CHEMICAL THERMODYNAMICS OF NICKEL · 2013-09-09 · of Inorganic and Analytical Chemistry,...

i

CHEMICAL THERMODYNAMICS OF

NICKEL

Heinz GAMSJÄGER (Chairman) Lehrstuhl für Physikalische Chemie

Montanuniversität Leoben Leoben (Austria)

Jerzy BUGAJSKI Lehrstuhl für Physikalische Chemie


Tamás GAJDA Department of Inorganic and Analytical Chemistry

University of Szeged Szeged (Hungary)

Robert J. LEMIRE Deep River

ON K0J 1P0 (Canada)

Wolfgang PREIS Lehrstuhl für Physikalische Chemie


Edited by Federico J. MOMPEAN (Series Editor and Project Co-ordinator)

Myriam ILLEMASSÈNE (Volume Editor) and Jane PERRONE OECD Nuclear Energy Agency, Data Bank

Issy-les-Moulineaux (France)

Err

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v

Preface

This volume is the sixth of the series "Chemical Thermodynamics" edited by the OECD Nuclear Energy Agency (NEA). It is a critical review of the Thermodynamic Data Base of nickel and its compounds initiated by the Management Board of the NEA Thermo-chemical Database Project Phase II (NEA TDB II).

The TDB Ni review team first met at the University of Leoben, Austria in January 1999. Three subsequent plenary meetings were held at NEA Headquarters at Issy-les-Moulineaux (France) in January 2001, October 2001 and in November 2002, two were held at Leoben in May 2002 and in November 2003. Smaller working sub-groups met at Luleå, Sweden in October 1999 and in April 2001. The Executive Group of the Management Board provided scientific assistance in the implementation of the NEA TDB Project Guidelines. Hans Wanner participated in meetings of the Review Team as the designated member of the Executive Group. At the NEA Data Bank the responsibility for the overall co-ordination of the Project was placed with Eric Östhols (from its initiation in 1999 to February 2000), with Stina Lundberg (from March 2000 to September 2000) and with Federico Mompean (since September 2000). Federico Mompean was in charge of the preparation of the successive drafts, updating the NEA thermodynamic database and editing the book to its present final form, with assistance of Myriam Illemassène, Jane Perrone, Katy Ben Said and Cristina Domènech-Ortí.

Originally Willis Forsling, Lars Gunneriusson (Luleå University of Technol-ogy, Sweden), Jerzy Bugajski, Heinz Gamsjäger, and Wolfgang Preis (University of Leoben, Austria) participated in the nickel project. In October 2001 Robert Lemire joined the Review Team. As member, chairman and peer reviewer of previous NEA TDB projects he contributed invaluable expertise to seeing the Ni review through to completion. In August 2002 time constraints and the pressure of other commitments forced Willis Forsling and Lars Gunneriusson to resign from the Review Team and Tamás Gajda, University of Szeged, Hungary, joined as a new member.

Although almost all of the members of the final review team contributed text and comments to several chapters, primary responsibility for the different chapters was divided as follows. Jerzy Bugajski prepared the sections on elemental nickel, silicon compounds and complexes, germanium compounds and complexes and, with the chairman, the section on boron compounds and complexes. Tamás Gajda drafted the sections for halide and pseudohalide complexes, the section for hydroxo complexes, and the sections on nitrato and thiocyanato complexes. Robert Lemire prepared Chapter VI, Discussion of auxiliary data selection, the sections on halates, sulphates, phosphorus compounds and complexes, arsenic compounds and complexes and, with the chairman, the section on solid halides (he also extensively reviewed several of the other sections, and revised the English throughout the text). Wolfgang Preis prepared the sections on the oxide, the sulphides and the hydrogensulphido complexes and, with the chairman,

PREFACE

vi

the section on Ni(III,IV) hydroxides (for the Leoben group he also implemented a data base for ionic strength corrections and calculations ensuring internal consistency of the selected values). The chairman drafted the sections on aqua ions, hydroxides and car-bonates.

The initial contributions from Willis Forsling and Lars Gunneriusson on halide and pseudohalide complexes and (together with Wolfgang Preis) on hydroxo complexes are gratefully acknowledged as they constituted the starting point for subsequent discus-sions within the Review Team.

While there is no need to repeat the general purpose of this review a side effect deserves mentioning. The selection of key values, e.g., for Ni2+(aq), revealed gaps in our knowledge which may stimulate rewarding projects on the experimental thermodynam-ics of nickel compounds.

Leoben, Austria, October 2004 Heinz Gamsjäger, Chairman

vii

Acknowledgements

For the preparation of this book, the authors have received financial support from the NEA TDB Phase II Project. The following organisations take part in the Project:

ANSTO, Australia ONDRAF/NIRAS, Belgium RAWRA, Czech Republic POSIVA, Finland ANDRA, France IPSN (now IRSN), France FZK, Germany JNC, Japan ENRESA, Spain SKB, Sweden SKI, Sweden HSK, Switzerland NAGRA, Switzerland PSI, Switzerland BNFL, UK NIREX, UK DoE, USA

Jerzy Bugajski, Heinz Gamsjäger and Wolfgang Preis would like to express their gratitude to the Institut für Physikalische Chemie of the Montanuniversität Leoben for having provided the infrastructure necessary for their contributions to this project.

Tamás Gajda gratefully acknowledges the technical support of the Department of Inorganic and Analytical Chemistry, University of Szeged.

Robert Lemire wishes to thank Atomic Energy of Canada Limited (AECL) for allowing him to participate in this project, and for allowing him to use the library and other AECL facilities at the Chalk River Laboratories.

The authors also thank Werner Sitte for his help with and comments on the section of silicon compounds and complexes. Heinz Gamsjäger thanks the library staff at the University of Leoben and the Austrian Central Library for Physics for their litera-ture searches and assistance in obtaining copies of many references, he also appreciates the help of Malcolm Rand, Alan Dinsdale, Heiko Kleykamp and Herbert Ipser for their

ACKNOWLEDGEMENTS

viii

advice concerning papers difficult to come by.

The unceasing efforts of Federico Mompean, coordinator for the TDB project during the time the main part of the draft of this book was assembled, edited for peer review and finally prepared for the press, is greatly appreciated. His work built on the earlier efforts of Erik Östhols and Stina Lundberg. The authors are indebted to Myriam Illemassène, Jane Perrone and Katy Ben Said who, with admirable alertness and scien-tific competence, transformed contributions of five authors, prepared in many different formats, into a consistent text with correctly numbered tables and figures. At the NEA Data Bank, Pierre Nagel and Eric Lacroix have provided excellent software and advice, which have eased the editorial and database work. Cynthia Picot, Solange Quarmeau and Amanda Costa from NEA Publications have provided considerable help in editing the present series. Their contributions and the support of many NEA staff members are highly appreciated.

The entire manuscript of this book has undergone a peer review by an inde-pendent international group of reviewers, according to the procedures in the TDB-6 Guideline, available from the NEA. The peer reviewers have seen and approved the modifications made by the authors in response to their comments. The peer review comment records may be obtained on request from the OECD Nuclear Energy Agency. The peer reviewers were:

Dr. Donald A. Palmer, Batelle, Oak Ridge National Laboratory, USA and

Prof. Wolfgang Voigt, University of Freiberg, Germany.

Their contributions are gratefully acknowledged.

ix

Note from the Chairman of the NEA TDB Project Phase II

The need to make available a comprehensive, internationally recognised and quality-assured chemical thermodynamic database that meets the modelling requirements for the safety assessment of radioactive waste disposal systems prompted the Radioactive Waste Management Committee (RWMC) of the OECD Nuclear Energy Agency (NEA) to launch in 1984 the Thermochemical Database Project (NEA TDB) and to foster its continuation as a semi-autonomous project known as NEA TDB Phase II in 1998.

The RWMC assigned a high priority to the critical review of relevant chemical thermodynamic data of inorganic species and compounds of the actinides uranium, neptunium, plutonium and americium, as well as the fission product technetium. The first four books in this series on the chemical thermodynamics of uranium, americium, neptunium and plutonium, and technetium originated from this initiative.

The organisation of Phase II of the TDB Project reflects the interest in many OECD/NEA member countries for a timely compilation of the thermochemical data that would meet the specific requirements of their developing national waste disposal pro-grammes.

The NEA TDB Phase II Review Teams, comprising internationally recognised experts in the field of chemical thermodynamics, exercise their scientific judgement in an independent way during the preparation of the review reports. The work of these Review Teams has also been subjected to further independent peer review.

Phase II of the TDB Project consisted of: (i) updating the existing, CODATA-compatible database for inorganic species and compounds of uranium, neptunium, plu-tonium, americium and technetium, (ii) extending it to include selected data on inor-ganic species and compounds of nickel, selenium and zirconium, (iii) and further adding data on organic complexes of citrate, oxalate, EDTA and iso-saccharinic acid (ISA) with uranium, neptunium, plutonium, americium, technetium, nickel, selenium, zirco-nium and some other competing cations.

The NEA TDB Phase II objectives were formulated by the 17 participating or-ganisations coming from the fields of radioactive waste management and nuclear regu-lation. The TDB Management Board is assisted for technical matters by an Executive Group of experts in chemical thermodynamics. In this second phase of the Project, the

NOTE FROM THE CHAIRMAN OF THE NEA TDB PROJECT PHASE II

x

NEA acts as coordinator, ensuring the application of the Project Guidelines and liaising with the Review Teams.

The present volume is the second one published within the scope of NEA TDB Phase II and contains a database for inorganic species and compounds of nickel. We trust that the efforts of the reviewers, the peer reviewers and the NEA Data Bank staff merit the same high recognition from the broader scientific community as received for previous volumes of this series.

Mehdi Askarieh United Kingdom Nirex limited Chairman of TDB Project Phase II Management Board On behalf of the NEA TDB Project Phase II Participating Organisations:

ANSTO, Australia ONDRAF/NIRAS, Belgium RAWRA, Czech Republic POSIVA, Finland ANDRA, France IPSN (now IRSN), France FZK, Germany JNC, Japan ENRESA, Spain SKB, Sweden SKI, Sweden HSK, Switzerland PSI, Switzerland BNFL, UK Nirex, UK DoE, USA

xi

Editor’s note

This is the sixth volume of a series of expert reviews of the chemical thermodynamics of key chemical elements in nuclear technology and waste management. This volume is devoted to the inorganic species and compounds of nickel. The tables contained in Chapters III and IV list the currently selected thermodynamic values within the NEA TDB Project. The database system developed at the NEA Data Bank, see Section II.6, assures consistency among all the selected and auxiliary data sets.

The recommended thermodynamic data are the result of a critical assessment of published information.. The values in the auxiliary data set, see Tables IV-1 and IV-2 have been adopted from CODATA key values or have been critically reviewed in this or earlier volumes of the series.

xii

How to contact the NEA TDB Project

Information on the NEA and the TDB Project, on-line access to selected data and computer programs, as well as many documents in electronic format are available at

www.nea.fr.

To contact the TDB project coordinator and the authors of the review reports, send comments on the TDB reviews, or to request further information, please send e−mail to [email protected]. If this is not possible, write to:

TDB project coordinator OECD Nuclear Energy Agency, Data Bank Le Seine-St. Germain 12, boulevard des Îles F-92130 Issy-les-Moulineaux FRANCE

The NEA Data Bank provides a number of services that may be useful to the reader of this book.

• The recommended data can be obtained via internet directly from the NEA Data Bank.

• The NEA Data Bank maintains a library of computer programs in vari-ous areas. This includes geochemical codes such as PHREEQE, EQ3/6, MINEQL, MINTEQ and PHRQPITZ, in which chemical thermody-namic data like those presented in this book are required as the basic in-put data. These computer codes can be obtained on request from the NEA Data Bank.

xiii

Contents

Preface v

Acknowledgement vii

Note from the chairman of the NEA TDB Project Phase II ix

Editor’s note xi

Part I Introductory material 1

I INTRODUCTION 3

I.1 Background .........................................................................................................3 I.2 Focus of the review .............................................................................................5 I.3 Review procedure and results............................................................................6

II STANDARDS, CONVENTIONS, AND CONTENTS OF THE TABLES 11

II.1 Symbols, terminology and nomenclature .......................................................11 II.1.1 Abbreviations ..............................................................................................11 II.1.2 Symbols and terminology............................................................................13 II.1.3 Chemical formulae and nomenclature .........................................................15 II.1.4 Phase designators.........................................................................................15 II.1.5 Processes .....................................................................................................17 II.1.6 Equilibrium constants ..................................................................................18

II.1.6.1 Protonation of a ligand...........................................................................18 II.1.6.2 Formation of metal ion complexes.........................................................19 II.1.6.3 Solubility constants................................................................................20 II.1.6.4 Equilibria involving the addition of a gaseous ligand............................21 II.1.6.5 Redox equilibria.....................................................................................22

II.1.7 pH................................................................................................................25 II.1.8 Order of formulae ........................................................................................27 II.1.9 Reference codes...........................................................................................28

II.2 Units and conversion factors............................................................................28 II.3 Standard and reference conditions .................................................................31

II.3.1 Standard state ..............................................................................................31 II.3.2 Standard state pressure ................................................................................32 II.3.3 Reference temperature.................................................................................35

CONTENTS

xiv

II.4 Fundamental physical constants......................................................................35 II.5 Uncertainty estimates .......................................................................................36 II.6 The NEA-TDB system ......................................................................................36 II.7 Presentation of the selected data .....................................................................38

Part II Tables of selected data 41

III SELECTED NICKEL DATA ...............................................................................43

IV SELECTED AUXILIARY DATA ........................................................................53

Part III Discussion of data selection 73

V DISCUSSION OF DATA SELECTION FOR NICKEL ..................................75

V.1 Elemental nickel........................................................................................................... 75

V.1.1 Nickel gas ....................................................................................................75 V.1.2 Nickel crystal...............................................................................................76 V.1.3 Nickel liquid ................................................................................................78

V.2 Simple nickel aqua ions ....................................................................................79 V.2.1 Ni2+ ..............................................................................................................79

V.2.1.1 Gibbs energy of formation of Ni2+ .........................................................80 V.2.1.1.1 Temperature coefficient of the standard electrode potential Ni2+ | Ni

..........................................................................................................82 V.2.1.2 Enthalpy of formation of Ni2+ ................................................................82 V.2.1.3 Partial molar entropy of Ni2+ .................................................................84

V.2.1.3.1 NiSO4·7H2O(cr) cycle.......................................................................84 V.2.1.3.2 NiCl2·6H2O(cr) cycle ........................................................................85 V.2.1.3.3 Entropy of Ni2+ from its f mGο∆ and f mH ο∆ ......................................86

V.2.1.4 Heat capacity of Ni2+ .............................................................................86 V.2.2 Other oxidation states ..................................................................................88

V.2.2.1 Ni3+ ........................................................................................................88 V.3 Oxygen and hydrogen compounds and complexes ........................................90

V.3.1 Aqueous nickel(II) hydroxo complexes.......................................................90 V.3.1.1 Hydroxo complexes in acidic or neutral solutions .................................90 V.3.1.2 Hydroxo complexes in alkaline solutions ..............................................99

V.3.2 Solid nickel oxides and hydroxides ...........................................................102 V.3.2.1 Ni(II) oxide ..........................................................................................102

V.3.2.1.1 Solubility measurements.................................................................107 V.3.2.2 Ni(II) hydroxides, Ni(OH)2(cr) ............................................................108

V.3.2.2.1 Crystallography and mineralogy of nickel hydroxide.....................108

CONTENTS

xv

V.3.2.2.1.1 β-Ni(OH)2 ...................................................................................108 V.3.2.2.1.2 α-Ni(OH)2 ...................................................................................109

V.3.2.2.2 Heat capacity and entropy of Ni(OH)2(cr) ......................................109 V.3.2.2.3 Thermodynamic analysis of solubility data of Ni(OH)2(s) .............111

V.3.2.3 Ni(III, IV) hydroxides..........................................................................115 V.3.2.3.1 β-NiOOH ........................................................................................116 V.3.2.3.2 γ-NiOOH.........................................................................................118

V.4 Group 17 (halogen) compounds and complexes...........................................119 V.4.1 Nickel halide compounds ..........................................................................119

V.4.1.1 Solid nickel fluoride, NiF2(cr) .............................................................119 V.4.1.1.1 Enthalpy of formation of NiF2(cr) ..................................................120

V.4.1.2 Solid nickel chloride, NiCl2(cr) ...........................................................122 V.4.1.2.1 Low-temperature heat capacity and entropy of NiCl2(cr) ...............122 V.4.1.2.2 The high-temperature heat capacity of NiCl2(cr)............................124 V.4.1.2.3 Enthalpy of formation of NiCl2(cr).................................................124 V.4.1.2.4 Gibbs energy of formation of NiCl2(cr) ..........................................127

V.4.1.3 Hydrated NiCl2 solids ..........................................................................128 V.4.1.3.1 Gibbs energy of formation of NiCl2·6H2O......................................128 V.4.1.3.2 Enthalpy of formation of NiCl2·6H2O.............................................128 V.4.1.3.3 The entropy and heat capacity of NiCl2·6H2O(cr) ..........................129 V.4.1.3.4 Gibbs energy and enthalpy of formation of nickel chloride

tetrahydrate .....................................................................................130 V.4.1.3.5 The entropy and heat capacity of nickel chloride dihydrate ...........131 V.4.1.3.6 Gibbs energy and enthalpy of formation of nickel chloride dihydrate

........................................................................................................131 V.4.1.4 Solid nickel bromide, NiBr2(cr) ...........................................................132

V.4.1.4.1 Low-temperature heat capacity and entropy of NiBr2(cr)...............132 V.4.1.4.2 The high-temperature heat capacity of NiBr2(cr)............................134 V.4.1.4.3 Enthalpy of formation of NiBr2(cr).................................................134

V.4.1.5 Hydrated solid, NiBr2·xH2O.................................................................135 V.4.1.6 Solid nickel iodide, NiI2(cr) .................................................................136

V.4.1.6.1 Entropy of NiI2(cr)..........................................................................136 V.4.1.6.2 Heat capacity of NiI2(cr).................................................................136 V.4.1.6.3 Enthalpy of formation of NiI2(cr) ...................................................137

V.4.2.7 Solid nickel iodate compounds ............................................................137 V.4.2 Aqueous nickel halide complexes .............................................................140

V.4.2.1 Introduction..........................................................................................140 V.4.2.2 Solution structural studies....................................................................141 V.4.2.3 Aqueous Ni(II) - fluoro complexes......................................................142 V.4.2.4 Aqueous Ni(II) - chloro complexes .....................................................146 V.4.2.5 Aqueous Ni(II) – bromo complexes.....................................................153

V.4.2.5.1 Determination of the Ni2+ – Br– ion interaction coefficient ............156

CONTENTS

xvi

V.4.2.6 Aqueous Ni(II) – iodine complexes .....................................................157 V.4.2.6.1 Aqueous Ni(II) – iodo complexes...................................................157

V.4.3 Determination of the Ni2+ – 4ClO− ion interaction coefficient..................158 V.5 Group 16 (chalcogen) compounds and complexes .......................................161

V.5.1 Sulphur compounds and complexes ..........................................................161 V.5.1.1 Nickel sulphides...................................................................................161

V.5.1.1.1 Solid nickel sulphides .....................................................................161 V.5.1.1.1.1 Crystallography and mineralogy of nickel sulphides ..................161

V.5.1.1.1.1.1 Heazlewoodite, Ni3S2(cr) ...................................................161 V.5.1.1.1.1.2 α-Ni7S6...............................................................................161 V.5.1.1.1.1.3 Pentlandite, (Fe, Ni)9S8(cr) ................................................161 V.5.1.1.1.1.4 Godlevskite, Ni9S8(cr)........................................................162 V.5.1.1.1.1.5 Millerite, β-NiS..................................................................162 V.5.1.1.1.1.6 Polydymite, Ni3S4(cr).........................................................162 V.5.1.1.1.1.7 Vaesite, NiS2(cr) ................................................................163

V.5.1.1.1.2 Ni – S phase diagram ..................................................................163 V.5.1.1.1.2.1 Ni – S melt .........................................................................166 V.5.1.1.1.2.2 β1-Ni3S2 and β2-Ni4S3 ........................................................167 V.5.1.1.1.2.3 Ni3S2(cr) (heazlewoodite) ..................................................167 V.5.1.1.1.2.4 Ni7S6(cr) and Ni9S8(cr).......................................................168 V.5.1.1.1.2.5 NiS(α) and NiS(β) .............................................................170 V.5.1.1.1.2.6 NiS2(cr), vaesite .................................................................173 V.5.1.1.1.2.7 Ni3S4(cr), polydymite.........................................................174 V.5.1.1.1.2.8 Ni4.5Fe4.5S8(cr), pentlandite ................................................174

V.5.1.1.1.3 Discussion of selected thermodynamic properties of nickel sulphides......................................................................................174

V.5.1.1.2 Solubility of NiS(s) and aqueous nickel hydrogen sulphido species........................................................................................................177

V.5.1.2 Nickel sulphates...................................................................................181 V.5.1.2.1 Aqueous nickel (II) sulphato complexes.........................................181

V.5.1.2.1.1 Enthalpy of formation of NiSO4(aq) ...........................................188 V.5.1.2.1.2 Heat capacity values....................................................................190

V.5.1.2.2 Solid nickel sulphates .....................................................................190 V.5.1.2.2.1 NiSO4·7H2O(cr)...........................................................................190

V.5.1.2.2.1.1 α-NiSO4·6H2O ...................................................................191 V.5.1.2.2.1.2 Other hydrated nickel sulphate solids ................................192 V.5.1.2.2.1.3 NiSO4(cr) ...........................................................................193

V.5.2 Selenium compounds and complexes........................................................195 V.5.3 Tellurium compounds and complexes .......................................................195

V.5.3.1 Nickel tellurides...................................................................................195 V.5.3.1.1 NiTe0.775(γ1).....................................................................................195 V.5.3.1.2 NiTe0.85(γ2) ......................................................................................195

CONTENTS

xvii

V.6 Group 15 compounds and complexes............................................................196 V.6.1 Nitrogen compounds and complexes.........................................................196

V.6.1.1 Solid nickel nitrates .............................................................................196 V.6.1.1.1 Ni(NO3)2·9H2O(cr) .........................................................................196 V.6.1.1.2 Ni(NO3)2·6H2O(cr) .........................................................................196 V.6.1.1.3 Ni(NO3)2·4H2O(cr), Ni(NO3)2·3H2O(cr), Ni(NO3)2·2H2O(cr) ........198 V.6.1.1.4 Ni(NO3)2(anhydrous) ......................................................................200

V.6.1.2 Aqueous Ni(II)-nitrato complexes .......................................................200 V.6.1.2.1 Determination of the Ni2+ – 3NO− ion interaction coefficient........203

V.6.2 Phosphorus compounds and complexes ....................................................204 V.6.2.1 Solid nickel phosphorus compounds....................................................204

V.6.2.1.1 Nickel phosphides...........................................................................204 V.6.2.1.2 Nickel phosphates ...........................................................................204

V.6.2.2 Aqueous nickel phosphorus species.....................................................205 V.6.2.2.1 Simple nickel phosphato complexes ...............................................205 V.6.2.2.2 Aqueous nickel diphosphato complexes .........................................207

V.6.3 Arsenic compounds ...................................................................................210 V.6.3.1 Nickel arsenides...................................................................................210

V.6.3.1.1 NiAs(cr) ..........................................................................................210 V.6.3.1.2 Ni11As8(cr) ......................................................................................211 V.6.3.1.3 Ni5As2(cr) .......................................................................................212 V.6.3.1.4 NiAs2(cr).........................................................................................212

V.6.3.2 Nickel arsenates ...................................................................................212 V.6.3.2.1 Solid nickel arsenates......................................................................212 V.6.3.2.2 Aqueous nickel arsenato complexes ...............................................213

V.6.3.3 Nickel arsenites....................................................................................214 V.6.3.3.1 Solid nickel arsenites ......................................................................214

V.7 Group 14 compounds and complexes............................................................214 V.7.1 Carbon compounds and complexes ...........................................................214

V.7.1.1 Nickel carbonates.................................................................................214 V.7.1.1.1 Solid nickel carbonates ...................................................................214

V.7.1.1.1.1 NiCO3(cr) ....................................................................................214 V.7.1.1.1.1.1 Crystallography and mineralogy of nickel carbonate.........214 V.7.1.1.1.1.2 Heat capacity and entropy of NiCO3(cr) ............................216 V.7.1.1.1.1.3 Thermodynamic analysis of solubility data on gaspéite ....217

V.7.1.1.1.2 NiCO3·5.5H2O(cr) ....................................................................................... 218 V.7.1.1.1.2.1 Crystallography and mineralogy of nickel carbonate hydrate .

...........................................................................218 V.7.1.1.1.2.2 Heat capacity and entropy of NiCO3·5.5H2O(cr) ............219 V.7.1.1.1.2.3 Thermodynamic analysis of solubility data on hellyerite ..219

V.7.1.1.2 Aqueous nickel carbonato complexes.............................................223

CONTENTS

xviii

V.7.1.2 Nickel cyanides....................................................................................226 V.7.1.2.1 Aqueous Ni(II) cyano complexes ...................................................226

V.7.1.2.1.1 Complexes in neutral and alkaline solution.................................226 V.7.1.2.1.2 Protonated complexes .................................................................231

V.7.1.3 Nickel thiocyanates..............................................................................231 V.7.1.3.1 Aqueous nickel thiocyanato complexes ..........................................231

V.7.2 Silicon compounds and complexes............................................................238 V.7.2.1 Solid nickel silicates ............................................................................238

V.7.2.1.1 Nickel orthosilicate Ni2SiO4(cr) .....................................................238 V.7.2.1.1.1 Crystal structure and phase transitions ........................................238 V.7.2.1.1.2 Crystallography and mineralogy of nickel orthosilicate..............239 V.7.2.1.1.3 Heat capacity and standard entropy.............................................239 V.7.2.1.1.4 Enthalpy of formation .................................................................241 V.7.2.1.1.5 Gibbs energy of formation ..........................................................243

V.7.3 Germanium compounds and complexes ....................................................245 V.7.3.1 Solid nickel germanates .......................................................................245

V.7.3.1.1 Nickel orthogermanate Ni2GeO4(cr) ...............................................245 V.7.3.1.1.1 Crystal structure of nickel orthogermanate .................................245 V.7.3.1.1.2 Enthalpy of formation from NiO(cr) and GeO2(cr).....................245

V.8 Group 13 compounds and complexes............................................................245 V.8.1 Boron compounds and complexes .............................................................245

V.8.1.1 Solid nickel borates..............................................................................245 V.8.1.1.1 Crystal structure..............................................................................245 V.8.1.1.2 Thermodynamic data for NiO·2B2O3(s)..........................................246 V.8.1.1.3 Solubility of Ni(BO2)2·4H2O(s) ......................................................246

V.8.1.2 Aqueous nickel borato complexes .......................................................247

VI DISCUSSION OF AUXILIARY DATA SELECTION ...............................249

VI.1 Group 15 auxiliary species .............................................................................249 VI.1.1 Additional consistent auxiliary data for As compounds ............................249

VI.2 Other auxiliary species ...................................................................................249 VI.2.1 Auxiliary data for KCl(cr), KBr(cr) and KI(cr).........................................249

Part IV Appendices 251

A Discussion of selected references .......................................................................253

B Ionic strength corrections ..................................................................................443

B.1 The specific ion interaction equations...........................................................445 B.1.1 Background ...............................................................................................445

CONTENTS

xix

B.1.2 Ionic strength corrections at temperatures other than 298.15 K ................451 B.1.3 Estimation of ion interaction coefficients..................................................453

B.1.3.1 Estimation from mean activity coefficient data ...................................453 B.1.3.2 Estimations based on experimental values of equilibrium constants at

different ionic strengths .......................................................................453 B.1.4 On the magnitude of ion interaction coefficients.......................................457

B.2 Ion interaction coefficients versus equilibrium constants for ion pairs .....458 B.3 Tables of ion interaction coefficients.............................................................458

C Assigned uncertainties ......................................................................................471

C.1 The general problem.......................................................................................471 C.2 Uncertainty estimates in the selected thermodynamic data. .......................473 C.3 One source datum ...........................................................................................474 C.4 Two or more independent source data .........................................................475

C.4.1 Discrepancies.............................................................................................477 C.5 Several data at different ionic strengths .......................................................479

C.5.1 Discrepancies or insufficient number of data points..................................481 C.6 Procedures for data handling ........................................................................483

C.6.1 Correction to zero ionic strength ...............................................................483 C.6.2 Propagation of errors .................................................................................485 C.6.3 Rounding ...................................................................................................486 C.6.4 Significant digits........................................................................................487

Bibliography ..............................................................................................................489

List of authors............................................................................................................585

List of Figures

Figure II-1: Standard order of arrangement of the elements and compounds based on the periodic classification of the elements ...........................................27

Figure V-1: The standard molar heat capacity of nickel as a function of temperature. ..............................................................................................................77

Figure V-2: Standard electrode potential of Ni | NiSO4 | Hg2SO4 | Hg at 25°C. .....81 Figure V-3: Estimation of the Ni3+ | Ni2+ standard electrode potential. ..................89 Figure V-4: Extrapolation to Im = 0 of the experimental data for Reaction (V.23) in

NaClO4 media.......................................................................................96 Figure V-5: SIT analysis of the experimental data for Reaction (V.23) in chloride

media, including 10 1,1*log οb ((V.23), 298.15 K) obtained in NaClO4....96

Figure V-6: Extrapolation to Im = 0 of the experimental data for Reaction (V.25) in perchlorate media. ................................................................................98

Figure V-7: Solubility of Ni(OH)2(s) in alkaline solutions at 298.15 K. ..............100 Figure V-8: Experimental results of heat capacity measurements on nickel oxide

from 3.2 to 1809.7 K. ........................................................................104 Figure V-9: Comparison between the temperature dependence of the Gibbs energy

of formation for NiO obtained from electrochemical as well as chemical reduction/oxidation equilibrium measurements and the prediction based on the present selection of thermodynamic properties for NiO. .........106

Figure V-10: The heat capacity of Ni(OH)2(cr) as a function of temperature .........110 Figure V-11: Solubility constant of Ni(OH)2 as a function of temperature. ............114 Figure V-12: β-Ni(OH)2 | β-NiOOH system. Variation of reversible potential with the

oxidation state of Ni at 298.15 K and I = 0. ......................................116 Figure V-13: α-Ni(OH)2 | γ-NiOOH system. Variation of reversible potential with the

oxidation state of Ni at 298.15 K and I = 0. ......................................118 Figure V-14: Molar heat capacity of anhydrous NiCl2 between 2 and 30 K............123 Figure V-15: Heat capacity of anhydrous NiCl2(cr) ................................................123 Figure V-16: r m

ο∆ G –function of Reaction (V.59)....................................................125 Figure V-17: Third-law analysis of data for Reaction (V.59). . ...............................126 Figure V-18: Gibbs energy of formation of NiCl2(cr) versus temperature. ............127 Figure V-19: Values from the heat capacity studies of Stuve et al. [78STU/FER] and

White and Staveley [82WHI/STA]. ...................................................133

LIST OF FIGURES

xxii

Figure V-20: Extrapolation to Im = 0 of the experimental data for Reaction (V.70) in NaClO4 media. ...................................................................................143

Figure V-21: Extrapolation to Im = 0 of the experimental enthalpies for Reaction (V.70) in NaClO4 media. ....................................................................145

Figure V-22: Extrapolation to I = 0 of the 10 1log b values reported in [75LIB/TIA] for Reaction (V.71) in Ni(II) perchlorate solution. ............................149

Figure V-23: Extrapolation to I = 0 of the accepted experimental data for Reaction (V.71) in NaClO4 media. ....................................................................150

Figure V-24: Extrapolation to I = 0 of the 10 1log b values reported in [78LIB/KOW] for Reaction (V.73) in Ni(II) perchlorate media.................................155

Figure V-25: Plot of 10log ±γ versus molality of aqueous NiBr2 solutions at 298.15 K............................................................................................................157

Figure V-26: Plot of osmotic coefficient versus molality of aqueous NiBr2 solutions at 298.15 K. ...........................................................................................158

Figure V-27: Osmotic coefficient plotted as a function of molality of aqueous Ni(ClO4)2 solutions at 298.15 K. ........................................................160

Figure V-28: Plot of 10log ±γ versus molality of aqueous Ni(ClO4)2 solutions at 298.15 K. ............................................................................................160

Figure V-29: Phase diagram for the binary system Ni – S with T plotted versus mole fraction of sulphur. ............................................................................164

Figure V-30: Section of the calculated phase diagram for the composition range 0.30 < xS < 0.55. . ...............................................................................165

Figure V-31: Gibbs energy of reaction for (V.79) plotted versus temperature ........168 Figure V-32: Partial pressure of sulphur for various phase equilibria in the binary

system Ni – S plotted against temperature. .......................................170 Figure V-33: Gibbs energy function for the stoichiometric composition of NiAs-type

Ni(II) sulphide, α-NiS, plotted versus temperature. ..........................173 Figure V-34: Plot of equilibrium partial pressure of sulphur versus temperature ....176 Figure V-35: Distribution of nickel sulphide complexes as a function of total molality

of nickel (II) in aqueous solutions saturated with H2S, T = 298.15 K, I = 1.00 mol·kg–1 NaCl. ...........................................................................179

Figure V-36: Solubility of Ni(II) sulphides in aqueous solutions saturated with H2S as a function of initial concentration of hydrochloric acid at 298.15 K. .181

Figure V-37: The variation of molar conductances of aqueous nickel sulphate solutions with concentration at 25°C..................................................186

Figure V-38: Values of 10 1log K ο for nickel sulphate association at different temperatures .......................................................................................189

LIST OF FIGURES

xxiii

Figure V-39: Measured water vapour pressures over hydrated Ni(NO3)2·6H2O(cr) [37SAN], [72AUF/CAR] and Ni(NO3)2·4H2O(cr) [72AUF/CAR2]..199

Figure V-40: Extrapolation to I = 0 of the constants for Reaction (V.96) in lithium perchlorate media based on the data in [73FED/SHM]. . ...................202

Figure V-41: Plot of 10log γ ± versus molality of aqueous Ni(NO3)2 solutions at 298.15 K.............................................................................................203

Figure V-42: Osmotic coefficient plotted as a function of molality of aqueous Ni(NO3)2 solutions at 298.15 K. . .......................................................204

Figure V-43: Temperature dependence of hellyerite solubility. ..............................220 Figure V-44: Ionic strength dependence of hellyerite solubility... ...........................221 Figure V-45: Thermodynamic analysis of hellyerite solubility. ..............................222 Figure V-46: Three-dimensional predominance diagram for the system

Ni2+–CO2–H2O. ..................................................................................223 Figure V-47: calculated solubility of NiCO3(cr) in dilute sodium carbonate solutions

under atmospheric conditions.............................................................225 Figure V-48: Extrapolation to I = 0 of experimental formation constants of

24Ni(CN) − (V.113) determined in NaClO4 medium (insert shows the

corresponding plot of data reported for KClO4 medium). ..................229 Figure V-49: Extrapolation to I = 0 of experimental equilibrium constants for

Reaction (V.114) determined in NaClO4 medium..............................229 Figure V-50: Extrapolation to Im = 0 of the accepted formation constants reported for

the formation of NiSCN+ in perchlorate media. .................................235 Figure V-51: Extrapolation to Im = 0 of the accepted formation constants reported for

the formation of Ni(SCN)2(aq) in perchlorate media. ........................236 Figure V-52: Extrapolation to Im = 0 of the accepted formation constants reported for

the formation of 3Ni(SCN)− in perchlorate media..............................236 Figure V-53: The standard molar heat capacity of Ni2SiO4-olivine as a function of

temperature.........................................................................................240 Figure V-54: Standard Gibbs energy of Reaction (V.121) as a function of

temperature.........................................................................................244 Figure A-1: Solubility of precipitated nickel carbonate.........................................268 Figure A-2: High temperature enthalpy of anhydrous NiCl2. ................................274 Figure A-3: Plot of E (vs. SHE) versus log10([NiSO4]/m)......................................276 Figure A-4: Extrapolation to Im = 0 (using the SIT) of the formation constants for the

species NiSCN+ as reported in [58YAT/KOR] . ................................290 Figure A-5: Extrapolation of activity coefficients to saturated solution at 25°C. .295 Figure A-6: Extrapolation of osmotic coefficients to saturated solution at 25°C...295

LIST OF FIGURES

xxiv

Figure A-7: Plot of 10 1log b (A.17) reported in [61LIS/WIL] vs. 1/T...................298 Figure A-8: 10 1log b – Im plot obtained from the potentiometric data (t = 35°C)

reported for the Ni2+ – SCN– system in [61MOH/DAS]. ...................299 Figure A-9: Solubility product of Ni(BO2)2·4H2O(s) at 22°C from data reported in

[61SHC]. ............................................................................................300 Figure A-10: Nickel and cobalt borate complexes from data in [61SHC]. ..............301 Figure A-11: Temperature dependence of 10 1,1

*log b values derived from [63BOL/JAU].....................................................................................306

Figure A-12: Temperature dependence of 10 4,4*log b values derived from

[63BOL/JAU].....................................................................................307 Figure A-13: Extrapolation to I = 0 of the equilibrium constants 10 1,1

*log b (A.23) reported in [64PER] at 20°C, using the SIT. ......................................313

Figure A-14: Plot of 1,1*ln οb (A.23) recalculated from [64PER] vs. 1/T...................314

Figure A-15: Temperature dependence of the association constant for the reaction 24Ni(CN) − + CN– 3

5Ni(CN) − reported in [60MCC/JON]. ...........320 Figure A-16: Molar enthalpy of solution of NiSO4·6H2O(cr) at 25°C. . ..................325 Figure A-17: Temperature dependence of 10 1log b values of NiSCN+ complex

reported in [68MAL/TUR]. ................................................................333 Figure A-18: Heats of reaction determined from the calorimetric data of [69IZA/EAT]

by using the 298.15 K values of 10 1log K from [73KAT]. ................336 Figure A-19: Extrapolation to I = 0 of the formation constant of NiCl+ using Equation

(A.40) and the experimental data published in [70HAL/VAN]. ........343 Figure A-20: Temperature dependence of 10 1log b and 10 2log b values for the

Ni(II) – (SCN)– system (I = 0.2 M) as reported in [71TUR/MAL]. ...351 Figure A-21: Heats of dilution of hydrated nickel nitrate salts based on the

experimental heats of solution and an assumed heat of dilution to infinite dilution of 520 J·mol–1 for a solution with a solvent:salt ratio of 8000:1.................................................................................................352

Figure A-22: Experimental [73FED/SHM] and calculated solubility of nickel(II) in the aqueous phase with increasing nitrate concentrations. .................358

Figure A-23: Experimental [73FED/SHM] and calculated values of the solubility of Ni(IO3)2·3H2O in aqueous LiClO4 solutions.......................................358

Figure A-24: Variation of formation constants from Perlmutter-Hayman and Secco [73PER/SEC] with ionic strength.......................................................363

Figure A-25: Temperature dependence of 10 2log K for the reaction 24Ni(CN) − + H+

3Ni(HCN)(CN)− at different ionic strengths using NaCl. .............367

LIST OF FIGURES

xxv

Figure A-26: Extrapolation to I = 0 of experimental equilibrium constants for the reaction 3Ni(HCN)(CN)− + H+ 2 2Ni(HCN) (CN) (aq) at different temperatures.. .....................................................................................367

Figure A-27: r mH ο∆ (298 K) of reaction of B2O3(cr) with MO or M2O. ..................370 Figure A-28: Heat capacity of NiS2 plotted versus temperature. .............................382 Figure A-29: SIT-plot of recalculated data from [80MIL/BUG], [65BUR/LIL2] and

[71OHT/BIE], and the selected 10 4,4*log οb ((A.70), 298.15 K) ..........391

Figure A-30: Solubility of NiO in acidic solutions as a function of initial molality of HCl. . ..................................................................................................394

Figure A-31: Solubility of NiO in basic solutions as a function of initial molality of NaOH. ...............................................................................................395

Figure A-32: Values of ,mpC (NiBr2, cr) as a function of temperature as reported by White and Staveley [82WHI/STA]. ...................................................400

Figure A-33: Plot of logarithm of the oxygen partial pressure versus temperature for the phase equilibria α−NiS / β2-Ni4S3 and β2-Ni4S3 / NiO, respectively, at

2SOp = 1 atm. .................................................................................407

Figure A-34: Logarithm of the oxygen partial pressure plotted as a function of temperature for the equilibrium NiS(α) / NiO at p(SO2) = 1 atm.......408

Figure A-35: Plot of logarithm of the oxygen partial pressure versus temperature for the phase equilibria NiS(α) / β2-Ni4S3 and β2-Ni4S3 / NiO, respectively, at

2SOp = 1 atm ...................................................................................419

Figure A-36: Heat capacity data for Ni3S2 as a function of temperature..................421 Figure A-37: Heat capacity of “Ni7S6” consisting of a mixture of Ni3S2 and Ni9S8 as a

function of temperature. .....................................................................427 Figure A-38: Heat capacity of NiS as a function of temperature .............................429 Figure A-39: Logarithm of stability constants of nickel bisulphide complexes in

seawater (NaCl solutions) plus the Debye-Hückel term for ionic strength correction plotted as a function of ionic strength, T = 298.15 K.. ......432

Figure A-40: Solubility of theophrastite as determined by the pH variation method.. ..............................................................................................................440

Figure A-41: Ionic strength dependence of theophrastite solubility. .......................440

Figure B-1: Plot of 10 1,log 4m D+b versus Im for reaction (B.12), at 25°C and 1 bar.. ...........................................................................................455

xxvii

List of Tables

Table II-1: Abbreviations for experimental methods .................................................11

Table II-2: Symbols and terminology. .......................................................................13 Table II-3: Abbreviations used as subscripts of ∆ to denote the type of chemical

process. ....................................................................................................17 Table II-4: Unit conversion factors ............................................................................28 Table II-5: Factors for the conversion of molarity, cB, to molality, mB, of a substance

B, in various media at 298.15 K ..............................................................30 Table II-6: Reference states for some elements at the reference temperature of

298.15 K and standard pressure of 0.1 MPa. ...........................................31 Table II-7: Fundamental physical constants...............................................................36

Table III-1: Selected thermodynamic data for nickel compounds and complexes ......44 Table III-2: Selected thermodynamic data for reaction involving nickel compounds and

complexes ................................................................................................48 Table III-3: Selected temperature coefficients for heat capacities of nickel ...............52

Table IV-1: Selected thermodynamic data for auxiliary compounds and complexes..55 Table IV-2: Selected thermodynamic data for reactions involving auxiliary compounds

and complexes..........................................................................................68

Table V-1: The temperature coefficients of the heat capacity function for Ni(cr) .....78 Table V-2: Enthalpy of Reaction (V.3) ......................................................................83 Table V-3: Partial molar heat capacity values for Ni(II) salts in aqueous solution at

298.15 K...................................................................................................88 Table V-4: Experimental equilibrium data (logarithmic values) for the hydrolysis of

Ni(II) in acidic or neutral solutions. .........................................................93 Table V-5: Smoothed ,mpCο data and calculated values for the entropy of NiO. .....103 Table V-6: Literature data for 10 ,0

*log sK ο (Ni(OH)2)................................................113 Table V-7: Thermodynamic properties of Ni(OH)2(cr)............................................115 Table V-8: Experimental stability constants (logarithmic values) of the species NiF+. .

...............................................................................................................142 Table V-9: Experimental enthalpy values for the Reaction (V.70). .........................144

LIST OF TABLES

xxviii

Table V-10: Experimental formation constants for the NiCl+ species. ......................147 Table V-11: Experimental equilibrium constants for the Reaction (V.72).................152 Table V-12: Reaction enthalpies reported for Reaction (V.71)..................................153 Table V-13: Experimental formation constants of the NiBr+ complex. .....................154 Table V-14: Reaction enthalpies reported for Reaction (V.73)..................................156 Table V-15: Gibbs energy functions for the binary system Ni – S relative to metallic

nickel, Ni(s), and gaseous sulphur, S2(g). ..............................................165 Table V-16: Temperatures and enthalpies of transition for β-NiS α-NiS. ...........172 Table V-17: Temperatures for invariant three-phase equilibria in the binary system

Ni - S. .....................................................................................................175 Table V-18: Heat capacity functions of nickel sulphides. ..........................................176 Table V-19: Standard enthalpy of formation of nickel sulphides at 298.15 K. ..........177 Table V-20: Standard entropy of nickel sulphides at 298.15 K..................................177 Table V-21: Heat capacity of nickel sulphides at 298.15 K. ......................................177 Table V-22: Recalculated stability constants and SIT parameters for nickel sulphide

complexes at T = 298.15 K and I = 0. ....................................................178 Table V-23: Solubility constants for Ni(II) sulphide at T = 298.15 K and I = 0. .......180 Table V-24: Experimental values for the association constants and other

thermodynamic quantities for nickel sulphate. ......................................183 Table V-25: Results of reanalysis of NiSO4(aq) conductance data using a form of the

Lee-Wheaton equation that is compatible with the SIT procedure. .......185 Table V-26: Experimental values for the first association constant of nickel sulphate

determined in the presence of supporting electrolytes ..........................187 Table V-27: Formation constants of the 3NiNO+ complex (ion pair)........................201 Table V-28: Experimental values for the association constants and other

thermodynamic quantities for nickel phosphate complexes...................206 Table V-29: Experimental values for the association constants and other

thermodynamic quantities for nickel pyrophosphate complexes. ..........209 Table V-30: Experimental equilibrium data for the Ni(II) cyanide system................228 Table V-31: Experimental enthalpy values for the Reaction (V.113) and (V.114). ...230 Table V-32: Experimental equilibrium data for the Ni(II) – thiocyanate system. ......233 Table V-33: Experimental reaction enthalpy values for the formation of Ni(II)-

thiocyanate complexes. ..........................................................................237

Table VI-1: Enthalpy of formation values for KCl(cr), KBr(cr), and KI(cr) at 298.15 K.................................................................................................250

LIST OF TABLES

xxix

Table A-1: Equilibrium constants of Reaction (A.3) [21BER/CRU], [24CRU]. .....258 Table A-2: Equilibrium constants of NiCl2(cr) + H2(g) Ni(cr) + 2 HCl(g).........260 Table A-3: Electrode potential of the cell: Ni|NiSO4(m)|Hg2SO4|Hg.......................262 Table A-4: Standard electrode potential of Ni2+|Ni at 25°C. ....................................276 Table A-5: Equilibrium constants of Reaction (A.9)................................................278 Table A-6: Equilibrium constants of Reaction (A.9) [53SAN] ................................279 Table A-7: Enthalpy of formation values calculated from equilibrium gas

compositions. .........................................................................................282 Table A-8: Equilibrium constants of Reaction (A.10) [54SHC/TOL] .....................282 Table A-9: Enthalpy of solution of NiCl2(cr) in H2O(l) at 298.15 K [58MUL]. ......289 Table A-10: 10 1log b versus Im. .................................................................................304 Table A-11: Recalculated constants. ..........................................................................306 Table A-12: Comparison between experimental data [63LIN/LAF] and calculated

values of log10[p(H2S)/p(H2)] for the coexistence of Ni3+xS2 with metallic nickel......................................................................................................308

Table A-13: Reported and recalculated 10 1,1*log οb ....................................................314

Table A-14: Dataset. ..................................................................................................342 Table A-15: Experimental results for the enthalpy of formation for various nickel

sulphides determined by calorimetric measurements.............................345 Table A-16: Gibbs energy functions for various nickel sulphides. ............................385 Table A-17: Auxiliary data used in the recalculation of NaBr heat of formation values

with the data from [78STU/FER]...........................................................386 Table A-18: Corrected formation constants. ..............................................................390 Table A-19: Stability constants for the complexes NiOH+, Ni(OH)2 and 3Ni(OH)− .393 Table A-20: Stability constants 10 3,1

*log οb for 3Ni(OH)− , I = 0.................................395 Table A-21: Enthalpies and entropies of reaction from Skeaff et al.. ........................404 Table A-22: Enthalpies of formation for solid nickel sulphides at 298.15 K. ............406 Table A-23: Enthalpy of solution of NiCl2 in H2O at 298.15 K [90EFI/FUR]...........417 Table A-24: f mH ο∆ (NiCO3, 298.15 K) from decomposition data [91TAR/FAZ]......422 Table A-25: Enthalpy of solution of NiCl2 in H2O at 298.15 K [95MAN/KOR].......430 Table A-26: Recommended formation constants ( ,

*n mοb ) and reaction enthalpies for the

hydroxo complexes formed in the acidic region, at 298.15 K, mNi2+ + nH2O(l) 2Ni (OH) m n

m n− + nH+ .........................................................435

Table A-27: Recommended equilibrium constants ( , ,*

s n mK ο ) and reaction enthalpies for the hydroxo complexes formed in the alkaline region, at 298.15 K, mNi(OH)2(cr) + nOH– 2Ni (OH) n

m m n−

+ ............................................435

LIST OF TABLES

xxx

Table A-28: Recommended ion interaction coefficients ε(j, k). ................................436 Table A-29: Standard thermodynamic functions of the aqueous species. ..................436 Table B-1: Water activities

2H Oa at 2968.15 K for the most common ionic media at various concentrations applying Pitzer’s ion interaction approach and the interaction parameters given in [91PIT]. ...............................................448

Table B-2: Debye–Hückel constants as a function of temperature at a pressure of 1 bar below 100°C and at the steam saturated pressure for t ≥ 100°C.....452

Table B-3: The preparation of the experimental equilibrium constants for the extrapolation to I = 0 with the specific ion interaction method at 25°C and 1 bar, according to reaction (B.12).........................................................454

Table B-4: Ion interaction coefficients ( , )j kε (kg·mol–1) for cations j with k = Cl–, 4ClO− and 3NO− , taken from Ciavatta [80CIA], [88CIA] unless indicated

otherwise. ...............................................................................................460 Table B-5: Ion interaction coefficients ( , )j kε (kg·mol–1) for cations j with k = Li, Na

and K, taken from Ciavatta [80CIA], [88CIA] unless indicated otherwise................................................................................................................467

Table B-6: Ion interaction coefficients (1, , )j kε and (2, , )j kε for cations j with k = Cl–, 4ClO− and 3NO− (first part), and for anions j with k = Li+, Na+ and K+ (second part), according to the relationship ε = ε1 + ε2 log10Im. The data are taken from Ciavatta [80CIA], [88CIA]. .................................................470

Table B-7: SIT interaction coefficient ε(j,k) kg · mol–1 for neutral species, j, with k, electroneutral combination of ions.........................................................470

Table C-1: Details of the calculation of equilibrium constant corrected to I = 0, using (C.19) .....................................................................................................483

3

Chapter I

I Introduction

I.1 Background The modelling of the behaviour of hazardous materials under environmental conditions is among the most important applications of natural and technical sciences for the pro-tection of the environment. In order to assess, for example, the safety of a waste deposit, it is essential to be able to predict the eventual dispersion of its hazardous components in the environment (geosphere, biosphere). For hazardous materials stored in the ground or in geological formations, the most probable transport medium is the aqueous phase. An important factor is therefore the quantitative prediction of the reactions that are likely to occur between hazardous waste dissolved or suspended in ground water, and the surrounding rock material, in order to estimate the quantities of waste that can be transported in the aqueous phase. It is thus essential to know the relative stabilities of the compounds and complexes that may form under the relevant conditions. This infor-mation is often provided by speciation calculations using chemical thermodynamic data. The local conditions, such as ground water and rock composition or temperature, may not be constant along the migration paths of hazardous materials, and fundamental thermodynamic data are the indispensable basis for dynamic modelling of the chemical behaviour of hazardous waste components.

In the field of radioactive waste management, the hazardous material consists to a large extent of actinides and fission and activation products. Two isotopes of nickel (59Ni and 63Ni) pertain to the latter category. The scientific literature on thermodynamic data, mainly on equilibrium constants and redox potentials in aqueous solution, has been contradictory in a number of cases, especially in actinide chemistry. A critical and comprehensive review of the available literature is necessary in order to establish a reli-able thermochemical database that fulfils the requirements for rigorous modelling of the behaviour of the actinides, fission and activation products in the environment.

4 1 Introduction

The International Atomic Energy Agency (IAEA) in Vienna published special issues with compilations of physicochemical properties of compounds and alloys of elements important in reactor technology: Pu, Nb, Ta, Be, Th, Zr, Mo, Hf and Ti be-tween 1966 and 1983. In 1976, IAEA also started the publication of the series “The Chemical Thermodynamics of Actinide Elements and Compounds”, oriented towards nuclear engineers and scientists. This international effort has resulted in the publication of several volumes, each concerning the thermodynamic properties of a given type of compounds for the entire actinide series. These reviews cover the literature approxi-mately up to 1984. The latest volume in this series appeared in 1992, under Part 12: The Actinide Aqueous Inorganic Complexes [92FUG/KHO]. Unfortunately, data of impor-tance for radioactive waste management (for example, Part 10: The Actinide Oxides, or data on fission and activation products) is lacking in the IAEA series.

The Radioactive Waste Management Committee (RWMC) of the OECD Nu-clear Energy Agency recognised the need for an internationally acknowledged, high-quality thermochemical database for application in the safety assessment of radioactive waste disposal, and undertook the development of the NEA Thermochemical Data Base (TDB) project [85MUL], [88WAN], [91WAN]. The RWMC assigned a high priority to the critical review of relevant chemical thermodynamic data of compounds and com-plexes for this area containing the actinides uranium, neptunium, plutonium and ameri-cium, as well as the fission product technetium. The first four books in this series on the chemical thermodynamics of uranium [92GRE/FUG], americium [95SIL/BID], techne-tium [99RAR/RAN] and neptunium and plutonium [2001LEM/FUG] originated from this initiative. Simultaneously with the NEA’s TDB project, other reviews on the physi-cal and chemical properties of actinides appeared, including the book by Cordfunke et al. [90COR/KON2], the series edited by Freeman et al. [84FRE/LAN], [85FRE/LAN], [85FRE/KEL], [86FRE/KEL], [87FRE/LAN], [91FRE/KEL], the two volumes edited by Katz et al. [86KAT/SEA], and Part 12 by Fuger et al. [92FUG/KHO] within the IAEA review series mentioned above.

In 1998, Phase II of the TDB Project (TDB-II) was started to provide for the further needs of the radioactive waste management programs by updating the existing database and applying the TDB review methodology to other elements (occurring in waste as fission or activation products) and to simple organic complexes. In TDB-II the overall objectives are set by a Management Board, integrated by representatives of 17 organisations from the field of radioactive waste management. These participating or-ganisations, together with the NEA, provide financial support for TDB-II. The TDB-II Management Board is assisted in technical matters by a group of experts in chemical thermodynamics (the Executive Group). The NEA acts in this phase as Project Co-ordinator ensuring the implementation of the Project Guidelines and liaising with the Review Teams. The present volume, the sixth in the series, is the second one to be pub-lished within this second phase of the TDB Project. A previous volume dealt with an update on the chemical thermodynamics of uranium, neptunium, plutonium, americium

I.2 Focus of the review 5

and technetium [2003GUI/FAN]. Three additional volumes are in preparation as this book goes to press dealing with the inorganic species and compounds of selenium, the inorganic species and compounds of zirconium and with compounds and complexes formed by the organic ligands oxalate, citrate, EDTA and iso-saccharinic acid with U, Np, Pu, Am, Tc, Zr, Ni and Se.

I.2 Focus of the review The first and most important step in the modelling of chemical reactions is to decide whether they are controlled by chemical thermodynamics or kinetics, or possibly by a combination of the two. This also applies to the modelling of more complex chemical systems and processes, such as waste repositories of various kinds, the processes de-scribing transport of toxic materials in ground and surface water systems, the global geochemical cycles, etc.

As outlined in the previous section, the focus of the critical review presented in this report is on the thermodynamic data of nickel relevant to the safety assessment of radioactive waste repositories in the geosphere. This includes the release of waste com-ponents from the repository into the geosphere (i.e., its interaction with the waste con-tainer and the other near-field materials) and their migration through the geological formations and the various compartments of the biosphere. As ground waters and pore waters are the transport media for the waste components, the knowledge of the thermo-dynamics of the corresponding elements in waters of various compositions is of funda-mental importance.

The present review therefore puts much weight on the assessment of the low-temperature thermodynamics of nickel in aqueous solution and makes independent analyses of the available literature in this area. The standard method used for the analy-sis of ionic interactions between components dissolved in water (see Appendix B) al-lows the general and consistent use of the selected data for modelling purposes, regard-less of the type and composition of the ground water, within the ionic strength limits given by the experimental data used for the data analyses in the present review.

The interactions between solid compounds, such as the rock materials, and the aqueous solution and its components are as important as the interactions within the aqueous solution, because the solid materials in the geosphere control the chemistry of the ground water, and they also contribute to the overall solubilities of key elements. The present review therefore also considers the chemical behaviour of solid compounds that are likely to occur, or to be formed, under geological and environmental conditions. It is, however, difficult to assess the relative importance of the solid phases for perform-ance assessment purposes, particularly since their interactions with the aqueous phase are in many cases known to be subject to quantitatively unknown kinetic constraints. Furthermore, in some circumstances sorption of aqueous ions at mineral-water inter-faces may be a more important factor in determining migration of nickel than dissolu-

6 1 Introduction

tion and precipitation phenomena.

This book contains a summary and a critical review of the thermodynamic data on compounds and complexes containing nickel (cf. Chapter V), as reported in the available chemical literature up to mid-2002, but a few more recent references are also included. A comparatively large number of primary references are discussed separately in Appendix A.

Although the focus of this review is on nickel, it is necessary to use data on a number of other species during the evaluation process that lead to the recommended data. These so-called auxiliary data are taken both from the publication of CODATA Key Values [89COX/WAG] and from the evaluation of additional auxiliary data in the uranium and other volumes of this series [92GRE/FUG], [99RAR/RAN], [2003GUI/FAN], [2005OLI/NOL] and their use is recommended by this review. Care has been taken that all the selected thermodynamic data at standard conditions (cf. Sec-tion II.3) and 298.15 K are internally consistent. For this purpose, special software has been developed at the NEA Data Bank that is operational in conjunction with the NEA-TDB data base system, cf. Section II.6. In order to maintain consistency in the applica-tion of the values selected by this review, it is essential to use these auxiliary data when calculating equilibrium constants involving nickel compounds and complexes.

This review does not include any compounds and complexes containing or-ganic ligands or species in non-aqueous solvents. Nickel alloy systems have not been reviewed and treatment of gaseous species is very limited.

I.3 Review procedure and results The objective of the present review is the critical compilation of a database for the inor-ganic species of nickel. This aim is achieved by an assessment of all sources of thermo-dynamic data published until mid-2002. This assessment is performed in order to decide on the most reliable values that can be recommended. Experimental measurements pub-lished in the scientific literature are the main source for the selection of recommended data. Previous reviews are not neglected, but form a valuable source of critical informa-tion on the quality of primary publications.

When necessary, experimental source data are re-evaluated by using chemical models that are either found to be more realistic than those used by the original author, or are consistent with side-reactions discussed in another section of the review (for ex-ample, data on carbonate and hydroxide solubilities might need to be re-interpreted to take into account crystal structure and particle size of the phases actually investigated). Re-evaluation of literature values might be also necessary to correct for known system-atic errors (for example, if the junction potentials are neglected in the original publica-tion) or to make extrapolations to standard state conditions (I = 0) by using the specific ion interaction (SIT) equations (cf. Appendix B). For convenience, these SIT equations are referred to in some places in the text as “the SIT”. In order to ensure that consistent

I.3 Review procedure and results 7

procedures are used for the evaluation of primary data, a number of guidelines have been developed. They have been updated and improved since 1987, and their most re-cent versions are available at the NEA [99WAN], [99WAN/OST], [2000OST/WAN], [2000GRE/WAN], [2000WAN/OST]. Some of these procedures are also outlined in this volume, cf. Chapter II, Appendix B, and Appendix C. Parts of these sections, which were also published in earlier volumes [92GRE/FUG], [95SIL/BID], [99RAR/RAN], [2001LEM/FUG], [2003GUI/FAN] have been revised in this review.

Once the critical review process in the NEA TDB project is completed, the resulting manuscript is reviewed independently by qualified experts nominated by the NEA. The independent peer review is performed according to the procedures outlined in the TDB-6 guideline [99WAN]. The purpose of the additional peer review is to receive an independent view of the judgements and assessments made by the primary reviewers, to verify assumptions, results and conclusions, and to check whether the relevant litera-ture has been exhaustively considered. The independent peer review is performed by persons having technical expertise in the subject matter to be reviewed, to a degree at least equivalent to that needed for the original review.

The thermodynamic data selected in the present review (see Chapters III and IV) refer to the reference temperature of 298.15 K and to standard conditions, cf. Sec-tion II.3. For the modelling of real systems it is, in general, necessary to recalculate the standard thermodynamic data to non-standard state conditions. For aqueous species a procedure for the calculation of the activity factors is thus required. This review uses the approximate specific ion interaction method (SIT) for the extrapolation of experimental data to the standard state in the data evaluation process, and in some cases this requires the re-evaluation of original experimental values (solubilities, emf data, etc.). For maximum consistency, this method, as described in Appendix B, should always be used in conjunction with the selected data presented in this review.

The thermodynamic data selected in this review are provided with uncertainties representing the 95% confidence level. As discussed in Appendix C, there is no unique way to assign uncertainties, and the assignments made in this review are to a large ex-tent based on the subjective choice by the reviewers, supported by their scientific and technical experience in the corresponding area.

The quality of thermodynamic models cannot be better than the quality of the data on which they are based. The quality aspect includes both the numerical values of the thermodynamic data used in the model and the “completeness” of the chemical model used, e.g., the inclusion of all the relevant dissolved chemical species and solid phases. For the user it is important to consider that the selected data set presented in this review (Chapters III and IV) is certainly not “complete” with respect to all the conceiv-able systems and conditions; there are gaps in the information. The gaps are pointed out in the various sections of Part III, and this information may be used as a basis for the

8 1 Introduction

assignment of research priorities. The following example for a rewarding project in ex-perimental thermodynamics must suffice:

Precise measurements of the enthalpies of dissolution in aqueous media ac-cording to the reactions:

NiX2 (cr) Ni2+ + 2 X– (where X = Cl, Br, I) and

NiCl2·6H2O (cr) Ni2+ + 2 Cl– + 6 H2O(l)

as well as low temperature heat capacity measurements on NiCl2·6H2O(cr) would lead to the important key values f mH ο∆ (Ni2+, 298.15K) and mS ο (Ni2+, 298.15K) by thermo-dynamic cycles.

The results of experimental work done on formation of weak complexes are of-ten ambiguous. In the TDB assessments, the intent is to provide an interpretation that is consistent with the experimental data. This is not necessarily a unique interpretation, but must be one that can be used within the SIT framework. It can be an interpretation with or without an association constant, but it must be clear which is to be used. We have attempted to state the assumptions at each stage so that a reader can:

• reconstruct the arguments,

• reconstruct the calculations and results,

• with the final selected values (Tables III-1 and III-2) and interaction coeffi-cients reported in Tables B-4 and B-5, be able to regenerate the original ex-perimental data within the uncertainties of the selected values.

In some cases, an association constant and related interaction coefficients have been used where it might have been possible to “make do” with a different set of inter-action coefficients. We make no apologies for this so long as the calculation procedure is well defined and consistent.

For example, in several hydrolysis studies, allowance is made for complexation of Ni2+ with chloride, using an association constant, the calculated activity of water. An interaction coefficient of NiOH+ with Cl- is then derived as part of the same calculation that defines a value for the first hydrolysis constant for Ni2+. Conversely, association between Ni2+ and Cl– could have been neglected, an interaction coefficient for Ni2+ with Cl– then would have been specified, along with a different interaction coefficient for NiOH+ with Cl–. In theory, there should be no effect on the value of the selected hy-drolysis constant, 1K ο , at I = 0.

I.3 Review procedure and results 9

Our inherent assumptions about weak complexes can be stated as follows:

• A weak complex forms between Ni2+ and some singly charged anion X–,

• It is assumed that there is no complexation between Ni2+ and 4ClO− ,

• If a weak complex forms between Ni2+ and X–, the interaction coefficient ε(Ni2+, X–) is arbitrarily set to a “reasonable value”, in some cases a value similar or identical to the interaction coefficient ε(Ni2+, 4ClO− ),

• All other interaction between Ni2+ and X– is accounted for by the association constant ( K ο = a(NiX+) / a(Ni2+) a(X–)) and the interaction coefficient ε(NiX+, X–) (and, if necessary, ε(NiX+, 4ClO− ) also is specified).

The ε(Ni2+,X–) values are unlikely to be identical, and estimated uncertainties have been increased for values obtained from experiments in which a considerable part of the background electrolyte was replaced by the complex-forming anion. In some cases it has been necessary to make very arbitrary assumptions in generation of associa-tion constants for weak complexes. If used consistently, these should have little or no impact on geochemical modelling using the TDB database, or on regeneration of the experimental values.

When the experimental data are sparse, or the experiments were not well de-signed, the estimated uncertainties in some association constant values are large. If we omitted these association constants, and only generated interaction coefficients, the propagated uncertainties would not be reduced. They would simply reappear in a differ-ent guise.

In most cases this does not matter whatsoever, and a good argument can be made for eliminating the extraneous constants. However, modellers have found that when using a properly assembled database, there is only one thing worse than a species that has been included with a large uncertainty, and that is a species, for which there is qualitative or semi-quantitative evidence, but for which the database compiler has de-clined to propose a value. Therefore, there is a tendency for modellers to augment any database with extra values. In our view, it is better to supply association constants (with their inherent large uncertainties) than to allow the introduction of inconsistent “secon-dary” values.

11

Chapter II

II Standards, Conventions, and Contents of the TablesEquation Section 2

Equation Section 2

This chapter outlines and lists the symbols, terminology and nomenclature, the units and conversion factors, the order of formulae, the standard conditions, and the fundamental physical constants used in this volume. They are derived from international standards and have been specially adjusted for the TDB publications.

II.1 Symbols, terminology and nomenclature

II.1.1 Abbreviations Abbreviations are mainly used in tables where space is limited. Abbreviations for meth-ods of measurement are listed in Table II-1.

Table II-1: Abbreviations for experimental methods

AIX Anion exchange AES Atomic Emission Spectroscopy CAL Calorimetry CHR Chromatography CIX Cation exchange COL Colorimetry CON Conductivity COR Corrected COU Coulometry CRY Cryoscopy DIS Distribution between two phases DSC Differential Scanning Calorimetry DTA Differential Thermal Analysis EDS Energy Dispersive Spectroscopy EM Electromigration

(Continued on next page)

12 II. Standards, Conventions, and Contents of the Tables

Table II-1: (continued)

EMF Electromotive force, not specified EPMA Electron Probe Micro Analysis EXAFS Extended X-ray Absorption Fine Structure FTIR Fourier Transform Infra Red IDMS Isotope Dilution Mass-Spectroscopy IR Infrared GL Glass electrode ISE-X Ion selective electrode with ion X stated IX Ion exchange KIN Rate of reaction LIBD Laser Induced Break Down MVD Molar Volume Determination NMR Nuclear Magnetic Resonance PAS Photo Acoustic Spectroscopy POL Polarography POT Potentiometry PRX Proton relaxation QH Quinhydrone electrode RED Emf with redox electrode REV Review SEM Scanning Electron Microscopy SP Spectrophotometry SOL Solubility TC Transient Conductivity TGA Thermo Gravimetric Analysis TLS Thermal Lensing Spectrophotometry TRLFS Time Resolved Laser Fluorescence Spectroscopy UV Ultraviolet VLT Voltammetry XANES X-ray Absorption Near Edge Structure XRD X-ray Diffraction ? Method unknown to the reviewers

Other abbreviations may also be used in tables, such as SHE for the standard hydrogen electrode or SCE for the saturated calomel electrode. The abbreviation NHE has been widely used for the “normal hydrogen electrode”, which is by definition iden-tical to the SHE. It should nevertheless be noted that NHE customarily refers to a stan-dard state pressure of 1 atm, whereas SHE always refers to a standard state pressure of 0.1 MPa (1 bar) in this review.

II.1 Symbols, terminology and nomenclature 13

II.1.2 Symbols and terminology The symbols for physical and chemical quantities used in the TDB review follow the recommendations of the International Union of Pure and Applied Chemistry, IUPAC [79WHI], [93MIL/CVI]. They are summarised in Table II-2.

Table II-2: Symbols and terminology.

Symbols and terminology

length l

height h

radius r

diameter d

volume V

mass m

density (mass divided by volume) ρ

time t

frequency ν

wavelength λ internal transmittance (transmittance of the medium itself, disregarding boundaryor container influence)

Ti

internal transmission density, (decadic absorbance): log10(1/Ti) A

molar (decadic) absorption coefficient: B/A c l ε

relaxation time τ

Avogadro constant NA

relative molecular mass of a substance(a) Mr

thermodynamic temperature, absolute temperature T

Celsius temperature t

(molar) gas constant R

Boltzmann constant k

Faraday constant F

(molar) entropy mS

(molar) heat capacity at constant pressure ,mp

C

(molar) enthalpy mH

(molar) Gibbs energy mG

chemical potential of substance B µB

pressure p

partial pressure of substance B: xB p pB

fugacity of substance B fB

(Continued next page)



Symbols and terminology

fugacity coefficient: fB/pB γf,B amount of substance(b) n mole fraction of substance B xB

molarity or concentration of a solute substance B (amount of B divided by the volume of the solution) (c)

cB, [B]

molality of a solute substance B (amount of B divided by the mass of the solvent) (d)

mB

factor for the conversion of molarity to molality of a solution: B B/ cm mean ionic molality (e), + +(ν ν ) ν νm m m− −+

± + −= m± activity of substance B aB activity coefficient, molality basis: B B/a m γB activity coefficient, concentration basis: B B/a c yB mean ionic activity (e), B

+ +(ν ν ) ν νa a a a− −+± + −= = a±

mean ionic activity coefficient (e), + +(ν ν ) ν ν− −+± + −γ = γ γ γ±

osmotic coefficient, molality basis φ

ionic strength: 2 21 1 or 2 2

i im i i c i iI m z I c z= =∑ ∑ I

SIT ion interaction coefficient between substance B1 and substance B2 ε(B1, B2) stoichiometric coefficient of substance B (negative for reactants, positive for products)

νB

general equation for a chemical reaction B B0 ν B= ∑ equilibrium constant (f) K charge number of an ion B (positive for cations, negative for anions) zB charge number of a cell reaction n electromotive force E

+

1

10 HpH= log /(mol kg )[ ]a −− ⋅ pH electrolytic conductivity κ superscript for standard state(g) °

(a) ratio of the average mass per formula unit of a substance to 112 of the mass of an atom of nuclide 12C.

(b) cf. Sections 1.2 and 3.6 of the IUPAC manual [79WHI]. (c) This quantity is called “amount-of-substance concentration” in the IUPAC manual [79WHI]. A solu-

tion with a concentration equal to 0.1 3mol dm−⋅ is called a 0.1 molar solution or a 0.1 M solution. (d) A solution having a molality equal to 0.1 1mol kg −⋅ is called a 0.1 molal solution or a 0.1 m solution. (e) For an electrolyte

+ν νN X−

which dissociates into + ν ( ν ν )± −= + ions, in an aqueous solution with molality m, the individual cationic molality and activity coefficient are +( ν )m m+ = and

+ ( )γ /a m+ += . A similar definition is used for the anionic symbols. Electrical neutrality requires that

+ +ν νz z− −= . (f) Special notations for equilibrium constants are outlined in Section II.1.6. In some cases, cK is used to

indicate a constant in molar units, and mK a constant in molal units. (g) See Section II.3.1.


II.1.3 Chemical formulae and nomenclature This review follows the recommendations made by IUPAC [71JEN], [77FER], [90LEI] on the nomenclature of inorganic compounds and complexes, except for the following items:

• The formulae of coordination compounds and complexes are not enclosed in square brackets [71JEN] (Rule 7.21). Exceptions are made in cases where square brackets are required to distinguish between coordinated and uncoordi-nated ligands.

• The prefixes “oxy–” and “hydroxy–” are retained if used in a general way, e.g., “gaseous uranium oxyfluorides”. For specific formula names, however, the IUPAC recommended citation [71JEN] (Rule 6.42) is used, e.g., “uranium(IV) difluoride oxide” for 2UF O(cr).

An IUPAC rule that is often not followed by many authors [71JEN] (Rules 2.163 and 7.21) is recalled here: the order of arranging ligands in coordination com-pounds and complexes is the following: central atom first, followed by ionic ligands and then by the neutral ligands. If there is more than one ionic or neutral ligand, the alpha-betical order of the symbols of the ligating atoms determines the sequence of the lig-ands. For example, 2 2 3 3(UO ) CO (OH) − is standard, 2 2 3 3(UO ) (OH) CO− is non-standard and is not used.

Abbreviations of names for organic ligands appear sometimes in formulae. Following the recommendations by IUPAC, lower case letters are used, and if neces-sary, the ligand abbreviation is enclosed within parentheses. Hydrogen atoms that can be replaced by the metal atom are shown in the abbreviation with an upper case “H”, for example: 3H edta− , Am(Hedta)(s) (where edta stands for ethylenediaminetetraacetate).

II.1.4 Phase designators Chemical formulae may refer to different chemical species and are often required to be specified more clearly in order to avoid ambiguities. For example, 4UF occurs as a gas, a solid, and an aqueous complex. The distinction between the different phases is made by phase designators that immediately follow the chemical formula and appear in paren-theses. The only formulae that are not provided with a phase designator are aqueous ions. They are the only charged species in this review since charged gases are not con-sidered. The use of the phase designators is described below.

• The designator (l) is used for pure liquid substances, e.g., 2H O(l) .

• The designator (aq) is used for undissociated, uncharged aqueous species, e.g., 4U(OH) (aq) , 2CO (aq) . Since ionic gases are not considered in this re-view, all ions may be assumed to be aqueous and are not designed with (aq). If


a chemical reaction refers to a medium other than 2H O (e.g., 2D O , 90% etha-nol/10% 2H O ), then (aq) is replaced by a more explicit designator, e.g., “(in

2D O )” or “(sln)”. In the case of (sln), the composition of the solution is de-scribed in the text.

• The designator (sln) is used for substances in solution without specifying the actual equilibrium composition of the substance in the solution. Note the dif-ference in the designation of 2H O in Eqs.(II.2) and (II.3). 2H O(l) in Reaction (II.2) indicates that 2H O is present as a pure liquid, i.e., no solutes are present, whereas Reaction (II.3) involves a HCl solution, in which the thermodynamic properties of 2H O(sln) may not be the same as those of the pure liquid

2H O(l) . In dilute solutions, however, this difference in the thermodynamic properties of 2H O can be neglected, and 2H O(sln) may be regarded as pure

2H O(l) .

Example:

2 2 2UO Cl (cr) + 2 HBr(sln) UOBr (cr) + 2HCl(sln) (II.1)

2 2 2 2 2 2 2UO Cl 3H O(cr) UO Cl H O(cr) + 2 H O(l)⋅ ⋅ (II.2)

3 2 2 2UO (γ) + 2 HCl(sln) UO Cl (cr) + H O(sln) (II.3)

• The designators (cr), (am), (vit), and (s) are used for solid substances. (cr) is used when it is known that the compound is crystalline, (am) when it is known that it is amorphous, and (vit) for glassy substances. Otherwise, (s) is used.

• In some cases, more than one crystalline form of the same chemical composi-tion may exist. In such a case, the different forms are distinguished by separate designators that describe the forms more precisely. If the crystal has a mineral name, the designator (cr) is replaced by the first four characters of the mineral name in parentheses, e.g., 2SiO (quar) for quartz and 2SiO (chal) for chalced-ony. If there is no mineral name, the designator (cr) is replaced by a Greek let-ter preceding the formula and indicating the structural phase, e.g., α-UF5, β-UF5.

Phase designators are also used in conjunction with thermodynamic symbols to define the state of aggregation of a compound to which a thermodynamic quantity re-fers. The notation is in this case the same as outlined above. In an extended notation (cf. [82LAF]) the reference temperature is usually given in addition to the state of aggrega-tion of the composition of a mixture.


Example:

+f m (Na , 298.15 K)Gο∆ standard molar Gibbs energy of formation

of aqueous +Na at 298.15 K

m 2 4 2(UO (SO ) 2.5H O, cr, 298.15 K)S ο ⋅ standard molar entropy of 2 4 2UO (SO ) 2.5H O(cr)⋅ at 298.15 K

,m 3(UO , , 298.15 K)pCο α standard molar heat capacity of α-UO3 at 298.15 K

f m 2(HF, sln, HF 7.8H O)H∆ ⋅ enthalpy of formation of HF diluted 1:7.8 with water.

II.1.5 Processes Chemical processes are denoted by the operator ∆ , written before the symbol for a property, as recommended by IUPAC [82LAF]. An exception to this rule is the equilib-rium constant, cf. Section II.1.6. The nature of the process is denoted by annotation of the ∆, e.g., the Gibbs energy of formation, f mG∆ , the enthalpy of sublimation, sub mH∆ , etc. The abbreviations of chemical processes are summarised in Table II-3.

Table II-3: Abbreviations used as subscripts of ∆ to denote the type of chemical process.

Subscript of ∆ Chemical process

at separation of a substance into its constituent gaseous atoms (atomisation) dehyd elimination of water of hydration (dehydration) dil dilution of a solution f formation of a compound from its constituent elements fus melting (fusion) of a solid hyd addition of water of hydration to an unhydrated compound mix mixing of fluids r chemical reaction (general) sol process of dissolution sub sublimation (evaporation) of a solid tr transfer from one solution or liquid phase to another trs transition of one solid phase to another vap vaporisation (evaporation) of a liquid

The most frequently used symbols for processes are f G∆ and f H∆ , the Gibbs

energy and the enthalpy of formation of a compound or complex from the elements in their reference states (cf. Table II-6).


II.1.6 Equilibrium constants The IUPAC has not explicitly defined the symbols and terminology for equilibrium constants of reactions in aqueous solution. The NEA has therefore adopted the conven-tions that have been used in the work Stability Constants of Metal ion Complexes by Sillén and Martell [64SIL/MAR], [71SIL/MAR]. An outline is given in the paragraphs below. Note that, for some simple reactions, there may be different correct ways to in-dex an equilibrium constant. It may sometimes be preferable to indicate the number of the reaction to which the data refer, especially in cases where several ligands are dis-cussed that might be confused. For example, for the equilibrium:

M + L M Lm qm q (II.4)

both ,q mb and b (II.4) would be appropriate, and ,q mb (II.4) is accepted, too. Note that, in general, K is used for the consecutive or stepwise formation constant, and β is used for the cumulative or overall formation constant. In the following outline, charges are only given for actual chemical species, but are omitted for species containing general symbols (M, L).

II.1.6.1 Protonation of a ligand

+1H + H L H L r r− 1, +

1

[H L] = [H ][H L]

rr

r

K−

(II.5)

+ H + L H L rr 1, +

[H L] = [H ] [L]

rr rb (II.6)

This notation has been proposed and used by Sillén and Martell [64SIL/MAR], but it has been simplified later by the same authors [71SIL/MAR] from 1,rK to rK . This re-view retains, for the sake of consistency, cf. Eqs.(II.7) and (II.8), the older formulation of 1,rK .

For the addition of a ligand, the notation shown in Eq.(II.7) is used.

1HL + L HL q q− 1

[HL ] =

[HL ][L]q

qq

K−

(II.7).

Eq.(II.8) refers to the overall formation constant of the species H Lr q .

+ H + L H L r qr q , +

[H L ] =

[H ] [L]r q

q r r qb (II.8).

In Eqs.(II.5), (II.6) and (II.8), the second subscript r can be omitted if r = 1, as shown in Eq.(II.7).


Example:

+ 3 24 4H + PO HPO− −

24

1,1 1 + 34

[HPO ] = [H ][PO ]

−

−=β β

+ 34 2 42 H + PO H PO− − 2 4

1,2 + 2 34

[H PO ] = [H ] [PO ]

−

−β

II.1.6.2 Formation of metal ion complexes

1ML + L ML q q− 1

[ML ] =

[ML ][L]−

qq

q

K (II.9)

M + L ML qq [ ][ ][ML ]

= M L

qq qβ (II.10)

For the addition of a metal ion, i.e., the formation of polynuclear complexes, the follow-ing notation is used, analogous to Eq.(II.5):

1M + M L M L m m− [ ]

[ ][ ]1,1

M L =

M M Lm

mm

K−

(II.11).

Eq.(II.12) refers to the overall formation constant of a complex M L . m q

M + L M L m qm q [ ] [ ],

[M L ] =

M Lm q

q m m qβ (II.12)

The second index can be omitted if it is equal to 1, i.e., ,q mb becomes qβ if m = 1. The formation constants of mixed ligand complexes are not indexed. In this case, it is necessary to list the chemical reactions considered and to refer the constants to the corresponding reaction numbers.

It has sometimes been customary to use negative values for the indices of the protons to indicate complexation with hydroxide ions, OH− . This practice is not adopted in this review. If OH− occurs as a reactant in the notation of the equilibrium, it is treated like a normal ligand L, but in general formulae the index variable n is used instead of q. If 2H O occurs as a reactant to form hydroxide complexes, 2H O is consid-ered as a protonated ligand, HL, so that the reaction is treated as described below in Eqs.(II.13) to (II.15) using n as the index variable. For convenience, no general form is used for the stepwise constants for the formation of the complex MmLqHr. In many ex-


periments, the formation constants of metal ion complexes are determined by adding a ligand in its protonated form to a metal ion solution. The complex formation reactions thus involve a deprotonation reaction of the ligand. If this is the case, the equilibrium constant is supplied with an asterisk, as shown in Eqs.(II.13) and (II.14) for mononu-clear and in Eq.(II.15) for polynuclear complexes.

+1ML + HL ML + Hq q−

+

1

* ML H =

ML HL−

qq

q

K (II.13)

+M + HL ML + Hqq q [ ][ ]

+* ML H

= M HL

q

qq qb (II.14)

+ M + HL M L + Hm qm q q [ ] [ ]

+

,* M L H

= M HL

q

m qq m m q

b (II.15)

Example:

2+ + +2 2UO + HF(aq) UO F H+

+ +2

1 1 2+2

* * UO F H = =

UO HF (aq)K

b

2+ + +2 2 2 3 53 UO + 5 H O(l) (UO ) (OH) + 5 H

5+ +2 3 5

5,3 32+2

* (UO ) (OH) H =

UO

b

Note that an asterisk is only assigned to the formation constant if the proto-nated ligand that is added is deprotonated during the reaction. If a protonated ligand is added and coordinated as such to the metal ion, the asterisk is to be omitted, as shown in Eq.(II.16).

M + H L M(H L) r r qq [ ][ ]M(H L)

= M H L

r qq q

r

b (II.16)

Example:

2+2 2 4 2 2 4 3UO + 3 H PO UO (H PO )− − 2 2 4 3

3 32+2 2 4

UO (H PO ) =

UO H PO

−

−

b

II.1.6.3 Solubility constants Conventionally, equilibrium constants involving a solid compound are denoted as “sol-ubility constants” rather than as formation constants of the solid. An index “s” to the equilibrium constant indicates that the constant refers to a solubility process, as shown in Eqs.(II.17) to (II.19).


M L (s) M + L a b a b [ ] [ ],0 = M La bsK (II.17).

,0sK is the conventional solubility product, and the subscript “0” indicates that the equilibrium reaction involves only uncomplexed aqueous species. If the solubility constant includes the formation of aqueous complexes, a notation analogous to that of Eq.(II.12) is used:

M L (s) M L + La b m qm mb qa a

−

( )

, , = M L L−

mb qa

s q m m qK (II.18).

Example:

+2 2 2UO F (cr) UO F + F− +

,1,1 ,1 2= = UO F F s sK K − .

Similarly, an asterisk is added to the solubility constant if it simultaneously in-volves a protonation equilibrium:

+M L (s) + H M L + HLa b m qm mb mbq qa a a

− −

( )

, ,( )+

* M L HL =

H

−

−

mb qa

m qs q m mb q

a

K (II.19)

Example:

+ 2+4 2 2 4 2 4 2U(HPO ) 4H O(cr) + H UHPO + H PO + 4 H O(l)−⋅

2+4 2 4

,1,1 ,1 +* * UHPO H PO

= = Hs sK K

−

.

II.1.6.4 Equilibria involving the addition of a gaseous ligand A special notation is used for constants describing equilibria that involve the addition of a gaseous ligand, as outlined in Eq.(II.20).

1ML + L(g) ML q q− 1 L

p,

ML = ML

q

qqK p−

(II.20)

The subscript “p” can be combined with any other notations given above.

Example:

2 2CO (g) CO (aq) [ ]

2

2p

CO

CO (aq) K

p=


2+ 6 +2 2 2 2 3 3 63 UO + 6 CO (g) + 6 H O(l) (UO ) (CO ) 12 H− +

2

126 +2 3 3 6

p,6,3 32+ 62 CO

* (UO ) (CO ) H =

UO p

−

b

2 +2 3 2 2 2 3 2UO CO (cr) + CO (g) + H O(l) UO (CO ) 2 H− +

2

22 +2 3 2

p, s, 2CO

* UO (CO ) H = K

p

−

In cases where the subscripts become complicated, it is recommended that K or β be used with or without subscripts, but always followed by the equation number of the equilibrium to which it refers.

II.1.6.5 Redox equilibria Redox reactions are usually quantified in terms of their electrode (half cell) potential, E, which is identical to the electromotive force (emf) of a galvanic cell in which the elec-trode on the left is the standard hydrogen electrode, SHE1, in accordance with the “1953 Stockholm Convention” [93MIL/CVI]. Therefore, electrode potentials are given as re-duction potentials relative to the standard hydrogen electrode, which acts as an electron donor. In the standard hydrogen electrode, 2H (g) is at unit fugacity (an ideal gas at unit pressure, 0.1 MPa), and +H is at unit activity. The sign of the electrode potential, E, is that of the observed sign of its polarity when coupled with the standard hydrogen elec-trode. The standard electrode potential, Eο , i.e., the potential of a standard galvanic cell relative to the standard hydrogen electrode (all components in their standard state, cf. Section II.3.1, and with no liquid junction potential) is related to the standard Gibbs energy change r mGο∆ and the standard (or thermodynamic) equilibrium constant K ο as outlined in Eq.(II.21).

r m1 R = = lnF F

TE G Kn n

ο ο ο− ∆ (II.21)

and the potential, E, is related to Eο by:

= (R / F) lni iE E T n aο − ν∑ (II.22).

1 The definitions of SHE and NHE are given in Section II.1.1.


For example, for the hypothetical galvanic cell:

Pt H2(g, p = 1 bar) HCl(aq, +Ha = 1, 2Hf = 1)

Fe(ClO4)2 2+Fe(aq, = 1)a

Fe(ClO4)3 3+Fe(aq, = 1)a Pt (II.23)

where denotes a liquid junction and a phase boundary, the reaction is:

3+ 2+ +2

1Fe + H (g) Fe + H2

(II.24)

Formally Reaction (II.24) can be represented by two half cell reactions, each involving an equal number of electrons, (designated “ e− ”), as shown in the following equations:

3+ 2+Fe + e Fe− (II.25)

+2

1 H (g) H + e2

− . (II.26)

The terminology is useful, although it must be emphasised that “ e− ” here does not represent the hydrated electron.

The equilibrium constants of the two half cell reactions may be written:

K ο (II.25) = 2+

3+

Fe

Fe e

a

a a −⋅ (II.27)

K ο (II.26) = +

2

H e

H1

a a

f−⋅

= (II.28)

The “Stockholm Convention” implies that the constant K ο (II.26) = 1, which also defines the numerical value of ea − to 1 for the standard hydrogen electrode. In ad-dition, r mGο∆ (II.26) = 0, r mH ο∆ (II.26) = 0, r mS ο∆ (II.26) = 0 by definition, at all tempera-tures, and therefore r mGο∆ (II.25) = r mGο∆ (II.24).

The equilibrium constant of the more general cell reaction:

Ox + 12

n H2(g) Red + n H+ (II.29)

and of the half cell reaction

Ox + n e− Red (II.30)

together with K ο (II.26) yield

K ο (II.29) = K ο (II.30)× K ο (II.26)n Red Red

Ox Oxe e e

1 1( /( (SHE)=1))n n

a aa aa a a− − −

= = (II.31)


The activity scale is thus fixed by the above convention. Further

E(II.29) = Eo(II.29) Red10

Ox

R ln(10) logF

aTn a

−

(II.32)

Combination of Eqs.((II.21), (II.31) and (II.32)) yields the relationship between ea − and the electrode potential vs. SHE:

– log10 ea − = FR ln(10)

ET

(II.29) (II.33)

The splitting of redox reactions into two half cell reactions by introducing the symbol “ e− ” is highly useful even though the e− notation does not in any way refer to solvated electrons. When calculating the equilibrium composition of a chemical system, both “ e− ”, and +H can be chosen as components and they can be treated numerically in a similar way: equilibrium constants, mass balances, etc. may be defined for both. How-ever, while +H represents the hydrated proton in aqueous solution, the above equations use only the activity of “ e− ”, and never the concentration of “ e− ”. Concentration to activity conversions (or activity coefficients) are never needed for the electron (cf. Ap-pendix B, Example B.3).

In the literature on geochemical modelling of natural waters, it is customary to represent the “electron activity” of an aqueous solution with the symbol “pe” or “pε”( 10 e

log a −= − ) by analogy with pH ( +10 Hlog a= − ), and the redox potential of an

aqueous solution relative to the standard hydrogen electrode is usually denoted by either “Eh” or “ HE ” (see for example [82DRE], [84HOS], [86NOR/MUN], [96STU/MOR2],). The two representations are interrelated via Eq.(II.33) yielding the expression

HFpe

R ln(10)E

T= .

The symbol E ο′ is used to denote the so-called “formal potential” [74PAR]. The formal (or “conditional”) potential can be regarded as a standard potential for a particular medium in which the activity coefficients are independent (or approximately so) of the reactant concentrations [85BAR/PAR] (the definition of E ο′ parallels that of “concentration quotients” for equilibria). Therefore, from

R = lnF i iTE E c

nο′ − ν∑ (II.34)

E ο′ is the potential E for a cell when the ratio of the concentrations (not the activities) on the right-hand side and the left-hand side of the cell reaction is equal to unity, and

R = F iTE E

nο ο′ − ν∑ ln γi = r m

FG

n∆

− (II.35)

where the iγ are the molality activity coefficients and is m( )iic , the ratio of molality

to molarity (cf. Section II.2). The medium must be specified.


II.1.7 pH Because of the importance that potentiometric methods have in the determination of equilibrium constants in aqueous solutions, a short discussion on the definition of “pH” and a simplified description of the experimental techniques used to measure pH will be given here. For a comprehensive account see [2002BUC/RON].

The acidity of aqueous solutions is often expressed in a logarithmic scale of the hydrogen ion activity. The definition of pH as:

+ + +10 10H H HpH = log = log ( )a m− − γ

can only be strictly used in the limiting range of the Debye-Hückel equation (that is, in extremely dilute solutions). In practice the use of pH values requires extra assumptions on the values for single ion activities. In this review values of pH are used to describe qualitatively ranges of acidity of experimental studies, and the assumptions described in Appendix B are used to calculate single ion activity coefficients.

The determination of pH is often performed by emf measurements of galvanic cells involving liquid junctions [69ROS], [73BAT]. A common setup is a cell made up of a reference half cell (e.g. Ag(s)|AgCl(s) in a solution of constant chloride concentra-tion), a salt bridge, the test solution, and a glass electrode (which encloses a solution of constant acidity and an internal reference half cell):

Pt(s) Ag(s) AgCl(s) KCl(aq) salt

bridge test

solution KCl(aq) AgCl(s) Ag(s) Pt(s)

a b (II.36)

where stands for a glass membrane.

The emf of such a cell (assuming Nernstian behaviour of the glass electrode) is given by:

+ jH* R = ln +

FTE E a E+

where *E is a constant, and jE is the liquid junction potential. The purpose of the salt bridge is to minimise the junction potential in junction “b”, while keeping constant the junction potential for junction “a”. Two methods are most often used to reduce and con-trol the value of jE . An electrolyte solution of high concentration (the “salt bridge”) is a requirement of both methods. In the first method, the salt bridge is a saturated (or nearly saturated) solution of potassium chloride. A problem with a bridge of high potassium concentration, is that potassium perchlorate might precipitate1 inside the liquid junction when the test solution contains a high concentration of perchlorate ions.

1 KClO4(cr) has a solubility of ≈ 0.15 M in pure water at 25°C


In the other method the salt bridge contains the same high concentration of the same inert electrolyte as the test solution (for example, 3 M NaClO4). However, if the concentration of the background electrolyte in the salt bridge and test solutions is re-duced, the values of jE are dramatically increased. For example, if both the bridge and the test solution have 4[ClO ]− = 0.1 M as background electrolyte, the dependence of the liquid junction at “b” on acidity is jE ≈ – 440 × [H+] mV·dm3·mol–1 at 25°C [69ROS] (p.110), which corresponds to a correction at pH = 2 of ≥ 0.07 pH units.

Because of the problems in eliminating the liquid junction potentials and in defining individual ionic activity coefficients, an “operational” definition of pH is given by IUPAC [93MIL/CVI]. This definition involves the measurement of pH differences between the test solution and standard solutions of known pH and similar ionic strength (in this way similar values of +H

γ and jE cancel each other when emf values are sub-tracted).

In order to deduce the stoichiometry and equilibrium constants of complex formation reactions and other equilibria, it is necessary to vary the concentrations of reactants and products over fairly large concentration ranges under conditions where the activity coefficients of the species are either known, or constant. Only in this manner is it possible to use the mass balance equations for the various components together with the measurement of one or more free concentrations to obtain the information desired [61ROS/ROS], [90BEC/NAG], [97ALL/BAN], p. 326 – 327. For equilibria involving hydrogen ions, it is necessary to use concentration units, rather than hydrogen ion activ-ity. For experiments in an ionic medium, where the concentration of an “inert” electro-lyte is much larger than the concentration of reactants and products we can ensure that, as a first approximation, their trace activity coefficients remain constant even for mod-erate variations of the corresponding total concentrations. Under these conditions of fixed ionic strength the free proton concentration may be measured directly, thereby defining it in terms of – log10[H+] rather than on the activity scale as pH, and the value of – log10[H+] and pH will differ by a constant term, i.e., +10 H

log γ . Equilibrium con-stants deduced from measurements in such ionic media are therefore conditional con-stants, because they refer to the given medium, not to the standard state. In order to compare the magnitude of equilibrium constants obtained in different ionic media it is necessary to have a method for estimating activity coefficients of ionic species in mixed electrolyte systems to a common standard state. Such procedures are discussed in Ap-pendix B.

Note that the precision of the measurement of – log10[H+] and pH is virtually the same, in very good experiments, ± 0.001. However, the accuracy is generally con-siderably poorer, depending in the case of glass electrodes largely on the response of the electrode (linearity, age, pH range, etc.), and to a lesser extent on the calibration method employed.


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II.1.8 Order of formulae To be consistent with CODATA, the data tables are given in “Standard Order of Ar-rangement” [82WAG/EVA]. This scheme is presented in Figure II-1 below, and shows the sequence of the ranks of the elements in this convention. The order follows the ranks of the elements.

For example, for uranium, this means that, after elemental uranium and its monoatomic ions (e.g., 4+U ), the uranium compounds and complexes with oxygen would be listed, then those with hydrogen, then those with oxygen and hydrogen, and so on, with decreasing rank of the element and combinations of the elements. Within a class, increasing coefficients of the lower rank elements go before increasing coeffi-cients of the higher rank elements. For example, in the U–O–F class of compounds and complexes, a typical sequence would be 2UOF (cr) , 4UOF (cr) , 4UOF (g) , 2UO F(aq) ,

+2UO F , 2 2UO F (aq) , 2 2UO F (cr) , 2 2UO F (g) , 2 3UO F− , 2

42UO F − , 2 3 6U O F (cr) , etc. [92GRE/FUG]. Formulae with identical stoichiometry are in alphabetical order of their designators. An exception is made for ammonium salts, which are included in group 1 before the alkali metals.

Figure II-1: Standard order of arrangement of the elements and compounds based on the periodic classification of the elements (from [82WAG/EVA]).


II.1.9 Reference codes The references cited in the review are ordered chronologically and alphabetically by the first two authors within each year, as described by CODATA [87GAR/PAR]. A refer-ence code is made up of the final two digits of the year of appearance (if the publication is not from the 20th century, the year will be put in full). The year is followed by the first three letters of the surnames of the first two authors, separated by a slash.

If there are multiple reference codes, a “2” will be added to the second one, a “3” to the third one, and so forth. Reference codes are always enclosed in square brack-ets.

II.2 Units and conversion factors Thermodynamic data are given according to the Système International d'unités (SI units). The unit of energy is the joule. Some basic conversion factors, also for non-thermodynamic units, are given in Table II-4.

Table II-4: Unit conversion factors

To convert from

(non-SI unit symbol)

to

(SI unit symbol)

multiply by

ångström (Å) metre (m) 1×10–10 (exactly)

standard atmosphere (atm) pascal (Pa) 1.01325×105 (exactly)

bar (bar) pascal (Pa) 1×105 (exactly)

thermochemical calorie (cal) joule (J) 4.184 (exactly)

entropy unit 1 1e.u. cal K mol− −⋅ ⋅ 1 1J K mol− −⋅ ⋅ 4.184 (exactly)

Since a large part of the NEA-TDB project deals with the thermodynamics of aqueous solutions, the units describing the amount of dissolved substance are used very frequently. For convenience, this review uses “M” as an abbreviation of “ 3mol dm−⋅ ” for molarity, c, and, in Appendices B and C, “m” as an abbreviation of “ 1mol kg−⋅ ” for molality, m. It is often necessary to convert concentration data from molarity to molality and vice versa. This conversion is used for the correction and extrapolation of equilib-rium data to zero ionic strength by the specific ion interaction theory, which works in molality units (cf. Appendix B). This conversion is made in the following way. Molality is defined as Bm moles of substance B dissolved in 1 kilogram of pure water. Molarity is defined as Bc moles of substance B dissolved in B( )c M−ρ kilogram of pure water, where ρ is the density of the solution in kg·dm–3 and MB the molar weight of the solute in kg·mol–1.

From this it follows that:

BB

B

=

cmc M−ρ

.

II.2 Units and conversion factors 29

When the ionic strength is kept high and constant by an inert electrolyte, I, the ratio mB /cB can be approximated by:

B

B I I

1

mc c M

≈−ρ

where cI is the concentration of the inert electrolyte in mol·dm–3 and MI its molar mass in kg·mol–1.

Baes and Mesmer [76BAE/MES], (p.439) give a table with conversion factors (from molarity to molality) for nine electrolytes and various ionic strengths. Conversion factors at 298.15 K for twenty one electrolytes, calculated using the density equations reported by Söhnel and Novotný [85SOH/NOV], are reported in Table II-5.

Example:

4 4

4 4

3 3

1.00 M NaClO 1.05 m NaClO1.00 M NaCl 1.02 m NaCl4.00 M NaClO 4.95 m NaClO6.00 M NaNO 7.55 m NaNO

It should be noted that equilibrium constants need also to be converted if the concentration scale is changed from molarity to molality or vice versa. For a general equilibrium reaction, B B0 = ν B∑ , the equilibrium constants can be expressed either in molarity or molality units, cK or mK , respectively:

10 B 10 B

B

10 B 10 BB

log ν log

log ν log

c

m

K c

K m

= ∑

= ∑

With B B( / )m c = , or 10 B 10 B (log log )m c− = log10 , the relationship between cK and mK becomes very simple, as shown in Eq.(II.37).

10 10 BB

log log + νm cK K= ∑ log10 (II.37)

B B ν∑ is the sum of the stoichiometric coefficients of the reaction, cf. Eq. (II.53) and the values of are the factors for the conversion of molarity to molality as tabulated in Table II-5 for several electrolyte media at 298.15 K. In the case of very dilute solutions, these factors are approximately equal to the reciprocal of the density of the pure solvent. Then, if the solvent is water, molarity and molality may be used inter-changeably, and c mK K≈ . The differences between the values in Table II-5 and the values listed in the uranium NEA-TDB review [92GRE/FUG] (p.23) are found at the highest concentrations, and are no larger than ± 0.003 dm3·kg–1, reflecting the accuracy expected in this type of conversion. The uncertainty introduced by the use of Eq.(II.37) in the values of 10log mK will be no larger than ± 0.001 B B ν∑ .

30 II. Standards, Conventions and Contents of the Tables

Table II-5: Factors for the conversion of molarity, cB, to molality, mB, of a substance B, in various media at 298.15 K (calculated from densities in [85SOH/NOV])

= mB / cB (dm3 of solution per kg of H2O)

c (M) HClO4 NaClO4 LiClO4 NH4ClO4 Ba(ClO4)2 HCl NaCl LiCl

0.10 1.0077 1.0075 1.0074 1.0091 1.0108 1.0048 1.0046 1.0049 0.25 1.0147 1.0145 1.0141 1.0186 1.0231 1.0076 1.0072 1.0078 0.50 1.0266 1.0265 1.0256 1.0351 1.0450 1.0123 1.0118 1.0127 0.75 1.0386 1.0388 1.0374 1.0523 1.0685 1.0172 1.0165 1.0177 1.00 1.0508 1.0515 1.0496 1.0703 1.0936 1.0222 1.0215 1.0228 1.50 1.0759 1.0780 1.0750 1.1086 1.1491 1.0324 1.0319 1.0333 2.00 1.1019 1.1062 1.1019 1.2125 1.0430 1.0429 1.0441 3.00 1.1571 1.1678 1.1605 1.3689 1.0654 1.0668 1.0666 4.00 1.2171 1.2374 1.2264 1.0893 1.0930 1.0904 5.00 1.2826 1.3167 1.1147 1.1218 1.1156 6.00 1.3547 1.4077 1.1418 1.1423

c (M) KCl NH4Cl MgCl2 CaCl2 NaBr HNO3

NaNO3 LiNO3

0.10 1.0057 1.0066 1.0049 1.0044 1.0054 1.0056 1.0058 1.0059 0.25 1.0099 1.0123 1.0080 1.0069 1.0090 1.0097 1.0102 1.0103 0.50 1.0172 1.0219 1.0135 1.0119 1.0154 1.0169 1.0177 1.0178 0.75 1.0248 1.0318 1.0195 1.0176 1.0220 1.0242 1.0256 1.0256 1.00 1.0326 1.0420 1.0258 1.0239 1.0287 1.0319 1.0338 1.0335 1.50 1.0489 1.0632 1.0393 1.0382 1.0428 1.0478 1.0510 1.0497 2.00 1.0662 1.0855 1.0540 1.0546 1.0576 1.0647 1.0692 1.0667 3.00 1.1037 1.1339 1.0867 1.0934 1.0893 1.1012 1.1090 1.1028 4.00 1.1453 1.1877 1.1241 1.1406 1.1240 1.1417 1.1534 1.1420 5.00 1.2477 1.1974 1.1619 1.1865 1.2030 1.1846 6.00 1.2033 1.2361 1.2585 1.2309

c (M) NH4NO3 H2SO4 Na2SO4 (NH4)2SO4 H3PO4 Na2CO3 K2CO3 NaSCN

0.10 1.0077 1.0064 1.0044 1.0082 1.0074 1.0027 1.0042 1.0069 0.25 1.0151 1.0116 1.0071 1.0166 1.0143 1.0030 1.0068 1.0130 0.50 1.0276 1.0209 1.0127 1.0319 1.0261 1.0043 1.0121 1.0234 0.75 1.0405 1.0305 1.0194 1.0486 1.0383 1.0065 1.0185 1.0342 1.00 1.0539 1.0406 1.0268 1.0665 1.0509 1.0094 1.0259 1.0453 1.50 1.0818 1.0619 1.0441 1.1062 1.0773 1.0170 1.0430 1.0686 2.00 1.1116 1.0848 1.1514 1.1055 1.0268 1.0632 1.0934 3.00 1.1769 1.1355 1.2610 1.1675 1.1130 1.1474 4.00 1.2512 1.1935 1.4037 1.2383 1.1764 1.2083 5.00 1.3365 1.2600 1.3194 1.2560 1.2773 6.00 1.4351 1.3365 1.4131 1.3557

II.3 Standard and reference conditions 31

II.3 Standard and reference conditions

II.3.1 Standard state A precise definition of the term “standard state” has been given by IUPAC [82LAF]. The fact that only changes in thermodynamic parameters, but not their absolute values, can be determined experimentally, makes it important to have a well-defined standard state that forms a base line to which the effect of variations can be referred. The IUPAC [82LAF] definition of the standard state has been adopted in the NEA-TDB project. The standard state pressure, pο = 0.1 MPa (1 bar), has therefore also been adopted, cf. Sec-tion II.3.2. The application of the standard state principle to pure substances and mix-tures is summarised below. It should be noted that the standard state is always linked to a reference temperature, cf. Section II.3.3.

• The standard state for a gaseous substance, whether pure or in a gaseous mixture, is the pure substance at the standard state pressure and in a (hypothetical) state in which it exhibits ideal gas behaviour.

• The standard state for a pure liquid substance is (ordinarily) the pure liquid at the standard state pressure.

• The standard state for a pure solid substance is (ordinarily) the pure solid at the standard state pressure.

• The standard state for a solute B in a solution is a hypothetical liquid solution, at the standard state pressure, in which = Bm mο = 1 mol·kg–1, and in which the ac-tivity coefficient Bγ is unity.

It should be emphasised that the use of superscript, ο , e.g., in f mH ο∆ , implies that the compound in question is in the standard state and that the elements are in their reference states. The reference states of the elements at the reference temperature (cf. Section II.3.3) are listed in Table II-6.

Table II-6: Reference states for some elements at the reference temperature of 298.15 K and standard pressure of 0.1 MPa [82WAG/EVA], [89COX/WAG], [91DIN].

O2 gaseous Al crystalline, cubic H2 gaseous Zn crystalline, hexagonal He gaseous Cd crystalline, hexagonal Ne gaseous Hg liquid Ar gaseous Cu crystalline, cubic Kr gaseous Ag crystalline, cubic Xe gaseous Fe crystalline, cubic, bcc F2 gaseous Tc crystalline, hexagonal Cl2 gaseous V crystalline, cubic




Br2 liquid Ti crystalline, hexagonal I2 crystalline, orthorhombic Am crystalline, dhcp S crystalline, orthorhombic Pu crystalline, monoclinic Se crystalline, hexagonal (“black”) Np crystalline, orthorhombic Te crystalline, hexagonal U crystalline, orthorhombic N2 gaseous Th crystalline, cubic P crystalline, cubic (“white”) Be crystalline, hexagonal As crystalline, rhombohedral (“grey”) Mg crystalline, hexagonal Sb crystalline, rhombohedral Ca crystalline, cubic, fcc Bi crystalline, rhombohedral Sr crystalline, cubic, fcc C crystalline, hexagonal (graphite) Ba crystalline, cubic Si crystalline, cubic Li crystalline, cubic Ge crystalline, cubic Na crystalline, cubic Sn crystalline, tetragonal (“white”) K crystalline, cubic Pb crystalline, cubic Rb crystalline, cubic B β, crystalline, rhombohedral Cs crystalline, cubic

II.3.2 Standard state pressure The standard state pressure chosen for all selected data is 0.1 MPa (1 bar) as recom-mended by the International Union of Pure and Applied Chemistry IUPAC [82LAF].

However, the majority of the thermodynamic data published in the scientific literature and used for the evaluations in this review, refer to the old standard state pres-sure of 1 “standard atmosphere” (= 0.101325 MPa). The difference between the ther-modynamic data for the two standard state pressures is not large and lies in most cases within the uncertainty limits. It is nevertheless essential to make the corrections for the change in the standard state pressure in order to avoid inconsistencies and propagation of errors. In practice the parameters affected by the change between these two standard state pressures are the Gibbs energy and entropy changes of all processes that involve gaseous species. Consequently, changes occur also in the Gibbs energies of formation of species that consist of elements whose reference state is gaseous (H, O, F, Cl, N, and the noble gases). No other thermodynamic quantities are affected significantly. A large part of the following discussion has been taken from the NBS tables of chemical ther-modynamic properties [82WAG/EVA], see also Freeman [84FRE].

The following expressions define the effect of pressure on the properties of all substances:

= (1 )pT

H VV T V Tp T

∂ ∂ − = − α ∂ ∂ (II.38)

II.3 Standard and reference conditions 33

2

2 = p

pT

C VTp T

∂ ∂− ∂ ∂

(II.39)

= = pT

S VVp T

∂ ∂ − α − ∂ ∂ (II.40)

= T

G Vp

∂ ∂

(II.41)

where 1 p

VV T

∂ α ≡ ∂ (II.42)

For ideal gases, V = RT / p and α = R / pV = 1 /T. The conversion equations listed below (Eqs. (II.43) to (II.50)) apply to the small pressure change from 1 atm to 1 bar (0.1 MPa). The quantities that refer to the old standard state pressure of 1 atm are assigned the superscript ( ) ,atm while those that refer to the new standard state pressure of 1 bar carry the superscript ( )bar .

For all substances the changes in the enthalpy of formation and heat capacity are much smaller than the experimental accuracy and can be disregarded. This is exactly true for ideal gases.

(bar) (atm)

f f( ) ( ) = 0H T H T∆ − ∆ (II.43) (bar) (atm)( ) ( ) = 0p pC T C T− (II.44)

For gaseous substances, the entropy difference is:

(atm)

(bar) (atm)(bar)( ) ( ) = R ln = R ln 1.01325pS T S T

p

−

1 1 = 0.1094 J K mol− −⋅ ⋅ (II.45)

This is exactly true for ideal gases, as follows from Eq.(II.40) with α = R / pV.The entropy change of a reaction or process is thus dependent on the number of moles of gases involved:

(atm)

(bar) (atm)r r (bar) = δ R ln pS S

p

∆ − ∆ ⋅

1 1 = δ 0.1094 J K mol− −× ⋅ ⋅ (II.46)

where δ is the net increase in moles of gas in the process.


Similarly, the change in the Gibbs energy of a process between the two stan-dard state pressures is:

(atm)(bar) (atm)

r r (bar) = δ R ln pG G Tp

∆ − ∆ − ⋅

1 = δ 0.03263 kJ mol−− ⋅ ⋅ at 298.15 K. (II.47)

Eq.(II.47) applies also to (bar) (atm)f f G G∆ − ∆ , since the Gibbs energy of for-

mation describes the formation process of a compound or complex from the reference states of the elements involved:

. (bar) (atm) 1f f = δ 0.03263 kJ molG G −∆ − ∆ − × ⋅ at 298.15 K. (II.48).

The changes in the equilibrium constants and cell potentials with the change in the standard state pressure follow from the expression for Gibbs energy changes, Eq.(II.47)

(bar) (atm)(bar) (atm) r r

10 10

(atm)

(bar) (atm)

10 (bar)

log log = R ln 10

ln δ = δ log

ln10

G GK KT

pp p

p

∆ − ∆− −

= ⋅ ⋅

δ 0.005717= ⋅ (II.49)

(bar) (atm)(bar) (atm) r r

(atm)

(bar)

= F

R ln δ

F

G GE En

pTp

n

∆ − ∆− −

= ⋅

0.0003382 δ V n

= ⋅ at 298.15K (II.50).

It should be noted that the standard potential of the hydrogen electrode is equal to 0.00 V exactly, by definition.

+ 2

def1H e H (g) 0.00V2 =E− ο+ (II.51).

This definition will not be changed, although a gaseous substance, 2H (g) , is involved in the process. The change in the potential with pressure for an electrode po-tential conventionally written as:

+Ag + e Ag (cr)−

II.4 Fundamental physical constants 35

should thus be calculated from the balanced reaction that includes the hydrogen elec-trode,

+ +2

1Ag + H (g) Ag (cr) + H2

Here δ = – 0.5 Hence, the contribution to δ from an electron in a half cell reaction is the same as the contribution of a gas molecule with the stoichiometric coefficient of 0.5. This leads to the same value of δ as the combination with the hydrogen half cell.

Example: + 2+ (bar) (atm)

2

(bar) (bar2 2 10 10

Fe(cr) + 2 H Fe + H (g) = 1 = 0.00017 V

CO (g) CO (aq) = 1 log log

E E

K K

δ −

δ − − )

(bar) (atm) 13 2 2 r r

(bar) (atm)2 2 4 f f

= 0.0057

5 3NH (g) + O NO(g) + H O(g) = 0.25 = 0.008 kJ mol

4 21

Cl (g) + 2 O (g) + e ClO = 3 = 2

G G

G G

−

− −

−

δ ∆ − ∆ − ⋅

δ − ∆ − ∆ 10.098 kJ mol−⋅

II.3.3 Reference temperature The definitions of standard states given in Section II.3 make no reference to fixed tem-perature. Hence, it is theoretically possible to have an infinite number of standard states of a substance as the temperature varies. It is, however, convenient to complete the definition of the standard state in a particular context by choosing a reference tempera-ture. As recommended by IUPAC [82LAF], the reference temperature chosen in the NEA-TDB project is T = 298.15 K or t = 25.00°C. Where necessary for the discussion, values of experimentally measured temperatures are reported after conversion to the IPTS-68 [69COM]. The relation between the absolute temperature T (K, kelvin) and the Celsius temperature t (°C) is defined by t = ( )oT T− where oT = 273.15 K.

II.4 Fundamental physical constants To ensure the consistency with other NEA TDB Reviews, the fundamental physical constants are taken from a publication by CODATA [86COD]. Those relevant to this review are listed in Table II-7. Note that updated values of the fundamental constants can be obtained from CODATA, notably through its Internet site. In most cases, recal-culation of the NEA TDB database entries with the updated values of the fundamental constants will not introduce significant (with respect to their quoted uncertainties) ex-cursions from the current NEA TDB selections.


Table II-7: Fundamental physical constants. These values have been taken from CODATA [86COD]. The digits in parentheses are the one-standard-deviation uncer-tainty in the last digits of the given value.

Quantity Symbol Value Units

speed of light in vacuum c 299 792 458 m·s–1

permeability of vacuum µ0 4π×10–7 = 12.566 370 614… 10–7 N·A–2 permittivity of vacuum є ο 1/µ0 c2 = 8.854 187 817… 10–12 C2·J–1·m–1

Planck constant h 6.626 0755(40) 10–34 J·s elementary charge e 1.602 177 33(49) 10–19C Avogadro constant NA 6.022 1367(36) 1023 mol–1

Faraday constant F 96 485.309(29) C·mol–1

molar gas constant R 8.314 510(70) J·K–1·mol–1

Boltzmann constant, R/NA k 1.380 658(12) 10–23 J·K–1

Non-SI units used with SI:

electron volt, (e/C) J eV 1.602 177 33(49) 10–19 J atomic mass unit, u 1.660 5402(10) 10–27 kg

12

u

11u = ( )12

m m C=

II.5 Uncertainty estimates One of the principal objectives of the NEA-TDB development effort is to provide an idea of the uncertainties associated with the data selected in the reviews. In general the uncertainties should define the range within which the corresponding data can be repro-duced with a probability of 95%. In many cases, a full statistical treatment is limited or impossible due to the availability of only one or few data points. Appendix C describes in detail the procedures used for the assignment and treatment of uncertainties, as well as the propagation of errors and the standard rules for rounding.

II.6 The NEA-TDB system A data base system has been developed at the NEA Data Bank that allows the storage of thermodynamic parameters for individual species as well as for reactions. The structure of the data base system allows consistent derivation of thermodynamic data for individ-ual species from reaction data at standard conditions, as well as internal recalculations of data at standard conditions. If a selected value is changed, all the dependent values will be recalculated consistently. The maintenance of consistency of all the selected data, including their uncertainties (cf. Appendix C), is ensured by the software devel-oped for this purpose at the NEA Data Bank. The literature sources of the data are also stored in the data base.

The following thermodynamic parameters, valid at the reference temperature of 298.15 K and at the standard pressure of 1 bar, are stored in the data base:

II.6 The NEA-TDB system 37

f mο∆ G the standard molar Gibbs energy of formation from the elements in

their reference states (kJ·mol–1)

f mH ο∆ the standard molar enthalpy of formation from the elements in their reference states (kJ·mol–1)

mS ο the standard molar entropy (J·K–1·mol–1)

,mpCο the standard molar heat capacity (J·K–1·mol–1).

For aqueous neutral species and ions, the values of f mGο∆ , f mH ο∆ , mS ο and ,mpCο correspond to the standard partial molar quantities, and for individual aqueous

ions they are relative quantities, defined with respect to the aqueous hydrogen ion, ac-cording to the convention [89COX/WAG] that f mH ο∆ ( +H , T) = 0 and that mS ο ( +H , T) = 0. Furthermore, for an ionised solute B containing any number of different cations and anions:

f m + f m f m+

(B , aq) = (cation, aq) + (anion, aq)H H Hο ο ο± −

−∆ ν ∆ ν ∆∑ ∑

m + m m+

(B , aq) = (cation, aq) + (anion, aq)S S Sο ο ο± −

−ν ν∑ ∑

As the thermodynamic parameters vary as a function of temperature, provision is made for including the compilation of the coefficients of empirical temperature func-tions for these data, as well as the temperature ranges over which they are valid. In many cases the thermodynamic data measured or calculated at several temperatures were published for a particular species, rather than the deduced temperature functions. In these cases, a non-linear regression method is used in this review to obtain the most significant coefficients of the following empirical function for a thermodynamic pa-rameter, X:

2 1 2

3 3

( ) = ln ln

+ + + .

X X X X X X X

XX X X

X T a b T c T d T e T f T g T Tih T j T k TT

− −

−

+ ⋅ + ⋅ + ⋅ + ⋅ + ⋅ + ⋅

⋅ ⋅ + ⋅ (II.52)

Most temperature variations can be described with three or four parameters, a, b and e being the ones most frequently used. In the present review, only ,m ( )p TCο , i.e., the thermal functions of the heat capacities of individual species, are considered and stored in the data base. They refer to the relation:

2 1 2 1/ 2,m ( ) = pC T a b T c T d T e T i Tο − − −+ × + × + × + × + ×

(where the subindices for the coefficients have been dropped) and are listed in Table III–3.

The pressure dependence of thermodynamic data has not been the subject of critical analysis in the present compilation. The reader interested in higher temperatures and pressures, or the pressure dependency of thermodynamic functions for geochemical


applications, is referred to the specialised literature in this area, e.g., [82HAM], [84MAR/MES], [88SHO/HEL], [88TAN/HEL], [89SHO/HEL], [89SHO/HEL2], [90MON], [91AND/CAS].

Selected standard thermodynamic data referring to chemical reactions are also compiled in the data base. A chemical reaction “r”, involving reactants and products ‘B”, can be abbreviated as:

BB

0 Br= ν∑ (II.53)

where the stoichiometric coefficients Brν are positive for products, and negative for re-

actants. The reaction parameters considered in the NEA-TDB system include:

10log rK ο the equilibrium constant of the reaction, logarithmic

r mGο∆ the molar Gibbs energy of reaction (kJ·mol–1)

r mH ο∆ the molar enthalpy of reaction (kJ·mol–1)

r mS ο∆ the molar entropy of reaction (J·K–1·mol–1)

r ,mpCο∆ the molar heat capacity of reaction (J·K–1·mol–1)

The temperature functions of these data, if available, are stored according to Eq.(II.52).

The equilibrium constant, rK ο , is related to r mGο∆ according to the following relation:

r m10log =

R ln(10)rGK

T

οο ∆

−

and can be calculated from the individual values of f m (B)Gο∆ (for example, those given in Table III–1 and Table IV–1), according to:

10 B f mB

1log = (B)R ln(10)

rrK G

Tο ο− ν ∆∑ (II.54).

II.7 Presentation of the selected data The selected data are presented in Chapters III and IV. Unless otherwise indicated, they refer to standard conditions (cf. Section II.3) and 298.15K (25°C) and are provided with an uncertainty which should correspond to the 95% confidence level (see Appendix C).

Chapters III contains a table of selected thermodynamic data for individual compounds and complexes of nickel (Table III–1), a table of selected reaction data (Table III–2) for reactions concerning nickel species and a table containing selected thermal functions of the heat capacities of individual species of nickel (Table III–3). The selection of these data is discussed in Chapter V.

II.7 Presentation of the selected data 39

Chapter IV contains, for auxiliary compounds and complexes that do not con-tain nickel, a table of the thermodynamic data for individual species (Table IV–1) and a table of reaction data (Table IV–2). Most of these values are the CODATA Key Values [89COX/WAG]. The selection of the remaining auxiliary data is discussed in [92GRE/FUG], [99RAR/RAN] and [2001LEM/FUG].

All the selected data presented in Table III–1, Table III–2, Table IV–1 and Table IV–2 are internally consistent. This consistency is maintained by the internal con-sistency verification and recalculation software developed at the NEA Data Bank in conjunction with the NEA-TDB data base system, cf. Section II.6. Therefore, when us-ing the selected data for nickel species, the auxiliary data of Chapter IV must be used together with the data in Chapter III to ensure internal consistency of the data set.

It is important to note that Table III–2 and Table IV–2 include only those spe-cies for which the primary selected data are reaction data. The formation data derived there from and listed in Table III–1 are obtained using auxiliary data, and their uncer-tainties are propagated accordingly. In order to maintain the uncertainties originally assigned to the selected data in this review, the user is advised to make direct use of the reaction data presented in Table III–2 and Table IV–2, rather than taking the derived values in Table III–1 and Table IV–1 to calculate the reaction data with Eq.(II.54). The later approach would imply a twofold propagation of the uncertainties and result in re-action data whose uncertainties would be considerably larger than those originally as-signed.

The thermodynamic data in the selected set refer to a temperature of 298.15 K (25.00°C), but they can be recalculated to other temperatures if the corresponding data (enthalpies, entropies, heat capacities) are available [97PUI/RAR]. For example, the temperature dependence of the standard reaction Gibbs energy as a function of the stan-dard reaction entropy at the reference temperature ( 0T = 298.15 K), and of the heat ca-pacity function is:

0 0

r ,mr m r m 0 r ,m r m 0

(T)( ) = ( ) + (T) d ( ) d ,

T T ppT T

CG T H T C T T S T T

T

οο ο ο ο

∆∆ ∆ ∆ − ∆ +

∫ ∫

and the temperature dependence of the standard equilibrium constant as a function of the standard reaction enthalpy and heat capacity is:

0 0

r m 010 10 0

0

r ,mr ,m

( ) 1 1log ( ) = log ( ) R ln(10)

( ) 1 1 ( ) d + d ,R ln(10) R ln(10)

T Tp

pT T

H TK T K TT T

C TC T T T

T T

οο ο

οο

∆− −

∆

− ∆∫ ∫


where R is the gas constant (cf. Table II-7).

In the case of aqueous species, for which enthalpies of reaction are selected or can be calculated from the selected enthalpies of formation, but for which there are no selected heat capacities, it is in most cases possible to recalculate equilibrium constants to temperatures up to 100 to 150°C, with an additional uncertainty of perhaps about 1 to 2 logarithmic units, due to neglecting the heat capacity contributions to the temperature correction. However, it is important to observe that “new” aqueous species, i.e., species not present in significant amounts at 25°C and therefore not detected, may be significant at higher temperatures, see for example the work by Ciavatta et al. [87CIA/IUL]. Addi-tional high-temperature experiments may therefore be needed in order to ascertain that proper chemical models are used in the modelling of hydrothermal systems. For many species, experimental thermodynamic data are not available to allow a selection of pa-rameters describing the temperature dependence of equilibrium constants and Gibbs energies of formation. The user may find information on various procedures to estimate the temperature dependence of these thermodynamic parameters in [97PUI/RAR]. The thermodynamic data in the selected set refer to infinite dilution for soluble species. Ex-trapolation of an equilibrium constant K, usually measured at high ionic strength, to K ο at I = 0 using activity coefficients γ, is explained in Appendix B. The corresponding Gibbs energy of dilution is:

dil m r m r m = G G Gο∆ ∆ − ∆ (II.55)

r = R lnT γ ±− ∆ (II.56)

Similarly dil mS∆ can be calculated from ln ±γ and its variations with T, while:

2dil m r= R ( ln ) pH T

Tγ ±

∂∆ ∆∂

(II.57)

depends only on the variation of γ with T, which is neglected in this review, when no data on the temperature dependence of γ’s are available. In this case the Gibbs energy of dilution dil mG∆ is entirely assigned to the entropy difference. This entropy of reaction is calculated using the r m r m r m= G H T Sο ο ο∆ ∆ − ∆ equation, the above assumption dil mH∆ = 0, and dil mG∆ .

41

Part II

Tables of selected data

43

Chapter III

III Selected nickel data

This chapter presents the chemical thermodynamic data set for nickel species that has been selected in this review. Table III–1 contains the recommended thermodynamic data of the nickel compounds and species, Table III–2 the recommended thermody-namic data of chemical equilibrium reactions by which the nickel compounds and com-plexes are formed, and Table III–3 the temperature coefficients of the heat capacity data of Table III–1 where available.

The species and reactions in the tables appear in standard order of arrange-ment. Table III–2 contains information only on those reactions for which primary data selections are made in Chapter V of this review. These selected reaction data are used, together with data for key nickel species and auxiliary data selected in this review, to derive the corresponding formation data in Table III–1. The uncertainties associated with values for key nickel species and the auxiliary data are in some cases substantial, leading to comparatively large uncertainties in the formation quantities derived in this manner.

The values of r mGο∆ for many reactions are known more accurately than would be calculated directly from the uncertainties of the f mGο∆ values in Table III–1 and aux-iliary data. The inclusion of a table for reaction data (Table III–2) in this report allows the use of equilibrium constants with total uncertainties that are based directly on the experimental accuracies. This is the main reason for including both Table III–1 and Table III–2.

The selected thermal functions of the heat capacities, listed in Table III–3, refer to the relation ,mpCο (T) = a + b × T + c × T 2 + d × T –1 + i × T -1/2

A detailed discussion of the selection procedure is presented in Chapter V. It may be noted that this chapter contains data on more species or compounds than are present in the tables of Chapter III. The main reasons for this situation are the lack of information for a proper extrapolation of the primary data to standard conditions in some systems and lack of solid primary data in others.

A warning: The addition of any aqueous species and their data to this internally consistent data base can result in a modified data set, which is no longer rigorous and can lead to erroneous results. The situation is similar when gases or solids are added.

III. Selected nickel data

44

Table III–1: Selected thermodynamic data for nickel compounds and complexes. All ionic species listed in this table are aqueous species. Unless noted otherwise, all data refer to the reference temperature of 298.15 K and to the standard state, i.e., a pressure of 0.1 MPa and, for aqueous species, infinite dilution (I = 0). The uncertainties listed below each value represent total uncertainties and correspond in principle to the statisti-cally defined 95% confidence interval. Values obtained from internal calculation, cf. footnotes (a) and (b), are rounded at the third digit after the decimal point and may therefore not be exactly identical to those given in Part V. Systematically, all the values are presented with three digits after the decimal point, regardless of the significance of these digits. The data presented in this table are available on computer media from the OECD Nuclear Energy Agency.

f mGο∆ f mH ο∆ mS ο ,mpCο Compound

(kJ · mol–1) (kJ · mol–1) (J · K–1 · mol–1) (J · K–1 · mol–1) Ni(cr)

0.000 ±0.000

0.000±0.000

29.870±0.200

26.070±0.100

(c)

Ni(g)

384.686 ±8.400

(a) 430.100±8.400

182.190±0.080

23.360±0.100

Ni(l)

14.035 ±0.272

(a) 17.500±0.250

41.490±0.300

Ni2+

– 45.773 ±0.771

(b) – 55.012±0.878

(a) – 131.800±1.400

(d) – 46.100±7.500

NiO(cr)

– 211.660 ±0.422

(a) – 239.700±0.400

38.400±0.400

44.400±0.100

(c)

NiOH+

– 228.458 ±1.111

(b) – 287.042±1.914

(b) – 64.046±6.454

(b)

β-Ni(OH)2

– 457.100 ±1.400

– 542.300±1.500

80.000±0.800

82.010±0.300

Ni(OH)3 –

– 590.519

±9.735 (b)

Ni2OH3+

– 268.181 ±5.913

(b) – 349.955±6.252

(b) – 242.636±27.917

(b)

Ni4(OH)44+

– 974.567

±3.203 (b) – 1173.370

±10.601 (b) – 137.004

±34.126 (b)

NiF+

– 335.458 ±1.132

(b) – 380.862±3.193

(b) – 86.360±10.305

(b)

NiF2(cr)

– 609.852 ±8.001

(a) – 657.300±8.000

73.520±0.400

63.210±2.000

(c)

NiCl+

– 177.447 ±3.512

(b)



45

Table III–1 (continued)


(kJ · mol–1) (kJ · mol–1) (J · K–1 · mol–1) (J · K–1 · mol–1) NiCl2(cr)

– 258.743 ±0.154

(a) – 304.900±0.110

98.140±0.300

71.670±0.300

(c)

NiCl2. 2 H2O(cr)

– 754.382 ±1.256

(b) – 913.374±1.653

(a) 186.200±3.600

230.000±15.000

NiCl2. 4 H2O(cr)

– 1234.946 ±1.040

(b) – 1514.048±9.179

(b) 249.861±30.980

(b)

NiCl2. 6 H2O(cr)

– 1713.669 ±0.846

(b) – 2104.700±1.800

340.964±6.673

(a)

NiBr2(cr)

– 195.408 ±2.421

(a) – 213.500±2.400

121.400±1.000

75.400±1.000

(c)

NiI2(cr)

– 94.360 ±0.898

(a) – 96.420±0.840

139.100±1.000

77.280±1.000

(c)

β-Ni(IO3)2

– 323.735 ±1.743

(b) – 487.112±4.216

(b) 213.496±14.075

(b)

Ni(IO3)2. 2 H2O(cr)

– 802.068 ±1.832

(b) – 1087.672±5.175

(b) 270.058±17.403

(b)

α-NiS

– 87.792 ±1.007

(a) – 88.100±1.000

60.890±0.340

49.760±0.300

(c)

β-NiS

– 91.330 ±1.007

(a) – 94.000±1.000

52.970±0.330

47.080±0.300

(c)

NiS2(cr)

– 123.832 ±7.396

(a) – 128.000±7.000

80.000±8.000

Ni3S2(cr)

– 211.172 ±1.624

(a) – 217.200±1.600

133.500±0.700

118.200±0.800

(c)

Ni9S8(cr)

– 746.803 ±9.255

(a) – 760.000±9.000

481.000±7.000

407.000±5.000

(c)

NiHS+

– 63.098 ±2.524

(b)

NiSO4(cr)

– 762.688 ±1.572

(a) – 873.280±1.570

101.300±0.200

97.600±0.100

(c)

α-NiSO4. 6 H2O

– 2225.467 ±1.059

(a) – 2683.817±1.014

(b) 334.450±1.000

327.900±1.000

β-NiSO4. 6 H2O

– 2224.915 ±1.815

(a) – 2677.287±1.034

(b) 354.500±5.000

NiSO4. 7 H2O(cr)

– 2462.699 ±0.929

(b) – 2977.327±0.978

(b) 378.950±1.000

(b) 364.600±4.000



46



(kJ · mol–1) (kJ · mol–1) (J · K–1 · mol–1) (J · K–1 · mol–1) NiSO4(aq)

– 803.191 ±0.893

(b) – 958.692±1.260

(b) – 49.326±3.135

(b)

Ni(NO3)2. 2 H2O(cr)

– 1012.732±1.555

(b)

Ni(NO3)2. 4 H2O(cr)

– 1620.392±1.441

(b)

Ni(NO3)2. 6 H2O(cr)

– 2214.452±1.571

(b) 463.000±2.000

NiNO3+

– 159.421

±5.775 (b)

NiHPO4(aq)

– 1159.167 ±1.820

(b)

NiP2O72–

– 2031.107

±4.843 (b)

HNiP2O7 –

– 2064.270

±4.766 (b)

NiAs(cr)

– 66.583 ±4.606

(a) – 70.820±4.240

50.760±6.000

50.800±4.000

(c)

NiAs2(cr)

– 90.100±8.000

Ni5As2(cr)

– 237.638 ±21.914

(a) – 244.600±20.000

196.200±30.000

Ni11As8(cr)

– 715.758 ±39.338

(a) – 743.000±35.000

518.000±60.000

490.000±40.000

(c)

Ni3(AsO3)2(cr, hyd)

– 1253.627 ±9.251

(b)

Ni3(AsO4)2. 8 H2O(cr)

– 3491.557 ±8.823

(b) – 4179.000±20.000

540.771±73.331

(a)

NiHAsO4(aq)

– 776.918 ±4.426

(b)

NiCO3(cr)

– 636.416 ±1.292

(b) – 713.320±1.424

(a) 85.400±2.000

90.300±4.100

(c)

NiCO3. 5.5 H2O(cr)

– 1920.881 ±1.004

(b) – 2312.992±3.147

(a) 311.100±10.000

405.400±50.000

NiCO3(aq)

– 597.646 ±2.441

(b)



47



(kJ · mol–1) (kJ · mol–1) (J · K–1 · mol–1) (J · K–1 · mol–1) Ni(CN)4

2–

449.601 ±10.129

(b) 353.688±14.745

(b) 245.032±36.559

(b)

Ni(CN)53–

626.244

±12.938 (b) 490.638

±19.449 (b) 278.786

±51.082 (b)

NiSCN+

36.596 ±4.080

(b) 9.588±6.463

(b) 7.543±25.373

(b)

Ni(SCN)2(aq)

124.273 ±8.047

(b) 76.788±11.348

(b) 137.802±46.516

(b)

Ni(SCN)3 –

215.089

±12.069 (b) 145.188

±15.645 (b) 261.556

±66.172 (b)

Ni2SiO4(oliv)

– 1288.441 ±5.002

(a) – 1396.000±5.000

128.100±0.200

123.200±0.500

(c)

(a) Value calculated internally using equation f m f m m,ii

G H T Sο ο ο∆ = ∆ − ∑ .

(b) Value calculated internally from reaction data (see Table III–2).

(c) Temperature coefficients of this function are listed in Table III–3.

(d) Value calculated internally from reaction data for a different key species.


48

Table III–2: Selected thermodynamic data for reactions involving nickel compounds and complexes. All ionic species listed in this table are aqueous species. Unless noted otherwise, all data refer to the reference temperature of 298.15 K and to the standard state, i.e., a pressure of 0.1 MPa and, for aqueous species, infinite dilution (I = 0). The uncertainties listed below each value represent total uncertainties and correspond in principal to the statistically defined 95% confidence interval. Values obtained from in-ternal calculation, cf. footnote (a), are rounded at the third digit after the decimal point and may therefore not be exactly identical to those given in Part V. Systematically, all the values are presented with three digits after the decimal point, regardless of the sig-nificance of these digits. The data presented in this table are available on computer me-dia from the OECD Nuclear Energy Agency.

Species Reaction log 10Ko ∆r G m

o ∆r H m o ∆r S m

o (kJ . mol–1) (kJ . mol–1) (J . K–1 . mol–1)

Ni2+ Ni(cr) Ni2+ + 2 e – 8.019

±0.135 (b) – 45.773

±0.771

NiOH+ H2O(l) + Ni2+ H+ + NiOH+ – 9.540

±0.140 54.455

±0.799 53.800

±1.700 – 2.196

±6.300 (a)

Ni(OH)3 – 3H2O(l) + Ni2+ 3H+ + Ni(OH)3

– – 29.200

±1.700 166.675

±9.704

Ni2OH3+ H2O(l) + 2Ni2+ H+ + Ni2OH3+ – 10.600

±1.000 60.505

±5.708 45.900

±6.000 – 48.986

±27.776 (a)

Ni4(OH)44+ 4H2O(l) + 4Ni2+ 4H+ + Ni4(OH)4

4+ – 27.520

±0.150 157.085

±0.856 190.000

±10.000 110.396

±33.663 (a)

NiF+ F – + Ni2+ NiF+ 1.430

±0.080 – 8.163

±0.457 9.500

±3.000 59.240

±10.178 (a)

NiCl+ Cl – + Ni2+ NiCl+ 0.080

±0.600 – 0.457

±3.425

NiCl2. 2H2O(cr) NiCl2. 4H2O(cr) 2H2O(g) + NiCl22H2O(cr) – 4.099

±0.123 23.400

±0.700



49

Table III–2 (continued) Species Reaction log 10Ko ∆r G m



NiCl2. 4H2O(cr) NiCl2. 6H2O(cr) 2H2O(g) + NiCl24H2O(cr) – 3.777

±0.105 (c) 21.560

±0.600 107.000

±9.000 286.567

±30.253 (a)

NiCl2. 6H2O(cr) 2Cl – + 6H2O(l) + Ni2+ NiCl2. 6H2O(cr) – 3.045

±0.014 (c) 17.380

±0.080

β– Ni(IO3)2 2IO3 – + Ni2+ β– Ni(IO3)2

4.430±0.020

– 25.287±0.114

7.300±4.000

109.296±13.422

(a)

Ni(IO3)2. 2H2O(cr) 2H2O(l) + 2IO3 – + Ni2+ Ni(IO3)2. 2H2O(cr)

5.140±0.100

– 29.339±0.571

– 21.600±5.000

25.958±16.879

(a)

NiHS+ HS – + Ni2+ NiHS+ 5.180

±0.200 – 29.568

±1.142

α– NiSO4. 6H2O 6H2O(l) + Ni2+ + SO42– α– NiSO4. 6H2O

– 4.485±0.200

β– NiSO4. 6H2O α– NiSO4. 6H2O β– NiSO46H2O 6.530

±0.200

NiSO4. 7H2O(cr) 7H2O(l) + Ni2+ + SO42– NiSO4. 7H2O(cr)

2.267±0.019

– 12.940±0.110

– 12.167±0.226

2.593±0.843

(a)

NiSO4(aq) Ni2+ + SO42– NiSO4(aq)

2.350±0.030

– 13.414±0.171

5.660±0.810

63.974±2.777

(a)

Ni(NO3)2. 2H2O(cr) 2H2O(l) + 2NO3 – + Ni2+ Ni(NO3)2. 2H2O(cr)

27.640±1.000


– 8.360±0.800



50





– 30.760±1.000

NiNO3+ NO3

– + Ni2+ NiNO3+

0.500±1.000

– 2.854±5.708

NiHPO4(aq) HPO42– + Ni2+ NiHPO4(aq)

3.050±0.090

– 17.410±0.514

NiP2O72– Ni2+ + P2O7

4– NiP2O72–

8.730±0.250

– 49.831±1.427

30.600±10.000

269.768±33.880

(a)

HNiP2O7 – HP2O7

3– + Ni2+ HNiP2O7 –

5.140±0.250

– 29.339±1.427

47.900±15.000

259.062±50.537

(a)

Ni3(AsO3)2(cr, hyd) 2H2O(l) + 2HAsO2(aq) + 3Ni2+ 6H+ + Ni3(AsO3)2(cr, hyd) – 28.700

±0.700 163.821

±3.996

Ni3(AsO4)2. 8H2O(cr) 2AsO43– + 8H2O(l) + 3Ni2+ Ni3(AsO4)2. 8H2O(cr)

28.100±0.500

– 160.396±2.854

NiHAsO4(aq) HAsO42– + Ni2+ NiHAsO4(aq)

2.900±0.300

– 16.553±1.712

NiCO3(cr) CO2(g) + H2O(l) + Ni2+ 2H+ + NiCO3(cr) – 7.160

±0.180 40.870

±1.027

NiCO3. 5.5H2O(cr) CO2(g) + 6.5H2O(l) + Ni2+ 2H+ + NiCO3. 5.5H2O(cr) – 10.630

±0.100 60.676

±0.571

NiCO3(aq) CO32– + Ni2+ NiCO3(aq)

4.200±0.400

– 23.974±2.283



51




Ni(CN)42– 4CN – + Ni2+ Ni(CN)4

2– 30.200

±0.120 – 172.383

±0.685 – 180.700

±4.000 – 27.896

±13.611 (a)

Ni(CN)53– 5CN – + Ni2+ Ni(CN)5

3– 28.500

±0.500 – 162.679

±2.854 – 191.100

±8.000 – 95.324

±28.489 (a)

NiSCN+ Ni2+ + SCN – NiSCN+ 1.810

±0.040 – 10.332

±0.228 – 11.800

±5.000 – 4.925

±16.788 (a)

Ni(SCN)2(aq) Ni2+ + 2SCN – Ni(SCN)2(aq) 2.690

±0.070 – 15.355

±0.400 – 21.000

±8.000 – 18.935

±26.866 (a)

Ni(SCN)3 – Ni2+ + 3SCN – Ni(SCN)3

– 3.020

±0.180 – 17.238

±1.027 – 29.000

±10.000 – 39.449

±33.717 (a)

(a) Value calculated internally using r m r m r mG H T Sο ο ο∆ = ∆ − ∆ .

(b) Value calculated from a selected standard potential.

(c) Value of 10log K ο calculated internally from r mGο∆ .


52

Table III–3: Selected temperature coefficients for heat capacities marked with (c) in Table III–1, according to the form 2 1 2 1/ 2

,m ( )pC T a bT cT dT eT iTο − − −= + + + + + . The functions are valid between the temperatures Tmin and Tmax (in K). The notation E±nm indicates the power of 10. Units for ,mpCο are J · K–1 · mol–1.

Compound a b c d / i* e Tmin Tmax

Ni(cr) 3.36472E+01 – 2.47530E– 02 4.35458E– 05 0 – 3.61955E+05 298 450

Ni(cr) – 4.93235E+01 1.676480E–01– 7.48518E–05 3.76214E+06 450 600

Ni(cr) 1.05023E+05 – 2.29221E+02 1.40772E–01 0 – 6.52809E+09 600 631

Ni(cr) 5.29406E+04 – 1.65828E+02 1.29936E–01 0 0 631 640

Ni(cr) 1.82638E+02 – 4.20407E–01 2.90521E– 04 0 0 640 690

Ni(cr) 1.23677E+01 2.22800E–02 4.45300E– 06 0 2.46766E+06 690 1728

NiO(cr) 4.11064E+03 – 5.30241E+00 3.52061E– 03 5.30393E+04(i) 2.43067E+07 298 519

NiO(cr) – 8.776E+00 4.2232E–02 –7.5267E– 06 7.87250E+02(i) 3.6067E+06 519 1800

NiF2(cr) 6.55050E+01 1.53770E– 02 0 0 – 6.11603E+05 250 1450

NiCl2(cr) 7.37949E+01 1.22269E– 02 0 0 – 5.12775E+05 298 1281

NiBr2(cr) 7.23163E+01 1.46270E– 02 0 0 – 1.13805E+05 298 1200

NiI2(cr) 7.35775E+01 1.64250E– 02 0 0 – 1.05775E+05 200 550

α-NiS 4.66760E+01 1.99807E– 02 0 0 – 2.55499E+05 298 660

α-NiS 5.00783E+01 1.49007E– 02 0 0 – 3.36916E+05 660 1000

β-NiS 5.40080E+01 – 1.43700E– 02 3.22820E– 05 0 – 4.90363E+05 298 660

Ni3S2(cr) 1.65794E+02 – 1.03160E– 01 1.20000E– 04 0 – 2.44462E+06 298 835

Ni9S8(cr) 1.07925E+03 – 2.05994E+00 1.98000E– 03 0 – 2.08396E+07 298 640

NiSO4(cr) 1.18934E+02 3.67610E– 02 0 0 – 2.86810E+06 298 1000

NiAs(cr) 3.65400E+01 4.18700E– 02 0 0 1.58000E+05 298 800

Ni11As8(cr) 3.80120E+02 3.67270E– 01 0 0 1.80000E+03 298 800

NiCO3(cr) 8.87010E+01 3.89100E– 02 0 0 – 1.23400E+06 298 700

Ni2SiO4(oliv) 1.47616E+02 4.86860E– 02 – 1.50143E– 05 0 – 3.34214E+06 270 1570

* : One single column is used for the two coefficients d and i. The coefficient con-cerned is indicated in parentheses behind each value, when any of the two values is different from zero.

53

Chapter IV

IV Selected auxiliary data

This chapter presents the chemical thermodynamic data for auxiliary compounds and complexes which are used within the NEA’s TDB project. Most of these auxiliary spe-cies are used in the evaluation of the recommended nickel data in Table III–1, Table III–2, and Table III–3. It is therefore essential to always use these auxiliary data in con-junction with the selected data for nickel. The use of other auxiliary data can lead to inconsistencies and erroneous results.

The values in the tables of this chapter are either CODATA Key Values, taken from [89COX/WAG], or were evaluated within the NEA’s TDB project, as described in the corresponding chapters of the uranium review [92GRE/FUG], the technetium re-view [99RAR/RAN], the Update review [2003GUI/FAN], the selenium review [2005OLI/NOL] and the present Review.

Table IV–1 contains the selected thermodynamic data of the auxiliary species and Table IV–2 the selected thermodynamic data of chemical reactions involving auxil-iary species. The reason for listing both reaction data and entropies, enthalpies and Gibbs energies of formation is, as described in Chapter III, that uncertainties in reaction data are often smaller than the derived mS ο , f mH ο∆ and f mGο∆ , due to uncertainty accu-mulation during the calculations.

All data in Table IV–1 and Table IV–2 refer to a temperature of 298.15 K, the standard state pressure of 0.1 MPa and, for aqueous species and reactions, to the infinite dilution standard state (I = 0).

The uncertainties listed below each reaction value in Table IV–2 are total un-certainties, and correspond mainly to the statistically defined 95% confidence interval. The uncertainties listed below each value in Table IV–1 have the following signifi-cance:


54

• for CODATA values from [89COX/WAG], the ± terms have the meaning: “it is probable, but not at all certain, that the true values of the thermodynamic quantities differ from the recommended values given in this report by no more than twice the ± terms attached to the recommended values”.

• for values from [92GRE/FUG], [99RAR/RAN], [2003GUI/FAN], the on-going NEA TDB review on Selenium (to be published) and the present Review, the ± terms are derived from total uncertainties in the corresponding equilibrium constant of reaction (cf. Table IV–2), and from the ± terms listed for the nec-essary CODATA key values.

CODATA [89COX/WAG] values are available for CO2(g), 3HCO− , 23CO − ,

2 4H PO− and 24HPO − . From the values given for f mH ο∆ and mS ο the values of f mGο∆

and, consequently, all the relevant equilibrium constants and enthalpy changes can be calculated. The propagation of errors during this procedure, however, leads to uncer-tainties in the resulting equilibrium constants that are significantly higher than those obtained from experimental determination of the constants. Therefore, reaction data for CO2(g), 3HCO− , 2

3CO − , which were absent form the corresponding Table IV–2 in [92GRE/FUG], are included in this volume to provide the user of selected data for spe-cies of nickel (cf. Chapter III) with the data needed to obtain the lowest possible uncer-tainties on reaction properties.

Note that the values in Table IV–1 and Table IV–2 may contain more digits than those listed in either [89COX/WAG] or in the discussions on auxiliary data selec-tion of NEA TDB reviews, because the data in the present chapter are retrieved directly from the computerised data base and rounded to three digits after the decimal point throughout.


55

Table IV–1: Selected thermodynamic data for auxiliary compounds and complexes adopted in the NEA TDB project, including the CODATA Key Values [89COX/WAG]. All ionic species listed in this table are aqueous species. Unless noted otherwise, all data refer to 298.15 K and a pressure of 0.1 MPa and, for aqueous spe-cies, a standard state of infinite dilution (I = 0). The uncertainties listed below each value represent total uncertainties and correspond in principle to the statistically defined 95% confidence interval. Values in bold typeface are CODATA Key Values and are taken directly from Ref. [89COX/WAG] without further evaluation. Values obtained from internal calculation, cf. footnotes (a) and (b), are rounded at the third digit after the decimal point. Systematically, all the values are presented with three digits after the decimal point, regardless of the significance of these digits. The reference listed for each entry in this table indicates the NEA TDB Review where the corresponding data have been adopted as NEA TDB Auxiliary data. The data presented in this table are available on computer media from the OECD Nuclear Energy Agency.

f mGο∆ f mH ο∆ mS ο ,mpCο Compound and NEA TDB Review (kJ · mol–1) (kJ · mol–1) (J · K–1 · mol–1) (J · K–1 · mol–1) O(g) [92GRE/FUG]

231.743 ±0.100

(a) 249.180±0.100

161.059±0.003

21.912±0.001

O2(g) [92GRE/FUG]

0.000 0.000 205.152

±0.005 29.378

±0.003

H(g) [92GRE/FUG]

203.276 ±0.006

(a) 217.998±0.006

114.717±0.002

20.786±0.001

H+ [92GRE/FUG]

0.000 0.000 0.000 0.000

H2(g) [92GRE/FUG]

0.000 0.000 130.680

±0.003 28.836

±0.002

OH – [92GRE/FUG]

– 157.220 ±0.072

(a) – 230.015±0.040

– 10.900±0.200

H2O(g) [92GRE/FUG]

– 228.582 ±0.040

(a) – 241.826±0.040

188.835±0.010

33.609±0.030

H2O(l) [92GRE/FUG]

– 237.140 ±0.041

(a) – 285.830±0.040

69.950±0.030

75.351±0.080

H2O2(aq) [92GRE/FUG]

– 191.170±0.100

He(g) [92GRE/FUG]

0.000 0.000 126.153

±0.002 20.786

±0.001

Ne(g) [92GRE/FUG]

0.000 0.000 146.328

±0.003 20.786

±0.001



56

Table IV–1: (continued)

f mGο∆ f mH ο∆ mS ο ,mpCο Compound and NEA TDB Review (kJ · mol–1) (kJ · mol–1) (J · K–1 · mol–1) (J · K–1 · mol–1) Ar(g) [92GRE/FUG]

0.000 0.000 154.846

±0.003 20.786

±0.001

Kr(g) [92GRE/FUG]

0.000 0.000 164.085

±0.003 20.786

±0.001

Xe(g) [92GRE/FUG]

0.000 0.000 169.685

±0.003 20.786

±0.001

F(g) [92GRE/FUG]

62.280 ±0.300

(a) 79.380±0.300

158.751±0.004

22.746±0.002

F – [92GRE/FUG]

– 281.523 ±0.692

(a) – 335.350±0.650

– 13.800±0.800

F2(g) [92GRE/FUG]

0.000 0.000 202.791

±0.005 31.304

±0.002

HF(aq) [92GRE/FUG]

– 299.675 ±0.702

– 323.150±0.716

88.000±3.362

(a)

HF(g) [92GRE/FUG]

– 275.400 ±0.700

(a) – 273.300±0.700

173.779±0.003

29.137±0.002

HF2 –

[92GRE/FUG] – 583.709

±1.200 – 655.500

±2.221 92.683

±8.469 (a)

Cl(g) [92GRE/FUG]

105.305 ±0.008

(a) 121.301±0.008

165.190±0.004

21.838±0.001

Cl – [92GRE/FUG]

– 131.217 ±0.117

(a) – 167.080±0.100

56.600±0.200

Cl2(g) [92GRE/FUG]

0.000 0.000 223.081

±0.010 33.949

±0.002

ClO – [92GRE/FUG]

– 37.669 ±0.962

ClO2 –

[92GRE/FUG] 10.250

±4.044

ClO3 –

[92GRE/FUG] – 7.903

±1.342 (a) – 104.000

±1.000 162.300

±3.000

ClO4 –

[92GRE/FUG] – 7.890

±0.600 (a) – 128.100

±0.400 184.000

±1.500

HCl(g) [92GRE/FUG]

– 95.298 ±0.100

(a) – 92.310±0.100

186.902±0.005

29.136±0.002

HClO(aq) [92GRE/FUG]

– 80.023 ±0.613

HClO2(aq) [92GRE/FUG]

– 0.938 ±4.043



57


f mGο∆ f mH ο∆ mS ο ,mpCο Compound and NEA TDB Review (kJ · mol–1) (kJ · mol–1) (J · K–1 · mol–1) (J · K–1 · mol–1) Br(g) [92GRE/FUG]

82.379 ±0.128

(a) 111.870±0.120

175.018±0.004

20.786±0.001

Br – [92GRE/FUG]

– 103.850 ±0.167

(a) – 121.410±0.150

82.550±0.200

Br2(aq) [92GRE/FUG]

4.900 ±1.000

Br2(g) [92GRE/FUG]

3.105 ±0.142

(a) 30.910±0.110

245.468±0.005

36.057±0.002

Br2(l) [92GRE/FUG]

0.000 ±0.000

0.000±0.000

152.210±0.300

BrO – [92GRE/FUG]

– 32.095 ±1.537

BrO3 –

[92GRE/FUG] 19.070

±0.634 (a) – 66.700

±0.500 161.500

±1.300

HBr(g) [92GRE/FUG]

– 53.361 ±0.166

(a) – 36.290±0.160

198.700±0.004

29.141±0.003

HBrO(aq) [92GRE/FUG]

– 81.356 ±1.527

(b)

I(g) [92GRE/FUG]

70.172 ±0.060

(a) 106.760±0.040

180.787±0.004

20.786±0.001

I – [92GRE/FUG]

– 51.724 ±0.112

(a) – 56.780±0.050

106.450±0.300

I2(cr) [92GRE/FUG]

0.000 ±0.000

0.000±0.000

116.140±0.300

I2(g) [92GRE/FUG]

19.323 ±0.120

(a) 62.420±0.080

260.687±0.005

36.888±0.002

IO3 –

[92GRE/FUG] – 126.338

±0.779 (a) – 219.700

±0.500 118.000

±2.000

HI(g) [92GRE/FUG]

1.700 ±0.110

(a) 26.500±0.100

206.590±0.004

29.157±0.003

HIO3(aq) [92GRE/FUG]

– 130.836 ±0.797

S(cr)(orthorhombic) [92GRE/FUG]

0.000 ±0.000

0.000±0.000

32.054±0.050

22.750±0.050

S(g) [92GRE/FUG]

236.689 ±0.151

(a) 277.170±0.150

167.829±0.006

23.674±0.001

S2– [92GRE/FUG]

120.695 ±11.610



58


f mGο∆ f mH ο∆ mS ο ,mpCο Compound and NEA TDB Review (kJ · mol–1) (kJ · mol–1) (J · K–1 · mol–1) (J · K–1 · mol–1) S2(g) [92GRE/FUG]

79.686 ±0.301

(a) 128.600±0.300

228.167±0.010

32.505±0.010

SO2(g) [92GRE/FUG]

– 300.095 ±0.201

(a) – 296.810±0.200

248.223±0.050

39.842±0.020

SO32–

[92GRE/FUG] – 487.472

±4.020

S2O32–

[92GRE/FUG] – 519.291

±11.345

SO42–

[92GRE/FUG] – 744.004

±0.418 (a) – 909.340

±0.400 18.500

±0.400

HS – [92GRE/FUG]

12.243 ±2.115

(a) – 16.300±1.500

67.000±5.000

H2S(aq) [92GRE/FUG]

– 27.648 ±2.115

(a) – 38.600±1.500

126.000±5.000

H2S(g) [92GRE/FUG]

– 33.443 ±0.500

(a) – 20.600±0.500

205.810±0.050

34.248±0.010

HSO3 –

[92GRE/FUG] – 528.684

±4.046

HS2O3 –

[92GRE/FUG] – 528.366

±11.377

H2SO3(aq) [92GRE/FUG]

– 539.187 ±4.072

HSO4 –

[92GRE/FUG] – 755.315

±1.342 (a) – 886.900

±1.000 131.700

±3.000

HSO4 –

[2001LEM/FUG] – 755.315

±1.342 – 886.900

±1.000 131.700

±3.000

Te(cr) [92GRE/FUG]

0.000 ±0.000

0.000±0.000

49.221±0.050

25.550±0.100

TeO2(cr) [2003GUI/FAN]

– 265.996 ±2.500

(a) – 321.000±2.500

69.890±0.150

60.670±0.150

(c)

N(g) [92GRE/FUG]

455.537 ±0.400

(a) 472.680±0.400

153.301±0.003

20.786±0.001

N2(g) [92GRE/FUG]

0.000 ±0.000

0.000±0.000

191.609±0.004

29.124±0.001

N3 –

[92GRE/FUG] 348.200

±2.000 275.140

±1.000 107.710

±7.500 (a)

NO3 –

[92GRE/FUG] – 110.794

±0.417 (a) – 206.850

±0.400 146.700

±0.400



59


f mGο∆ f mH ο∆ mS ο ,mpCο Compound and NEA TDB Review (kJ · mol–1) (kJ · mol–1) (J · K–1 · mol–1) (J · K–1 · mol–1) HN3(aq) [92GRE/FUG]

321.372 ±2.051

260.140±10.050

147.381±34.403

NH3(aq) [92GRE/FUG]

– 26.673 ±0.305

– 81.170±0.326

109.040±0.913

NH3(g) [92GRE/FUG]

– 16.407 ±0.350

(a) – 45.940±0.350

192.770±0.050

35.630±0.005

NH4+

[92GRE/FUG] – 79.398

±0.278 (a) – 133.260

±0.250 111.170

±0.400

P(am)(red) [92GRE/FUG]

– 7.500±2.000

P(cr)(white, cubic) [92GRE/FUG]

0.000 0.000 41.090

±0.250 23.824

±0.200

P(g) [92GRE/FUG]

280.093 ±1.003

(a) 316.500±1.000

163.199±0.003

20.786±0.001

P2(g) [92GRE/FUG]

103.469 ±2.006

(a) 144.000±2.000

218.123±0.004

32.032±0.002

P4(g) [92GRE/FUG]

24.419 ±0.448

(a) 58.900±0.300

280.010±0.500

67.081±1.500

PO43–

[92GRE/FUG] – 1025.491

±1.576 – 1284.400

±4.085 – 220.970

±12.846

P2O74–

[92GRE/FUG] – 1935.503

±4.563

HPO42–

[92GRE/FUG] – 1095.985

±1.567 (a) – 1299.000

±1.500 – 33.500

±1.500

H2PO4 –

[92GRE/FUG] – 1137.152

±1.567 (a) – 1302.600

±1.500 92.500

±1.500

H3PO4(aq) [92GRE/FUG]

– 1149.367 ±1.576

– 1294.120±1.616

161.912±2.575

HP2O73–

[92GRE/FUG] – 1989.158

±4.482

H2P2O72–

[92GRE/FUG] – 2027.117

±4.445

H3P2O7 –

[92GRE/FUG] – 2039.960

±4.362

H4P2O7(aq) [92GRE/FUG]

– 2045.668 ±3.299

– 2280.210±3.383

274.919±6.954

As(cr) [92GRE/FUG]

0.000 ±0.000

0.000±0.000

35.100±0.600

24.640±0.500



60


f mGο∆ f mH ο∆ mS ο ,mpCο Compound and NEA TDB Review (kJ · mol–1) (kJ · mol–1) (J · K–1 · mol–1) (J · K–1 · mol–1) AsO2

– [92GRE/FUG]

– 350.022 ±4.008

(a) – 429.030±4.000

40.600±0.600

AsO43–

[92GRE/FUG] – 648.360

±4.008 (a) – 888.140

±4.000 – 162.800

±0.600

As2O5(cr) [92GRE/FUG]

– 782.449 ±8.016

(a) – 924.870±8.000

105.400±1.200

116.520±0.800

As4O6(cubic) [92GRE/FUG]

– 1152.445 ±16.032

(a) – 1313.940±16.000

214.200±2.400

191.290±0.800

As4O6(monoclinic) [92GRE/FUG]

– 1154.009 ±16.041

(a) – 1309.600±16.000

234.000±3.000

As4O6(g) This Review

– 1092.716 ±16.116

(a) – 1196.250±16.000

408.600±6.000

HAsO2(aq) [92GRE/FUG]

– 402.925 ±4.008

(a) – 456.500±4.000

125.900±0.600

H2AsO3 –

[92GRE/FUG] – 587.078

±4.008 (a) – 714.790

±4.000 110.500

±0.600

H3AsO3(aq) [92GRE/FUG]

– 639.681 ±4.015

(a) – 742.200±4.000

195.000±1.000

HAsO42–

[92GRE/FUG] – 714.592

±4.008 (a) – 906.340

±4.000 – 1.700

±0.600

H2AsO4 –

[92GRE/FUG] – 753.203

±4.015 (a) – 909.560

±4.000 117.000

±1.000

H3AsO4(aq) [92GRE/FUG]

– 766.119 ±4.015

(a) – 902.500±4.000

184.000±1.000

(As2O5)3. 5 H2O(cr) [92GRE/FUG]

– 4248.400±24.000

Sb(cr) [92GRE/FUG]

0.000 ±0.000

0.000±0.000

45.520±0.210

25.260±0.200

Bi(cr) [2001LEM/FUG]

0.000 ±0.000

0.000±0.000

56.740±0.420

25.410±0.200

C(cr) [92GRE/FUG]

0.000 ±0.000

0.000±0.000

5.740±0.100

8.517±0.080

C(g) [92GRE/FUG]

671.254 ±0.451

(a) 716.680±0.450

158.100±0.003

20.839±0.001

CO(g) [92GRE/FUG]

– 137.168 ±0.173

(a) – 110.530±0.170

197.660±0.004

29.141±0.002

CO2(aq) [92GRE/FUG]

– 385.970 ±0.270

(a) – 413.260±0.200

119.360±0.600



61


f mGο∆ f mH ο∆ mS ο ,mpCο Compound and NEA TDB Review (kJ · mol–1) (kJ · mol–1) (J · K–1 · mol–1) (J · K–1 · mol–1) CO2(g) [92GRE/FUG]

– 394.373 ±0.133

(a) – 393.510±0.130

213.785±0.010

37.135±0.002

CO32–

[92GRE/FUG] – 527.900

±0.390 (a) – 675.230

±0.250 – 50.000

±1.000

HCO3 –

[92GRE/FUG] – 586.845

±0.251 (a) – 689.930

±0.200 98.400

±0.500

CN – [2005OLI/NOL]

166.939 ±2.519

(b) 147.350±3.541

(b) 101.182±8.475

(b)

HCN(aq) [2005OLI/NOL]

114.368 ±2.517

(b) 103.750±3.536

(b) 131.271±8.440

(b)

HCN(g) [2005OLI/NOL]

119.517 ±2.500

(a) 129.900±2.500

201.710±0.100

SCN – [92GRE/FUG]

92.700 ±4.000

76.400±4.000

144.268±18.974

(a)

Si(cr) [92GRE/FUG]

0.000 0.000 18.810

±0.080 19.789

±0.030

Si(g) [92GRE/FUG]

405.525 ±8.000

(a) 450.000±8.000

167.981±0.004

22.251±0.001

SiO2(α– quartz) [92GRE/FUG]

– 856.287 ±1.002

(a) – 910.700±1.000

41.460±0.200

44.602±0.300

SiO2(OH)22–

[92GRE/FUG] – 1175.651

±1.265 – 1381.960

±15.330 – 1.488

±51.592

SiO(OH)3 –

[92GRE/FUG] – 1251.740

±1.162 – 1431.360

±3.743 88.024

±13.144

Si(OH)4(aq) [92GRE/FUG]

– 1307.735 ±1.156

(b) – 1456.960±3.163

(b) 189.973±11.296

(b)

Si2O3(OH)42–

[92GRE/FUG] – 2269.878

±2.878

Si2O2(OH)5 –

[92GRE/FUG] – 2332.096

±2.878

Si3O6(OH)33–

[92GRE/FUG] – 3048.536

±3.870

Si3O5(OH)53–

[92GRE/FUG] – 3291.955

±3.869

Si4O8(OH)44–

[92GRE/FUG] – 4075.179

±5.437

Si4O7(OH)53–

[92GRE/FUG] – 4136.826

±4.934



62


f mGο∆ f mH ο∆ mS ο ,mpCο Compound and NEA TDB Review (kJ · mol–1) (kJ · mol–1) (J · K–1 · mol–1) (J · K–1 · mol–1) SiF4(g) [92GRE/FUG]

– 1572.773 ±0.814

(a) – 1615.000±0.800

282.760±0.500

73.622±0.500

Ge(cr) [92GRE/FUG]

0.000 0.000 31.090

±0.150 23.222

±0.100

Ge(g) [92GRE/FUG]

331.209 ±3.000

(a) 372.000±3.000

167.904±0.005

30.733±0.001

GeO2(tetragonal) [92GRE/FUG]

– 521.404 ±1.002

(a) – 580.000±1.000

39.710±0.150

50.166±0.300

GeF4(g) [92GRE/FUG]

– 1150.018 ±0.584

(a) – 1190.200±0.500

301.900±1.000

81.602±1.000

Sn(cr) [92GRE/FUG]

0.000 0.000 51.180

±0.080 27.112

±0.030

Sn(g) [92GRE/FUG]

266.223 ±1.500

(a) 301.200±1.500

168.492±0.004

21.259±0.001

Sn2+ [92GRE/FUG]

– 27.624 ±1.557

(a) – 8.900±1.000

– 16.700±4.000

SnO(tetragonal) [92GRE/FUG]

– 251.913 ±0.220

(a) – 280.710±0.200

57.170±0.300

47.783±0.300

SnO2(cassiterite, tetragonal) [92GRE/FUG]

– 515.826 ±0.204

(a) – 577.630±0.200

49.040±0.100

53.219±0.200

Pb(cr) [92GRE/FUG]

0.000 0.000 64.800

±0.300 26.650

±0.100

Pb(g) [92GRE/FUG]

162.232 ±0.805

(a) 195.200±0.800

175.375±0.005

20.786±0.001

Pb2+ [92GRE/FUG]

– 24.238 ±0.399

(a) 0.920±0.250

18.500±1.000

PbSO4(cr) [92GRE/FUG]

– 813.036 ±0.447

(a) – 919.970±0.400

148.500±0.600

B(cr) [92GRE/FUG]

0.000 0.000 5.900

±0.080 11.087

±0.100

B(g) [92GRE/FUG]

521.012 ±5.000

(a) 565.000±5.000

153.436±0.015

20.796±0.005

B2O3(cr) [92GRE/FUG]

– 1194.324 ±1.404

(a) – 1273.500±1.400

53.970±0.300

62.761±0.300

B(OH)3(aq) [92GRE/FUG]

– 969.268 ±0.820

(a) – 1072.800±0.800

162.400±0.600

B(OH)3(cr) [92GRE/FUG]

– 969.667 ±0.820

(a) – 1094.800±0.800

89.950±0.600

86.060±0.400



63


f mGο∆ f mH ο∆ mS ο ,mpCο Compound and NEA TDB Review (kJ · mol–1) (kJ · mol–1) (J · K–1 · mol–1) (J · K–1 · mol–1) BF3(g) [92GRE/FUG]

– 1119.403 ±0.803

(a) – 1136.000±0.800

254.420±0.200

50.463±0.100

Al(cr) [92GRE/FUG]

0.000 0.000 28.300

±0.100 24.200

±0.070

Al(g) [92GRE/FUG]

289.376 ±4.000

(a) 330.000±4.000

164.554±0.004

21.391±0.001

Al3+ [92GRE/FUG]

– 491.507 ±3.338

(a) – 538.400±1.500

– 325.000±10.000

Al2O3(corundum) [92GRE/FUG]

– 1582.257 ±1.302

(a) – 1675.700±1.300

50.920±0.100

79.033±0.200

AlF3(cr) [92GRE/FUG]

– 1431.096 ±1.309

(a) – 1510.400±1.300

66.500±0.500

75.122±0.400

Tl+ [99RAR/RAN]

– 32.400 ±0.300

Zn(cr) [92GRE/FUG]

0.000 0.000 41.630

±0.150 25.390

±0.040

Zn(g) [92GRE/FUG]

94.813 ±0.402

(a) 130.400±0.400

160.990±0.004

20.786±0.001

Zn2+ [92GRE/FUG]

– 147.203 ±0.254

(a) – 153.390±0.200

– 109.800±0.500

ZnO(cr) [92GRE/FUG]

– 320.479 ±0.299

(a) – 350.460±0.270

43.650±0.400

Cd(cr) [92GRE/FUG]

0.000 0.000 51.800

±0.150 26.020

±0.040

Cd(g) [92GRE/FUG]

77.230 ±0.205

(a) 111.800±0.200

167.749±0.004

20.786±0.001

Cd2+ [92GRE/FUG]

– 77.733 ±0.750

(a) – 75.920±0.600

– 72.800±1.500

CdO(cr) [92GRE/FUG]

– 228.661 ±0.602

(a) – 258.350±0.400

54.800±1.500

CdSO4. 2.667 H2O(cr) [92GRE/FUG]

– 1464.959 ±0.810

(a) – 1729.300±0.800

229.650±0.400

Hg(g) [92GRE/FUG]

31.842 ±0.054

(a) 61.380±0.040

174.971±0.005

20.786±0.001

Hg(l) [92GRE/FUG]

0.000 ±0.000

0.000±0.000

75.900±0.120

Hg2+ [92GRE/FUG]

164.667 ±0.313

(a) 170.210±0.200

– 36.190±0.800



64


f mGο∆ f mH ο∆ mS ο ,mpCο Compound and NEA TDB Review (kJ · mol–1) (kJ · mol–1) (J · K–1 · mol–1) (J · K–1 · mol–1) Hg2

2+ [92GRE/FUG]

153.567 ±0.559

(a) 166.870±0.500

65.740±0.800

HgO(montroydite, red) [92GRE/FUG]

– 58.523 ±0.154

(a) – 90.790±0.120

70.250±0.300

Hg2Cl2(cr) [92GRE/FUG]

– 210.725 ±0.471

(a) – 265.370±0.400

191.600±0.800

Hg2SO4(cr) [92GRE/FUG]

– 625.780 ±0.411

(a) – 743.090±0.400

200.700±0.200

Cu(cr) [92GRE/FUG]

0.000 0.000 33.150

±0.080 24.440

±0.050

Cu(g) [92GRE/FUG]

297.672 ±1.200

(a) 337.400±1.200

166.398±0.004

20.786±0.001

Cu2+ [92GRE/FUG]

65.040 ±1.557

(a) 64.900±1.000

– 98.000±4.000

CuCl(g) [2003GUI/FAN]

76.800±10.000

CuSO4(cr) [92GRE/FUG]

– 662.185 ±1.206

(a) – 771.400±1.200

109.200±0.400

Ag(cr) [92GRE/FUG]

0.000 0.000 42.550

±0.200 25.350

±0.100

Ag(g) [92GRE/FUG]

246.007 ±0.802

(a) 284.900±0.800

172.997±0.004

20.786±0.001

Ag+ [92GRE/FUG]

77.096 ±0.156

(a) 105.790±0.080

73.450±0.400

AgCl(cr) [92GRE/FUG]

– 109.765 ±0.098

(a) – 127.010±0.050

96.250±0.200

Ti(cr) [92GRE/FUG]

0.000 0.000 30.720

±0.100 25.060

±0.080

Ti(g) [92GRE/FUG]

428.403 ±3.000

(a) 473.000±3.000

180.298±0.010

24.430±0.030

TiO2(rutile) [92GRE/FUG]

– 888.767 ±0.806

(a) – 944.000±0.800

50.620±0.300

55.080±0.300

TiCl4(g) [92GRE/FUG]

– 726.324 ±3.229

(a) – 763.200±3.000

353.200±4.000

95.408±1.000

Th(cr) [92GRE/FUG]

0.000 ±0.000

0.000±0.000

51.800±0.500

26.230±0.050

Th(g) [92GRE/FUG]

560.745 ±6.002

(a) 602.000±6.000

190.170±0.050

20.789±0.100



65


f mGο∆ f mH ο∆ mS ο ,mpCο Compound and NEA TDB Review (kJ · mol–1) (kJ · mol–1) (J · K–1 · mol–1) (J · K–1 · mol–1) ThO2(cr) [92GRE/FUG]

– 1169.238 ±3.504

(a) – 1226.400±3.500

65.230±0.200

Be(cr) [92GRE/FUG]

0.000 0.000 9.500

±0.080 16.443

±0.060

Be(g) [92GRE/FUG]

286.202 ±5.000

(a) 324.000±5.000

136.275±0.003

20.786±0.001

BeO(bromellite) [92GRE/FUG]

– 580.090 ±2.500

(a) – 609.400±2.500

13.770±0.040

25.565±0.100

Mg(cr) [92GRE/FUG]

0.000 0.000 32.670

±0.100 24.869

±0.020

Mg(g) [92GRE/FUG]

112.521 ±0.801

(a) 147.100±0.800

148.648±0.003

20.786±0.001

Mg2+ [92GRE/FUG]

– 455.375 ±1.335

(a) – 467.000±0.600

– 137.000±4.000

MgO(cr) [92GRE/FUG]

– 569.312 ±0.305

(a) – 601.600±0.300

26.950±0.150

37.237±0.200

MgF2(cr) [92GRE/FUG]

– 1071.051 ±1.210

(a) – 1124.200±1.200

57.200±0.500

61.512±0.300

Ca(cr) [92GRE/FUG]

0.000 0.000 41.590

±0.400 25.929

±0.300

Ca(g) [92GRE/FUG]

144.021 ±0.809

(a) 177.800±0.800

154.887±0.004

20.786±0.001

Ca2+ [92GRE/FUG]

– 552.806 ±1.050

(a) – 543.000±1.000

– 56.200±1.000

CaO(cr) [92GRE/FUG]

– 603.296 ±0.916

(a) – 634.920±0.900

38.100±0.400

42.049±0.400

CaF(g) [2003GUI/FAN]

– 302.118 ±5.104

– 276.404±5.100

229.244±0.500

33.671±0.500

CaCl(g) [2003GUI/FAN]

– 129.787 ±5.001

– 103.400±5.000

241.634±0.300

35.687±0.010

Sr(cr) [92GRE/FUG]

0.000 ±0.000

0.000±0.000

55.700±0.210

Sr2+ [92GRE/FUG]

– 563.864 ±0.781

(a) – 550.900±0.500

– 31.500±2.000

SrO(cr) [92GRE/FUG]

– 559.939 ±0.914

(a) – 590.600±0.900

55.440±0.500

SrCl2(cr) [92GRE/FUG]

– 784.974 ±0.714

(a) – 833.850±0.700

114.850±0.420



66


f mGο∆ f mH ο∆ mS ο ,mpCο Compound and NEA TDB Review (kJ · mol–1) (kJ · mol–1) (J · K–1 · mol–1) (J · K–1 · mol–1) Sr(NO3)2(cr) [92GRE/FUG]

– 783.146 ±1.018

(a) – 982.360±0.800

194.600±2.100

Ba(cr) [92GRE/FUG]

0.000 ±0.000

0.000±0.000

62.420±0.840

Ba(g) [2003GUI/FAN]

152.852 ±5.006

185.000±5.000

170.245±0.010

20.786±0.001

Ba2+ [92GRE/FUG]

– 557.656 ±2.582

(a) – 534.800±2.500

8.400±2.000

BaO(cr) [92GRE/FUG]

– 520.394 ±2.515

(a) – 548.100±2.500

72.070±0.380

BaF(g) [2003GUI/FAN]

– 349.569 ±6.705

– 324.992±6.700

246.219±0.210

34.747±0.300

BaCl2(cr) [92GRE/FUG]

– 806.953 ±2.514

(a) – 855.200±2.500

123.680±0.250

Li(cr) [92GRE/FUG]

0.000 0.000 29.120

±0.200 24.860

±0.200

Li(g) [92GRE/FUG]

126.604 ±1.002

(a) 159.300±1.000

138.782±0.010

20.786±0.001

Li+ [92GRE/FUG]

– 292.918 ±0.109

(a) – 278.470±0.080

12.240±0.150

Na(cr) [92GRE/FUG]

0.000 0.000 51.300

±0.200 28.230

±0.200

Na(g) [92GRE/FUG]

76.964 ±0.703

(a) 107.500±0.700

153.718±0.003

20.786±0.001

Na+ [92GRE/FUG]

– 261.953 ±0.096

(a) – 240.340±0.060

58.450±0.150

NaF(cr) [2001LEM/FUG]

– 546.327 ±0.704

(a) – 576.600±0.700

51.160±0.150

NaCl(cr) [2001LEM/FUG]

– 384.221 ±0.147

– 411.260±0.120

72.150±0.200

50.500

NaNO3(cr) [2003GUI/FAN]

– 467.580±0.410

K(cr) [92GRE/FUG]

0.000 0.000 64.680

±0.200 29.600

±0.100

K(g) [92GRE/FUG]

60.479 ±0.802

(a) 89.000±0.800

160.341±0.003

20.786±0.001

K+ [92GRE/FUG]

– 282.510 ±0.116

(a) – 252.140±0.080

101.200±0.200



67


f mGο∆ f mH ο∆ mS ο ,mpCο Compound and NEA TDB Review (kJ · mol–1) (kJ · mol–1) (J · K–1 · mol–1) (J · K–1 · mol–1) KCl(cr) This Review

– 436.461±0.129

KBr(cr) This Review

– 393.330±0.188

KI(cr) This Review

– 329.150±0.137

Rb(cr) [92GRE/FUG]

0.000 0.000 76.780

±0.300 31.060

±0.100

Rb(g) [92GRE/FUG]

53.078 ±0.805

(a) 80.900±0.800

170.094±0.003

20.786±0.001

Rb+ [92GRE/FUG]

– 284.009 ±0.153

(a) – 251.120±0.100

121.750±0.250

Cs(cr) [92GRE/FUG]

0.000 0.000 85.230

±0.400 32.210

±0.200

Cs(g) [92GRE/FUG]

49.556 ±1.007

(a) 76.500±1.000

175.601±0.003

20.786±0.001

Cs+ [92GRE/FUG]

– 291.456 ±0.535

(a) – 258.000±0.500

132.100±0.500

CsCl(cr) [2001LEM/FUG]

– 413.807 ±0.208

(a) – 442.310±0.160

101.170±0.200

52.470

CsBr(cr) [2001LEM/FUG]

– 391.171 ±0.305

– 405.600±0.250

112.940±0.400

52.930

(a) Value calculated internally using f m f m m,i

iG H T Sο ο ο∆ = ∆ − ∑ .

(b) Value calculated internally from reaction data (see Table IV–2).


68

Table IV–2: Selected thermodynamic data for reactions involving auxiliary compounds and complexes used in the evaluation of thermodynamic data for the NEA TDB Project data. All ionic species listed in this table are aqueous species. Unless noted otherwise, all data refer to 298.15 K and a pressure of 0.1 MPa and, for aqueous species, a standard state of infinite dilution (I = 0). The uncertainties listed below each value represent total uncertainties and correspond in principle to the statistically defined 95% confidence interval. Systematically, all the values are presented with three digits after the decimal point, regardless of the significance of these digits. The reference listed for each entry in this table indicates the NEA TDB Review where the corresponding data have been adopted as NEA TDB Auxiliary data.The data presented in this table are available on computer media from the OECD Nuclear Energy Agency.

Reaction Species and NEA TDB Review 10log K ο

r mGο∆ r mH ο∆ r mS ο∆ (kJ · mol–1) (kJ · mol–1) (J · K–1 · mol–1)

HF(aq) F – + H+ HF(aq) [92GRE/FUG] 3.180

±0.020 – 18.152

±0.114 12.200

±0.300 101.800

±1.077 (a)

HF2 – F – + HF(aq) HF2

– [92GRE/FUG] 0.440

±0.120 – 2.511

±0.685 3.000

±2.000 18.486

±7.090 (a)

ClO – HClO(aq) ClO – + H+ [92GRE/FUG] – 7.420

±0.130 42.354

±0.742 19.000

±9.000 – 78.329

±30.289 (a)

ClO2 – HClO2(aq) ClO2

– + H+ [92GRE/FUG] – 1.960

±0.020 11.188

±0.114

HClO(aq) Cl2(g) + H2O(l) Cl – + H+ + HClO(aq) [92GRE/FUG] – 4.537

±0.105 25.900

±0.600

HClO2(aq) H2O(l) + HClO(aq) 2H+ + HClO2(aq) + 2 e – [92GRE/FUG] – 55.400

±0.700 (b) 316.230

±3.996

BrO – HBrO(aq) BrO – + H+ [92GRE/FUG] – 8.630

±0.030 49.260

±0.171 30.000

±3.000 – 64.600

±10.078 (a)

HBrO(aq) Br2(aq) + H2O(l) Br – + H+ + HBrO(aq) [92GRE/FUG] – 8.240

±0.200 47.034

±1.142



69

Table IV–2 (continued)



HIO3(aq) H+ + IO3 – HIO3(aq)

[92GRE/FUG] 0.788±0.029

– 4.498±0.166

S2– HS – H+ + S2– [92GRE/FUG] – 19.000

±2.000 108.450

±11.416

SO32– H2O(l) + SO4

2– + 2 e – 2OH – + SO32–

[92GRE/FUG] – 31.400±0.700

(b) 179.230±3.996

S2O32– 3H2O(l) + 2SO3

2– + 4 e – 6OH – + S2O32–

[92GRE/FUG] – 39.200±1.400

(b) 223.760±7.991

H2S(aq) H2S(aq) H+ + HS – [92GRE/FUG] – 6.990

±0.170 39.899

±0.970 (a)

HSO3 – H+ + SO3

2– HSO3 –

[92GRE/FUG] 7.220±0.080

– 41.212±0.457

66.000±30.000

359.590±100.630

(a)

HS2O3 – H+ + S2O3

2– HS2O3 –

[92GRE/FUG] 1.590±0.150

– 9.076±0.856

H2SO3(aq) H+ + HSO3 – H2SO3(aq)

[92GRE/FUG] 1.840±0.080

– 10.503±0.457

16.000±5.000

88.891±16.840

(a)

HSO4 – H+ + SO4

2– HSO4 –

[92GRE/FUG] 1.980±0.050

– 11.302±0.285

HN3(aq) H+ + N3 – HN3(aq)

[92GRE/FUG] 4.700±0.080

– 26.828±0.457

– 15.000±10.000

39.671±33.575

(a)

NH3(aq) NH4+ H+ + NH3(aq)

[92GRE/FUG] – 9.237±0.022

52.725±0.126

52.090±0.210

– 2.130±0.821

(a)



70




HNO2(aq) H+ + NO2 – HNO2(aq)

[92GRE/FUG] 3.210±0.160

– 18.323±0.913

– 11.400±3.000

23.219±10.518

(a)

PO43– HPO4

2– H+ + PO43–

[92GRE/FUG] – 12.350±0.030

70.494±0.171

14.600±3.800

– 187.470±12.758

(a)

P2O74– HP2O7

3– H+ + P2O74–

[92GRE/FUG] – 9.400±0.150

53.656±0.856

H2PO4 – H+ + HPO4

2– H2PO4 –

[92GRE/FUG] 7.212±0.013

– 41.166±0.074

– 3.600±1.000

126.000±3.363

(a)

H3PO4(aq) H+ + H2PO4 – H3PO4(aq)

[92GRE/FUG] 2.140±0.030

– 12.215±0.171

8.480±0.600

69.412±2.093

(a)

HP2O73– H2P2O7

2– H+ + HP2O73–

[92GRE/FUG] – 6.650±0.100

37.958±0.571

H2P2O72– H3P2O7

– H+ + H2P2O72–

[92GRE/FUG] – 2.250±0.150

12.843±0.856

H3P2O7 – H4P2O7(aq) H+ + H3P2O7

– [92GRE/FUG] – 1.000

±0.500 5.708

±2.854

H4P2O7(aq) 2H3PO4(aq) H2O(l) + H4P2O7(aq) [92GRE/FUG] – 2.790

±0.170 15.925

±0.970 22.200

±1.000 21.045

±4.673 (a)

CO2(aq) H+ + HCO3 – CO2(aq) + H2O(l)

[92GRE/FUG] 6.354±0.020

– 36.269±0.114

(a)

CO2(g) CO2(aq) CO2(g) [92GRE/FUG] 1.472

±0.020 – 8.402

±0.114 (a)



71




HCO3 – CO3

2– + H+ HCO3 –

[92GRE/FUG] 10.329±0.020

– 58.958±0.114

(a)

CN – HCN(aq) CN – + H+ [2005OLI/NOL] – 9.210

±0.020 52.571

±0.114 43.600

±0.200 – 30.089

±0.772 (a)

HCN(aq) HCN(g) HCN(aq) [2005OLI/NOL] 0.902

±0.050 – 5.149

±0.285 – 26.150

±2.500 – 70.439

±8.440 (a)

SiO2(OH)22– Si(OH)4(aq) 2H+ + SiO2(OH)2

2– [92GRE/FUG] – 23.140

±0.090 132.080

±0.514 75.000

±15.000 – 191.460

±50.340 (a)

SiO(OH)3 – Si(OH)4(aq) H+ + SiO(OH)3

– [92GRE/FUG] – 9.810

±0.020 55.996

±0.114 25.600

±2.000 – 101.950

±6.719 (a)

Si(OH)4(aq) 2H2O(l) + SiO2(quar) Si(OH)4(aq) [92GRE/FUG]$ – 4.000

±0.100 22.832

±0.571 25.400

±3.000 8.613

±10.243 (a)

Si2O3(OH)42– 2Si(OH)4(aq) 2H+ + H2O(l) + Si2O3(OH)4

2– [92GRE/FUG] – 19.000

±0.300 108.450

±1.712

Si2O2(OH)5 – 2Si(OH)4(aq) H+ + H2O(l) + Si2O2(OH)5

– [92GRE/FUG] – 8.100

±0.300 46.235

±1.712

Si3O6(OH)33– 3Si(OH)4(aq) 3H+ + 3H2O(l) + Si3O6(OH)3

3– [92GRE/FUG] – 28.600

±0.300 163.250

±1.712


3– [92GRE/FUG] – 27.500

±0.300 156.970

±1.712



72





4– [92GRE/FUG] – 36.300

±0.500 207.200

±2.854


3– [92GRE/FUG] – 25.500

±0.300 145.560

±1.712

(a) Value calculated internally using r m r m r mG H T Sο ο ο∆ = ∆ − ∆ .

(b) Value calculated from a selected standard potential.

(c) Value of 10log K ο calculated internally from r mGο∆ .

73

Part III

Discussion of data selection

75

Chapter VEquation Section 22

Discussion of data selection for nickel

V.1 Elemental nickel V.1.1 Nickel gas In this review, it is assumed that Ni forms an ideal gas phase. Our recommendations are based on the critically evaluated data given in [73HUL/DES], [76MAH/PAN] and [98CHA]. Gaseous ions e.g., Ni+(g), Ni2+(g) and Ni–(g) also evaluated in [98CHA], [76MAH/PAN] and [73HUL/DES] are not comprehensively treated by this review.

An interest for the interpretation of spectra of highly-ionised atoms has arisen from the investigation of hot plasmas generated to achieve nuclear fusion. These activi-ties have produced a substantial increase in spectroscopic information and the prepara-tion of new compilations of energy levels. The energy levels of nickel atoms and its ions, as derived from analyses of atomic spectra, have been published by Corliss and Sugar [81COR/SUG]. Based on this publication, the thermodynamic data for Ni+(g) and Ni–(g) have been calculated by [98CHA]. The source of data for the calculation of the enthalpy of formation, the heat capacity and the entropy for Ni2+(g) ions were the en-ergy levels deduced from NiII and NiIII spectra. Such calculations were carried out by Mah and Pankratz [76MAH/PAN].

The enthalpy of formation for Ni(g) was calculated by [98CHA] by analysis of the vapour pressure data over liquid nickel obtained by [57MOR/ZEL]. This calcula-tions gave vap mH ο∆ (Ni, 298.15 K) = (412.29 ± 0.50) kJ·mol–1 (3rd law) and (415.57 ± 5.15) kJ·mol–1 (2nd law).

Using the 3rd law vap mH ο∆ (Ni, 298.15 K) and f mH ο∆ (Ni, l, 298.15 K) the en-thalpy of formation f mH ο∆ (Ni, g, 298.15 K) was calculated [98CHA]. The value

f mH ο∆ (Ni, g, 298.15 K) = (430.1 ± 8.4) kJ·mol–1

obtained by [98CHA] is the same as the one chosen by [76MAH/PAN] and [73HUL/DES]. This value is also selected in the present review.

V.1 Elemental nickel 76

Rutner and Haury [74RUT/HAU] determined the “best value” of the heat of vaporisation and sublimation of nickel by a statistical treatment of data available in the literature and their own results obtained by the Langmuir technique. However, the dif-ference of the entropy of sublimation determined by the 2nd and 3rd law analysis − (10.5 ± 16.3) J·K–1·mol–1, being a criterion of the precision of measurements, was rather high. Therefore, Chase [98CHA] has not used results of [74RUT/HAU] for the calculation of the enthalpy of formation for Ni(g). Other vapour pressure studies were discussed by Hultgren et al. [73HUL/DES].

The source of data for the calculation of the heat capacity and entropy of Ni(g), tabulated in [98CHA], [76MAH/PAN] and [73HUL/DES], were the calculations of electronic energy levels and quantum weights done by [70ROT] and [68MOO/MER]. All tabulated entropy values agree within 0.004 J·K–1·mol–1, therefore, the present re-viewers select the values:

mS ο (Ni, g, 298.15 K) = (182.19 ± 0.08) J·K–1·mol–1

,mpCο (Ni, g, 298.15 K) = (23.36 ± 0.10) J·K–1·mol–1

which are the same as those recommended by [98CHA]. With the selections of f mH ο∆ and mS ο the Gibbs energy can be calculated to be:

f mGο∆ (Ni, g, 298.15 K) = (384.7 ± 8.4) kJ·mol–1.

V.1.2 Nickel crystal Ni metal has a fcc (A1) structure existing up to the melting point of the solid at (1726 ± 4) K [73HUL/DES]. Ferromagnetic Ni is frequently designated Ni(α) and para-magnetic Ni, Ni(β) [60KEL].

The melting point selected by Hultgren et al. [73HUL/DES] and also retained by this review, Tfus = (1726 ± 4) K, is based on the following determinations:

[66VOL/KOH]: 1725 K [63GEO/FER]: 1726 K [54ORI/JON]: (1725 ± 4) K.

The value of the Curie point, Tc = 631 K, is taken from the study of Connelly et al. [71CON/LOO] using a calorimetric method which permits continuous observation of

,mpC versus T, and from the study of Vollmer et al. [66VOL/KOH] using a high tem-perature adiabatic calorimeter.

Pure Ni metal is defined as the nickel reference phase. As such, its Gibbs en-ergy of formation and enthalpy of formation are zero by definition at 298.15 K and 0.1 MPa.

The source of data for the entropy determination of nickel has been the tem-perature dependence of the heat capacities. The measurements of the molar heat of


nickel in the temperature range 1.1 to 19.0 K were performed by Keesom and Clark [35KEE/CLA]. The standard molar entropy of Ni determined by Busey and Giauque [52BUS/GIA2], using measurements of the heat capacity from 15 to 300 K, was equal to 29.86 J·K–1·mol–1.

In the present review, the selected value of the standard molar entropy of crys-talline nickel, Ni(cr), is based on the above determination:

mS ο (Ni, cr, 298.15 K) = (29.87 ± 0.20) J·K–1·mol–1.

This is essentially the same as the values given in the NIST-JANAF Thermo-chemical Tables [98CHA] (29.87 ± 0.21) J·K–1·mol–1, [76MAH/PAN] (29.87 ± 0.08) J·K–1·mol–1, [73HUL/DES] (29.87 ± 0.08) J·K–1·mol–1 and in [77BAR/KNA] 29.87 J·K−1·mol–1.

In the literature, numerous heat capacity studies of Ni metal, especially in the vicinity of the Curie point, are reported. The studies upon which our adopted values are based are as follows: [36BRO/WIL], [38SYK/WIL], [55KRA/WAR], [65PAW/STA] , [66VOL/KOH] and [71CON/LOO]. The experimental results of these studies and a polynomial fitting curve derived for this review are shown in Figure V-1.

Figure V-1: The standard molar heat capacity of nickel as a function of temperature. The data are taken from (+)[36BRO/WIL], ( )[38SYK/WIL], ( )[55KRA/WAR], ( )[65PAW/STA], ( )[66VOL/KOH] and ( )[71CON/LOO]. The dotted line corre-sponds to the thermal heat capacity function given by [77BAR/KNA]. The fitted curve (solid line) refers to the thermal heat capacity function recommended by this review.

200 400 600 800 1000 1200 1400 1600 1800

20

22

24

26

28

30

32

34

36

38

40

T / K

C° p,m

/ J·K

–1·m

ol–1


The heat capacity calculated using the temperature coefficients given by Barin and Knacke [77BAR/KNA] and the experimental results disagree conspicuously in the temperature range 640 < T/K < 1728 (Figure V-1). Based on the fitted curve, this review selects the heat capacity of Ni at 298.15 K:

,mpCο (Ni, cr, 298.15 K) = (26.07 ± 0.10) J·K–1·mol–1

and the temperature coefficients of the heat capacity function: ,mpCο (T) = a + bT + cT2 + eT–2, which are listed in Table V-1. The recommended value for the heat capacity at 298.15 K is consistent with the values selected by the US Bureau of Mines [76MAH/PAN], 25.98 J·K–1·mol–1, the US Geological Survey [95ROB/HEM], 25.99 J·K–1·mol–1, the value selected by Hultgren et al. [73HUL/DES], 26.07 J·K–1·mol–1, the NIST-JANAF Thermochemical Tables [98CHA], 25.987 J·K–1·mol–1, and also with the [77BAR/KNA] compilation, 26.09 J·K–1·mol–1. The uncertainty assigned to the selected value of ,mpCο (Ni, cr, 298.15 K) was calculated using the experimental data of [36BRO/WIL] and [52BUS/GIA] obtained from measurements in the vicinity of 298 K.

The temperature coefficients of the heat capacity function for Ni derived for the present review also allow calculation of heat capacities which coincide very well with the experimental data in the temperature range 640 < T/K < 1728, as shown in Figure V-1.

Table V-1: The temperature coefficients of the heat capacity function for Ni(cr):

,mpCο (T) = a + bT + cT 2 + eT –2. The notation E±nn indicates the power of 10, units for ,mpCο are J·K–1·mol–1. The sources of experimental data used for the determination of

these coefficients are given in Figure V-1.

a b c e Tmin/K Tmax/K

3.36472 E+01 – 2.47530E–02 4.35458E–05 – 3.61955E+05 298 450 – 4.93235E+01 1.676480E–01 – 7.48518E–05 3.76214E+06 450 600

1.05023E+05 – 2.29221E+02 1.40772E–01 – 6.52809E+09 600 631 5.29406E+04 – 1.65828E+02 1.29936E–01 0 631 640 1.82638E+02 – 4.20407E–01 2.90521E– 04 0 640 690 1.23677E+01 2.22800E–02 4.45300E– 06 2.46766E+06 690 1728

V.1.3 Nickel liquid The enthalpy of formation of Ni(l) can be calculated from that of Ni(cr) by adding the enthalpy of fusion and the difference in enthalpy, ( mH (1728 K) – mH ο (298.15 K)) be-tween the solid and liquid phase. The value

f mH ο∆ (Ni, l, 298.15 K) = (17.50 ± 0.25) kJ·mol–1

given by [98CHA] is essentially the same as that recommended by the present review.


The enthalpy of fusion fus mH ο∆ = (17.15 ± 0.42) kJ·mol–1, used for the calcula-tion of the enthalpy of formation, is based on the calorimetric study of Vollmer et al. [66VOL/KOH] fus mH ο∆ = (16.90 ± 0.25) kJ·mol–1, and Geoffray et al. [63GEO/FER]

fus mH ο∆ = (17.47 ± 0.23) kJ·mol–1. An another reported value, (18.43 ± 0.27) kJ·mol–1 obtained by the electrical explosion study i.e., by rapid heating using a current of high density [71LEB/SAV] was significantly different and, therefore, disregarded. Another study performed using a similar experimental technique i.e., fast pulse heating gave

fus mH ο∆ = (18.02 ± 0.54) kJ·mol–1 [93OBE/KAS]. Although this value is higher than the calorimetric one, the respective error limits overlap.

The enthalpy of fusion retained by this review is identical with that given in the NIST-JANAF Tables [98CHA]. The compilations [73HUL/DES] and [77BAR/KNA] adopted the value determined by [63GEO/FER].

The heat capacity of nickel in the liquid region (1728 – 1822 K) was measured by Vollmer et al. [66VOL/KOH] using an adiabatic high temperature calorimeter with an inert gas atmosphere, ,mpCο (Ni, l) = 38.99 J·K–1·mol–1 was obtained. Another study yielding the liquid phase heat capacity is that of Geoffray et al. [63GEO/FER],

,mpCο (Ni, l) = 43.10 J·K–1·mol–1. In this review, the constant value independent of the temperature, ,mpCο (Ni, l) = (41.1 ± 2.1) J·K–1·mol–1 is based on these two studies. The uncertainty was calculated assuming that the uncertainty in both studies was 2%, i.e., the same as reported by [66VOL/KOH]. The thermochemical tables [98CHA] and [76MAH/PAN] adopted the heat capacity of nickel in the liquid phase, 38.911 J·K−1·mol–1, and 38.91 J·K–1·mol–1, respectively which is close to the one measured by [66VOL/KOH].

The entropy at 298.15 K can be calculated in a manner analogous to the one used for the enthalpy of formation of nickel in the liquid region. The value selected by this review:

mS ο (Ni, l, 298.15 K) = (41.49 ± 0.30) J·K–1·mol–1

is almost identical with the one chosen by NIST-JANAF Thermochemical Tables [98CHA].

The selections above yield:

f mGο∆ (Ni, l, 298.15 K) = (14.04 ± 0.27) kJ·mol–1.

V.2 Simple nickel aqua ions V.2.1 Ni2+ In every review of thermodynamic data care should be taken that the same and simulta-neously the most reliable auxiliary quantities are employed. Consequently CODATA values (see NEA-TDB auxiliary data) were taken whenever they had been available. The standard enthalpy of formation and the partial molar entropy of the bivalent nickel

V.2 Simple nickel aqua ions 80

ion have been re-evaluated in the present review. The recommendations of recent re-views by Plyasunova et al. [98PLY/ZHA] as well as Archer [99ARC] (see Sections V.2.1.2 and V.2.1.3) have been taken into account.

V.2.1.1 Gibbs energy of formation of Ni2+

In principle the standard Gibbs energy of formation of Ni2+ can be obtained directly from potentiometric data. Plyasunova et al. [98PLY/ZHA] maintain, however, that no accurate value for the standard electrode potential of Ni2+ | Ni according to:

Ni2+ + H2(g) Ni(cr) + 2H+ (V.1)

can be produced by emf measurements, because the truly reversible equilibrium appears to be unattainable. Thus, in Plyasunova et al.’s review the results given by Murata [28MUR], Haring and Vanden Bosche [29HAR/BOS], Colombier [34COL2], as well as Carr and Bonilla [52CAR/BON] were simply averaged and taken as the result of poten-tiometric data on the Ni2+ | Ni couple. The present reviewers, however, re-evaluated these and other data [62LOP], [62VAG/UVA], see Appendix A.

From the point of view of thermodynamics the standard electrode potential of Ni2+ | Ni can be obtained most accurately from a cell without a liquid junction, such as that used by Haring and Vanden Bosche [29HAR/BOS] and discussed thoroughly by Archer [99ARC]. The former authors measured potentials of cells of the type:

Ni | NiSO4(m) | Hg2SO4 | Hg

at 25°C for 0.05 < m < 0.16 mol·kg–1. The recalculated mean value of E° (SIT model [97GRE/PLY2]) of the 26 cells measured was (0.8500 ± 0.0025) V, see Figure V-2. Combined with the standard potential of the Hg2SO4 | Hg electrode ((0.6128 ± 0.0031) V, [89COX/WAG]) a value of E°(Ni2+ | Ni, 298.15 K) = − (0.2372 ± 0.0040) V has been obtained which agrees almost perfectly with the NBS tables (– 0.236 V) [82WAG/EVA], [89BRA].

Attempts to measure the standard electrode potential of nickel using the similar cell:

Ni | NiCl2, solvent | Hg2Cl2 | Hg

failed when the solvent was water [29HAR/BOS], see also [61TAN2].

For the selected standard Gibbs energy of formation of Ni2+ an uncertainty of ± 0.77 kJ·mol–1 has to be taken into account,

f mGο∆ (Ni2+, 298.15 K) = – (45.77 ± 0.77) kJ·mol–1.


Figure V-2: Standard electrode potential of Ni | NiSO4 | Hg2SO4 | Hg at 25°C. (Thick solid curve: calculated according to [29HAR/BOS]; thin solid curves: corresponding error limits; ○ experimental data [29HAR/BOS]; thick dash curve: calculated using pre-viously selected thermodynamic quantities [82WAG/EVA], [89COX/WAG]).

0.04 0.06 0.08 0.10 0.12 0.14 0.16 0.18

0.952

0.956

0.960

0.964

0.968

0.972

: E/V = 0.858 − (RT/F) ln aNiSO4

E / V

m (NiSO4) / mol·kg−1

▬▬ : E/V = 0.8500 – (RT/F) ln aNiSO4


V.2.1.1.1 Temperature coefficient of the standard electrode potential Ni2+ | Ni

The formation reaction of Ni2+ is the reverse of the reaction of Equation (V.1):

Ni(cr) + 2H+ Ni2+ + H2(g), (V.2)

and the corresponding standard entropy may be calculated by:

f mS ο∆ (Ni2+, 298.15 K) = – n F (d E°(Ni2+ | Ni)/d T)298.

Bratsch [89BRA] recommended the value (d E°/d T)298 = (0.146 ± 0.010) mV·K–1 which leads to f mS ο∆ (Ni2+, 298.15 K) = – (28 ± 2) J·K–1·mol–1. According to Archer [99ARC] this value can be traced to the Ph.D. thesis of Muldrow [58MUL], who measured the enthalpy of solution of NiCl2(cr) in water and obtained f mH ο∆ (Ni2+, 298.15 K) = − (54.040 ± 0.800) kJ·mol–1. Combined with f mGο∆ (Ni2+, 298.15 K) from the NBS Tables [82WAG/EVA], (d E°(Ni2+ | Ni)/d T)298 = 0.1465 mV·K–1 follows [58MUL]. When Ko and Hepler [63KO/HEP] derived f mH ο∆ (Ni2+, 298.15 K) = – 53.14 kJ·mol–1 by combination of the enthalpies of formation and solution in water of NiCl2(cr) with the enthalpy of formation of Cl– and f mGο∆ (Ni2+, 298.15 K) from Carr and Bonilla [52CAR/BON], almost the same value of (d E°(Ni2+ | Ni)/d T)298 was ob-tained. From Reaction (V.2) with NEA-TDB auxiliary values and the data for Ni(cr), optimised in this review, the partial molar entropy of the nickel ion can then be calcu-lated to be: mS ο (Ni2+, 298.15 K) = – (128.9 ± 3.0) J·K–1·mol–1. This result necessarily agrees with the NBS value [82WAG/EVA]. It must be emphasised that (d E°(Ni2+ | Ni)/d T)298 [89BRA] was obtained from previously determined entropies and not independently by investigating the temperature dependence of E°(Ni2+ | Ni), as might be concluded from the review of Plyasunova et al. [98PLY/ZHA]. Consequently the selection of mS ο (Ni2+, 298.15 K) has been deferred to Section V.2.1.3.

V.2.1.2 Enthalpy of formation of Ni2+

A straightforward calculation of f mH ο∆ (Ni2+, 298.15 K) can be based on Reactions (V.3), (V.4), (V.5), see Equations (V.6), (V.7), (V.8):

NiCl2(cr) Ni2+ + 2Cl– (V.3)

NiBr2(cr) Ni2+ + 2Br– (V.4)

NiI2(cr) Ni2+ + 2I–. (V.5)

Thomsen [1883THO] measured the enthalpy of dissolution of NiCl2(cr) in wa-ter as early as the 19th century, however, the data reported in [58MUL], [90EFI/FUR] and [95MAN/KOR] have been evaluated, see Table V-2. Other data were considered to be of lower quality [70EFI/KUD], [75BAR/MAS], [75WEE/KOE]. As NiCl2(cr) is very hygroscopic, a systematic error may be introduced when moisture is not excluded effec-tively. It was not possible to decide which of the three results chosen is the most reli-able. Consequently the unweighted average was calculated, and the uncertainty (2σ)


assigned covers the whole range of expectancy, thus sol mH ο∆ (V.3) = − (83.9 ± 1.7) kJ·mol–1 at 298.15 K. For the calculation of f mH ο∆ (Ni2+) according to Equation (V.6),

f mH ο∆ (NiCl2, cr, 298.15 K) = – (304.90 ± 0.11) kJ·mol–1 has been selected in this re-view and f mH ο∆ (Cl–) = – (167.08 ± 0.10) kJ·mol–1 has been taken from CODATA [89COX/WAG] as adopted by the NEA-TDB reviews:

f mH ο∆ (Ni2+) = sol mH ο∆ (V.3) – 2 f mH ο∆ (Cl–) + f mH ο∆ (NiCl2, cr). (V.6)

This results in f mH ο∆ (Ni2+) = – (54.64 ± 1.71) kJ·mol–1.

Table V-2: Enthalpy of Reaction (V.3)

1sol m / kJ molH ο −− ∆ ⋅ References

(83.41 ± 1.40) [58MUL]

(84.89 ± 0.27) [90EFI/FUR]

(83.48 ± 0.57) [95MAN/KOR]

Efimov and Furkalyuk [90EFI/FUR] determined f mH ο∆ (Ni2+) =

− (55.18 ± 0.58), − (55.27 ± 0.50), – (56.22 ± 0.42) kJ·mol–1 with Equations (V.6), (V.7) and (V.8), using the enthalpies of nickel halide formation determined by [88EFI/EVD], [89EVD/EFI] and [90EFI/EVD]. In the present review the selected enthalpy of forma-tion values for f mH ο∆ (NiBr2, cr) and f mH ο∆ (NiI2, cr), – (213.5 ± 2.4) kJ·mol-1 and − (96.42 ± 0.84) kJ·mol–1 are used with the enthalpies of formation for the halide ions

f mH ο∆ (Br–) = – (121.41 ± 0.15) kJ·mol–1 and f mH ο∆ (I–) = − (56.78 ± 0.05) kJ·mol–1 taken from CODATA [89COX/WAG] as adopted by the NEA-TDB reviews:

f mH ο∆ (Ni2+) = sol mH ο∆ (V.4) – 2 f mH ο∆ (Br–) + f mH ο∆ (NiBr2, cr) (V.7)

f mH ο∆ (Ni2+) = sol mH ο∆ (V.5) – 2 f mH ο∆ (I–) + f mH ο∆ (NiI2, cr). (V.8)

This results in f mH ο∆ (Ni2+) = – (57.4 ± 2.4) kJ·mol–1 from Equation (V.7) and f mH ο∆ (Ni2+) = – (56.2 ± 0.9) kJ·mol–1 from Equation (V.8). Together with f mH ο∆ (Ni2+)

= – 52.10, – 51.85 kJ·mol–1 determined by Vasil’ev et al. [84VAS/VAS], [86VAS/DMI], there are five independent values which, giving equal weight to each one, result in a mean value of f mH ο∆ (Ni2+, 298.15 K) = – (54.4 ± 4.9) kJ·mol–1.

This value agrees with the one selected by Plyasunova et al. (– (54.1 ± 2.5) kJ·mol–1) from essentially the same sources.

Archer [99ARC] cited values taken, partly from the same, partly from different sources, resulting in – (54.2 ± 2.1) kJ·mol–1 which again agrees within the error limits with the earlier assessments [84VAS/VAS], [86VAS/DMI], [90EFI/FUR], [98PLY/ZHA]. As the uncertainty of the calorimetric determinations of f mH ο∆ (Ni2+, 298.15 K) is rather high (± 4.9 kJ·mol–1), the enthalpy of Reaction (V.2) has been calcu-lated from the Gibbs energy using the definition in Equation (V.9). For the calculation


of f mS ο∆ (Ni2+, 298.15 K) = r mS ο∆ (V.2) the selected values of mS ο (Ni, cr, 298.15 K), mS ο (Ni2+, 298.15 K), see Sections V.1.2 and V.2.1.3, as well as mS ο (H2, g, 298.15 K)

from the selected auxiliary data of Table IV-1 were used: H = G + T S (V.9)

f mH ο∆ (Ni2+, 298.15 K) = f mGο∆ (Ni2+, 298.15 K) + 298.15· r mS ο∆ (V.2)

f mH ο∆ (Ni2+, 298.15 K) = (– (45.77 ± 0.77) + 0.29815·(– (30.99 ± 1.41))) kJ·mol–1

f mH ο∆ (Ni2+, 298.15 K) = – (55.01 ± 0.88) kJ·mol–1.

This value is similar to the result obtained by the sol mH ο∆ measurements, but the uncertainty is much less.

V.2.1.3 Partial molar entropy of Ni2+

V.2.1.3.1 NiSO4·7H2O(cr) cycle

Dissolution: NiSO4·7H2O(cr) Ni2+ + 24SO − + 7H2O(l) (V.10)

The solubility of NiSO4·7H2O(cr) at 25°C has been taken from the authors rec-ommended by Linke and Seidell on p. 1219 [65LIN/SEI]. The mean value and error limits have been estimated by averaging the six different solubilities listed at 25°C [23VIL/BEN], [24TAN], [39ROH]:

40.6 g (NiSO4)/100g (H2O)

which yields: m (sat) = (2.62 ± 0.05) mol·kg–1 (NiSO4·7H2O(cr)).

The value of ln ±γ (sat) = – (3.326 ± 0.007) has been derived from the data of Robinson and Stokes [59ROB/STO]. The water activity can be calculated according to:

ln aw = – [M (H2O) · m(NiSO4) · 2 ·φ ] / 1000 (V.11)

where φ is the osmotic coefficient also listed by Robinson and Stokes [59ROB/STO]. The value of ln aw (sat) = – (0.0707 ± 0.0030) has been derived from those data. The error limits were estimated by inserting the higher and lower values of both m and φ into Equation (V.11).

The standard Gibbs energy of Reaction (V.10) can be calculated according to Equation (V.12), as was shown for example by Larson et al. [68LAR/CER]:

sol mGο∆ (V.10) = – RT{2·[ln m (sat) + lnγ ± (sat)] + 7·ln aw (sat)} (V.12)

which leads to the selection:

sol mGο∆ (V.10) = (12.940 ± 0.110) kJ·mol–1.

Stout et al. [66STO/ARC] determined the standard enthalpy of Reaction (V.13):

NiSO4·7H2O(cr) NiSO4·6H2O(α) + H2O(l) (V.13)


to be r mH ο∆ (V.13) = (7.682 ± 0.105) kJ·mol–1 ((1836 ± 25) cal·mol–1).

This value can be combined with the recalculated standard enthalpy of Reac-tion (V.14) determined by Goldberg et al. [66GOL/RID]:

NiSO4·6H2O(α) Ni2+ + 24SO − + 6H2O(l) (V.14)

r mH ο∆ (V.14) = (4.485 ± 0.200) kJ·mol–1 to result in:

sol mH ο∆ (V.10) = (12.167 ± 0.226) kJ·mol–1.

The entropy of Reaction (V.10) has been calculated by rearranging Equation (V.9):

sol mS ο∆ (V.10) = { sol mH ο∆ (V.10) – sol mGο∆ (V.10)}/298.15

sol mS ο∆ (V.10) = – (2.59 ± 0.84) J·K–1·mol–1.

The selected value for partial molar entropy of Ni2+ follows from Equation (V.15):

mS ο (Ni2+) = mS ο (NiSO4·7H2O, cr) – mS ο ( 24SO − ) – 7 mS ο (H2O, l) + sol mS ο∆ (V.10) (V.15)

mS ο (Ni2+, 298.15 K) = – (131.8 ± 1.4) J·K–1·mol–1

where mS ο (NiSO4·7H2O, cr) = (378.95 ± 1.00) J·K–1·mol–1 as assessed in the present review (see Section V.5.1.2.2.1) and the values for mS ο ( 2

4SO − ) and mS ο (H2O, l) are auxiliary values from Table IV-1.

V.2.1.3.2 NiCl2·6H2O(cr) cycle

Another possible cycle for determining the partial molar entropy of Ni2+ is based on the dissolution reaction:

NiCl2·6H2O(cr) Ni2+ + 2Cl– + 6H2O(l) (V.16)

Rard [92RAR2] calculated a value of sol mGο∆ (V.16) = – (17.38 ± 0.04) kJ·mol–1 from his precise isopiestic vapour-pressure and solubility experiments. Thomsen [06THO] measured r mH∆ (V.17) = – (85.06 ± 1.80) kJ·mol–1 (– 20.33 kcal·mol–1) for the reaction:

NiCl2(cr) + 6H2O(l) NiCl2·6H2O(cr) (V.17)

at 291 – 293 K. The uncertainty assigned in the present work to r mH∆ (V.17) accounts for the deviation of the actual measurements and correction to the standard temperature (298.15 K) [06THO], [98PLY/ZHA]. As discussed in Section V.2.1.2, based on experi-mental heat of solution measurements, sol mH ο∆ (V.3) = − (83.9 ± 1.7) kJ·mol–1 at 298.15

K. From Reaction (V.3) and (V.17) sol mH ο∆ (V.16) = (1.2 ± 2.5) kJ·mol–1 follows. The entropy of Reaction (V.16) can be calculated in a manner quite analogous to that used in Section V.2.1.3.1:


sol mS ο∆ (V.16) = { sol mH ο∆ (V.16) – sol mGο∆ (V.16)}/298.15

sol mS ο∆ (V.16) = (62.3 ± 8.3) J·K–1·mol–1.

If a reliable independent value were available for mS ο (NiCl2·6H2O, cr, 298.15 K), the partial molar entropy of Ni2+ would follow from Equation (V.18):

mS ο (Ni2+) = mS ο (NiCl2·6H2O, cr) – 2 mS ο (Cl–) – 6 mS ο (H2O, l) + sol mS ο∆ (V.16) (V.18)

For example, if mS ο (NiCl2·6H2O, cr) = 344.3 J·K–1·mol–1 from [63KO/HEP] is used with auxiliary data for mS ο (Cl–) and mS ο (H2O, l) from Table IV-1,

mS ο (Ni2+) = – (126.3 ± 8.3) J·K–1·mol–1.

However, in [63KO/HEP], the value of mS ο (NiCl2·6H2O, cr) was taken from unpublished heat capacity data, and the value accepted in the present review is not inde-pendent of the value selected for f mH ο∆ (Ni2+) (cf. Section V.4.1.3.3).

Although the entropy values of the NiSO4·7H2O(cr) and the NiCl2·6H2O(cr) cycles agree within the error limits, the former has been preferred, because it is more precise and has a more reliable experimental basis.

V.2.1.3.3 Entropy of Ni2+ from its f mGο∆ and f mH

ο∆

When the independently determined values for f mGο∆ (Ni2+, 298.15 K) = − (45.77 ± 0.77) kJ·mol–1 (see Section V.2.1.1) and f mH ο∆ (Ni2+, 298.15 K) = − (54.4 ± 4.9) kJ·mol–1 (see Section V.2.1.2) are employed, f mS ο∆ (Ni2+) can be calcu-lated by Equation (V.19):

f mS ο∆ (Ni2+) = [ f mH ο∆ (Ni2+) – f mGο∆ (Ni2+)]/298.15. (V.19)

The partial molar entropy of Ni2+ follows from Equation (V.20)

mS ο (Ni2+) = f mS ο∆ (Ni2+) – mS ο (H2, g) + mS ο (Ni, cr) (V.20)

mS ο (Ni2+) = – (131.4 ± 16.6) J·K–1·mol–1.

A comparison of Equations (V.15), (V.18), (V.20) shows that three completely independent paths have led to mS ο (Ni2+) values that agree with each other within the experimental uncertainty. The result of the Section V.2.1.3.1, however, is the most pre-cise one. Plyasunova et al. [98PLY/ZHA] selected mS ο (Ni2+, 298.15 K) = − (130 ± 3) J·K−1·mol–1, which is very close to the finally selected value given in Equation (V.15):

mS ο (Ni2+, 298.15 K) = – (131.8 ± 1.4) J·K–1·mol–1.

V.2.1.4 Heat capacity of Ni2+

Spitzer et al. [78SPI/SIN], Perron et al. [81PER/ROU], and Smith-Magowan et al. [82SMI/WOO] reported apparent molar heat capacity values for NiCl2(aq); values for Ni(ClO4)2(aq) were reported by Spitzer et al. [78SPI/SIN2] and Pan and Campbell [97PAN/CAM]. There are also literature values for Ni(NO3)2(aq) [79SPI/OLO],


NiBr2(aq) [79VOR/VAS] and NiSO4(aq) [85DRA/MAD]. Values of ,mpCο (Ni2+, T), rela-tive to the standard state ,mpCο (H+, T) = 0 J·K–1·mol–1 can be determined if the apparent molar heat capacities for the salts NiX2(aq, T) are extrapolated to I = 0, and values for the partial molar heat capacities of the corresponding acids, ,mpCο (HX, aq, T), are known. The measurements of Smith-Magowan et al. [82SMI/WOO] (50 to 300°C) were limited at each temperature, were carried out at a relatively high pressure (17.7 MPa), and are not useful in the selection of a value for ,mpCο (Ni2+, 298.15 K). Similarly, the measurements of Drakin et al. [85DRA/MAD] cannot be used because they were done at higher temperatures (50 to 90°C) and concentrations.

The procedures for an extrapolation of the electrolyte apparent molar heat ca-pacities to I = 0 differ considerably in the different references. For example, the appar-ent molar heat capacities for Ni(ClO4)2(aq) found by Spitzer et al. [78SPI/SIN2] are approximately 3 J·K–1·mol–1 less negative than those found by Pan and Camp-bell [97PAN/CAM]. Yet, the reported partial molar heat capacities for 25°C differ by 6 J·K–1·mol–1. Apparent molar heat capacities for low molalities (< 0.07 m) are needed for a reasonable extrapolation to I = 0. Unfortunately, uncertainties in the apparent molar heat capacities increase with decreasing electrolyte concentration (small uncertainties in the solute concentrations become significant). This is especially evident in the low-concentration results of Perron et al. [81PER/ROU], which are so badly scattered that the values as extrapolated to I = 0 must be considered unreliable.

The same problems in extrapolation are found to a somewhat lesser degree in calculations of values of the partial molar heat capacities for the (1:1) aqueous acids. Reanalysis and comparison of the apparent molar heat capacity data for inorganic acids is not within the scope of the present review. Here, the partial molar heat capacity val-ues ,mpCο (HNO3, aq, 298.15 K) = – 71.0 J·K–1·mol–1 from Patterson and Woolley [2002PAT/WOO], ,mpCο (HCl, aq, 298.15 K) = – 123.2 J·K–1·mol–1 [2002PAT/WOO], and ,mpCο (HClO4, aq, 298.15 K) = – 26.3 J·K–1·mol–1 from Oakes and Rai [2001OAK/RAI] (based on the experimental work of Hovey and Hepler [89HOV/HEP] and Lemire and Campbell [96LEM/CAM]) have been retained. No cor-rection has been applied for values obtained at different pressures to 0.6 MPa, as these corrections are markedly less than other uncertainties in the measurements.

The unweighted average of the values for ,mpCο (Ni2+, 298.15 K) in Table V-3 (neglecting the discrepant value from Perron et al. [81PER/ROU]) is – 46.1 J·K–1·mol–1 (relative to ,mpCο (H+) = 0 J·K–1·mol–1), which is very close to the value – 47.0 J·K−1·mol−1 from the most extensive set of experimental data [97PAN/CAM]. The se-lected value is:

,mpCο (Ni2+, 298.15 K) = – (46.1 ± 7.5) J·K–1·mol–1.

The unweighted average has been used, and the uncertainty of 7.5 J·K–1·mol–1 has been selected, because most of the uncertainty in the value of the partial molar heat


capacity arises from systematic experimental errors and the method of extrapolation to infinite dilution.

Over the temperature range of the measurements, the value of ,mpCο (Ni2+) shows only a very shallow maximum [97PAN/CAM]. Within the uncertainties of the measurements, the value for 298.15 K can be used over the temperature range between 298.15 and 358.15 K. The values of the apparent molar heat capacities for nickel salts depend strongly on the electrolyte concentrations [97PAN/CAM]1.

Table V-3: Partial molar heat capacity values for Ni(II) salts in aqueous solution at 298.15 K. salt ,mpCο (Ni-Salt, aq, 298.15 K)/

J·K–1·mol–1 ,mpCο (Ni2+, 298.15 K)/

J·K–1·mol–1 reference

Ni(ClO4)2(aq) – 99.6 – 47.0 [97PAN/CAM] – 93.1 – 40.5 [78SPI/SIN] NiCl2(aq) – 294.0 – 47.6 [78SPI/SIN] – 306 – 59.6(a) [81PER/ROU] Ni(NO3)2(aq) – 186.5 – 44.5 [79SPI/OLO] NiBr2 – 310 – 51(b) [79VOR/VAS] (a) The apparent molar heat capacities of Perron et al. [81PER/ROU] for low concentrations are too badly

scattered to give a reliable value for ,mpCο (Ni2+, 298.15 K). A second value for ,mpCο (NiCl2, aq, 298.15 K) from the same paper (– 295.3 J·K–1·mol–1) is based primarily on higher concentration values.

(b) Calculated using a value for ,mpCο (Br–, 298.15 K) based on the value for HCl(aq) [2002PAT/WOO] and ,mpCο (Br–, 298.15 K) – ,mpCο (Cl–, 298.15 K) = 6 J·K–1·mol–1

[96HEP/HOV], [96CRI/MIL].

V.2.2 Other oxidation states Nickel(III) and nickel(IV) complexes are well documented, but have not been treated in this review. Whereas a ‘transient’ Ni3+ species presumably exists, there is no aqua ion representative of Ni(IV).

V.2.2.1 Ni3+

Although the trivalent oxidation state of Ni can be stabilised by certain arrangements of donor ligands, no stable aqua ion of Ni(III) appears to exist in dilute aqueous solutions at ambient conditions [97RIC]. Vogel [68VOG] found strong evidence that a redox couple composed of tri- and divalent nickel can exist in potassium hydroxide melts con-taining a mass fraction of water w(H2O) = 0.15 at 498 and 523 K. The standard potential measured, E°(Ni3+ | Ni2+) = 1.168 and 1.174 V at 498 and 523 K, respectively, cannot be compared with the value estimated for 298 K in predominantly aqueous solutions.

1 After this review was completed, results from an extensive experimental study (from 278.15 to 393.15 K

at 0.35 MPa) were reported [2004BRO/MER] for aqueous Ni(NO3)2 solutions. The new apparent molar heat capacity values are markedly more negative than those reported by Spitzer et al. [79SPI/OLO].


Bhattacharya et al. [86BHA/MUK] synthesised the tris(2,2'-bipyridine 1,1'-dioxide) Ni(III) complex, 3

2 3Ni(bpyO ) + , and measured the potential of Reaction (V.21) in acetonitrile against the standard calomel electrode (SCE).

32 3Ni(bpyO ) + + e– 2

2 3Ni(bpyO ) + (V.21)

These authors observed a linear correlation between the M3+ | M2+ and 3

2 3M(bpyO ) + | 22 3M(bpyO ) + potentials (M being Cr, Fe, Mn, Co) versus the SCE in

aqueous and acetonitrile solutions, respectively. A correlation of similar quality is ob-tained when the standard electrode potentials of M3+ | M2+ [89BRA] are linearly fitted against the potentials measured by [86BHA/MUK], see Figure V-3. As the potential of

32 3Ni(bpyO ) + | 2

2 3Ni(bpyO ) + can be measured, the experimentally inaccessible standard electrode potential of Ni3+ | Ni2+ can be estimated as E°(Ni3+ | Ni2+, 298.15 K) = (2.56 ± 0.27) V. Twice the standard deviation of the regression analysis has been taken as an uncertainty limit. With this result and f mGο∆ (Ni2+, 298.15 K) = − (45.77 ± 0.77) kJ·mol–1 (see Section V.2.1.1) the Gibbs energy of formation of Ni3+ has been roughly estimated to be f mGο∆ (Ni3+, 298.15 K) = (202 ± 26) kJ·mol–1. Although this result is of general interest, it is not sufficiently well defined to be a selected value in this review.

Figure V-3: Estimation of the Ni3+ | Ni2+ standard electrode potential. Solid straight line: least square fit of SHE data; dash straight line: least square fit of SCE data; ▼ SHE po-tentials; ▲ SCE potentials ; extrapolated value of 298Eο (Ni3+ | Ni2+, SHE); extrapo-lated value of 298Eο (Ni3+ | Ni2+, SCE).

-1.2 -0.8 -0.4 0.0 0.4 0.8 1.2 1.6

-1.0

-0.5

0.0

0.5

1.0

1.5

2.0

2.5

3.0

E°2 98 / V(SCE), 32 3M(bpyO ) + ‌ 2

2 3M(bpyO ) +

E°2 9

8 / V

, M3+

‌ M2+

V.3 Oxygen and hydrogen compounds and complexes 90

V.3 Oxygen and hydrogen compounds and complexes V.3.1 Aqueous nickel(II) hydroxo complexes V.3.1.1 Hydroxo complexes in acidic or neutral solutions

The Ni(II) ion is a rather weak acid in aqueous solution. The hydrolysis reactions of Ni(II) may be defined by the following generalised equilibrium process:

mNi2+ + nH2O(l) 2Ni (OH) m nm n

− + nH+; ,*

n mb (V.22)

The formation of five water soluble hydroxo complexes is generally recog-nised: NiOH+, Ni(OH)2(aq), 3Ni(OH)− , Ni2OH3+, 4

4 4Ni (OH) + .

Several experimental difficulties may arise during the study of the hydrolysis of Ni(II): (i) nickel forms a relatively insoluble hydroxide, i.e., the hydroxide or basic salt precipitates even in solutions with a low degree of hydrolysis, and therefore the solutions have a low buffer capacity; (ii) because of the formation of oligomer hydroxo complexes, the mole fractions of different hydroxo species depend strongly on the ex-perimental solution concentrations; (iii) the kinetics of formation of oligomer species is relatively slow. The first mentioned problem can be avoided by using concentrated solu-tions, but this introduces complications (ii) and (iii). Consequently, the choice of ex-perimental methods and conditions is a critical factor in determining the accuracy of the hydrolysis constant values. The published data often refer to high ionic strengths and concentrations of Ni(II), while parallel measurements at different temperatures are usu-ally missing. For these reasons, a limited number of sets of serviceable experimental data are available to evaluate the thermodynamic properties of Ni(II) hydroxo com-plexes.

Only a few compilations and reviews have been published on the thermody-namic data of Ni(II) hydrolysis. Baes and Mesmer [76BAE/MES] have critically re-viewed the available experimental data, and provided recommended values for I = 0, 0.1, 1.0 and 3.0 mol·kg–1. A slightly different set of stability constants has been recom-mended by Smith and Martell [76SMI/MAR] in their compilation. In a recent review of Plyasunova et al. [98PLY/ZHA], critical evaluations, and in some cases re-evaluations, of literature data have been carried out, using the SIT approach. The values selected by these authors are given in Appendix A.

The hydrolysis of Ni(II) has mainly been studied by potentiometric titrations, although some solubility, kinetic and calorimetric studies have also been published.


In all cases when sufficient experimental data were reported, a re-evaluation of data was performed for the purposes of this review (see discussions in Appendix A). Such re-assessments were required for several reasons:

(i) Several background electrolytes ((Na)ClO4, (Li)ClO4, (Na)Cl, K(Cl) and (K)NO3) have been used to keep the anion concentration constant in experimen-tal solutions, and thereby to limit changes in activity coefficients. Generally, the cations of these background electrolytes were singly-charged, and in most stud-ies varying and relatively high Ni(II) concentrations were used (up to 1.5 M). In such cases, multiply-charged Ni(II) species were a substantial fraction of the to-tal cationic species in the solutions. In a few cases polynuclear Ni hydroxo com-plexes even predominated. Consequently, the constancy of the activity coeffi-cients is questionable due to both the medium effect and the changing ionic strength. Moreover, at high chloride or nitrate concentrations, the complexation of these anions with Ni(II) cannot be ignored. Therefore, a re-evaluation of data was performed to minimise the influence of the medium effects (often using only a portion of the reported data) and to take into account the presence of NiX+ (X– = Cl– or 3NO− ) complexes.

(ii) In publications that appeared before 1965, the experimental data were interpreted in terms of formation only of the mononuclear NiOH+ complex. However, after the work of Sillén and his co-workers [65BUR/LIL], the formation of

4+4 4Ni (OH) as the predominant hydrolytic species at [Ni2+]tot > 1 mM, became

widely accepted. Consequently, as pointed out in [98PLY/ZHA], almost all pre-viously reported values should be corrected. In such cases, re-evaluation was performed including oligonuclear species, especially 4+

4 4Ni (OH) .

(iii) In addition, although the experimental methods (e.g., pH-measurements) were already well established in the 1950s, the evaluation of experimental data has developed considerably since that time (in a great part of the cited works, graphi-cal evaluation was used). Therefore, re-assessment of data was performed to check (and if necessary complete) the proposed equilibrium model.

The formation of two main species is generally recognised in this pH region: NiOH+, and 4

4 4Ni (OH) + . Due to the presence of oligonuclear species, the onset of the hydrolysis reaction of Ni(II) is strongly concentration dependent. Under natural condi-tions above pH = 6.9, and at low Ni(II) concentrations (< 0.0001 M), only the dominant aqua complex and the species NiOH+ may be significant. In more concentrated solu-tions ([Ni2+]tot > 0.005 M), however, the tetramer 4

4 4Ni (OH) + is the major hydroxo complex before precipitation occurs. The formation of a minor species Ni2OH3+ has been suggested by several authors [65BUR/LIL], [65BUR/LIL2], [66BUR/IVA], [71BUR/ZIN], [71OHT/BIE], [78BUR/KAM], because by adding this complex to the model a better fit of the data could be achieved. Some other species were also postulated by different authors ( 2

2 2Ni (OH) + [69KEN2], 22 6Ni (OH) − [71BHA/SUB], 3

3 3Ni (OH) +


[66BUR/IVA]). However, the formation of these species has never been independently confirmed. Actually the re-evaluation of data reported in [66BUR/IVA] indicated the presence of 4

4 4Ni (OH) + rather than 33 3Ni (OH) + (see discussion in Appendix A).

The following equations describe the formation of the first mentioned three species:

Ni2+ + H2O(l) NiOH+ + H+ 1,1*b (V.23)

2Ni2+ + H2O(l) Ni2OH3+ + H+ 1,2*b (V.24)

4Ni2+ + 4H2O(l) 44 4Ni (OH) + + 4H+ 4,4

*b (V.25)

The reported and re-calculated hydrolysis constants are listed in Table V-4.

As mentioned above, several early studies reported the formation of NiOH+ as the only hydrolytic species, but the experiments were conducted at relatively high con-centrations of Ni(II), where a considerable amount of 4

4 4Ni (OH) + species can be ex-pected. To avoid its presence, the [Ni2+]tot should not exceed 0.001 M, and the pH should not exceed 8.6 to preclude precipitation. The studies of Perrin [64PER] and of Funahashi and Tanaka [69FUN/TAN] were the only ones in which conditions appropri-ate to the determination of a reliable formation constant for NiOH+ were used. The re-evaluation of the experimental data reported in [63BOL/JAU], [65BUR/LIL], [65BUR/LIL2], [71BUR/ZIN] resulted in acceptable 10 1,1

*log b values. In a few cases, a simple correction was applied to take into account the formation of the NiCl+ complex [71OHT/BIE], [80MIL/BUG]. However, some of the reported values cannot be cor-rected, as insufficient experimental details or data have been communicated [52CHA/COU], [52GAY/WOO], [53SCH/POL], [56CUT/KSA], [59ACH], [63SHA/DES]. Tremaine and LeBlanc have reported a sharp increase of the first hy-drolysis constant of Ni(II) between 150 and 300°C [80TRE/LEB2]. These data indicate that the formation of NiOH+ is more pronounced at higher temperatures. The mole frac-tion of NiOH+ never exceeds 15% at 25°C, but reaches 70% at 300°C. Although the reported data seem to be reliable, the temperature and pressure range used is well above of the scope of this review. The authors’ extrapolation to the standard state resulted in considerably different values from those determined at 25°C.


Table V-4: Experimental equilibrium data (logarithmic values) for the hydrolysis of Ni(II) in acidic or neutral solutions. Method Medium Im t /°C 10 ,

*log n mb reported(a)

10 ,*log n mb

retained(b) Reference

Ni2+ + H2O(l) NiOH+ + H+ gl KCl 0.1 30 – 9.4 – [52CHA/COU] gl NiCl2 → 0 25 – (10.64 ± 0.24) – [52GAY/WOO] gl NiCl2, NiSO4 ~ 0.08 25 – 9 – [53SCH/POL] gl Ni(ClO4)2 → 0 25 – 8.94 – [56CUT/KSA] gl KClO4, Ni(NO3)2 → 0 25 – (10.92 ± 0.12) – [59ACH] gl (Na)ClO4 0.25

1.05 1.05 1.05 1.05

25 25 30 40 50

– (9.76 ± 0.05) – – (9.65 ± 0.04) – (9.51 ± 0.04) – (9.32 ± 0.06)

– (9.84 ± 0.30)(c) – (9.71 ± 0.30)(c) – (9.72 ± 0.30)(c) – (9.62 ± 0.30)(c) – (9.66 ± 0.30)(c)

[63BOL/JAU]

gl (K)NO3 → 0 15 20 25 30 36 42

– (10.22 ± 0.04) – (10.05 ± 0.04) – (9.86 ± 0.03) – (9.75 ± 0.07) – (9.58 ± 0.06) – (9.43 ± 0.03)

– (10.22 ± 0.20)(d) – (10.07 ± 0.20)(d) – (9.80 ± 0.20)(d) – (9.74 ± 0.20)(d) – (9.59 ± 0.20)(d) – (9.38 ± 0.20)(d)

[64PER]

gl (Na)ClO4 3.50 25 – – (9.79 ± 0.12)(c) [65BUR/LIL] gl (Na)Cl 3.20 25 – – (9.53 ± 0.20)(c) [65BUR/LIL2] kin NaClO4 0.10 25 – (9.5 ± 0.1) – (9.5 ± 0.2) [69FUN/TAN] kin ? → 0 ? – (11.0 ± 0.2) – [71HOH/GEI] gl (Na)Cl 3.20 25 ≤ – 10.5 – (10.23 ± 0.20)(c) [71OHT/BIE] gl (Na)Cl 3.20 60 – – (8.45 ± 0.30)(c) [71BUR/ZIN] pH (sp) (Na)ClO4 0.82 25 – (8.93 ± 0.1) – [75CLA/KEP] sol HCl, NaOH → 0 150

200 250 300

– (8.05 ± 0.48) – (6.95 ± 0.22) – (6.03 ± 0.20) – (5.23 ± 0.36)

– – – –

[80TRE/LEB2]

gl (K)Cl

0.51 1.57

25 – (9.87 ± 0.40) – (10.06 ± 0.40)

– (9.81 ± 0.50)(c) – (9.93 ± 0.50)(c)

[80MIL/BUG]

sol NaClO4 0.01 25 – (8.35 ± 0.10) – [97MAT/RAI]



Table V-4 (continued)

Method Medium Im t /°C 10 ,*log n mb

reported(a) 10 ,

*log n mb retained(b)

Reference

2Ni2+ + H2O(l) Ni2OH3+ + H+ gl (Na)ClO4 3.50 25 – (10.0 ± 0.5) – (9.8 ± 0.5)(c) [65BUR/LIL] gl (Na)Cl 3.20 25 – (9.3 ± 0.2) – (8.95 ± 0.40)(c) [65BUR/LIL2] gl (Na)NO3 3.30 25 – (9.6 ± 0.2) – [66BUR/IVA] gl (Na)Cl 3.20 25 – (10.5 ± 0.5) – (10.0 ± 0.6)(c) [71OHT/BIE] gl (Na)Cl 3.20 60 – (8.5 ± 0.2) – (8.58 ± 0.60)(c) [71BUR/ZIN] gl (Na)Br 3.30 25 – (9.5 ± 0.1) – (8.82 ± 0.40)(c) [78BUR/KAM]

4Ni2+ + 4H2O(l) 44 4Ni (OH) + + 4H+

gl (Na)ClO4 0.25 1.05 1.05 1.05 1.05

25 25 30 40 50

– – – – –

– (27.34 ± 0.30)(c) – (27.66 ± 0.40)(c) – (27.10 ± 0.30)(c) – (26.39 ± 0.30)(c) – (25.24 ± 0.30)(c)

[63BOL/JAU]

gl (Na)ClO4 3.50 25 – (27.37 ± 0.02) – (27.36 ± 0.05)(c) [65BUR/LIL] gl (Na)Cl 3.20 25 – (28.42 ± 0.05) – (27.52 ± 0.20)(c) [65BUR/LIL2] gl (Na)ClO4 1.61 25 – (27.03 ± 0.06) – (27.00 ± 0.12) [69KOL/KIL] gl (Na)Cl 3.20 25 – (28.55 ± 0.02) – (27.57 ± 0.30)(c) [71OHT/BIE] gl (Na)Cl 3.20 60 – (25.33 ± 0.02) – (24.06 ± 0.30)(c) [71BUR/ZIN] gl (Li)ClO4 3.48 25 – (27.32 ± 0.08) – (27.25 ± 0.16) [73KAW/OTS] sp 0.8 M (Na)ClO4 0.82 25 – (27.00 ± 0.25) – (26.99 ± 0.40)(c) [75CLA/KEP] gl (Na)Br 3.3 25 – (28.18 ± 0.05) – (27.15 ± 0.30)(c) [78BUR/KAM] gl (K)Cl 0.51

1.03 1.57 2.13 3.31

25 – (28.04 ± 0.10) – (28.14 ± 0.06) – (28.28 ± 0.04) – (28.24 ± 0.04) – (28.58 ± 0.06)

– – – – –

[80MIL/BUG]

gl NiCl2 → 0 25 – (27.0 ± 0.1) – (27.0 ± 0.3) [91SPI/TRI]

(a) Reported values. (b) Accepted values corrected to molal scale. The accepted values reported in Appendix A are expressed on

the molar or molal scales, depending on which units were used originally by the authors. (c) Re-evaluated value, see Appendix A. (d) Corrected to I = 0 by means of the SIT.

A single least squares SIT analysis was done using all the accepted (appropri-ately weighted) constants for 25°C in perchlorate and chloride media (Figure V-4 and Figure V-5). The available data for chloride media are somewhat less consistent than those for perchlorate media. This analysis results in the selected value of:

10 1,1*log οb ((V.23), 298.15 K) = – (9.54 ± 0.14).


In this analysis, the value recalculated from Perrin [64PER] for I = 0 and 25ºC (weighted according to its estimated uncertainty) was also included (see Appendix A). From the slopes in Figure V-4 and Figure V-5, ∆ε( 4ClO− ) = – (0.088 ± 0.061) kg·mol–1 and ∆ε(Cl–) = – (0.063 ± 0.072) kg·mol–1 are calculated. Using the NEA-TDB values of 0.37, 0.17, 0.14 and 0.12 kg·mol–1 for ε(Ni2+, 4ClO− ), ε(Ni2+,Cl–), ε(H, 4ClO− ) and ε(H+,Cl–), respectively, leads to the selected values:

ε(NiOH+, 4ClO− ) = (0.14 ± 0.07) kg·mol–1

and

ε(NiOH+, Cl–) = – (0.01 ± 0.07) kg·mol–1 .

From the selected 10 1,1*log οb value, r mGο∆ ((V.23), 298.15 K) = (54.5 ± 0.8)

kJ·mol–1 can be derived. The Gibbs energy of formation of NiOH+ is calculated using the selected values for Ni2+ and H2O(l).

f mGο∆ (NiOH+, 298.15 K) = – (228.5 ± 1.1) kJ·mol–1.

When the discussion of Perrin [64PER] in Appendix A is also considered, the ε(NiOH+, X–) values (X = Cl–, 3NO− , 4ClO− ) show a larger than expected spread – 0.01, – 0.27, 0.14 kg·mol–1, though the value for ε(NiOH+, 3NO− ) is not a selected value in the present review.

The 10 1,1*log οb ((V.23), 298.15 K) value reported above is somewhat higher

than recommended by Baes and Mesmer (– (9.86 ± 0.03), [76BAE/MES]), but agrees well with the one obtained by [98PLY/ZHA] viz. – (9.50 ± 0.36).

From the temperature dependence of the first hydrolysis constant determined by Perrin [64PER], r mH ο∆ = (51.7 ± 3.4) kJ·mol–1 can be obtained. The re-evaluation of the data reported in [68ARN] yielded r mH∆ = (54.3 ± 2.0) kJ·mol–1 for 3.2 m NaCl me-dium. Combining the accepted hydrolysis constants of [65BUR/LIL2] and [71BUR/ZIN], r mH∆ = (59 ± 10) kJ·mol–1 can be calculated for the formation of the species NiOH+ in 3.2 m NaCl medium. From the 10 1,1

*log b values reported for 150 – 300°C in [80TRE/LEB2], r mH∆ = (87 ± 15) kJ·mol–1 can be derived, but this value was not considered in this review, as it is probably valid only in the temperature range cov-ered by the measurements. The temperature variation of 10 1,1

*log b values accepted in [63BOL/JAU] is small compared to the uncertainties in the values. Consequently, no enthalpy value has been derived from these constants. As the relatively high uncertain-ties of the remaining three data cannot be used to extract a reliable ionic strength de-pendence, it was assumed to be negligible. Assigning statistical weights to the data , the value of:

r mH ο∆ ((V.23), 298.15 K) = (53.8 ± 1.7) kJ·mol–1

is selected. The standard enthalpy of formation of NiOH+ is:

f mH ο∆ (NiOH+, 298.15 K) = – (287.0 ± 1.9) kJ·mol–1.


Figure V-4: Extrapolation to Im = 0 of the experimental data for Reaction (V.23) in NaClO4 media. Figure V-5: SIT analysis of the experimental data for Reaction (V.23) in chloride me-dia, including 10 1,1

*log οb ((V.23), 298.15 K) obtained in NaClO4.

0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0-10.5

-10.0

-9.5

-9.0

-8.5

log 10

β 1,1 +

2D

- lo

g 10a w

I / mol·kg−1

Ni2+ + H2O(l) NiOH+ + H+

0 1 2 3-10.5

-10.0

-9.5

-9.0

-8.5

log 10

β 1,1 +

2D

− lo

g 10a w

I / mol · kg−1

Ni2+ + H2O(l) NiOH+ + H+


The dinuclear Ni2OH3+ species is always a minor component, and its formation is considered to account for small deviations between the observed and calculated titra-tion curves. Therefore, the rather wide scatter of the reported values is not surprising. As stated in [76BAE/MES], the reported constants for this minor species should be con-sidered as the upper limits for its stability, if indeed such a species is present. In some cases, the accuracy of the reported experimental data was insufficient to take into ac-count the formation of Ni2OH3+ during the re-evaluation [63BOL/JAU], [75CLA/KEP], thus all accepted constants refer to high ionic strength. Therefore, the extrapolation to Im = 0 was made assuming ε(Ni2OH3+, 4ClO− ) = (0.50 ± 0.15) kg·mol–1, based on the estimated value for ε(Be2OH3+, 4ClO− ), see Table B-4, (ε(Ni2OH3+, 4ClO− ) = (0.59 ± 0.16) kg·mol–1 is reported in [98PLY/ZHA], based on the corrected data from [59ACH], [63BOL/JAU], [75CLA/KEP]). In this way, ∆ε((V.24), 4ClO− ) is estimated to be − (0.1 ± 0.2) kg·mol–1, which leads to the selection:

10 1,2*log οb ((V.24), 298.15K) = – (10.6 ± 1.0).

From this value, r mGο∆ ((V.24), 298.15 K) = (60.5 ± 5.7) kJ·mol–1 and

f mGο∆ (Ni2OH3+, 298.15 K) = – (268.2 ± 5.9) kJ·mol–1

can be derived.

Two reliable values are available concerning the enthalpy of reaction to form Ni2OH3+. From the re-evaluation of the calorimetric data reported in [68ARN] r mH∆ = (45.9 ± 1.6) kJ·mol–1 can be calculated. Combining the accepted hydrolysis constants of [65BUR/LIL2] and [71BUR/ZIN], r mH∆ = (20 ± 20) kJ·mol–1 can be obtained. Both values refer to 3.2 m NaCl medium. The more precise calorimetric value is accepted, but with an increased uncertainty, and negligible dependence on the ionic strength is assumed.

r mH ο∆ ((V.24), 298.15 K) = (45.9 ± 6.0) kJ·mol–1

is selected in this review, which leads to:

f mH ο∆ (Ni2OH3+, 298.15 K) = – (350.0 ± 6.3) kJ·mol–1.

At Ni(II) concentrations higher than 0.005 M, the 4+4 4Ni (OH) complex is

dominant in the acidic pH region. Most of the accepted values concerning the formation of the tetrameric hydroxo species have been re-evaluated [63BOL/JAU], [65BUR/LIL], [65BUR/LIL2], [71BUR/ZIN], [75CLA/KEP], [78BUR/KAM] and corrected consider-ing complexation with the background anion [65BUR/LIL2], [71BUR/ZIN], [71OHT/BIE], [78BUR/KAM].

The SIT analysis of the accepted constants for perchlorate medium (Figure V-6) resulted in the selected value of:

10 4,4*log οb ((V.25), 298.15 K) = – (27.52 ± 0.15).


From the slope in Figure V-6, ∆ε((V.25), 4ClO− ) = (0.16 ± 0.05) kg·mol–1 can be calculated. Using the selected values for ε(Ni2+, 4ClO− ) and ε(H+, 4ClO− ), ∆ε((V.25)

4ClO− ) leads to a value of ε( 44 4Ni (OH) + , 4ClO− ) = (1.08 ± 0.08) kg·mol–1.

In sodium chloride medium, only two values are accepted at 25°C [65BUR/LIL2], [71OHT/BIE]; both are valid for Im = 3.2 mol·kg–1. Combining them with the above selected 10 4,4

*log οb ((V.25), 298.15 K), ∆ε((V.25), Cl–) = (0.23 ± 0.07) kg·mol–1 can be derived, which leads to a value of ε( 4

4 4Ni (OH) + ,Cl–) = (0.43 ± 0.08) kg·mol–1. A similar treatment, using the value accepted for NaBr medium, results in ∆ε((V.25), Br–) = (0.11 ± 0.07) kg·mol–1. Since no value is currently available for ε(H+, Br–), only an estimate can be given for ε( 4

4 4Ni (OH) + ,Br–). Assuming ε(H+,Br–) = ε(H+,Cl–), ε( 4

4 4Ni (OH) + ,Br–) = (0.71 ± 0.12) kg·mol–1 can be estimated.

Figure V-6: Extrapolation to Im = 0 of the experimental data for Reaction (V.25) in per-chlorate media.

0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5-29.5

-29.0

-28.5

-28.0

-27.5

-27.0

-26.5

-26.0

log 10

β4,

4 − 4

D −

4 lo

g 10 a

w

I / mol·kg−1

4Ni2+ + 4H2O(l) 44 4Ni (OH) + + 4H+


From the selected 10 4,4*log οb value, r mGο∆ ((V.25), 298.15 K) = (157.1 ± 0.9)

kJ·mol–1 can be obtained. The Gibbs energy of formation of 44 4Ni (OH) + is

f mGο∆ ( 44 4Ni (OH) + , 298.15 K) = – (974.6 ± 3.2) kJ·mol–1.

From the temperature dependence of 10 4,4*log b values recalculated from the

data in [63BOL/JAU], r mH∆ = (174 ± 15) kJ·mol–1 can be derived for 1.0 M NaClO4 medium (see Appendix A). The re-evaluation of the calorimetric data reported in [68ARN] yielded r mH∆ = (192.9 ± 1.0) kJ·mol–1 for 3.2 m NaCl medium (see Appendix A). Combining the accepted hydrolysis constants of [65BUR/LIL2] and [71BUR/ZIN],

r mH∆ = (186 ± 6) kJ·mol–1 can be obtained for the formation of 44 4Ni (OH) + species in

3.2 m NaCl medium.

Based on these three values in two different media, and the probability that the most accurate result is the value determined calorimetrically in 3.2 m NaCl(aq), the value:

r mH ο∆ ((V.25), 298.15 K) = (190 ± 10) kJ·mol–1

is selected in this review, which leads to

f mH ο∆ ( 44 4Ni (OH) + , 298.15 K) = – (1173.4 ± 10.6) kJ·mol–1.

Electrostatic models such as the one presented in [94PLY/GRE] have been used to estimate the temperature dependence of formation constants and entropies of hydrolysis reactions of polynuclear hydroxo complexes, such as 4

4 4Ni (OH) + . Such semi-theoretical treatments are beyond the scope of this review.

V.3.1.2 Hydroxo complexes in alkaline solutions

Three further complexes have been reported to form in Ni(II) solutions above ca. pH = 9: Ni(OH)2(aq), 3Ni(OH)− and 2

4Ni(OH) − , based on the increasing solubility of Ni(OH)2(cr) in alkaline solutions. According to a re-evaluation, carried out in this re-view, only the hydrolysis constant of the second species (V.27) can be determined with reasonable certainty:

Ni2+ + 2H2O Ni(OH)2(aq) + 2H+; 2,1*b (V.26)

Ni2+ + 3H2O 3Ni(OH)− + 3H+; 3,1*b (V.27)

Ni2+ + 4H2O 24Ni(OH) − + 4H+; 4,1

*b . (V.28)

The values of these hydrolysis constants selected by [76BAE/MES] and [98PLY/ZHA], (see Appendix A) are based on a few experimental points of a single paper [49GAY/GAR]. These data can, however, be equally well described when only

3Ni(OH)− is assumed to be present.


-3.5 -3.0 -2.5 -2.0 -1.5 -1.0 -0.5 0.0 0.5 1.0 1.5-9

-8

-7

-6

-5

-4

-3

log 10

([Ni(I

I)]to

t/mol

·kg−1

)

log10([OH−]/mol·kg−1)

K0s,2,1, K

0s,3,1, ∆ε

K0s,3,1, ∆ε

K0s,3,1

exp. data [49GAY/GAR]

Three publications [49GAY/GAR], [80TRE/LEB2], [89ZIE/JON] report ther-modynamic properties for the Ni(OH)2(aq) and 3Ni(OH)− species. The hydrolysis con-stants for reactions (V.26) and (V.27) can be derived from solubility equilibria:

Ni(OH)2(cr) Ni(OH)2(aq) ,2,1sK ο (V.29)

Ni(OH)2(cr) + OH– 3Ni(OH)− ,3,1sK ο (V.30)

with

Ni(OH)2(cr) + 2H+ Ni2+ + 2H2O ,0*

sK ο (V.31)

and assuming that hydrolysis of Ni2+ to NiOH+ was essentially complete in the experi-mental solutions used.

Several hydrolysis models can be considered in interpreting nickel hydrolysis data obtained from experiments in basic solutions (see Figure V-7).

Figure V-7: Solubility of Ni(OH)2(s) in alkaline solutions at 298.15 K. • data from [49GAY/GAR], estimated error: ± 0.5 log10 units, dash-dotted line: log10[Ni(II)]tot =

10 ,3,1log sK ο + log10[OH–]ini, dotted line: log10 [Ni(II)]tot = 10 ,3,1log sK ο – ∆ε·[OH–]ini +

log10[OH–]ini; solid line: log10[Ni(II)]tot = ini·[OH ]10 ,2,1 ,3,1 inilog ( + [OH ] 10 )s sK K

−−∆εο ο −


Gayer and Garrett [49GAY/GAR] reported 10 ,0*log sK ο = (10.8 ± 0.15),

10 ,2,1log sK ο ≈ – 7 and 10 ,3,1log sK ο = – (4.2 ± 0.6). By combining these values, and the ionisation constant of water (pKw = 14.0), 10 2,1

*log οb = – (17.8 ± 1.0) and 10 3,1*log οb =

− (29.0 ± 0.8) can be derived for reactions (V.26) and (V.27), respectively. Re-evaluation of the solubility data reported by these authors in alkaline solutions, using the SIT approach, assuming formation of both 3Ni(OH)− and Ni(OH)2(aq) (but exclud-ing the solubility values obtained in 8 M, 10 M and 15 M NaOH(aq)), results in

10 ,2,1log sK ο ≈ – (7.20 ± 0.22), 10 ,3,1log sK ο = − (4.58 ± 0.11), and ∆ε = (0.84 ± 0.24) kg·mol–1. These equilibrium constants correspond to 10 2,1

*log οb = − (18.00 ± 0.30) and 10 3,1

*log οb = − (29.38 ± 0.30)1.

A re-evaluation of the solubility data, using the SIT approach, and assuming 3Ni(OH)− (but not Ni(OH)2(aq)) is formed in the dissolution process, results in

10 ,3,1log sK ο = − (4.40 ± 0.12), ∆ε = (0.66 ± 0.17) kg·mol–1. The equilibrium constant corresponds to 10 3,1

*log οb = − (29.20 ± 0.48). If activity coefficient effects are ne-glected, 10 ,3,1log sK ο = − (4.63 ± 0.22) and 10 3,1

*log οb = − (29.43 ± 0.48).

Tremaine and LeBlanc obtained, 10 2,1*log οb = − (20.26 ± 1.77) and 10 3,1

*log οb = – (29.78 ± 1.97) at 298.15 K, by extrapolation, based on the solubility of NiO in NaOH solutions measured between 150 and 300°C [80TRE/LEB2]. A re-evaluation revealed that their experimental data can also be explained with 3Ni(OH)− as the only solubility controlling species in alkaline solutions (see Appendix A). The revised ex-trapolation to T = 298.15 K resulted in 10 3,1

*log οb = – (30.9 ± 1.3), in reasonable agree-ment with [80TRE/LEB2], although Ni(OH)2 was neglected. Ziemniak et al. [89ZIE/JON] also measured the solubility of NiO between 18 and 289°C in alkaline solutions containing 2

4HPO − ion. A re-evaluation of their data by [95LEM] resulted in 10 2,1

*log οb = − (19.77 ± 1.00) and 10 3,1*log οb = – (30.9 ± 1.3). As the numerical value of

10 3,1*log οb remains within its error limits whether Ni(OH)2 is taken into account or not,

no thermodynamic quantities for the latter species are selected in this review. The solid line of Figure V-7 shows that a value of 10 ,2,1log sK ο = – (7.2 ± 0.2) is also consistent with the data. No thermodynamic quantities for the latter species are selected in this review, though – 7 can be tentatively assigned as its upper limit. The value for 10 3,1

*log οb ob-tained by the re-evaluation of data reported in [49GAY/GAR] has been selected, but with an increased uncertainty to reflect the lower values obtained from use of the higher temperature data of [80TRE/LEB2] and [89ZIE/JON]:

10 3,1*log οb ((V.27), 298.15 K) = – (29.2 ± 1.7).

1 Gayer and Garrett [49GAY/GAR] did not report on the characterisation of their solid phase, which may,

or may not, have been identical to β-Ni(OH)2, the solid discussed in Section V.3.2.2.1.1. Therefore, the value 10 ,0

*log sK ο = (10.8 ± 0.15) was used for these recalculations, rather than the value for β-Ni(OH)2,

10 ,0*log sK ο = (11.02 ± 0.20), selected in the present review (also see the discussion of [49GAY/GAR] in

Appendix A).


From this value, r mGο∆ ((V.27), 298.15 K) = (166.7 ± 9.7) kJ·mol–1 can be ob-tained, while the Gibbs energy of formation is

f mGο∆ ( 3Ni(OH)− , 298.15 K) = – (590.5 ± 9.7) kJ·mol–1.

From the temperature dependence of 10 2,1*log οb and 10 3,1

*log οb measured be-tween 150 and 300°C, r mH∆ (V.26) = (109 ± 15) kJ·mol–1 and r mH∆ (V.27) = (119 ± 15) kJ·mol–1 were obtained by [80TRE/LEB2]. From the re-evaluation of Ziem-niak's data [89ZIE/JON] reported in [95LEM], r mH ο∆ (V.26) = (94 ± 10) kJ·mol–1 and

r mH ο∆ (V.27) = (126 ± 15) kJ·mol–1 can be estimated. The re-evaluation of the data of Tremaine and LeBlanc results in r mH ο∆ ((V.27), 298.15 K) = (121.2 ± 6.5) kJ·mol–1 which agrees satisfactorily with the values given by [80TRE/LEB2] and [95LEM]. Al-though the quantities are not sufficiently accurate to allow selection of selected values for the standard reaction enthalpies in question (see above), they can be used as rough estimates.

We have not found any convincing evidence for formation of 24Ni(OH) − .

V.3.2 Solid nickel oxides and hydroxides V.3.2.1 Ni(II) oxide

Bunsenite (NiO) is extremely rare, and a member of the periclase group of minerals. It had been discovered as early as 1868 in Johanngeorgenstadt, Erzgebirge, Saxony, Ger-many in a hydrothermal Ni-U vein. It is translucent to opaque, and occurs as well-formed fine sized crystals (up to 3 mm). Its hardness is 5.5, the colour pistachio-green, the lustre adamantine, and the streak brownish black.

The lattice of NiO is face centred cubic, space group: Fm3m , with unit cell dimensions a0 = 4.177 Å, Z = 4, ρ(calc.) = 6.806 g·cm–3 (according to JCPDS-ICDD card No.4-835, [53SWA/TAT]), ρ(obs.) = 6.4 – 6.8 g·cm–3.

A very accurate technique for the determination of the standard entropy mS ο (298.15 K) of a solid crystalline compound is the integration of low-temperature

heat capacity data, ,mpCο , between 0 and 298.15 K:

mS ο (298.15 K) = 298.15

,m

0

dpCT

T

ο

∫ . (V.32)

In the case of nickel oxide several publications dealing with heat capacity measurements in the temperature range 3.2 – 477.8 K are available [40SEL/DEW], [57KIN], [74WHI], [79DUB/NAU], [93WAT]. Seltz et al. [40SEL/DEW] investigated the heat capacity from 68 to 298 K, King [57KIN] performed low-temperature meas-urements between 54 and 296 K, and DuBose and Naugle [79DUB/NAU] determined the specific heat of NiO from 14 to 280 K. Whereas bulk-like NiO powder samples were employed in all these investigations, White [74WHI] used single crystals for the determination of ,mpCο from 3.2 to 18.75 K. Furthermore, Watanabe [93WAT] measured


the heat capacity of single crystals of NiO over the temperature range from 135.7 to 477.8 K. The experimental data of all studies mentioned above coincide remarkably, as can be seen in Figure V-8. The heat capacity values between 0 and 40 K are perfectly described by a Debye – T 3 law. A fifth order polynomial fit to the experimental data in the range 40 – 477.8 K results in a ,mpCο function which can be easily integrated. The smoothed ,mpCο data as well as calculated entropy values are listed in Table V-5. The standard entropy of nickel oxide at 298.15 K, mS ο (NiO), is equal to 38.4 J·K–1·mol–1. This result is somewhat higher than the value (38.0 ± 0.2) J·K–1·mol–1 obtained by King [57KIN] which was adopted by Kelley and King [61KEL/KIN], Kellogg [69KEL], the NBS tables [82WAG/EVA], Knacke et al. [91KNA/KUB] and Robie and Hemingway [95ROB/HEM]. As our value for mS ο (NiO) is based on a simultaneous evaluation of five independent experimental studies of comparable accuracy, the selected standard entropy of NiO amounts to:

mS ο (NiO, cr, 298.15 K) = (38.4 ± 0.4) J·K–1·mol–1.

The uncertainty has been chosen, such that both values (the result of the pre-sent evaluation and the value of King [57KIN]) fall within the error limits.

Table V-5: Smoothed ,mpCο data and calculated values for the entropy of NiO.

T (K)

,mpCο (J·K–1·mol–1)

mS ο (NiO) (J·K–1·mol–1)

T (K)

,mpCο (J·K–1·mol–1)

mS ο (NiO) (J·K–1·mol–1)

10 0.02 0.006 160 27.09 15.865 20 0.13 0.047 170 28.93 17.563 30 0.45 0.159 180 30.68 19.266 40 0.96 0.376 190 32.32 20.969 50 2.98 0.800 200 33.85 22.666 60 5.13 1.536 210 35.29 24.353 70 7.37 2.494 220 36.63 26.026 80 9.66 3.627 230 37.87 27.682 90 11.97 4.898 240 39.03 29.318 100 14.28 6.279 250 40.10 30.933 110 16.57 7.748 260 41.11 32.526 120 18.82 9.286 270 42.04 34.095 130 21.00 10.880 280 42.93 35.640 140 23.12 12.514 290 43.76 37.161 150 25.15 14.179 298.15 44.41 38.383

300 44.55 38.658


Figure V-8: Experimental results of heat capacity measurements on nickel oxide from 3.2 to 1809.7 K. ▲ [40SEL/DEW], ■ [74WHI], ○ [57KIN], × [79DUB/NAU], [93WAT], + [90HEM], □ [58KIN/CHR]; solid line: polynomial fit (present assess-ment).

The heat content of nickel oxide was investigated at temperatures above 298.15 K by Tomlinson et al. [55TOM/DOM] as well as King and Christensen [58KIN/CHR]. The experimental data of these contributions are in close agreement. King and Christensen interpreted their results in terms of two phase transitions occur-ring at 525 K and 565 K, respectively. Hemingway [90HEM] reinvestigated the high-temperature heat capacity of bunsenite (NiO). Both the experimental data of [58KIN/CHR], [90HEM] and the low-temperature heat capacities are illustrated in Figure V-8. The data of Hemingway clearly demonstrate that in the high-temperature region only the antiferromagnetic-paramagnetic phase transition can be observed. The corresponding Néel temperature (the antiferromagnetic-paramagnetic transition tem-perature) was found to be equal to 519 K and the following polynomial fits were given [90HEM]:

519,m 298.15[ ]pCο (NiO, cr) = (4110.720 – 5.302412 (T/K) + 3.52061×10–3 (T/K)2 –

53039.297 (T/K)–0.5 + 2.43067×107 (T/K)–2) J·K–1·mol–1 (V.33)

0 200 400 600 800 1000 1200 1400 1600 1800

0

10

20

30

40

50

60

70

80

90

Co p,m /

J·mol

−1· K

−1

T / K

C° p,m

/ J·K

–1·m

ol–1


valid over the range from 298.15 to 519 K, and 1800

,m 519[ ]pCο (NiO, cr) = (– 8.776 + 4.2232×10–2 (T/K) – 7.5267×10–6 (T/K)2 + 787.25 (T/K)–0.5 + 3.6067×106 (T/K)–2) J·K–1·mol–1 (V.34)

valid from 519 to 1800 K. Equation (V.33) has to be modified somewhat in order to obtain a heat capacity function which is consistent with the value at 298.15 K listed in Table V-5:

519,m 298.15[ ]pCο (NiO, cr) = (4110.64 – 5.302412 (T/K) + 3.52061×10–3 (T/K)2 –

53039.297 (T/K)–0.5 + 2.43067×107 (T/K)–2) J·K–1·mol–1 (V.35)

valid over the range from 298.15 to 519 K. Equations (V.34) and (V.35) describe the temperature dependence of the heat capacity above 298.15 K satisfactorily. Moreover, Equation (V.34) is consistent with the data of King and Christensen [58KIN/CHR] above 846 K, see the solid line in Figure V-8. Therefore, the Reviewers have selected the heat capacity function Equations (V.35) and (V.34), where ,mpCο at 298.15 K amounts to:

,mpCο (NiO, cr, 298.15 K) = (44.4 ± 0.1) J·K–1·mol–1.

The uncertainty corresponds to the deviation between Equations (V.33) and (V.35) at 298.15 K.

Boyle et al. [54BOY/KIN] determined the enthalpy of formation of nickel ox-ide directly by means of combustion calorimetry, f mH ο∆ (NiO) = – (239.7 ± 0.4) kJ·mol−1. The application of a solid-state electrochemical cell to determine the Gibbs energy of formation of nickel oxide at high temperatures was originally introduced by Kiukkola and Wagner [57KIU/WAG]. By employing electrochemical cells of the type:

Reference electrode | YSZ or CSZ | NiO, Ni, Pt (V.36)

the temperature dependence of the Gibbs energy function of NiO was studied by a num-ber of authors, where yttria- or calcia-stabilised zirconia served as the solid electrolyte and the reference electrode consisted of Pt, O2(air) or Fe, FeOx (wuestite) or Cu, Cu2O [63RAP], [65TRE/SCH], [65STE/ALC], [67RIZ/BID], [68CHA/FLE], [69MOR/SAT], [70HUE/SAT], [75MOS/FIT2], [76BER2], [78IWA/FUJ], [79KEM/KAT], [80TRA/BRU], [84COM/PRA], [84RAY/PET], [86HOL/NEI], [93NEI/POW], [94KAL/FRA]. Some of these references are discussed in Appendix A. Basically, the results of all these experimental studies agree within the limits of uncertainty. Espe-cially, Charette and Flengas [68CHA/FLE], Berglund [76BER2], Comert and Pratt [84COM/PRA], Holmes et al. [86HOL/NEI] and O'Neill and Pownceby [93NEI/POW] performed highly reliable and precise electrochemical measurements of the Gibbs en-ergy of formation as a function of temperature, as shown in Figure V-9.

Equation (V.37) allows the calculation of the Gibbs energy of formation of NiO as a function of temperature:


f m f m f m

f ,mf ,m298.15 298.15

( ) = (298.15K) (298.15K) +

d dT T p

p

G T H T S

CC T T T

T

ο ο∆ ∆ − ∆

∆∆ −∫ ∫

(V.37)

The auxiliary data for O2(g) at 298.15 K can be found in the NEA TDB auxil-iary data set, heat capacities at higher temperatures were taken from CODATA [89COX/WAG] and from Robie and Hemingway [95ROB/HEM] in tabular and ana-lytical form, respectively. When the selected data for Ni(cr), the calorimetric value for the standard enthalpy of formation of NiO, according to [54BOY/KIN], the heat capac-ity function and the standard entropy of NiO (selected above) are used, the predicted temperature dependence of the Gibbs energy of formation agrees remarkably well with experimental data obtained from various high-temperature electrochemical measure-ments, see Figure V-9. Thus, we select for the standard enthalpy of formation of nickel oxide at 298.15 K:

f mH ο∆ (NiO, cr, 298.15 K) = – (239.7 ± 0.4) kJ·mol–1

and the corresponding derived value for the Gibbs energy of formation at 298.15 K is:

f mGο∆ (NiO, cr, 298.15 K) = – (211.66 ± 0.42) kJ·mol–1.

Figure V-9: Comparison between the temperature dependence of the Gibbs energy of formation for NiO obtained from electrochemical as well as chemical reduc-tion/oxidation equilibrium measurements and the prediction based on the present selec-tion of thermodynamic properties for NiO.

600 800 1000 1200 1400 1600 1800-190

-180

-170

-160

-150

-140

-130

-120

-110

-100

-90

-80

[26PEA/COO] [33WAT] [38BOG] [42FRI/WEI] [73RAU/GUE] [76BER2] [86HOL/ONE] [84COM/PRA] [68CHA/FLE] present review

∆ fG° m

(NiO

) / k

J·mol

−1

T / K

∆ fG

° m (N

iO) /

kJ·m

ol–1


In addition, equilibrium constants for the reduction of NiO with H2 and CO are reported in [26PEA/COO], [73RAU/GUE] and [33WAT], [38BOG], [42FRI/WEI], [64ALC/BEL], [67ANT/WAR], respectively. The equilibrium constants for the reac-tions:

NiO(cr) + H2(g) Ni(cr) + H2O(g)

NiO(cr) + CO(g) Ni(cr) + CO2(g)

can be transformed into the Gibbs energy of formation of NiO by using the pertinent auxiliary quantities for H2(g), H2O(g), CO(g) and CO2(g) listed in the CODATA tables [89COX/WAG] and Robie and Hemingway [95ROB/HEM]. The equilibrium data for the reduction/oxidation reactions given by [26PEA/COO], [33WAT], [42FRI/WEI], [64ALC/BEL], [67ANT/WAR], [73RAU/GUE] are in perfect agreement with calcu-lated values for the Gibbs energy of formation using the thermodynamic data selected in the present compilation, as can be deduced from Figure V-9 (the results of [64ALC/BEL], [67ANT/WAR] are discussed in Appendix A). However, the data of Bogatzki [38BOG] deviate somewhat from the calculated line at lower temperatures, probably because thermodynamic equilibrium was not attained below 1000 K.

The solid – gas equilibria:

NiO(cr) + H2(g) Ni(cr) + H2O(g)


were likewise studied by [21WOH/BAL], [32SKA/DAB] and [36KAP/SIL], respec-tively, who obtained equilibrium constants deviating by approximately one order of magnitude from those predicted according to the thermodynamic model of the present assessment. The presumably erroneous results of these early works may be caused by (i) insufficient equilibration of the solid material with the gas phase, (ii) deposition of car-bon by the Boudouard reaction, or (iii) inaccurate chemical analysis of the composition of the gas phase.

Also, the relative mass change of NiO was monitored as a function of the oxy-gen partial pressure at 1473 K by means of thermogravimetry [89KIT]. Hence, the de-composition of NiO into O2(g) and Ni was found to occur at log10(p(O2)/bar) = – 7.74 which concurs well with log10(p(O2)/bar) = – 7.84 calculated from the thermodynamic data of the present compilation.

V.3.2.1.1 Solubility measurements

Due to the kinetically inert nature of nickel oxide with respect to its dissolution in aque-ous media the solubility of NiO has been studied only at elevated temperatures so far [80TRE/LEB2], [89ZIE/JON]. These studies are not suitable for the calculation of any thermodynamic properties of NiO because of the high uncertainty of the measured solu-bilities compared to the high-temperature emf data and the low-temperature heat capac-


ity data discussed above. Moreover, the evaluation of the solubility experiments per-formed at hydrothermal conditions may cause an additional uncertainty for the solubil-ity constant of NiO owing to the lack of heat capacity functions for the ionic species including the hydroxo complexes. Thus, the calculated value for the solubility constant of NiO at 298.15 K, according to the reaction:

NiO(cr) + 2H+ Ni2+ + H2O(l),

derived from the thermodynamic data accepted in the present assessment, is :

10 ,0*log sK ο (NiO, cr, 298.15 K) = (12.48 ± 0.15).

V.3.2.2 Ni(II) hydroxides, Ni(OH)2(cr)

Recently the thermodynamic property values of Ni(II) hydroxide have been reviewed by Archer [99ARC] and critically evaluated by Plyasunova et al. [98PLY/ZHA]. The data of Ni(OH)2(cr) reported are rather uncertain and, as this solid phase is of considerable importance for the deposition and remobilisation of nickel, they have been re-evaluated on the basis of recent solubility measurements [2002GAM/WAL].

V.3.2.2.1 Crystallography and mineralogy of nickel hydroxide

V.3.2.2.1.1 β-Ni(OH)2

In a comprehensive review on bivalent metal hydroxides it was pointed out that the unit cell dimensions of synthetic β-Ni(OH)2 were determined by Lotmar and Feitknecht for the first time [36LOT/FEI], [77OSW/ASP]. Like most M(OH)2 hydroxides, the crystal structure is of the CdI2- or brucite- type, trigonal, space group: P 3 m1, with unit cell dimensions a = 3.126 Å, c = 4.605 Å, Z = 1 (according to JCPDS-ICDD card No. 14−117).

The natural occurrence of Ni(II) hydroxide in the Lord Brassey mine, Heazle-wood, Tasmania, was first reported by Williams [60WIL2]. Almost pure Ni(OH)2 was found in the Vermion region, northern Greece, whereas Mg-bearing (Ni, Mg)(OH)2 oc-curs in Unst, Shetland, Scotland [81MAR/ECO], [82LIV/BIS]. The X-ray powder dif-fraction patterns of the former mineral match closely with synthetic β-Ni(OH)2. It was named theophrastite after Theophrastos, the first Greek mineralogist, 373/372–288/287 B.C. [81MAR/ECO] (approved by IMA 1980). Theophrastite from the Vermion region is emerald green, translucent with vitreous luster and has a pale green streak. The Mohs hardness is 3.5. The specific gravity is ρ(obs.) = 4.0 and ρ(calc.) = 3.95 g·cm–3. It is a gangue mineral in ore consisting of magnetite (Fe3O4), chromite (FeCr2O4) and Ni-sulphides as minor component. Theophrastite is formed from Ni-bearing solutions be-tween 80 ≤ t(°C) ≤ 115 in alkaline moderately oxidising media [83ECO/MAR].


V.3.2.2.1.2 α-Ni(OH)2

Other varieties of crystallised divalent nickel hydroxide, α-Ni(OH)2 and α∗-Ni(OH)2, differ from the thermodynamically stable β-form by the presence of a layer of water molecules in the van der Waals gap. Bode et al. [66BOD/DEH] proposed an α−3Ni(OH)2·2H2O unit cell (a = 3.08 Å, c = 8.09 Å), whereas Braconnier et al. [84BRA/DEL]1 synthesised a compound with a stoichiometric composition of α∗−Ni(OH)2·0.75H2O and the following unit cell dimensions: a = 3.08 Å, c = 23.41 Å. In any case the intercalated crystal water of α-Ni(OH)2 causes an extension of the struc-ture along the crystallographic c-axis, whereas the other dimension remains similar to that of theophrastite. The α-, α∗- and β-forms of nickel hydroxide have the same trigo-nal structure. The β-form consists of well-ordered brucite type layers with a separation of 0.46 nm, whereas the hydrated α-form shows larger layer distances and turbostratic character as a result of intercalated water and distorted stacking, respectively. The α−form may also intercalate different anions so that the interlayer separation depends on the type as well as the extent of these ions and varies between 0.8 and 0.9 nm. By electron microscopy the very small crystallites appear as formless aggregates [84BRA/DEL]. Although it plays an important role in the charge/discharge cycle of nickel batteries no thermodynamic data can be definitively assigned to it (see Sections V.2.2 and V.3.2.3 about Ni(III, IV) ions and hydroxides). The natural occurrence of α−Ni(OH)2 has never been reported as it is probably too unstable to persist under ambi-ent conditions.

V.3.2.2.2 Heat capacity and entropy of Ni(OH)2(cr)

Sorai et al. [69SOR/KOS] measured the heat capacity of Ni(OH)2(cr) at low tempera-tures. These data have been used to determine the standard entropy, mS ο (298.15 K), which has been accepted for the selected thermodynamic data of nickel compounds [76MAH/PAN]. Sorai et al. [69SOR/KOS] investigated three samples, (iii), (ii), and (i), of Ni(OH)2(cr) with different particle sizes, and reported ,mpCο values up to 104 K and 300 K for the coarsest (iii) and the finer crystals (ii, i), respectively. As shown in Figure V-10 the heat capacity and consequently the entropy turned out to depend on the parti-cle size and thus on the specific surface area s.

When the excess surface heat capacity estimated by Sorai et al. between 80 and 130 K was assumed to remain constant up to 300 K values of ,mpCο (Ni(OH)2, cr, 298.15 K) = 81.7 J·K–1·mol–1 and mS ο (Ni(OH)2, cr, 298.15 K) = 79.4 J·K–1·mol–1 have been extrapolated for the coarsest material. When, however, the entropy increases line-arly with s, extrapolation to s = 6 m2·g–1 (according to Sorai et al. the specific surface area of the coarsest sample) results in mS ο (Ni(OH)2, cr, 298.15 K) = 80.0 J·K–1·mol–1, in

1 The latter phase has been called α∗ to emphasise that the composition and the intersheet distance are close

to those of the turbostratic α-Ni(OH)2, whereas the texture and structure are different.


perfect agreement with the value selected by Mah and Pankratz [76MAH/PAN]. This value has been retained, but the uncertainty was increased from ± 0.4 to ± 0.8:

mS ο (Ni(OH)2, β, 298.15 K) = (80.0 ± 0.8) J·K–1·mol–1. (V.38)

These increased error limits probably account also for the uncertainty intro-duced by Enoki and Tsujikawa’s slightly differing results for bulky Ni(OH)2 between 4.2 to 35 K [78ENO/TSU], see Figure V-10. In addition the interpolated ,mpCο value for sample (ii) at 298.15 K has been preferred to the extrapolated one of sample (iii). An uncertainty has been selected, so that both values fall within the error limits:

,mpCο (Ni(OH)2, β, 298.15 K) = (82.01 ± 0.30) J·K–1·mol–1.

Figure V-10: The heat capacity of Ni(OH)2(cr) as a function of temperature (▲ s(Ni(OH)2(iii)) = 6 m2·g–1 [69SOR/KOS]; ☼ s(Ni(OH)2(ii)) = 64 to 75 m2·g–1 estimated by the reviewer; × s(Ni(OH)2(i)) = 320 m2·g–1 [69SOR/KOS]; thick line: bulky Ni(OH)2 [78ENO/TSU]).

0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0 4.5 5.0

0

2

4

6

8

10

12

14

16

18

20

22

24

C p,m /

J·mol

−1 ·K

−1

ln (T / K)

C° p,m

/ J·K

–1·m

ol–1


V.3.2.2.3 Thermodynamic analysis of solubility data of Ni(OH)2(s)

For a methodological review of the experimental determination of solubilities of spar-ingly soluble ionic solids in aqueous media see [2003GAM/KON].

All auxiliary variables discussed in Section V.2.1 were used in the thermody-namic analysis of solubility data.

The determination of the solubility product of Ni(OH)2(cr) is in principle a straightforward method to evaluate the standard Gibbs energy of formation,

f mGο∆ (Ni(OH)2, cr). It presumes, however, that the solubility product (defined by Equa-tion (V.39)) or any appropriate solubility constant (as for example defined by Equation (V.40)) of the respective phase has indeed been determined.

Solubility product: β-Ni(OH)2 Ni2+ + 2OH– ,0sK (V.39)

Dissolution: β-Ni(OH)2 + 2H+ Ni2+ + 2H2O(l) ,0*

sK (V.40)

Standard thermodynamic quantities from solubility measurements on hydrox-ides can be obtained using Equations (V.41) – (V.45).

sol mGο∆ = sol mH ο∆ – T sol mS ο∆ = – R T ln ,0*

sK ο (V.41)

sol mH ο∆ = R T 2 (∂ln ,0*

sK ο /∂T) (V.42)

f mGο∆ (Ni(OH)2, β) = f mGο∆ (Ni2+) + 2 f mGο∆ (H2O, l) – sol mGο∆ (V.43)

f mH ο∆ (Ni(OH)2, β) = f mH ο∆ (Ni2+) + 2 f mH ο∆ (H2O, l) – sol mH ο∆ (V.44)

mS ο (Ni2+) = mS ο (Ni(OH)2, β) – 2 mS ο (H2O, l) + sol mS ο∆ (V.45)

As discussed in Section V.3.2.2.2, the entropy of nickel hydroxide was calori-metrically measured with sufficient accuracy [69SOR/KOS], [78ENO/TSU] that, in fact, the value of the entropy of dissolution contributes to the knowledge of the partial molar entropy mS ο (Ni2+), see Equation (V.45).

Actually most solubility data of Ni(OH)2(cr), reported so far and listed in Table V-6, suffer from an uncertainty in the physical state of the solid investigated [18ALM], [25BRI], [25WIJ], [38OKA], [43NAS], [49GAY/GAR], [50AKS/FIA], [53SCH/POL], [54DOB], [54FEI/HAR], [56CUT/KSA], [56JEN/PRA], [67BLA], [69MAK/SPI], [73KAW/OTS], [73NOV/COS], [80CHI/SAB], [96POU/DRE]. This is particularly sur-prising for Feitknecht and Hartmann [54FEI/HAR], because Lotmar and Feitknecht [36LOT/FEI] were the first to determine the structure of β-Ni(OH)2, and the concept that only well defined solid phases result in reliable solubility constants of metal oxides and hydroxides was developed by Schindler in the same laboratory [63SCH], [63FEI/SCH]. Moreover, the comprehensive review on Ni(II) hydroxide mentioned in Section V.3.2.2.1, is also based to a considerable extent on the work performed by Feitknecht and his co-workers [77OSW/ASP]. There it has been clearly pointed out that, apart from the well defined β-Ni(OH)2, a number of basic salts of changing com-


position exist. When nickel hydroxide is precipitated from aqueous NiCl2, Ni(NO3)2, or NiSO4 with NaOH or KOH solutions it is always contaminated with basic salts. The solubility of the latter varies depending on the anion and the molar ratio OH/Ni. This means that solubility studies on poorly defined nickel hydroxide or the respective basic salts are useless as an experimental basis to derive accurate thermodynamic functions of nickel hydroxide. They may, however, serve to find out relevant information concerning removal of Ni2+ from radioactive effluents.

Sometimes solubility products have to be rejected, because they were calcu-lated on the basis of quite inadequate experimental results, see Appendix A [18ALM], [53SCH/POL], [67BLA], [80CHI/SAB].

One of the most careful solubility studies reported so far was carried out by Mattigod et al. [97MAT/RAI], but according to their preparative method the respective results refer to a microcrystalline β-Ni(OH)2 that probably contained chloride. It is noteworthy that Poulson and Drever’s [96POU/DRE] study of a commercially available nickel hydroxide (Johnson Matthey Chemical Co.) resulted in a solubility constant which after being extrapolated to the same temperature ( 10 ,0

*log sK ο = 12.08) agrees within 0.18 log10-units with that of Mattigod et al. [97MAT/RAI].

A re-evaluation of f mGο∆ and f mH ο∆ of β-Ni(OH)2 has been based on the re-sults of [2002WAL/GAT] and [2002GAM/WAL].

In Figure V-11, the solubility constant 10 ,0*log sK , is plotted as a function of

temperature. It is quite obvious that after having rejected the most unreliable results [18ALM], [25WIJ], [67BLA], [80CHI/SAB], the maximum error limits according to Plyasunova et al. [98PLY/ZHA] turned out to be too pessimistic—albeit the usual ex-perimental error is exceeded by an order of magnitude. The results of several authors scatter closely around the curve adjusted to 10 ,0log sK ο of Mattigod et al. which seems to describe the solubility of finely divided β-Ni(OH)2, [25BRI], [54DOB], [56JEN/PRA], [96POU/DRE], [97MAT/RAI], whereas still higher solubilities presumably refer to amorphous Ni(OH)2 or even basic salts. Measurements carried out with carefully pre-pared coarse theophrastite led to lower solubilities, closer error limits, and a compara-tively reliable prediction of thermodynamic quantities [2002GAM/WAL].


Table V-6: Literature data(a) for 10 ,0*log sK ο (Ni(OH)2).

Medium I (mol·kg–1) T (K) 10 ,0*log sK ο (a) Method Remarks Reference

NiCl2, NaCl 0.030 – 0.065 291.15 12.20 H, titr. Ni(OH)2–xClx(s)? [25BRI]

Ni(NO3)2, NH4NO3

0.3 – 0.4 298.15 13.36 [OH–]→ [NH3]/[ +

4NH ]incipient ppt. [25WIJ]

Ni(NO3)2 dilute 298.15 13.50 pH, titr. incipient ppt. [38OKA]

NiCl2, KCl 0.002 – 2.05 298.15 12.92 H, titr. Ni(OH)2–xClx(s)? [43NAS]

NiCl2 0.038 – 0.15 298.15 10.78 pH, [Ni(II)] saturated? [49GAY/GAR]

NiCl2 0.038 – 0.15 308.15 10.78 pH, [Ni(II)] sat. at 35°C ? [49GAY/GAR]

NiSO4 0.006 – 5.37 291.15 13.40 H, QH(b), Sb(c) Ni(OH)2 pptd. [50AKS/FIA]

NiSO4 0.030 – 0.065 348.15 9.48 pH, titr. Figs. 3 and 5.1 [54DOB]

NiCl2 0 corr 298.15 13.30 pH, ? Ni(OH)2, fresh [54FEI/HAR]

NiCl2 0 corr 298.15 10.80 pH, ? Ni(OH)2, aged [54FEI/HAR]

NaClO4 1.05 293.15 12.56 pH, [Ni(II)] incipient ppt. [56CUT/KSA]

NiCl2 0.028 – 0.28 298.15 12.12 pH, [Ni(II)] Ni(OH)2 pptd. [56JEN/PRA]

CH3COONa 0.10 – 0.50 302.15 11.75 pH, [Ni(II)] Ni(OH)2 pptd. [56JEN/PRA]

NiCl2 0.001 – 4.7 M 298.15 12.46 pH, [Ni(II)] incipient ppt. [69MAK/SPI]

LiClO4 3.48 298.15 12.99 pH, Zmax(d) incipient ppt. [73KAW/OTS]

NaCl 0.56 298.15 12.49 pH, [Ni(II)] slow ppt. [73NOV/COS]

NaCl 0.56 298.15 10.89 pH, [Ni(II)] fast ppt. [73NOV/COS]

NaClO4 1.05 298.15 10.38 pH, [Ni(II)] fast ppt. [73NOV/COS]

NaNO3 0.02 – 0.2 295.15 12.24 pH, [Ni(II)] Johnson Matthey [96POU/DRE]

NaClO4 0.01 298.15 11.90 pH, [Ni(II)] X-ray,β- Ni(OH)2 [97MAT/RAI]

NaClO4 0.5 – 3.0, SIT 298.15 11.02 pH, [Ni(II)] X-ray, β- Ni(OH)2 [2002GAM/WAL]

(a) Whenever possible the entries of column 4 have been recalculated from the experimental data. (b) quinhydrone electrode. (c) antimony electrode. (d) maximum number of H+ set free per Ni.


Figure V-11: Solubility constant of Ni(OH)2 as a function of temperature. Thick solid curve: optimised from experimental data of Gamsjäger et al. [2002GAM/WAL]; thin solid curves: corresponding error limits; thick dash-dotted curve: calculated according to [98PLY/ZHA]; thin dash-dotted curves: corresponding maximum error limits; thick dash curve: calculated with f mH ο∆ (Ni(OH)2, cr) adjusted to coincide with the result of Mattigod et al. [97MAT/RAI]; ▲ [25BRI], ⊗ [38OKA], ∆ [43NAS], ∇ [49GAY/GAR], ■ [50AKS/FIA], ● [54DOB], × [54FEI/HAR] (aged), ♦ [54FEI/HAR] (freshly pre-pared), [56CUT/KSA], ∗ [56JEN/PRA], [56JEN/PRA], ○ [69MAK/SPI], ▼ [73KAW/OTS], + [73NOV/COS] (high value: slow, lower values: fast precipitation), ⊕ [96POU/DRE], □ [97MAT/RAI], [2002GAM/WAL].

In Table V-7 the thermodynamic properties of Ni(OH)2 obtained in [2002GAM/WAL] and selected by Plyasunova et al. are compared. The latter authors overlooked a direct calorimetric determination of mS ο of Ni(OH)2(cr) [69SOR/KOS], [76MAH/PAN] which has been discussed in Section V.3.2.2.2 and Appendix A ([69SOR/KOS], [78ENO/TSU]). For the optimisation algorithm applied, the calorimet-rically measured mS ο (β-Ni(OH)2) = 80.0 J·K–1·mol–1 (Equation (V.38)) and mS ο (Ni2+) = – 131.8 J·K–1·mol–1 (Equation (V.15)) were kept fixed. Based on solubility data, the

270 280 290 300 310 320 330 340 350 3606

7

8

9

10

11

12

13

14

log 10

* Ko s,0

T / K

log 1

0* K ° s ,

0

Ni(OH)2(s) + 2 H+ Ni2+ + 2 H2O(l)


precision and reliability of f mGο∆ and f mH ο∆ (β-Ni(OH)2) can be improved considera-bly, leading to the selections:

f mGο∆ (Ni(OH)2, β, 298.15 K) = – (457.1 ± 1.4) kJ·mol–1

f mH ο∆ (Ni(OH)2, β, 298.15 K) = – (542.3 ± 1.5) kJ·mol–1.

Table V-7: Thermodynamic properties of Ni(OH)2(cr).

Compound

10 ,0*log sK ο (V.40) sol mH ο∆

(kJ·mol–1) mS ο

(J·K–1·mol–1) Reference

Ni(OH)2(cr)(a) (10.5 ± 1.3) – (78.7 ± 10.2) (73 ± 10) [98PLY/ZHA] β-Ni(OH)2

(b) (11.02 ± 0.20)# – (84.36 ± 1.24)# [2002GAM/WAL] β-Ni(OH)2

(b) (80.0 ± 0.8) [69SOR/KOS], [76MAH/PAN]

(a) Uncertainties are taken from Plyasunova et al. [98PLY/ZHA]. (b) Uncertainties based on ∆ 10 ,0

*log sK = ± 0.2 (2σ), estimated from the experiments. # From selections in this review.

According to the present thermodynamic model the equilibrium temperature for:

β-Ni(OH)2 NiO(cr) + H2O(l)

amounts to T/K= (503 ± 31). The uncertainty has been obtained from error propagation of the estimated uncertainties for the corresponding enthalpy and entropy of reaction (T = 503 K): r mH ο∆ = (26.40 ± 1.55) kJ·mol–1; r mS ο∆ = (52.50 ± 0.89) J·K–1·mol–1. The lower value for the decomposition temperature of Ni(OH)2 found by [89ZIE/JON], 468 K, may be caused by a scarcely crystallised and, therefore, less stable Ni(OH)2.

V.3.2.3 Ni(III, IV) hydroxides

Nickel hydroxides of oxidation state two and higher have been used as the active mate-rial in the positive electrodes of several alkaline batteries for more than a hundred years [99MCB]. Bode et al. [66BOD/DEH] established the scheme depicted below.

charge → ← discharge

β-Ni(OH)2 β-NiOOH dehydratation ↑ ↓ overcharge

α-Ni(OH)2 γ-NiOOH

The charge and discharge cycles of nickel batteries involve two different pairs of solid phases. Oxidation of β-Ni(OH)2 produces β-NiOOH, oxidation of α-Ni(OH)2 produces γ-NiOOH. The end-products of these cycles are interconnected by dehydration and overcharge. For the crystallographic properties of Ni(II) hydroxides and Ni(III, IV) hydroxides see Section V.3.2.2.1 and Sections V.3.2.3.1, V.3.2.3.2, respectively.


V.3.2.3.1 β-NiOOH

Glemser and Einerhand [50GLE/EIN] synthesised β-NiOOH and determined the di-mensions of its hexagonal unit cell to be a0 = 2.81 Å, c0 = 4.84 Å. This Ni(III) oxide hydroxide is the primary oxidation product of electrodes containing β-Ni(OH)2.

Conway and Gileadi [62CON/GIL] measured the reversible potential of the β−NiOOH | β−Ni(OH)2 electrode as a function of the degree of oxidation in 1 molal KOH. The mercury | mercury oxide system, Hg | HgO, was used as reference electrode. Similar measurements were carried out by Barnard et al. [81BAR/RAN]. In this review the experimental data according to both teams of authors have been based on the stan-dard hydrogen electrode (SHE) and extrapolated to I = 0 with the SIT model [97GRE/PLY2]. In Figure V-12 the reversible potential of Reaction (V.46) has been plotted versus the oxidation state of nickel which is numerically equal to x (NiIII).

β-NiOOH + ½ H2(g) β-Ni(OH)2 (V.46)

Figure V-12: β-Ni(OH)2 | β-NiOOH system. Variation of reversible potential with the oxidation state of Ni at 298.15 K and I = 0. Solid curve: regular model; experimental data: ■ [62CON/GIL]; ○ [81BAR/RAN].

When nickel (II, III) hydroxides form regular solid solutions in the whole range of compositions between β-Ni(OH)2 and β-NiOOH the experimental data can be

0.0 0.2 0.4 0.6 0.8 1.0

1.28

1.30

1.32

1.34

1.36

1.38

1.40

1.42

1.44

E / V

x (Ni(III))


evaluated with Equation (V.47), where n = 1, A ( = exG∆ /[x·(1 – x)]) is the excess in-teraction parameter and x the mole fraction of Ni(III):

E = E° – R (1 )ln (1 2 )F FT x A x

n x n−

⋅ + ⋅ − (V.47)

A non-linear least squares analysis of these experimental data led to A = (3.4 ± 1.0) kJ·mol–1, E° = (1.361 ± 0.006) V and r mGο∆ (298.15 K) = – (131.3 ± 0.6) kJ·mol–1. Now the standard Gibbs energy of formation can be derived:

f mGο∆ (NiOOH, β, 298.15 K) = – (325.4 ± 6.0) kJ·mol–1 (V.48)

It must be emphasised, however, that potentiometric measurements are ther-modynamically comparable only when exactly the same phases are involved in the re-spective redox reactions. As shown in Figure V-11 the solubility constant and conse-quently the Gibbs energy of formation of β-Ni(OH)2 presumably depends on the particle size (compare thick solid and thick dash curves which differ by ≈ 0.8 log10K-units). Thus the total uncertainty in Equation (V.48) is composed of an experimental random error 2σ ≈ 1.5 kJ·mol–1 and a systematic error due to the solid phase used of ≈ 4.5 kJ·mol–1.

Jain et al. [98JAI/ELM] determined the standard potential of:

NiOOH(s) + H2O(l) + e– Ni(OH)2(s) + OH– (V.49)

against the silver | silver chloride electrode. When their result is recalculated versus SHE the following standard potential is obtained, E°(298.15 K) = 1.340 V. The Gibbs energy of NiOOH formation derived from this result differs by only 2 kJ·mol–1 from the Equation (V.48) value. It might be that not quite the same pair of phases have deter-mined the potentials measured by Barnard et al. [81BAR/RAN] and Jain et al. [98JAI/ELM] respectively.

Jain et al. measured the following value (∂ r∆ E°(V.50)/∂T)p = – 0.84 mV·K–1 for the temperature coefficient of the galvanic cell (V.50).

β-NiOOH + H2O(l) + Ag(cr) + Cl– β-Ni(OH)2 + AgCl(cr) + OH– (V.50)

This value is regarded to be in reasonable agreement with an earlier value of − 0.5 mV·K–1. On the basis of these data mS ο (NiOOH, β, 298.15 K) = 77.3 or 44.5 J·K−1·mol–1 is obtained, when the other standard entropies were taken from CODATA [89COX/WAG] as adopted by the NEA TDB Project and Equation (V.38). With the so-called “Latimer” entropy contributions listed in Tables XIII, XIV and XVI of [93KUB/ALC] mS ο (NiOOH, β) can be estimated. It is either 55.8 J·K–1·mol–1, when the OH– contribution for a divalent cation is accepted, or 62.6 J·K–1·mol–1, when NiO+ is considered to be a monovalent cation. To sum up the situation, the mean of these ob-served and estimated values (60 ± 27) J·K−1·mol–1, which seems too rough an estimate to be recommended.


V.3.2.3.2 γ-NiOOH

γ-NiOOH is produced either by oxidation of α-Ni(OH)2 or on overcharge of β-Ni(OH)2. This nickel oxide hydroxide was first prepared by Glemser and Einerhand [50GLE/EIN]. It crystallises in a rhombohedral system with cell dimensions of a0 = 2.8 Å, c0 = 20.65 Å and has a layer structure with a spacing of 7.2 Å between layers. γ-NiOOH always contains small quantities of water and alkali metal ions between the layers, whereas β-NiOOH does not. When γ-NiOOH is prepared by the method of Glemser and Einerhand its chemical composition is NiOOH·0.51H2O.

The standard potential measurements of Reaction (V.51) are difficult to inter-pret thermodynamically.

γ-NiOOH + ½ H2(g) α-Ni(OH)2 (V.51)

Clearly the γ-NiOOH | α-Ni(OH)2 electrode plays an important role in the op-eration of alkaline nickel batteries and has been investigated extensively, but thermody-namic data of α-Ni(OH)2 still remain unknown. Consequently, emf data of the above electrode measured against a suitable reference electrode contain contributions of both the γ- as well as the α-phases and thus do not suffice to determine the standard Gibbs energy of γ-NiOOH formation. Moreover, the γ-phase can be charged up to a formal oxidation number of z(Ni) ≤ 3.6. Balej and Divisek re-evaluated four series of experi-mental data [70MAC], [70TYS/KSE], [80BAR/RAN], [81BAR/RAN] implying the following model [93BAL/DIV], [93BAL/DIV2], [97BAL/DIV].

Figure V-13: α-Ni(OH)2 | γ-NiOOH system. Variation of reversible potential with the oxidation state of Ni at 298.15 K and I = 0. Solid curve: regular model; dotted curve [93BAL/DIV2]; experimental data: ■ [70MAC], ○ [70TYS/KSE], ▼ [80BAR/RAN], ● [81BAR/RAN].

0.0 0.2 0.4 0.6 0.8

1.20

1.25

1.30

1.35

1.40

E / V

x (Ni(IV))


The so-called “α-Ni(OH)2 | γ-NiOOH” system forms regular solid solutions of Ni(II) and Ni(IV) components in the whole range of compositions between α-Ni(OH)2 and NiO2·xH2O(s), without any participation of Ni(III) oxide hydroxides. Thus, Reac-tion (V.51) has to be modified.

NiO2·xH2O(s) + H2(g) α-Ni(OH)2 + xH2O(l) (V.52)

On the basis of Reaction (V.52), an equation analogous to (V.47) can be fitted to the experimental data, where z = 2, A ( = exG∆ /[x·(1 – x)]) is the regular excess inter-action parameter and x the mole fraction of Ni(IV). A non-linear least squares analysis of these experimental data led to A = (2.31 ± 1.10) kJ·mol–1, E° = (1.322 ± 0.004) V and

r mGο∆ (298.15 K) = – (255.1 ± 0.8) kJ·mol–1. For the calculation of f mGο∆ (NiO2·xH2O, s, 298.15 K) = (– 201.6 – 237.14 x) kJ·mol–1 the arbitrary assumption f mGο∆ (Ni(OH)2, α) = f mGο∆ (Ni(OH)2, β) has to be made.

To summarise the situation, it can be said that unless the equilibrium potential measurements of the system NiO2·xH2O | α-Ni(OH)2 are confirmed independently from the structural and stoichiometric point of view, the respective results do not qualify for thermodynamic recommendations. They can be used, however, for the calculation of E-pH diagrams [97BAL/DIV].

V.4 Group 17 (halogen) compounds and complexes V.4.1 Nickel halide compounds V.4.1.1 Solid nickel fluoride, NiF2(cr)

Catalano and Stout [54STO/CAT], [55CAT/STO], [55STO/CAT] reported low-temperature (adiabatic calorimetry) heat capacity measurements (12 K to 300 K) for this solid, and calculated mS ο (NiF2, cr, 298.15 K) = 73.60 J·K–1·mol–1. A λ-transition at (73.22 ± 0.05) K was reported, and a similar temperature was associated with anoma-lous behaviour in the magnetic anisotropy of a single crystal of NiF2 [54MAT/STO]. No heat capacity measurements for temperatures below 12 K have been found in the litera-ture. Binford and Hebert [70BIN/HEB] reported drop calorimetry results for tempera-tures between 381.9 K and 1461.7 K, and calculated a “best fit” curve to match their results to the earlier heat capacity measurements. The ,mpCο results of Catalano and Stout appear to be slightly higher than would merge smoothly with the drop calorimetry results.

In the present review, a common function is fit to the combined results of the two studies above 250 K. The heat capacity results are weighted using an uncertainty of 0.2 J·K–1·mol–1 for measurements between 250 and 270 K, and 0.3 J·K–1·mol–1 at higher temperatures. For weighting purposes, the enthalpy differences are assigned an uncer-tainty of 0.2 kJ·mol–1. The results are constrained so that m m( ) (298.15K)H T H ο− is

V.4 Group 17 (halogen) compounds and complexes 120

zero at 298.15 K (a similar procedure appears to have been used by Binford and Her-bert, but without the constraint at 298.15 K), and the resulting equation is:

m m[ ( /K) (298.15K)]H T H ο− / J·mol–1 = 65.505(T /K) + 0.0076887(T /K)2 + 611603 (T /K)–1 –22265.1 (V.53)

and therefore,

1450K,m 250K[ ]pCο (NiF2, cr)/ J·K–1·mol–1 = 65.505 + 0.015377(T /K) – 611603(T/K)–2 (V.54)

Based on Equation (V.54), the heat capacity of NiF2(cr) at 298.15 K is 63.21 J·K–1·mol–1, whereas the value reported by Catalano and Stout [55CAT/STO] was 64.06 J·K–1·mol–1. Equation (V.54) is accepted in the present review, and as a consequence the reported entropy value at 298.15 K [55CAT/STO] should also be adjusted slightly (by 0.08 J·K–1·mol–1).

Therefore, the selected values are:

mS ο (NiF2, cr, 298.15 K) = (73.52 ± 0.40) J·K–1·mol–1 and ,mpCο (NiF2, cr, 298.15 K) = (63.21 ± 2.00) J·K–1·mol–1,

where the uncertainties are estimated values.

V.4.1.1.1 Enthalpy of formation of NiF2(cr)

Several groups have determined values for f mH ο∆ (NiF2, cr) from enthalpies of reaction at temperatures between 500 K and 1000 K. Jellinek and Rudat [28JEL/RUD] investi-gated the reaction:

H2(g) + NiF2(cr) Ni(cr) + 2HF(g)

by a transpiration method (573 K to 773 K). Domange [37DOM] studied:

H2O(g) + NiF2(cr) NiO(cr) + 2HF(g)

from 773 K to 973 K, and Hood and Woyski [51HOO/WOY] studied:

NiCl2(cr) + 2HF(g) NiF2(cr) + 2HCl(g)

from 477 K to 830 K. Extrapolations based on these higher temperature experiments led to values of f mH ο∆ (NiF2, cr, 298.15 K) between – 650 kJ·mol–1 and − 670 kJ·mol−1 [67RUD/DEV].

The enthalpy of formation of NiF2(cr), was measured directly by bomb calo-rimetry at 298.15 K, and reported as – (657.72 ± 1.67) kJ·mol–1 by Rudzitis et al. [67RUD/DEV]. This complicated experimental work appears to have been carried out with great care.

In addition, Chattopadhyay and coworkers [75CHA/KAR] determined Gibbs energies of reaction from measurements of galvanic potentials for cells of the type


Pt|M’,M’F2|CaF2|NiF2,Ni|Pt between 850 and 1000 K (M’ = Fe, Co). However, use of these data to determine a value for f mH ο∆ (NiF2, cr, 298.15 K) would require assessment of an enthalpy of formation value for CoF2(cr) or FeF2(cr), and this falls outside the scope of the present review. Lofgren and McIver [66LOF/MCI] carried out measure-ments for the cell Mg,MgF2|CaF2|NiF2|Ni, and f mH ο∆ (NiF2, cr, 298.15 K) = − (657.0 ± 1.4) kJ·mol–1 has been calculated from a third-law analysis of their results (Appendix A).

As discussed in Appendix A, the early work of Jellinek and Rudat cannot be used, but the measurements in the two later high-temperature studies [37DOM], [51HOO/WOY] lead to inconsistent values of f mH ο∆ (NiF2, cr, 298.15 K), − (666.13 ± 0.64) kJ·mol–1 and – (652.76 ± 0.93) kJ·mol–1, respectively. Neither of these is in particularly good agreement with the calorimetric value of Rudzitis et al. [67RUD/DEV].

Brunetti and Piacente [96BRU/PIA] reported an extensive set of torsion bal-ance and Knudsen cell measurements for the determination of the enthalpy of sublima-tion of NiF2(cr) between 950 and 1250 K. The measurements were in reasonable agree-ment with, but somewhat lower than the sparse earlier sets of Farber et al. [58FAR/MEY] and Ehlert et al. [64EHL/KEN]. Pankratz [84PAN], and Pankratz et al. [84PAN/STU] combined the results from the earlier studies with vapour pressure results for higher temperatures [65CAN], and the enthalpy of formation value for NiF2(cr) at 298.15 K from Rudzitis et al. [67RUD/DEV], to determine the enthalpy of formation of NiF2(g). They also used spectroscopic data to estimate the thermal functions for the gas. Brunetti and Piacente [96BRU/PIA] indicated that the thermodynamic data from Pank-ratz et al. are inconsistent with their measurements, and attributed the problem to errors in either the enthalpy of formation (at 298.15 K) or the thermal functions of the solid (and would show better agreement if the value of f mH ο∆ (NiF2, cr, 298.15 K) obtained from the work of Hood and Woyski [51HOO/WOY] were used). The discrepancy re-ported by Brunetti and Piacente [96BRU/PIA] remains unresolved. The present review-ers could find no apparent substantive error in the experiments used to derive the differ-ent values for NiF2(cr), though confirmation of the function for ,mpCο (T) at temperatures above 300 K would be useful. Vogt [2001VOG] has reported further spectroscopic in-formation for NiF2(g), but recalculation of the thermal functions is beyond the scope of this review.

Some of the discrepancies may suggest problems with auxiliary data (e.g., f mH ο∆ (NiCl2, cr, 298.15 K), f mH ο∆ (NiO, cr, 298.15 K)). Also, it is possible that there is

some undetected systematic error in the work of Domange [37DOM], as his value is markedly more negative than any of the others. In the present review, the value:

f mH ο∆ (NiF2, cr, 298.15 K) = – (657.3 ± 8.0) kJ·mol–1

is selected. It is the average of the values based on the potentiometric work of Lofgren and McIver [66LOF/MCI] and the calorimetric work of Rudzitis et al. [67RUD/DEV],


but the uncertainty has been greatly increased to reflect the inconsistency with other values.

Together with the selected value for mS ο (NiF2, cr, 298.15 K), this selection yields:

f mGο∆ (NiF2, cr, 298.15 K) = – (609.9 ± 8.0) kJ·mol–1.

V.4.1.2 Solid nickel chloride, NiCl2(cr)

Anhydrous nickel chloride, NiCl2, has a layered CdCl2 structure (rhombohedral space group R3m [63FER/BRA]) of cubic close packed chloride ions with 1/2 octahedral holes filled by nickel ions. At T > 52 K a transition of the antiferromagnetic to the paramagnetic state occurs [52BUS/GIA2].

V.4.1.2.1 Low-temperature heat capacity and entropy of NiCl2(cr)

Trapeznikova et al. measured the heat capacity of anhydrous Ni(II) chloride in the range of 14 to 130 K [36TRA/SCH]. Their results were not confirmed by Busey and Giauque who took great care to prepare high purity NiCl2 and measured its heat capacity between 15 and 300 K [52BUS/GIA2]. Kostryukova assessed the singularities of the magnetic energy spectrum of NiCl2 by heat capacity measurements at liquid helium temperatures [68KOS], [69KOS], [75KOS]. In Figure V-14 the smoothed values of ,m /pC T are plot-ted versus T for temperatures up to 30 K. Figure V-14 clearly shows that the usual ex-trapolation to T = 0 with Debye's T 3 law is not applicable up to 15 K.

Thus, in the present review the standard entropy of NiCl2 was re-evaluated by integration of the original and smoothed ,mpC values listed in [52BUS/GIA2] and [75KOS] respectively. Two different linear spline functions, valid for the temperature range below and above the Néel temperature (TN = 52.32 K) were fitted to these data, shown in Figure V-15. The sum of these integrals resulted in the selected value:

mS ο (NiCl2, cr, 298.15 K) = (98.14 ± 0.30) J·mol–1·K–1.

This value is slightly higher than the one given in [52BUS/GIA2], which may be due to the more realistic extrapolation to T = 0 on the basis of Kostryukova's data [75KOS] as well as to the influence of different fitting functions employed by [52BUS/GIA2] (unknown) and in the present review. The uncertainty selected is greater than the difference arising from using the more complete set of fitting functions.

The standard heat capacity derived by interpolation of the fitted ,mpC –functions is essentially equal to the value given by Busey and Giauque (17.13 cal·mol−1·K–1) who stated that near 300 K heat leaks may have introduced an error of several tenths of a percent into their ,mpC –values. Thus the following value as well as its uncertainty are selected:

,mpCο (NiCl2, cr, 298.15K) = (71.67 ± 0.30) J·mol–1·K–1. (V.55)


Figure V-14: Molar heat capacity of anhydrous NiCl2 between 2 and 30 K. –– –– : smoothed values of [75KOS], –– –– : smoothed values of [52BUS/GIA2] Dotted line: T 3–law up to 2 K, dash line T 3–law up to 15 K, solid line: ,m /pC T = 0.0011715·T 2 + 0.0091211·T up to 3.2 K [75KOS].

Figure V-15: Heat capacity of anhydrous NiCl2(cr). : experimental data from [52BUS/GIA2]. : smoothed data of [75KOS]. Dash line: linear spline T > TN, dotted line: linear spline T < TN

0 5 10 15 20 25 30 35

0.00

0.05

0.10

0.15

0.20

0.25

(C° p,

m ·

T −1

) / J·

mol

−1·K

−2

T / K

(C° p,m

·T –1

) / J·

K–2

·mol

–1

0 50 100 150 200 250 300 350

0.0

0.1

0.2

0.3

0.4

0.5

(Cp,

m ·

T −1

) / J·

mol

−1·K

−2

T / K

(C° p,m

·T –1

) / J·

K–2

·mol

–1


V.4.1.2.2 The high-temperature heat capacity of NiCl2(cr)

Coughlin conducted high temperature heat content measurements of nickel chloride over the temperature range from 298.15 to 1336 K [51COU]. The experimentally de-termined heat contents listed in [51COU] were fitted to the following function up to the highest temperature before any premelting was observed:

1 2 2m m

1

[ ( ) (298.15K)] / J·mol ·(( / K) 298.15 )

·((K / ) 298.15 ) ·(( / K) 298.15)

H T H a T

b Tc T

ο −

−

− = − +

− +−

(V.56)

With the standard heat capacity of NiCl2 according to Equation (V.55) as a constraint, c can be eliminated from Equation (V.56).

2m m,m

( ( ) (298.15))( ) 2· · · pp

H T HC T a T b T cT

οο − ∂ −

= = − + ∂

2,m (298.15) 2· ·(298.15) ·(298.15)pc C a bο −= − − (V.57)

From a least squares analysis of the m m( ( ) (298.15))H T H ο− data listed in [51COU] a ,m ( )pC Tο –function was obtained which is numerically slightly different from the one given by Coughlin, but the calculated ,mpCο –values differ only marginally. For the sake of consistency with the selected value of ,mpCο (NiCl2, cr, 298.15 K) the follow-ing ,m ( )pC Tο –function is selected:

1281,m 298[ ]pCο (NiCl2, cr) = 73.7949 + 1.22269×10–2 (T / K)

– 5.12775×105 · (T / K)–2 /J·mol–1·K–1 (V.58)

V.4.1.2.3 Enthalpy of formation of NiCl2(cr)

The enthalpy of formation of NiCl2 can be determined by a third-law analysis of the equilibrium studies on the reaction:

NiCl2(cr) + H2(g) Ni(cr) + 2 HCl(g) (V.59)

carried out by Berger and Crut, [21BER/CRU], Crut, [24CRU], Jellinek and Uloth, [26JEL/ULO], Sano, [37SAN2], [53SAN], Busey and Giauque, [53BUS/GIA], as well as Shchukarev et al. [54SHC/TOL]. An inspection of Figure V-16 reveals that the data of [53BUS/GIA], [53SAN], and [54SHC/TOL] appear to be the most reliable and best suited for this analysis. This was confirmed by linear fits through the data pairs T, − RT·ln(Kp /bar) which resulted in lower standard errors for [53SAN], [53BUS/GIA] and [54SHC/TOL] than for [24CRU] or [26JEL/ULO].

An alternative method consists in subjecting high-purity nickel metal to chlori-nation in a calorimetric bomb according to reaction:

Ni(cr) + Cl2(g) NiCl2 (cr) (V.60)

The advantage of the latter method is that the standard enthalpy of formation


r mH ο∆ (NiCl2, cr, 298.15 K) can be determined directly at the reference temperature, [84LAV/TIM].

A third independent method has been developed by Efimov et al. and is de-scribed in Appendix A [88EFI/EVD].

Figure V-16: r mο∆ G –function of Reaction (V.59). r mGο∆ = – RT·lnK°, linear fit of

experimental T; experimental data from: [53BUS/GIA], [54SHC/TOL], [37SAN2] [53SAN], [26JEL/ULO], [21BER/CRU] [24CRU]. r mGο∆ data: solid line [53BUS/GIA], dotted line [53SAN], dashed line [54SHC/TOL]

The third-law analyses of the equilibrium constants obtained by Busey and Gi-auque [53BUS/GIA], Sano [53SAN] and Shchukarev et al. [54SHC/TOL] for reaction (V.59) are summarised in Figure V-17. A combination of the standard enthalpy of HCl(g) formation recommended by CODATA [89COX/WAG] with the reaction en-thalpies of Figure V-17 resulted in f mH ο∆ (NiCl2, cr, 298.15 K) / kJ·mol–1 = − (305.01 ± 0.17) [53BUS/GIA]; − (305.5 ± 2.4) [54SHC/TOL], – (307.07 ± 0.56) [53SAN]. Busey and Giauque's data show the least uncertainty. The results of Shchuka-rev et al. led to a similar value with a significantly enhanced uncertainty. Moreover, Busey and Giauque's as well as Sano's values scatter around the mean, whereas Shchu-karev et al.'s results increase with increasing temperature.

550 600 650 700 750 800 850

-20

-10

0

10

20

30

∆ rG° m

/ kJ

·mol

−1

T / K

∆ rG

° m / k

J·mol

–1

NiCl2(cr) + H2(g) Ni(cr) + 2 HCl(g)


In principle equilibrium constants of other reactions such as:

NiO(cr) + 2 HCl(g) NiCl2(cr) + H2O(g)

can also be subjected to the third-law method, see Appendix A [77KAS/SAK].

Figure V-17: Third-law analysis of data for Reaction (V.59). Experimental data from [53BUS/GIA], [54SHC/TOL], and [53SAN]. Solid line: r mH ο∆ / kJ·mol–1 = (120.41 ± 0.10) [53BUS/GIA], dotted line: r mH ο∆ = (122.45 ± 0.56) kJ·mol–1, [53SAN], dash line: r mH ο∆ = (120.99 ± 2.4) kJ·mol–1 [54SHC/TOL], dash-dotted line: linear regression of [54SHC/TOL] data.

Lavut et al. [84LAV/TIM] determined the standard enthalpy of NiCl2 forma-tion very accurately f mH ο∆ (NiCl2, cr, 298.15 K) = – (304.78 ± 0.16) kJ·mol–1. This re-sult overlaps with the value obtained from the third-law analysis of Busey and Giau-que's data. Consequently the mean value of these investigations was selected in this review:

f mH ο∆ (NiCl2, cr, 298.15 K) = – (304.90 ± 0.11) kJ·mol–1. (V.61)

550 600 650 700 750 800 850

119

120

121

122

123

∆rH

° m /

kJ·m

ol−1

T / K

NiCl2(cr) + H2(g) Ni(cr) + 2 HCl(g)

∆rH

° m / k

J·mol

–1


V.4.1.2.4 Gibbs energy of formation of NiCl2(cr)

The Gibbs energy of formation of NiCl2(cr) can be calculated as a function of tempera-ture using the selected values of f mH ο∆ , mS ο , ,mpCο for NiCl2(cr) and for Ni(cr) as well as the respective CODATA values for Cl2(g) [89COX/WAG], adopted (Chapter VI) as NEA TDB auxiliary data (see Figure V-18).

Gee and Shelton measured the emf of solid electrolyte cells and thus obtained, for example, r mGο∆ of the displacement reaction (V.62) directly [76GEE/SHE].

Fe(cr) + NiCl2(cr) FeCl2(cr) + Ni(cr) (V.62)

The Gibbs energy of formation of FeCl2 appeared to be well established and thus the Fe, FeCl2 electrode was taken as the reference electrode. Figure V-18 shows that values measured in this manner and calculated f mGο∆ (NiCl2, cr) values agree with each other satisfactorily. This is considered to be a strong argument in favour of the selected thermodynamic quantities for NiCl2(cr) and Ni(cr).

Figure V-18: Gibbs energy of formation of NiCl2(cr) versus temperature. Values (straight line) and uncertainty (dotted lines) derived from [76GEE/SHE]. Dashed line calculated from the selected values of f mH ο∆ , mS ο , ,mpCο .

500 600 700 800 900-250

-240

-230

-220

-210

-200

-190

-180

-170

-160

∆ fG° m

(NiC

l 2, cr)

/kJ·m

ol−1

T / K

Ni(cr) + Cl2(g) NiCl2(cr)

∆fG

° m (N

iCl 2,

cr)

/ kJ

·mol

–1


V.4.1.3 Hydrated NiCl2 solids

Thermodynamic data have been reported for the nickel chloride hexahydrate, tetrahy-drate and dihydrate, and the existence of a heptahydrate has been reported at low tem-peratures (below – 30°C) [33BOY]. At 0.1 MPa, in contact with saturated aqueous solu-tions, the hexahydrate was reported [16DER/YNG] to be the stable form to approxi-mately 36.25°C, and dehydration of the tetrahydrate to the dihydrate occurs above 60°C [33BOY]. The lower hydrates can also be prepared by equilibration with aqueous HCl solutions at room temperature [23FOO].

V.4.1.3.1 Gibbs energy of formation of NiCl2·6H2O

Many solubility measurements for NiCl2·6H2O(cr) in water have been reported [33BOY], [37PEA/ECK], [79OIK], [79OJK/MAK], [79PET/SHE], [79PET/SHE2], [80PET/SHE], [83KIM/IMA], [85FIL/CHA], [86FIL/CHA], [87RAR3], [89LIL/TEP].

NiCl2·6H2O(cr) Ni2+ + 2Cl–+ 6H2O(l) (V.63)

Rard [92RAR2] has done an excellent assessment of the solubility of NiCl2·6H2O and the associated activity and osmotic coefficients of the saturated solu-tion at 298.15 K. Based primarily on results from [37PEA/ECK], [87RAR3] and [92RAR2], Rard calculated ((4.9155 ± 0.0022) mol·kg–1),

K ο (V.63) = (1108 ± 37) and r mGο∆ (V.63) = – (17.38 ± 0.08) kJ·mol–1.

Rard’s value is selected in the present review and, with the auxiliary data for Cl– (Chapter IV) and the selected value for f mGο∆ (Ni2+, 298.15 K) (Table III.1), the value:

f mGο∆ (NiCl2·6H2O, cr) = – (1713.7 ± 0.8) kJ·mol–1

is calculated.

V.4.1.3.2 Enthalpy of formation of NiCl2·6H2O

Surprisingly, the only reported direct measurement of the heat of solution of the hexa-hydrate in water appears to be that of Thomsen [1883THO] (4.85 kJ·mol–1 at 18.5°C [20BUR]). Archer [99ARC] has reported that an attempt to correct this value to 25°C and infinite dilution (cf. Appendix A) resulted in the value 0.14 kJ·mol–1. Ashcroft [70ASH] reported the heat of solution of NiCl2·6H2O(cr) in 1 M HCl(aq) as (6.10 ± 0.22) kJ·mol–1, and Gurrieri et al., [73GUR/CAL] reported (7.00 ± 0.13) kJ·mol−1. These two values do not agree within the authors’ (2σ) uncertainties, even though the final solution concentrations of the nickel salt were apparently quite similar in the experiments. Kakolowicz and Giera [75KAK/GIE] reported the heat of solution of NiCl2·6H2O(cr) in 4.36 M HCl(aq) as (21.32 ± 0.13) kJ·mol–1.


The final nickel concentrations in Thomsen’s experiments were quite high (ap-proximately 0.14 m), thus applying a proper heat of dilution correction to determine

r mH ο∆ (V.63) is problematic [99ARC]. The final nickel concentrations in the experi-ments of Ashcroft [70ASH] and Gurrieri et al. [73GUR/CAL] were about a factor of five lower, but to use these values a correction for transfer of the aqueous species from the dissolution medium to water is needed.

The enthalpy of solution measurements of Manin and Korolev [95MAN/KOR] for dissolution of anhydrous nickel chloride in HCl(aq) indicate that if the final solu-tions have nickel chloride concentrations of less than 0.01 M, the enthalpy of solution of anhydrous nickel chloride is approximately 4 kJ·mol–1 more positive than in water. This is in semi-quantitative agreement with results of the other measurements [1883THO], [70ASH], [73GUR/CAL]. Manin and Korolev [95MAN/KOR] also reported heats of solution of NiCl2(cr) in water as a function of the final solution concentration (for 0.0042 to 0.030 m). The variation in these values could be applied equally well to the heats of solution of the hydrated salt, but the experimental concentration range is not adequate to determine heats of dilution to I = 0 to better than ± 1 kJ·mol–1. As discussed in Appendix A, using a reasonable value for the enthalpy of dilution, the heat of solu-tion of NiCl2·6H2O(cr) to infinite dilution in water at 298.15 K can be estimated as (0.64 ± 1.80) kJ·mol–1. From this, f mH ο∆ (NiCl2·6H2O, cr) = – (2104.41 ± 2.01) kJ·mol–1 can be calculated.

Another approach is to compare the heats of solution of the anhydrous salt and the hydrate in the same medium (with appropriate minor corrections for changes in the heat of formation of water in the medium). However, except for Thomsen [1883THO], none of these researchers reported enthalpy of solution values for the anhydrous solid, NiCl2(cr) and the hexahydrate in the same medium. As discussed in Appendix A, the heat of hydration of NiCl2(cr) to NiCl2·6H2O(cr) at 298.15 K can be estimated as − (85.06 ± 1.80) kJ·mol–1, and f mH ο∆ (NiCl2·6H2O, cr) = – (2104.9 ± 1.8) kJ·mol–1.

In the present review, the weighted average of the two values based on the measurements of Thomsen [1883THO] is selected:

f mH ο∆ (NiCl2·6H2O, cr, 298.15 K) = – (2104.7 ± 1.8) kJ·mol–1,

but with an uncertainty somewhat larger than the statistical uncertainty because the val-ues are based, in part, on the same measurement.

V.4.1.3.3 The entropy and heat capacity of NiCl2·6H2O(cr)

Based on the selected values of f mGο∆ (NiCl2·6H2O, cr, 298.15 K) and f mH ο∆ (NiCl2·6H2O, cr, 298.15 K), the selected entropy is:

mS ο (NiCl2·6H2O, cr, 298.15 K) = (341.0 ± 6.7) J·K–1·mol–1.

Robinson and Friedberg [60ROB/FRI] reported measurements of the heat ca-pacity of NiCl2·6H2O(cr) between 1.6 K and 20.0 K, and found that the Néel tempera-


ture is at 5.34 K [73HAM/FRI]. Donaldson and Edmonds [65DON/EDM] extended the measurements to lower temperatures, with measurements between 0.7 K and 4 K. Simi-lar measurements were reported by Okaji and Watanabe [76OKA/WAT]. Although Ko and Hepler [63KO/HEP] reported mS ο (NiCl2·6H2O, cr, 298.15 K) = 344.3 kJ·mol–1, based on unpublished ,mpC data of Friedberg, the primary ,mpC data apparently were never published. In the present review, no heat capacity value at 298.15 K is selected for NiCl2·6H2O(cr).

V.4.1.3.4 Gibbs energy and enthalpy of formation of nickel chloride tetrahydrate

Measurements of water vapour pressures of mixtures of nickel chloride hydrates [16DER/YNG] and [40BEL], and over pure solids [93UVA/TIM] have been reported. Water vapour pressure values reported for the equilibrium:

NiCl2·6H2O(cr) NiCl2·4H2O(cr) + 2H2O(g) (V.64)

by Derby and Yngve [16DER/YNG] and the equation of Bell [40BEL] between 20 and 36°C are in good agreement, and may even be based on the same experimental results (cf. Appendix A), whereas Uvaliev et al. [93UVA/TIM] reported slightly lower values. The latter (isopiestic) work was focussed on determining transition pressures at constant temperature (25 and 35°C) rather than on the temperature dependence of the water va-pour pressure. The equilibrium water vapour pressure at 298.15 K is (1.30 ± 0.15) kPa, based on the average of the values reported by [16DER/YNG] (1.40 kPa) and Uvaliev et al. [93UVA/TIM] (1.19 kPa). The selected value is:

r mGο∆ (V.64) = (21.56 ± 0.60) kJ·mol–1.

The uncertainty has been assigned to span the uncertainties in these inconsis-tent results.

The enthalpy of reaction, determined only from the results of Derby and Yngve [16DER/YNG], is (102.4 ± 1.8) kJ·mol–1; and (107.8 ± 5.8) kJ·mol–1 if the values from [93UVA/TIM] are included. If the water vapour pressure is fixed at the selected value at 298.15 K, the fitted slope of

210 H Olog p against (T/K)–1 leads to a calculated enthalpy of reaction of (111.2 ± 5.6) kJ·mol–1. In the present review a median value spanning all these values is selected:

r mH ο∆ (V.64) = (107 ± 9) kJ·mol–1.

Using the values selected above for the hexahydrate, and auxiliary data from Chapter IV, yields:

f mGο∆ (NiCl2·4H2O, cr, 298.15 K) = – (1234.9 ± 1.0) kJ·mol–1

and

f mH ο∆ (NiCl2·4H2O, cr, 298.15 K) = – (1514.0 ± 9.2) kJ·mol–1.


From these and auxiliary values from Chapter IV,

mS ο (NiCl2·4H2O, cr, 298.15 K) = (250 ± 31) J·K–1·mol–1

is calculated.

Bell [40BEL] reported on measurements of the vapour pressure of D2O over mixtures of the tetradeuterate and hexadeuterate, and it was found that the ratio of equi-librium values of

2 2H O D O/p p was greater than 1.0 near 25°C.

V.4.1.3.5 The entropy and heat capacity of nickel chloride dihydrate

Polgar et al. [72POL/HER] measured the specific heat of NiCl2·2H2O(cr) for tempera-tures between 1.2 and 24.5 K. Peaks were observed at 6.31 and 7.26 K, and at lower temperatures the heat capacity values are not a simple function of T 3. The authors esti-mated that the contribution to the molar entropy between 0 and 1.2 K is 0.03 R. Using an adiabatic calorimeter, Juraitis et al. [90JUR/DOM] measured the heat capacity of the solid between 80 and 281 K, with emphasis on the effects of the structural transition near 220 K. It is not clear from these results (only reported graphically) whether the results from 250 to 281 K can be extrapolated smoothly to 298.15 K.

A series of polynomials were fitted to the values of ,m /pC T over the entire temperature range from 1.2 K to 281 K. The uncertainty in the extrapolation below 1.5 to 0 K is of the order of ± 0.2 J·K–1·mol–1. A single simple polynomial was fitted to the combined results from 20 to 25 K [72POL/HER] and 80 to 100 K [90JUR/DOM]. In the present review, it is assumed that there are no further anomalies in this temperature range, and the uncertainty contribution to mS ο (NiCl2·2H2O, cr) over this range is esti-mated as ± 2.0 J·K–1·mol–1. The uncertainties in the set of ,mpCο values above 240 K (see Appendix A) independently introduce an uncertainty of ± 3.0 J·K−1·mol–1 in the calcu-lated value of mS ο . In the present review,

mS ο (NiCl2·2H2O, cr, 298.15 K) = (186.2 ± 3.6) J·K–1·mol–1

is selected, where the uncertainty is an estimate. From the work of Juraitis et al. [90JUR/DOM] (see Appendix A),

,mpCο (NiCl2·2H2O, cr, 298.15 K) = (230 ± 15) J·K–1·mol–1

is selected in the present review. Again, the uncertainty is estimated.

V.4.1.3.6 Gibbs energy and enthalpy of formation of nickel chloride dihydrate

Derby and Yngve [16DER/YNG] and Uvaliev et al. [93UVA/TIM] both reported water vapour pressure measurements for the dehydration of nickel chloride tetrahydrate to the dihydrate.

NiCl2·4H2O(cr) NiCl2·2H2O(cr) + 2 H2O(l) (V.65)


The values of 210 H Olog p from the earlier study do vary monotonically with

1/T, though they are in fair agreement with the results from Uvaliev et al. [93UVA/TIM] at 25 and 35°C. In the present review, the average of the discordant val-ues from the two studies at 298.15 K is used to calculate a value for r mGο∆ (V.65), but no value for r mH ο∆ (V.65) is selected,

r mGο∆ (V.65) = (23.4 ± 0.7) kJ·mol–1.

Using the values selected above for the tetrahydrate, and auxiliary data from Chapter IV,

f mGο∆ (NiCl2·2H2O, cr, 298.15 K) = – (754.4 ± 1.3) kJ·mol–1.

Based on this, auxiliary data from Chapter IV, and mS ο (NiCl2·2H2O, cr, 298.15 K) from Section V.4.1.3.5,

f mH ο∆ (NiCl2·2H2O, cr, 298.15 K) = – (913.4 ± 1.7) kJ·mol–1.

V.4.1.4 Solid nickel bromide, NiBr2(cr)

There are two different crystallographic forms of paramagnetic NiBr2(cr) [34KET]. Samples prepared by sublimation, or by melting and slow cooling, have a CdCl2 struc-ture, rhombohedral space group R3m [33BIJ/NIE], [34KET], [80ADA/BIL]. Samples obtained by low-temperature dehydration of the hydrated salts or by recrystallisation of these solids from organic solvents have a “Wechselstruktur” [34KET]. Day et al. [76DAY/DIN] describe the structure as a disordered intergrowth of the CdCl2 and CdI2 lattices [33BIJ/NIE]. Both forms have been used in experiments to determine the chemical thermodynamic properties of anhydrous nickel bromide, often without a rec-ognition that two forms can exist. NiBr2(cr) orders to an antiferromagnetic form below 52 K, and to a helimagnetic structure below 23 K [83KAT/SUG].

V.4.1.4.1 Low-temperature heat capacity and entropy of NiBr2(cr)

There are two sets of low-temperature (8 to 300 K) heat capacity (adiabatic calorimetry) measurements [78STU/FER], [82WHI/STA] that are in good agreement between 30 and 250 K (as discussed in Appendix A). Both sets of results show a small excess heat ca-pacity between 40 and 55 K, near the Néel temperature. White and Staveley [82WHI/STA] also found a small, sharp peak near 20 K, and apparently this is associ-ated with the transition from an antiferromagnetic structure to an incommensurate phase with helimagnetic ordering [76DAY/DIN], [81ADA/BIL]. No corresponding peak was found in the earlier study [78STU/FER], and the reported heat capacities below 15 K decrease much more slowly with decreasing temperature than suggested by the results of [78STU/FER]. Unlike the results of White and Staveley, the values reported by Stuve et al. below 18 K do not fit well to a T 3 relationship (see Figure V-19). Above 250 K, the heat capacity values reported by White and Staveley increase rapidly with increasing temperature, and the authors suggested [82WHI/STA] that the sample may have been


contaminated by moisture (at 298.15 K, their value of ,mpCο (NiBr2, cr) is approximately 83 J·K–1·mol–1, as opposed to 75.4 J·K–1·mol–1 from Stuve et al.).

Stuve et al. [78STU/FER] integrated their heat capacity values to obtain mS ο (NiBr2, cr, 298.15 K) = (122.42 ± 0.25) J·K–1·mol–1. In the present review, the ex-

perimental ,mpC (T) values reported in both papers were integrated. The results of White and Staveley [82WHI/STA] are accepted to 250 K (an entropy contribution of 108.23 J·K−1·mol–1). The contribution to the entropy from 0 to 8 K (0.09 J·K–1·mol–1) was esti-mated by assuming ,mpCο /T varies as T 2 (Figure V-19), though there are no measure-ments to support this assumption. The results of Stuve et al. [78STU/FER] are accepted between 250 K and 298.15 K (an entropy contribution of 13.14 kJ·mol–1). The sum of these contributions provides the selected value:

mS ο (NiBr2, cr, 298.15 K) = (121.4 ± 1.0) J·K–1·mol–1 .

The uncertainty is an estimate based on the spread of the entropy contributions from the two studies in different temperature ranges, and considering that there may be an unknown contribution from the extrapolation from 8 to 0 K.

The reported ,mpCο (NiBr2, cr, 298.15 K) value of Stuve et al. [78STU/FER]

,mpCο (NiBr2, cr, 298.15 K) = (75.40 ± 1.00) J·K–1·mol–1

is selected in the present review with a markedly increased uncertainty.

Figure V-19: Values from the heat capacity studies of Stuve et al. [78STU/FER] and White and Staveley [82WHI/STA]. The curves (from bottom to top) are based on the equation ,mpCο (T) = bT 2 for b = 0.000500, 0.000750, 0.001000, 0.001350 and 0.002000. 0 2 4 6 8 10 12 14 16 18 20

0.0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

[78STU/FER] [82WHI/STA]

Cp/T

/ J·

K-2·m

ol-1

T / K

(C° p,m

·T–1

) /J·K

–2·m

ol–1

T / K


V.4.1.4.2 The high-temperature heat capacity of NiBr2(cr)

The fitted heat capacity equation reported by the same authors [78STU/FER] repro-duces all their high-temperature experimental m m( ) (298.15)H T H ο− values to better than 0.5% and is fixed to the 298.15 K value from the low-temperature measurements. The values from the equation deviate slightly from the experimental values below 298.15 K. If this heat capacity equation is used with the high-temperature equilibrium measurements of Shchukarev et al. [54SHC/TOL], the derived values of f mH ο∆ (NiBr2, cr, 298.15 K) vary with the measurement temperature (see the Appendix A discussion of [54SHC/TOL]). In the present review, this drift has been attributed to experimental difficulties in the work of Shchukarev et al., but it might also indicate problems with the drop calorimetry work of Stuve et al. It is also not clear whether the more stable (or the same) crystalline form of NiBr2 was used for all the measurements [34KET], [54SHC/TOL]. Nevertheless, in the present review, the high-temperature heat capacity equation of Stuve et al. is selected:

1200K,m 298.15K[ ]pCο (NiBr2, cr)/J·K–1·mol–1 = 72.3163 + 0.014627T/K – 113805 (T/K)–2.

V.4.1.4.3 Enthalpy of formation of NiBr2(cr)

Stuve et al. [78STU/FER] measured the heats of heat of solution of samples of NiBr2(cr), (possibly the “Wechselstruktur” form), NiSO4(cr) and H2SO4·6H2O(l) in 4.36 m HCl(aq) and, from these measurements, as described in Appendix A, f mH ο∆ (NiBr2, cr, 298.15 K) = − (211.214 ± 1.752) kJ·mol–1. Evdokimova and Efimov [89EVD/EFI] measured the heats of solution of Br(l), nickel and NiBr2(cr), (probably the “Wechsel-struktur” form) in an aqueous KBr, Br2, HBr mixture. From these measurements,

f mH ο∆ (NiBr2, cr) = − (211.93 ± 0.90) kJ·mol–1, 2σ uncertainties (Appendix A).

Several groups, Crut [24CRU], Jellinek and Uloth [26JEL/ULO2] and Shchu-karev et al. [54SHC/TOL] have measured the equilibrium constant for the reaction:

NiBr2(cr) + H2(g) Ni(cr) + 2HBr(g).

at temperatures between 623 K and 928 K. The most extensive, and least scattered set of values is that of [54SHC/TOL]. As discussed in Appendix A, these measurements can be used to calculate f mH ο∆ (NiBr2, cr) = – (215.02 ± 0.91) kJ·mol–1. Values based on the results of [24CRU] and [54SHC/TOL] are in rough agreement (– 214 to – 217 J·K−1·mol–1). A value of f mH ο∆ (NiBr2, cr) based on the sparse enthalpy of solution measurements of Paoletti [65PAO] would depend on use of the selected value of

f mH ο∆ (Ni2+, 298.15 K) (also see the discussion of [90EFI/FUR] in Section V.2.1.2 and Appendix A).

It is unclear whether the enthalpy of formation values based on the measure-ments of Stuve et al., Shchukarev et al. and Evdokimova and Efimov are for the identi-cal solid. The solid of Evdokimova and Efimov was almost certainly the “Wechselstruk-tur” form, and that of Shchukarev et al. was possibly the rhombohedral form, at least for


some of their measurements. This seems to be reflected in the less negative value of f mH ο∆ (NiBr2, cr) found by Evdokimova and Efimov. In the absence of other informa-

tion, the average of the more precise values of Evdokimova and Efimov and of Shchu-karev et al. is selected, and the uncertainty is selected to span the uncertainties of the these measurements:

f mH ο∆ (NiBr2, cr, 298.15 K) = – (213.5 ± 2.4) kJ·mol–1.

This value results in the following derived Gibbs energy of formation:

f mGο∆ (NiBr2, cr, 298.15 K) = – (195.4 ± 2.4) kJ·mol–1.

Based on this enthalpy of formation value and the selected value for f mH ο∆ (Ni2+, 298.15 K), the enthalpy of solution of NiBr2(cr) in water (at infinite dilu-

tion) is – (83.9 ± 2.6) kJ·mol–1. This value is less negative than, but in marginal agree-ment with the – (86.18 ± 0.40) kJ·mol–1 reported by [90EFI/FUR], and is also in agree-ment, within the uncertainty limits, with the rather poorly documented result of Paoletti [65PAO] (– 82.1 kJ·mol–1).

V.4.1.5 Hydrated solid, NiBr2·xH2O

Although several hydrated compounds have been reported, information on well-characterised samples is sparse. The isolation of the hexahydrate (over a desiccant at room temperature) has been reported [66GME], [77BHA/CAR]. Conversion of the hexahydrate to the trihydrate may occur near 30°C.

The low-temperature heat capacity of NiBr2·6H2O was reported by Spence et al. [59SPE/FOR] (between 4.0 and 13 K), and by Bhatia et al. [77BHA/CAR] (between 1.5 and 30 K). There is a λ–transition in the heat capacity measurements. Spence et al. indicate that the transition is at 6.5 K, whereas Bhatia et al. report that the transition is at (8.30 ± 0.02) K, the same temperature as the paramagnetic-antiferromagnetic transi-tion found in their magnetic susceptibility measurements. The two groups report similar values for the entropy contribution of the transition (9.62 J·K−1·mol–1 [59SPE/FOR], (9.41 ± 0.21) J·K–1·mol–1 [77BHA/CAR]). Because no values could be found for the heat capacity of NiBr2·6H2O for temperatures between 30 and 298.15 K, no value for

mS ο (NiBr2·6H2O, cr) is selected.

A value for f mH ο∆ (NiBr2·3H2O) has been reported in several standard sets of tables [36BIC/ROS], [52ROS/WAG], [82WAG/EVA]. This value has been traced to enthalpy of solution measurements of Crut [24CRU], [24CRU2] (from – 79.1 kJ·mol–1 for the enthalpy of hydration of the anhydrous salt [24CRU]). However, the stoichiome-try of the hydrated salt is not specified, and there are ambiguities in the reported values (see the discussion in Appendix A). In the present review, no chemical thermodynamic values are selected for the nickel bromide hydrates for 298.15 K.


V.4.1.6 Solid nickel iodide, NiI2(cr)

V.4.1.6.1 Entropy of NiI2(cr)

Worswick et al. [74WOR/COW] reported low-temperature (adiabatic calorimetry) heat capacity measurements (10 K to 300 K) for this solid, and calculated mS ο (NiI2, cr) = 138.7 J·K–1·mol–1. A λ-transition at 58.9 K was reported. There are only limited confir-matory heat capacity measurements in the literature. Based on magnetic susceptibility measurements, Van Uitert et al. [65UIT/WIL] reported a λ-transition as occurring at approximately 75 K. Billerey et al. [77BIL/TER] found heat-capacity transitions at 59.5 K and near 76 K, and these were later identified by Kuindersma et al. [81KUI/SAN] as a structural transition and a magnetic phase transition, respectively. The raw data from Billerey et al. are not available, but the contributions to the entropy from the two transitions are estimated to be 0.40 and 0.37 J·K–1·mol–1 (see Appendix A). Examination of the data from Worswick et al. indicates that their value for mS ο (NiI2, cr) included a contribution of 0.34 J·K–1·mol–1 as an excess entropy in the region 50 K to 85 K. There are no reported heat capacity values at temperatures below 10 K, though the magnetic behaviour of the solid is very complex at low temperatures [81KUI/SAN], [90PAS/TAY], [92FRE/FAL]. The entropy, adjusted for the contribution of the transi-tion at 76 K, is 139.1 J·K–1·mol–1, and this value is accepted in the present review. A rather large uncertainty of ± 1.0 J·K–1·mol–1 is assigned to the selected value:

mS ο (NiI2, cr, 298.15 K) = (139.1 ± 1.0) J·K–1·mol–1 .

V.4.1.6.2 Heat capacity of NiI2(cr)

Worswick et al. [74WOR/COW] reported low-temperature (adiabatic calorimetry) heat capacity measurements (10 to 300 K) and differential scanning calorimetry (DSC) re-sults from 300 to 550 K. The equation (V.66) (200 to 550 K) is not in the standard form generally used for solids (e.g., [93KUB/ALC], p. 166, Equation (116)), however, the original DSC data are unavailable.

550K,m 200K[ ]pCο (NiI2, cr) / J·K–1·mol–1 = 65.5 + 0.0515 T/K – 0.0000397 (T/K)2 (V.66)

Equation (V.67), in a more standard form, was calculated to fit the adiabatic calorimetry results from 201.55 to 297.39 K and values from equation (V.66) at higher temperatures (as discussed in [74WOR/COW]). Equation (V.67) closely reproduces the values from the authors’ equation (within 0.5 J·K–1·mol–1 between 300 and 550 K, and within the experimental uncertainties at lower temperatures). Equation (V.67) is ac-cepted in the present review for the specified temperature range, but this equation should not be used at higher temperatures. NiI2(cr) is known to decompose at tempera-tures slightly above 700 K.

550K,m 200K[ ]pCο (NiI2, cr) / J·K–1·mol–1 = 73.5775 + 0.016425T – 1.057753×105 (T/K)–2 (V.67)

The value of ,mpCο (NiI2, cr, 298.15 K) based on equation (V.67),


,mpCο (NiI2, cr, 298.15 K) = (77.28 ± 1.00) J·K–1·mol–1

is selected in the present review with a markedly increased uncertainty (see Appen-dix A).

V.4.1.6.3 Enthalpy of formation of NiI2(cr)

Values for the enthalpy of formation of NiI2(cr) have been reported (or can be calcu-lated) based on enthalpy of solution measurements [65PAO/SAB], [90EFI/EVD] and from high temperature equilibrium measurements [80OPP/TOS]. The work of Opper-mann and Toschew [80OPP/TOS] considers many of the difficulties in determining values from the rather complex Ni-I system at higher temperatures, and the problems in doing a proper extrapolation to determine thermodynamic values at 298.15 K. However, as discussed in Appendix A, this process can only be carried out within very large un-certainty bounds, by using the selected ,m ( )pC Tο equation for NiI2(cr), and drawing ten-tative conclusions related to the thermodynamic functions for NiI2(g) and Ni2I4(g). The earlier measurements of Jellinek and Uloth [26JEL/ULO2] suffer from many of the same problems, but are also too sparse to be properly re-interpreted. The value reported by Bartovská et al. [74BAR/CER] is based, in part, on estimated values.

A value of f mH ο∆ (NiI2, cr) based on the enthalpy of solution measurements of Paoletti et al. [65PAO/SAB] or Efimov and Furkalyuk [90EFI/FUR] would depend on use of the selected value of f mH ο∆ (Ni2+, 298.15 K). However, experimental results re-ported by Efimov et al. [88EFI/EVD], Efimov and Evdokimova [90EFI/EVD] provide a value of f mH ο∆ (NiI2, cr, 298.15 K) that is independent of the value for f mH ο∆ (Ni2+, 298.15 K). From their work (see Appendix A), the value,

f mH ο∆ (NiI2, cr, 298.15 K) = – (96.42 ± 0.84) kJ·mol–1

is selected in the present review. This value is in agreement with – (91 ± 8) kJ·mol–1 based on Oppermann and Toschew [80OPP/TOS] (see Appendix A). This value results in the following derived Gibbs energy of formation:

f mGο∆ (NiI2, cr, 298.15 K) = – (94.4 ± 0.9) kJ·mol–1.

Based on the selected values of f mH ο∆ (NiI2, cr, 298.15 K) and f mH ο∆ (Ni2+, 298.15 K), the enthalpy of solution (at infinite dilution) of NiI2(cr) in water is − (71.8 ± 1.2) kJ·mol–1, in reasonable agreement with the rather poorly documented result of Paoletti [65PAO] (– 70.5 kJ·mol–1) for slightly higher ionic strength. The value of – (73.38 ± 0.12) kJ·mol–1 reported by Efimov and Furkalyuk [90EFI/FUR] is some-what more negative.

V.4.1.7 Solid nickel iodate compounds

Pracht et al. [97PRA/LAN] have summarised and examined phase relationships in the Ni(IO3)2-H2O system. They concluded that, as reported in earlier papers [73COR], [73NAS/SHI], two tetrahydrates, a dihydrate and a higher hydrate, perhaps a decahy-


drate, may have regions of metastability or stability near 300 K. Meusser [01MEU] measured the solubilities of several different hydrated nickel iodates, and the most sta-ble form near room temperature was the “β-dihydrate”, which is the Ni(IO3)2·2H2O of later workers. The solubility results of Meusser [01MEU] and Cordfunke [73COR] for the dihydrate are quite similar, and extrapolation [73COR] and interpolation [01MEU] to 298.15 K provide values of – (5.17 ± 0.06) and – (5.17 ± 0.04) for 10 ,0log sK ο ((V.68), 298.15 K) (see Appendix A). This agreement is rather surprising as neither worker re-ported a measurement at 298.15 K, and there are doubts concerning whether equilibrium was attained, especially in the work of Cordfunke [73COR].

Ni(IO3)2·2H2O(cr) Ni2+ + 2 3IO− + 2H2O(l) (V.68)

Fedorov et al. [73FED/SHM] measured the solubility of a solid with the re-ported composition of a trihydrate, Ni(IO3)2·3H2O, in aqueous lithium perchlo-rate/nitrate mixtures at 298.15 K and ionic strengths from 0.5 to 4.0 M. Their solubility results near 25°C, on extrapolation to I =0, are similar to those from the other two stud-ies. Based on the values obtained from experiments using solutions without nitrate, a value of the solubility product of Ni(IO3)2·3H2O, 10 ,0log sK ο = – (5.09 ± 0.16) is calcu-lated (see Appendix A).

It is clear from all the studies that the kinetics of the interconversion of the hy-drates is slow at room temperature. The evidence for a true trihydrate is slim, yet might be the result of catalysed precipitation of a more stable form that is difficult to synthe-size in other media. It might also reflect insufficient periods of equilibration, or that the solid was not properly handled prior to analysis. There does not seem to be any thermo-dynamic information (vapour pressure measurements) available on the interconversion of the hydrates.

In the present review the solubilities of Fedorov et al. [73FED/SHM] are as-sumed to relate to the dihydrate rather than the “trihydrate”. The three values for the dihydrate at 298.15 K are then in marginal agreement within the statistical (2σ) uncer-tainties. The average from the three studies is selected:

10 ,0log sK ο (V.68) = − (5.14 ± 0.10),

where the uncertainty of 0.10 in 10 ,0log sK ο is assigned to reflect uncertainty in the crys-talline form.

From the solubility product,

r mGο∆ ((V.68), 298.15 K) = (29.34 ± 0.57) kJ·mol–1.

Using this selected value and the auxiliary data in Table IV-1,

f mGο∆ (Ni(IO3)2·2H2O, cr, 298.15 K) = – (802.1 ± 1.8) kJ·mol–1.

The effective enthalpy of the dissolution reaction can be calculated from the temperature dependence of the solubility products, provided that the heat capacity of


reaction is small over the experimental temperature range. For a dissolution reaction of an ionic salt, it is unlikely that the latter condition is met; nevertheless, as shown in Ap-pendix A, the values based on the data of Meusser [01MEU] and Cordfunke [73COR] do not appear to change significantly if the temperature range used for calculation of

r mH∆ (V.68) is changed. On this basis, the average of the calculated temperature-averaged values for the enthalpy of reaction (V.68), 21.4 kJ·mol–1 (281 to 323 K) [01MEU] and 21.7 kJ·mol–1 (302 to 323 K) [73COR], is selected for 298.15 K, but with a slightly increased estimated uncertainty:

r mH ο∆ ((V.68), 298.15 K) = (21.6 ± 5.0) kJ·mol–1.

Using this value and the auxiliary data in Table IV-1,

f mH ο∆ (Ni(IO3)2·2H2O, cr, 298.15 K) = – (1087.7 ± 5.2) kJ·mol–1

mS ο (Ni(IO3)2·2H2O, cr, 298.15 K) = (270.1 ± 17.4) J·K–1·mol–1

are calculated.

The heat capacity of Ni(IO3)2·2H2O(cr) has been measured [64CHA/BOE], [72KLA/DOK] between 1 and 30 K, but no higher temperature measurements appear to have been reported.

There are two forms of anhydrous nickel iodate, and there is some evidence [73NAS/SHI], [97PRA/LAN] that the yellow (hexagonal) β form is more stable than the green α form. The measurements of Meusser for the solubility of an anhydrous nickel iodate, probably the β-Ni(IO3)2 [73NAS/SHI], at four temperatures leads to

10 ,0log sK ο ((V.69), Ni(IO3)2, β, 298.15 K) = – (4.43 ± 0.02), and an average value of r mH∆ (V.69) = – (7.3 ± 0.7) kJ·mol–1 (303.15 K to 363.15 K). It is clear that at 298.15 K

the dihydrate is less soluble than the anhydrous salt. In the absence of other data, the value of the solubility product is selected, but with an increased uncertainty:

β-Ni(IO3)2 Ni2+ + 2 3IO− (V.69)

r mGο∆ ((V.69), 298.15 K) = (25.3 ± 0.1) kJ·mol–1.


f mGο∆ (Ni(IO3)2, β, 298.15 K) = – (323.7 ± 1.7) kJ·mol–1.

The enthalpy of reaction value is somewhat suspect because it is an average value from measurements over a rather large temperature range (not including 298.15 K), and no attempt has been made to include heat capacity values in the data analysis. Nevertheless, the value is selected here with an increased estimated uncer-tainty.

r mH ο∆ ((V.69), 298.15 K) = – (7.3 ± 4.0) kJ·mol–1.



f mH ο∆ (Ni(IO3)2, β, 298.15 K) = – (487.1 ± 4.2) kJ·mol–1

mS ο (Ni(IO3)2, β, 298.15 K) = (213.5 ± 14.1) J·K–1 mol–1

are selected.

V.4.2 Aqueous nickel halide complexes V.4.2.1 Introduction

Halide anions, with the exception of fluoride, form rather unstable complexes with Ni(II) in aqueous solution. This is mostly due to the strong hydration of Ni(II). Thus, water can efficiently compete with the essentially electrostatic Ni(II)-halide interaction. Consequently, high and varying excesses of ligand anions over Ni(II) have been used to assess the stability of the complexes formed. Under such conditions a large part of the medium ions are replaced by the complex-forming electrolyte, causing substantial change in the activity coefficients, and thus, the constant medium principle is violated in these measurements. According to Harned's rule, the logarithms of the activity coeffi-cients in a mixture of two electrolytes vary linearly with their mole fractions. This varia-tion, which depends mainly on the hydration of the ions involved, introduces an inher-ent and electrolyte dependent uncertainty into the determined stability constants of weak complexes [89BJE]. The hypothesis generally made in the earlier volumes of this series, was that in all studies of weak complexation, to a first approximation, the activity coef-ficients do not change significantly as the weak complex-forming anion substitutes for perchlorate, provided that the total ionic strength is maintained constant. In the present review the same procedure is accepted, emphasising the approximative validity of this hypothesis. Therefore, if considerable medium changes occurred during the experi-ments, we assign higher uncertainty to the accepted formation constants than was pub-lished in the original report, regardless of the quality of the experimental work. Fur-thermore, as it is almost impossible to distinguish between a medium effect and the formation of higher complexes, for lack of solid evidence, only NiX+ species are ac-cepted in this review, if a substantial medium effect can be assumed.

In some cases mixed or ‘semi-thermodynamic’ formation constants are re-ported [36JOB], [71PAA/HUM], [72RET/HUM], [88BJE], [90BJE]. In these sets of measurements, the formation of weak complexes was studied at low concentrations of the metal ion and up to the highest possible concentrations of the highly soluble salts (NX) of the complex-forming anions (X–). The ‘semi-thermodynamic’ formation con-stants can be written as:

[ ][ ] [ ]1

NiXNiX X

nn

n

K− ±

=⋅ ⋅γ

where γ± is the molar activity coefficient of the complex forming electrolyte (NX) in molar units [88BJE], [90BJE]. Others [71PAA/HUM], [72RET/HUM] used the so called complete equilibrium constant ( 'nK ), in which the water is also taken into con-sideration in the equilibrium process,


[NiXn–1(H2O)7–n]3–n + X– [NiXn(H2O)6–n]2–n + H2O(l)

( )( ) [ ]

2 w6'

1 2 7

NiX H O

NiX H O Xn n

nn n

aK −

− ±−

⋅ = ⋅ ⋅ γ

Although these constants probably provide a better description of the given system in the case of a high and varying excess of complex-forming anion than do those determined using the constant medium principle, their chemical meaning is not clear [90BEC/NAG]. These ‘semi-thermodynamic’ formation constants are not equivalent to the value extrapolated to zero ionic strength using the SIT, except for the case when the ratio ( ) ( )2 1 26 7[NiX H O ] [NiX H O ]/

n nn nγ γ

−− −= 1 over the whole range of the ionic strength studied.

Besides the weakness of complex formation, the simultaneous presence of in-ner- and outer-sphere complexes, especially in the cases of chloride and bromide, fur-ther complicates the determination of the correct solution speciation. Using most of the experimental methods the measured formation constant is the sum of the formation con-stants of two types of complexes, 2 5 inner[NiL(H O) ] and 2 6 outer[Ni(H O) L] :

exp inner outer2 5 inner 2 6 outer1 1 1

2 6

[NiL(H O) ] + [Ni(H O) L] =[Ni(H O) ] [L]

= +⋅

β β β

However, the d-d transitions of Ni(II) are certainly more influenced by inner- than outer-sphere complexation. In the case of weak complexes, inner-sphere interaction can be expected to dominate only at markedly reduced water activity, i.e., in case of concentrated solutions. Consequently, a visible spectrophotometric study, using high concentrations of the complex-forming anion, mostly provides information on inner-sphere complexes, and it is therefore not strictly comparable with e.g., potentiometric results obtained at considerably lower ligand concentrations.

V.4.2.2 Solution structural studies

A number of NMR [68LIN/APR], [79WEI/HER], [79WEI/MUL], [83GOW/DOD], EXAFS [80LAG/FON], [83LIC/PAS], [99HOF/DAR], solution X-ray [80CAM/LIC], [82WAK/ICH], [82MAG/PAS], [85MAG/MOR], [99WAI/KOU], neutron diffraction [83NEI/END] and neutron scattering [2002CHI/SIM] studies are available concerning the structure of concentrated NiCl2 and NiBr2 solutions. In presence of 2 M NiCl2 and 4 − 6 M LiCl/HCl, solution X-ray [82MAG/PAS] and EXAFS [83LIC/PAS] studies clearly demonstrated the formation of inner-sphere chloro complexes, although the av-erage number of Ni2+ − Cl– contacts are relatively low (nCl– = 0.7-1.4). In stoichiometric solutions the situation is less clear. Some authors [80LAG/FON], [80CAM/LIC], [83NEI/END], [99HOF/DAR] have detected the formation of ion pairs only, but the majority of reports agree with the presence of inner-sphere complexes at higher concen-trations. The first solid evidence in favour of the existence of inner-sphere [Ni(H2O)5Cl]+ complex in stoichiometric NiCl2 solutions was the detection of slow chloride exchange by 35Cl NMR at – 5 and + 1°C [83GOW/DOD]. This observation


demonstrates the presence of inner-sphere complexes, since outer-sphere interaction is unlikely to produce slow ligand-exchange. Solution X-ray diffraction studies of concen-trated NiBr2 solutions [82WAK/ICH], [85MAG/MOR] confirmed the above observa-tions, since the very large backscattering amplitude of a Br scatterer results in better identification of Ni2+ − Br– contacts than the Ni2+ − Cl– ones. Although in [99HOF/DAR] the authors report the formation of ion pairs at 298 K, the temperature dependence of XAFS spectra clearly indicate the formation of tetrahedral [NiBr4]2– complex above 698 K. The average number of Ni2+ − X– contacts reported in most of the publications mentioned above allowed the calculation of the stability constants of inner-sphere com-plexes. Although the applied ionic strengths are generally very high, some of these con-stants are reported in Sections V.4.2.4 and V.4.2.5.

V.4.2.3 Aqueous Ni(II) - fluoro complexes

The reported stability constants for Reaction (V.70) are listed in Table V-8.

Ni2+ + F– NiF+ (V.70)

Table V-8: Experimental stability constants (logarithmic values) of the species NiF+.

Method Medium t (°C) 10 1log K reported

10 1log K (a) accepted

Reference

ise-F 0.05 M (CH3)4NClO4 25 (1.32 ± 0.05) (1.32 ± 0.30) [83SOL/BON] ise-F 0.1 M NaClO4 25 (1.02 ± 0.12) (1.02 ± 0.20) [81KUL/BLO] kin 0.1 M NaClO4 25 (1.1 ± 0.1) (1.10 ± 0.20) [69FUN/TAN] ise-F 0.25 M NaClO4 25 (0.91 ± 0.10) (0.91 ± 0.30) Mironov et al.(b) ise-F 0.5 M NaClO4 25 (0.76 ± 0.08) (0.75 ± 0.20) Mironov et al.(b) ise-F 1.0 M NaClO4 25 (0.68 ± 0.09) (0.66 ± 0.20) Mironov et al.(b) ise-F 1.0 M NaClO4 25 (0.34 ± 0.04) (0.32 ± 0.30) [72BON/HEF] pol 1.0 M Na(ClO4, F) 25 (0.48 ± 0.05) (0.46 ± 0.30) [71BON] qh 1.0 M NaClO4 20

25 (0.66 ± 0.05) (0.64 ± 0.10)

(0.66 ± 0.10)(c)

[56AHR/ROS]

ise-F 2.0 M NaClO4 25 (0.65 ± 0.06) (0.61 ± 0.20) [81KUL/BLO] ise-F 3.0 M NaClO4 25 (0.76 ± 0.06) (0.69 ± 0.20) Mironov et al.(b) (a) Corrected to molal scale. The accepted values reported in Appendix A are expressed on the

molar or molal scales, depending on which units were used originally by the authors. (b) Average value from the studies reported by Mironov’s group [83AVR/BLO],

[81KUL/BLO], [76KUL/BLO]. (c) Corrected to 298.15 K using the selected value of r mH ο∆ .

The majority of data in Table V-8 were obtained in NaClO4 solutions using a fluoride selective electrode, but some pH-metric [56AHR/ROS], kinetic [69FUN/TAN] and polarographic [71BON] results are included, too. The SIT analysis of the experi-mental stability constants is depicted in Figure V-20.


Figure V-20: Extrapolation to Im = 0 of the experimental data for Reaction (V.70) in NaClO4 media. Experimental data from: : [56AHR/ROS], : [69FUN/TAN], : [71BON], : [72BON/HEF], : [81KUL/BLO], : [83SOL/BON], : [76KUL/BLO], [81KUL/BLO], [83AVR/BLO].

The value given by Arhland and Rosengren [56AHR/ROS] determined at 293.15 K has been corrected to 298.15 K using the enthalpy of reaction selected below. The constant determined in 0.05 M tetraethylammonium perchlorate [83SOL/BON] was also considered (but with a somewhat higher uncertainty). Although there are notable differences between the values published by different authors, the data in Figure V-20 show acceptable consistency. We should note, however, that most of these data were reported by the Mironov’s group [83AVR/BLO], [81KUL/BLO], [76KUL/BLO]. Only two of the five constants reported by other laboratories are consistent with the data of Mironov et al. The weighted linear regression using 11 data points yielded the selected value of:

10 1log οb ((V.70), 298.15 K) = (1.43 ± 0.08).

This value agrees well with that published in [81KUL/BLO] calculated by means of the Vasil’ev equation. The resulting ∆ε(V.70) value is – (0.049 ± 0.060) kg·mol–1. Using the selected values for ε(Ni2+, 4ClO− ) and ε(Na+, F–), ∆ε(V.70) leads to a value of ε(NiF+, 4ClO− ) = (0.34 ± 0.08) kg·mol–1.

0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.00.8

1.0

1.2

1.4

1.6

1.8

2.0

log 10

β1 +

4D

I / mol·kg−1

Ni2+ + F– NiF+


From the above formation constant r mGο∆ ((V.70), 298.15 K) = – (8.2 ± 0.5) kJ·mol–1 can be calculated. The Gibbs energy of formation of NiF+ is derived using the selected values for Ni2+ and F–. The selected value is:

f mGο∆ (NiF+, 298.15 K) = – (335.5 ± 1.1) kJ·mol–1.

The formation of higher complexes ( 2NiF nn

− , n > 1) is not reported in the litera-ture, not even in presence of more than a thousand-fold excess of fluoride over Ni(II) [71BON].

Reaction enthalpies for the formation of the NiF+ complex are collected in Table V-9. All data in Table V-9 except those of [74ARU] were determined in Mi-ronov’s laboratory. In order to determine r mH ο∆ ((V.70), 298.15 K), the SIT equation adapted to enthalpy format [97GRE/PLY2] has been used (cf. Figure V-21). The values at I = 0.1 and 0.25 M [81KUL/BLO] were not considered, as the amount of NiF+ formed is probably too small to yield reliable enthalpies. The value reported by Aruga was also neglected for the SIT, due to the different ionic medium used. The analysis of the remaining four points resulted in the selected standard reaction enthalpy:

r mH ο∆ (V.70) = (9.5 ± 3.0) kJ·mol–1

and ∆εL(V.70) = – (5 ± 35) × 10–5 kg·K–1·mol–1. The former value is fundamentally dif-ferent from that reported in [81KUL/BLO] ( r mH ο∆ ((V.70), 298.15 K) = – (2.9 ± 2.5) kJ·mol–1; the method of extrapolation to I = 0 is not reported). It agrees well, however, with the general observation that the stabilities of most aqueous monofluoride com-plexes are governed predominantly by electrostatic interactions and thus are entropy controlled [74HEF], [90GRZ].

Table V-9: Experimental enthalpy values for Reaction (V.70).

Method Medium t /°C r mH∆ (kJ·mol–1)reported

r mH∆ (kJ·mol–1)accepted

Reference

cal 0.1 M NaClO4 25 0 [81KUL/BLO] cal 0.25 M NaClO4 25 (3.3 ± 1.6) [81KUL/BLO] cal 0.5 M NaClO4 25 (6.7 ± 0.8) (6.7 ± 2.0) [81KUL/BLO] cal 0.5 M self(a) 25 (8.1 ± 0.2)(b) (4.4 ± 2.0)(c) [74ARU] cal 1.0 M NaClO4 25 (7.9 ± 1.2) (7.9 ± 2.0) [81KUL/BLO] cal 2.0 M NaClO4 25 (7.9 ± 1.6) (7.9 ± 2.0) [81KUL/BLO] cal 3.0 M NaClO4 25 (6.5 ± 2.0) (6.5 ± 2.0) Mironov et. al.(d)

(a) The solution of 0.17 M Ni(NO3)2 and 0.5 M NaF were mixed in different ratios. (b) Calculated from 10 1log b = 0.31, which was derived from the datum in [72BON/HEF]. (c) Recalculated data using the selected value for 10 1log οb ((V.70), 298.15 K) and ∆ε (V.70). (d) Average value from the studies reported by Mironov’s group [76KUL/BLO] and [81KUL/BLO].


From the above r mH ο∆ value, the following selected standard enthalpy of for-mation can be derived for NiF+:

f mH ο∆ (NiF+, 298.15 K) = – (380.9 ± 3.2) kJ·mol–1.

The above selections result in:

mS ο (NiF+, 298.15 K) = – (86.4 ± 10.3) J·K–1·mol–1.

Figure V-21: Extrapolation to Im = 0 of the experimental enthalpies for Reaction (V.70) in NaClO4 media from [76KUL/BLO] and [81KUL/BLO] (squares) and [74ARU] (tri-angle). The data marked with open square were not considered in the SIT treatment, since the amount of NiF+ formed is too small. According to [97GRE/PLY2],

Ψ(Im) = 23 ( )

4 1 1.5L m

m

A IzI

∆− ×

+

where ∆(z2) is the stoichiometric sum of the charge numbers squared of the reactants (in the present case – 4), AL is the Debye-Hückel parameter for the enthalpy, at 298.15 K and 1 bar AL = 1.986 kJ·kg1/2·mol–3/2.

0

2

4

6

8

10

12

14

16

0 0.5 1 1.5 2 2.5 3 3.5 4

I / mol·kg−1

∆ rH

m−Ψ

(Im

) / k

J·mol

−1

Ni2+ + F– NiF+


V.4.2.4 Aqueous Ni(II) - chloro complexes

The formation of Ni(II)-chloro complexes has been the subject of numerous investiga-tions using several experimental methods (cf. Table V-10). Since in most of these stud-ies the composition of the ionic medium was varied over a wide range, and no attention was paid to the medium effect, the literature reports on the complex formation are rather divergent concerning both the reported formation constants and stoichiometry of the formed complexes. Beside the formation of NiCl+ and NiCl2(aq) complexes, the pres-ence of other species is also proposed in aqueous solution at room temperature: e.g., Ni2Cl3+ [60LIS/ROS], 3NiCl− [64GRI/LIB] or 2

4NiCl − [68AND/KHA], [71KLY/GLE], [85SAP/CHI]. This reflects the fact that conventional methods are not always applicable to determine the stability of weak metal complexes if a high concentration of the ligand is needed to achieve a sufficient degree of complex formation. The majority of the ex-periments have shown that no anionic complexes are formed even at very high chloride concentrations. On the other hand, the formation of Ni2Cl3+ reported in [60LIS/ROS] was reconsidered and finally rejected in a subsequent paper of the authors [66KEN/LIS]. Therefore in this review only the formation constants for NiCl+ and NiCl2(aq) complexes are considered. The formation constants reported for Reaction (V.71),

Ni2+ + Cl– NiCl+ (V.71)

are listed in Table V-10. For reasons mentioned in Appendix A we do not consider the data reported in [57KIV/LUO], [58TRE], [62TRI/CAL], [65MOR/REE], [74GRA/WIL] and [76MUR/KUR]. Most of the accepted data are, however, also sub-ject to substantial experimental errors, due to the medium effect, and in such cases we assigned significantly higher uncertainty to the selected constants than reported in the original publications. The data in [75LIB/TIA] are free of significant medium effects, and this is the only data set where the systematic errors can be assumed identical for each point. Therefore, these data were used to determine the ion interaction coefficient between NiCl+ and 4ClO− , in spite of the fact that the applied ionic strength (Im = 3 – 9 m) is well above of the recommended range for the SIT analysis.

The plot ( 10 1log b + 4D) versus Im is depicted in Figure V-22. The weighted linear regression of data in Figure V-22 results in the values of 10 1log οb (V.71) = − (0.37 ± 0.27) and ∆ε(V.71) = – (0.073 ± 0.040) kg·mol–1. Using the selected values for ε(Ni2+, 4ClO− ), ε(Ni2+,Cl–), (in the absence of data to allow calculation of ε(Ni2+,Cl−) while also considering association; see, for example, Section V.6.1.2), ∆ε(V.71) leads to a value of:

ε(NiCl+, 4ClO− ) = (0.47 ± 0.06) kg·mol–1.

There is some uncertainty concerning this calculation. In [75LIB/TIA] Ni(ClO4)2 was used as a constant ionic medium (with a chloride concentration of ~ 0.01 m), therefore, ∆ε can be calculated as ∆ε = ε(NiCl+, 4ClO− ) – ε(Ni2+,Cl–) – ε(Ni2+, 4ClO− ).


In Appendix B.2 of previous TDB reviews it is mentioned that since the formation of chloride complexes is taken into account, the value of ε(Ni2+, Cl–) as listed, for example, in Appendix B of [92GRE/FUG], [2001LEM/FUG] should be replaced by ε(NiCl+, 4ClO− ) in all calculations when chloride is part of the ionic medium. This listed value of ε(Ni2+,Cl–) was obtained [80CIA] by neglecting association of nickel and chlo-ride. Following this process, and using ε(Ni2+, 4ClO− ) = (0.37 ± 0.03) kg·mol–1 (see Sec-tion V.4.3), from the above determined ∆ε = – (0.073 ± 0.04), ε(NiCl+, 4ClO− ) = (0.667 ± 0.06) kg·mol–1 can be calculated. This value of ε(NiCl+, 4ClO− ) is too high, taking into account the relatively accurate value for ε(NiF+, 4ClO− ).

In this review, ε(NiCl+, 4ClO− ) = (0.47 ± 0.06) kg mol–1 is used to determine ∆ε in other media to calculate 10 1log οb values from the rest of the data in Table V-10.

According to Appendix B.2, the same treatment should be applied, e.g., for 3NO− , i.e., ε(Mx+, 3NO− ) should be replaced by ε(Mx+, 4ClO− ) if nitrate is part of the me-

dium, but this recommendation is often disregarded.

Using the remaining experimental values for NaClO4 media (from Table V-10) and ε(Ni2+, 4ClO− ) = 0.37 to determine ε(NiCl+, 4ClO− ), from the equation ∆ε = ε(NiCl+, 4ClO− ) – ε(Na+, Cl–) – ε(Ni2+, 4ClO− ), ε(NiCl+, 4ClO− ) = 0.51 kg·mol–1 can be obtained. This may support the value calculated using ε(Ni2+, Cl–) from the data in [75LIB/TIA], however, this is not particularly convincing because most of the remain-ing data in the other papers have relatively low accuracy, due to substantial medium effects.

Table V-10: Experimental formation constants (logarithmic values) for the NiCl+ spe-cies.

Method Ionic medium t (°C) 10 1log b (a) 10 1log b (b) 10 1log οb (c) Reference

pol 2 M Na(ClO4, Cl) 25 – (0.25 ± 0.15) [57KIV/LUO]

cix 1.5 M (Ca, Ni) (NO3, Cl)

? – (0.66 ± 0.01) [58TRE]

cry 0.048 m KClO4 – 0.15 0.62 [59KEN]

cry 0.258 m KClO3 – 0.8 0.23 (0.23 ± 0.30) (0.79 ± 0.40) [59KEN]

ise-Cl 2 M (Na, Ni)ClO4 12 25 40

– 0.17 – 0.15

– (0.25 ± 0.20)(d) – (0.21 ± 0.20)(d) – (0.18 ± 0.20)(d)

(0.84 ± 0.30) (0.88 ± 0.30) (0.91 ± 0.30)

[60LIS/ROS]

cry 1.3 M K(NO3/Cl) – 2 (0.38 ± 0.08) (0.34 ± 0.30)(d) (1.14 ± 0.40) [62FAU/CRE]

sp 1.5 M Na(ClO4, Cl) ? – (0.85 ± 0.25) [62TRI/CAL]

sp 5.7 M H(Cl, ClO4) 25 – (0.5 ± 0.2) – (0.62 ± 0.40) (0.16 ± 0.50) [63NET/DRO]

cix 3 M H(ClO4, Cl) 25 – (0.64 ± 0.06) – (0.51 ± 0.50)(d) (0.35 ± 0.60) [64GRI/LIB]

cix 0.69 H(ClO4, Cl) 20 (0.23 ± 0.03) [65MOR/REE]




Method Ionic medium t (°C) 10 1log b (a) 10 1log b (b) 10 1log οb (c) Reference

sp 10 M Li(NO3, Cl) 30 – (1.02 ± 0.05) – (1.17 ± 0.50) [66FLO]

sol 4 M Na(ClO4, Cl) 25 – 0.2(e) – (0.71 ± 0.30)(d) (0.68 ± 0.40) [70HAL/VAN] (0.18 ± 0.50)(f)

sp 3 M Li(ClO4, Cl) 25 – (0.57 ± 0.03) – (0.63 ± 0.40) (0.36 ± 0.50) [70MIR/MAK]

pol 1 M Na(ClO4, Cl) 25 – (0.22 ± 0.08) – (0.24 ± 0.40) (0.66 ± 0.50) [71BON]

sp 6 M (Na, Li)(ClO4, Cl)0 corr(g)

25 – (0.5 ± 0.1) – (1.30 ± 0.09)

[71PAA/HUM]

kin 1 M Na(ClO4, Cl) 25 (0.07 ± 0.08) (0.00 ± 0.40)(d) (0.90 ± 0.50) [73HUT/HIG]

ise-Cl 1 M NaClO4 + 0.1 M HClO4

25 (0.00 ± 0.06) – (0.02 ± 0.30) (0.90 ± 0.40) [74BIX/LAR]

ise-Cl 3 M Na(ClO4,Cl) 25 (0.69 ± 0.04) [74GRA/WIL]

ise-Cl

1 m Ni(ClO4)2 1.5 m Ni(ClO4)2 2 m Ni(ClO4)2 2.5 m Ni(ClO4)2 3 m Ni(ClO4)2

25

– 1.1 – 1.1 – 1.0 – 0.92 – 0.80

– (1.10 ± 0.20) – (1.10 ± 0.20) – (1.00 ± 0.20) – (0.92 ± 0.20) – (0.80 ± 0.20)

– (0.37 ± 0.27) [75LIB/TIA]

cix 1 M Na(ClO4, Cl) 25 0.0 [76MUR/KUR]

sp/gl 3 M Na(ClO4, Cl) 25 – (0.50 ± 0.13) – (0.57 ± 0.40) (0.68 ± 0.50) [78FOR]

nmr 4 m KCl 25 (h) – (1.26 ± 1.00)(d) (0.18 ± 1.00) [79WEI/HER]

sp 0 corr(i) 25 – (1.90 ± 0.10) [88BJE]

sol

6 m Na(ClO4, Cl) 0 corr(j)

25

– (0.69 ± 0.08) – 0.83

– (0.69 ± 0.20)

(0.65 ± 0.30) [89IUL/POR]

(a) Reported values. (b) Accepted values corrected to molal scale. The accepted values reported in Appendix A are expressed on

the molar or molal scales, depending on which units were used originally by the authors. (c) Corrected to I = 0 by means of the SIT, using the selected ion interaction coefficients. (d) Re-evealuated value, see Appendix A. (e) The authors suggested solely the formation of NiCl2(aq) complex, but they also calculated the stabil-

ity constant of NiCl+ assuming its unique formation. (f) Re-evaluated value taking into account of the effect of medium changes, see discussion on

[70HAL/VAN]. (g) ‘Semi-thermodynamic’ constant, in which the activity of chloride (using γ± of the complex forming

electrolyte) and water were taken into account. (h) The average number of bound chloride per Ni(II) is reported. (i) ‘Semi-thermodynamic’ constant, in which the activity of chloride (using γ± of the complex forming

electrolyte) was taken into account. (j) Corrected to I = 0 by means of the SIT, using estimated ion interaction coefficients.


Figure V-22: Extrapolation to I = 0 of the 10 1log b values reported in [75LIB/TIA] for Reaction (V.71) in Ni(II) perchlorate solution. The solid line is calculated with

10 1log οb = – (0.37 ± 0.27), and ∆ε = – (0.073 ± 0.040).

In [89IUL/POR] the authors proposed ε(NiCl+, 4ClO− ) = 0.202 kg·mol–1, esti-mated from ε(Ni2+, 4ClO− ) = 0.375 kg·mol–1 and ε(Na+, Cl–) using Equation (B.22) (see Appendix B). Although Iuliano’s value is considerably lower, we recommend the value derived in the present review for ε(NiCl+, 4ClO− ), since it is based on experimental evi-dence.

All accepted constants for NaClO4 media are within the ionic strength range Im = 0 – 6 m. The SIT representation of these values is depicted in Figure V-23. Using the ion interaction coefficient ε(Ni2+, 4ClO− ) together with the above determined value of ε(NiCl+, 4ClO− ), ∆ε = (0.07 ± 0.10) kg·mol–1 can be calculated for Reaction (V.71) in NaClO4 medium. From the plot of experimental data in Figure V-23 a very similar value, ∆ε = (0.11 ± 0.06) kg·mol–1 can be obtained. However, this value must be treated with caution, due to the already mentioned medium change, which affects most of the data points in Figure V-23. Therefore, ∆ε((V.71), NaClO4) = (0.07 ± 0.10) kg·mol–1 was used to calculate the 10 1log οb values for NaClO4 media given in Table V-10. The SIT extrapolation resulted in somewhat lower 10 1log οb values in LiClO4 (∆ε((V.71),

-0.8

-0.6

-0.4

-0.2

0

0.2

0.4

0.6

0.8

1

0 1 2 3 4 5 6 7 8 9 10

I / mol·kg−1

log 1

0 β1 +

4D

Ni2+ + Cl– NiCl+


LiClO4) = (0.0 ± 0.1) kg·mol–1) and HClO4 (∆ε((V.71), HClO4) = – (0.04 ± 0.10) kg·mol−1) media.

Figure V-23: Extrapolation to I = 0 of the accepted experimental data for Reaction (V.71) in NaClO4 media (see text).

These extrapolated values are considerably higher than the value determined from the data in [75LIB/TIA]. This is probably a result of the disregarded medium changes. Obviously, the attribution of the global effect (complex formation and changes in activity coefficients) to complex formation alone results in higher constants.

The well documented experimental results in [70HAL/VAN] offered the pos-sibility to use the SIT to assess the medium effect caused by the replacement of the original background electrolyte (see Appendix A). In this work the background medium was changed gradually from 4 M NaClO4 to 2 M NaClO4/2 M NaCl. As explained in the discussion on [70HAL/VAN], the modified SIT treatment, in which the changes in ionic medium were taken into account, resulted in a value approximately half a unit lower in the 10 1log οb value (0.20 ± 0.40) than obtained by neglecting this effect.

The medium changes, caused by the replacement of 4ClO− by Cl– ions, also were accounted for by SIT in [89IUL/POR]. Due to the different ε(NiCl+, 4ClO− ) value

0 1 2 3 4 5 6 7

0.0

0.5

1.0

1.5

log 10

β1 +

4D

I / mol·kg−1

Ni2+ + Cl– NiCl+


used (see Appendix A), however, the authors reported a considerably lower 10 1log οb than was recalculated in this review from their data.

Bjerrum [88BJE] reported a ‘semi-thermodynamic’ formation constant ( '

10 1log οb ), which is much lower than any other 10 1log οb value in Table V-10. Due to the unknown ratio of 2 +Ni NiCl

/γ γ+ (see Appendix A), this constant is not compatible with the assumptions of the SIT model. Since Bjerrum applied a spectrophotometric method using the d-d transitions of Ni(II) and very high chloride concentrations, his constant probably refers to the formation of the inner-sphere +

2 5NiCl(H O) complex ( ,inner

10 1log οb ). Libus and Tialowska estimated a similar value for ,inner10 1log οb (– 2.0,

[75LIB/TIA]). Bjerrum’s value is also within the uncertainty range of the ‘semi-thermodynamic’ formation constant derived from the data reported in [79WEI/HER] ( '

10 1log οb = – (1.03 ± 1.00), see Appendix A). However, the NMR relaxation measure-ments in [79WEI/HER] also demonstrate the incompatibility of '

10 1log οb as defined by Bjerrum (see [88BJE] and [89BJE]) and the 10log οb values extracted by means of the SIT analysis. Using the recommended ion interaction coefficients to calculate ∆ε(V.71) for KCl medium, 10 1log οb ((V.71), 298.15 K) = (0.16 ± 1.00) is derived from the ex-perimental data in [79WEI/HER], which is considerably higher than Bjerrum’s

'10 1log οb value.

Due to the medium effect, most of the constants listed in Table V-10 represent only the upper limits of the true values. Therefore, 10 1log οb ((V.71), 298.15 K) is ob-tained as the average of the following two data: – (0.37 ± 0.27) (determined from the data in [75LIB/TIA]) and (0.52 ± 0.38) (average of the remaining 10 1log οb values re-ported in Table V-10 for 298 K in (H/Li/Na)ClO4 and KCl media). This treatment re-sulted in the selected value:

10 1log οb ((V.71), 298.15 K) = (0.08 ± 0.60).

From this value, r mGο∆ ((V.71), 298.15 K) = – (0.5 ± 3.4) kJ·mol–1 can be de-rived. The Gibbs energy of formation of NiCl+ is calculated using the selected values for Ni2+ and Cl–.

f mGο∆ (NiCl+, 298.15 K) = – (177.4 ± 3.5) kJ·mol–1.

Several authors reported equilibrium constants for Reaction (V.72), which are collected in Table V-11:

NiCl+ + Cl– NiCl2(aq). (V.72)

For reasons mentioned in Appendix A, and since in cases in which there are considerable changes in the ionic medium it is impossible to distinguish between a me-dium effect and the formation of higher complexes, we do not consider the data reported in [57KIV/LUO], [65MOR/REE] and [70HAL/VAN]. ‘Semi-thermodynamic’ constants are reported for Reaction (V.72) in [71PAA/HUM] and [88BJE] (see Appendix A), based on a spectrophotometric method using the d-d transitions of Ni(II) and very high


chloride concentrations ([Cl–] ≥ 9 M). Therefore, these constants most probably refer to the formation of the inner-sphere complex NiCl2(H2O)4(aq). Iuliano et al. used solubil-ity measurements to determine the equilibrium constant for Reaction (V.72). They re-ported a higher value for K2 than for K1, which is rather unlikely considering the low affinity of Ni(II) for chloride. Therefore, this review does not find it justified to select a recommended value for K2 (V.72).

Table V-11: Experimental equilibrium constants (logarithmic values) for Reaction (V.72).

Method Ionic medium t /°C 10 2log K (a) Reference

pol 2 M Na(ClO4, Cl) 25 (0.20 ± 0.15) [57KIV/LUO] cix 0.69 H(ClO4, Cl) 20 – (0.27 ± 0.05) [65MOR/REE] sol(b) 4 M Na(ClO4, Cl) 25 – (0.89 ± 0.01) [70HAL/VAN] sp 0 corr(c) 25 – (2.8 ± 0.3) [71PAA/HUM] sp 0 corr(d) 25 – (2.3 ± 0.6) [88BJE] sol 6 M Na(ClO4, Cl)

0 corr(e) 25 – (0.5 ± 0.2)

– 0.3 [89IUL/POR]

(a) Reported values. (b) Corrected to I = 0 by means of the SIT, using the selected ion interaction coefficients. (c) ‘Semi-thermodynamic’ constant, in which the activity of chloride (using γ± of the complex form-

ing electrolyte) and water were taken into account. (d) ‘Semi-thermodynamic’ constant, in which the activity of chloride (using γ± of the complex form-

ing electrolyte) was taken into account. (e) Corrected to I = 0 by means of the SIT, using estimated ion interaction coefficients.

Higher chloro complexes 2NiCl nn

− with n = 3 – 4 readily form in several non-aqueous solutions [59GIL/NYH]. As was mentioned previously, such complexes are unlikely to be present in detectable concentrations in aqueous solution at room tempera-ture. The 10 4log b in aqueous solution was estimated to be – 14 in [76BIA/CAR] (see Appendix A). Accepting this value, the calculated concentration of 2

4NiCl − is negligi-ble, even in highly concentrated chloride solutions. In aqueous solutions such com-plexes can be only expected at elevated temperature in melts [66ANG/GRU] or under hydrothermal conditions [96UCH/GOR]. The only exception is reported in [69GRI/SCA]. Using a concentrated aqueous solution of (CH3)4NCl ( > 5 M) as back-ground electrolyte, the formation of tetrahedral 2

4NiCl − complex was detected by UV-VIS spectrophotometry at as low a temperature as 50°C. This exceptional behaviour was attributed to the specific effect of tetramethylammonium cation, which reduces the hydration of chloride anion enhancing its coordination ability [69GRI/SCA].


Only a few studies have reported reaction enthalpies for the formation of NiCl+ species (Table V-12), and such data are not available for NiCl2(aq). Only the experi-mental results reported in [66KEN/LIS] for 298 K could be re-evaluated using

10 1log οb ((V.71), 298.15 K) and ∆ε((V.71), NaClO4), therefore no value is selected in the present review.

Table V-12: Reaction enthalpies reported for Reaction (V.71).

Method Medium t (°C) r mH∆ (kJ·mol–1)(a)r mH∆ (kJ·mol–1)(b) Reference

cal 2 M (Na, Ni)ClO4 25 (2.1 ± 0.2)(c) (8.2 ± 2.5)(d) [66KEN/LIS]

cal 2 M (Na, Ni)ClO4 40 (1.8 ± 0.2)(e) [66KEN/LIS]

cal 3 M Li(ClO4/Cl) 25 (9.6 ± 0.4) [74BLO/RAZ]

(a) Reported value. (b) Accepted value. (c) Calculated using 10 1log b = – 0.17 (25°C) and 10 1log b = – 0.15 (40°C) from [60LIS/ROS]. (d) Re-evaluated value, see Appendix A. (e) Calculated using 10 1log b = – 0.57 from [70MIR/MAK].

V.4.2.5 Aqueous Ni(II) – bromo complexes

Only a few studies are available concerning the complex formation between Ni(II) and bromide ions. Most of the studies reported the formation of a monobromo complex (NiBr+), but at very high concentrations of the bromide ion the formation of bis- [36JOB], [72RET/HUM], [90BJE] and tetrakis- bromo complexes [36JOB], [68AND/KHA2] are also reported. The graphical data presented in [36JOB] were later re-evaluated in [72RET/HUM], and were found to be consistent with the formation of NiBr+ and NiBr2(aq) complexes (see comments on [36JOB]). The qualitative work in [68AND/KHA2] using spectrophotometry may well suggest formation of 5 – 10% of the tetrakis - bromo complex, but only in extremely concentrated (18 M) HBr solutions (see Appendix A). Nevertheless, the assessment of values for formation of negatively charged complexes in aqueous solutions is not justified.

The available quantitative information on the formation of NiBr+, according to Reaction (V.73), is collected in Table V-13.

Ni2+ + Br– NiBr+ (V.73)

There are large differences between the values listed in Table V-13. This is the consequence of the inherent difficulties that occur when formation constants of weak complexes are determined (see also Sections V.4.2.1 and V.4.2.4). The values reported in [63NET/DRO], [70MIR/MAK] and [73HUT/HIG] are of low precision, due to a con-siderable medium effect. The ‘semi-thermodynamic’ formation constant reported in [90BJE] probably gives an exact description of complex formation in 3.5 – 11 M LiBr,


but this constant is not compatible with the assumptions of the SIT model (see also Sec-tions V.4.2.1 and V.4.2.4 and [89BJE]). The most reliable data set is given in [78LIB/KOW]. The SIT representation of the data reported in [78LIB/KOW] is de-picted in Figure V-24. Although the ionic strengths in the experiments of [78LIB/KOW] are beyond the optimal range for SIT, the experiments are relatively free of interfering medium effects. Therefore, the estimation of ε(NiBr+, 4ClO− ) has been based on this data set. This probably leads to a more accurate value than reliance on the other experi-mental values in Table V-13, or on Equation (B.22). From the plot in Figure V-24, ∆ε(V.73) = – (0.046 ± 0.09) kg·mol–1 and 10 1log οb = − (0.5 ± 0.6) are obtained. Using the recommended values for ε(Ni2+, Br–) (in the absence of data to allow calculation of ε(Ni2+, Br–), while also considering association; see Section V.4.2.5.1) and ε(Ni2+, 4ClO− ),

ε(NiBr+, 4ClO− ) = (0.59 ± 0.10) kg·mol–1

can be derived. Using this unexpectedly large value, it is possible to estimate the ∆ε values for all background electrolytes listed in Table V-13 (∆ε((V.73), NaClO4) = (0.19 ± 0.11) kg·mol–1, ∆ε((V.73), HClO4) = (0.08 ± 0.11) kg·mol–1, ∆ε((V.73), LiClO4) = (0.07 ± 0.11) kg·mol–1, ∆ε((V.73), KCl) = (0.22 ± 0.11) kg·mol–1). These coefficients were used to calculate the 10 1log οb values given in Table V-13.

Table V-13: Experimental formation constants (logarithmic values) of the NiBr+ complex.

Method Ionic medium t /°C 10 1log b (a) 10 1log b (b) 10 1log οb (c) Reference

ise-Br 2 M (Na, Ni)ClO4 25(d) – (0.12 ± 0.02) – (0.16 ± 0.40) (1.20 ± 0.50) [61LIS/WIL] sp 5.7 M H(Br, ClO4) 25 – (0.30 ± 0.22) – (0.42 ± 0.50) (1.28 ± 0.60) [63NET/DRO] sp 3 M Li(ClO4, Br) 25 – (0.82 ± 0.03) – (0.89 ± 0.50) (0.35 ± 0.60) [70MIR/MAK] kin 1 M Na(ClO4, Br) 25 – (0.09 ± 0.04) – (0.11 ± 0.50) (0.89 ± 0.60) [73HUT/HIG] ise-Br

1.5 m Ni(ClO4)2 2 m Ni(ClO4)2 2.5 m Ni(ClO4)2 3 m Ni(ClO4)2

25

– 1.3 – 1.3 – 1.3 – 1.16

– (1.30 ± 0.30)– (1.30 ± 0.30)– (1.30 ± 0.30)– (1.16 ± 0.30)

– (0.50 ± 0.60) [78LIB/KOW]

nmr 4 m KBr 25 – (1.10 ± 0.60) (0.80 ± 0.60) [79WEI/HER] sp 0 corr(e) 25 – (2.37 ± 0.08) [90BJE]

(a) Reported values. (b) Accepted values corrected to the molal scale. The accepted values reported in Appendix A are ex-

pressed on the molar or molal scales, depending on which units were used originally by the authors. (c) Corrected to I = 0 by means of the SIT, using the recommended ion interaction coefficients. (d) Data for several temperatures, see Appendix A. (e) ‘Semi-thermodynamic’ constant, in which the activity of bromide (using γ± of the complex

forming electrolyte) was taken into account.


As already mentioned in previous sections, experiments in which there was a considerable medium effect lead to formation constants that represent only upper limits for the true values. Therefore, the weighted average of the following two values is cal-culated to assess the most probable formation constant at infinite dilution: − 0.5 ( 10 1log οb calculated from the data in [78LIB/KOW], weight = 2) and 0.90 (average of the remaining 10 1log οb values in Table V-13, weight = 1). This treatment yielded

10 1log οb ((V.73), 298.15 K) = – (0.03 ± 1.30). The uncertainty range was assigned to cover all the reported constants in Table V-13. Taking into account this large uncer-tainty, the above value can not be recommended, but can be used as the most probable value, until more precise data are published for Reaction (V.73).

Figure V-24: Extrapolation to I = 0 of the 10 1log b values reported in [78LIB/KOW] for Reaction (V.73) in Ni(II) perchlorate media.

The formation of the dibromo complex NiBr2(aq) is dominant only at very high bromide concentrations (8 – 12 M LiBr or HBr). Therefore, only ‘semi-thermodynamic’ stability constants ( '

10 2log K ο ) are available for this species. The reported values are as

-2

-1.5

-1

-0.5

0

0.5

1

1.5

2

0 1 2 3 4 5 6 7 8 9 10

I / mol·kg−1

log 1

0 β1 +

4D

Ni2+ + Br– NiBr+


follows: '10 2log K ο = – (2.97 ± 0.13) [90BJE], '

10 2log K ο = – (4.25 ± 0.06) [72RET/HUM] and '

10 2log K ο = – 3.7 (calculated from the data in [36JOB] by [72RET/HUM]). These data are not directly comparable since in the calculation of the last two values the activ-ity of water was also taken into account. Thus these values refer to the following equi-librium:

2 5NiBr(H O)+ + Br– NiBr2(H2O)4(aq) + H2O(l).

The higher value reported in [90BJE] is reasonable when the very low water activity in concentrated LiBr or HBr solutions is taken into account.

Two quantitative results are available for the enthalpy of Reaction (V.73) (Table V-14). The lack of experimental details and some unidentified experimental er-rors (see Appendix A), however, do not allow selection of accepted values for

r mH∆ (V.73).

Table V-14: Reaction enthalpies reported for Reaction (V.73).

Method Medium t /°C r mH∆ / kJ·mol–1 (a) Reference

ise-Br 2 M (Na, Ni)ClO4 25 (8 ± 5)(b) [61LIS/WIL]

cal 3 M Li(ClO4/Br) 25 (11.7 ± 0.4)(c) [74BLO/RAZ]

(a) Reported value. (b) Calculated using the reported temperature dependence of 10 1log b (V.73), see Appendix A. (c) Calculated using 10 1log b = – 0.82 from [70MIR/MAK].

V.4.2.5.1 Determination of the Ni2+ – Br– ion interaction coefficient

Because the value of 10 1log οb ((V.73), 298.15 K) given above is low and rather uncer-tain, the description of the Ni(II)-bromo system in terms of ion-ion interactions may be in many cases more straightforward. Therefore, the ion interaction coefficient ε(Ni2+,Br−) has also been derived in this review from the osmotic and mean activity co-efficients of NiBr2 solutions, using Equations (V.74) and (V.75), see Section V.4.3.

The only relevant experimental data are reported in [78LIB/MAC]. Goldberg et al. reported recommended values for activity and osmotic coefficients of NiBr2 solu-tions [79GOL/NUT], based on the data in [78LIB/MAC]. In [79GOL/NUT] a corrected data set for the experimentally determined osmotic coefficients was given. Reference was made to a personal communication from Z. Libus that provided a corrected version of the erroneous tables in [78LIB/MAC]. Therefore, only the activity coefficients were taken from [78LIB/MAC].


Plots of 10log γ ± and φ as a function of the molality, m, of NiBr2 (m = Im/3) are depicted in Figure V-25 and Figure V-26, respectively. The experimental data can be reasonably described in the concentration range of 0 – 3 m (Im = 0 – 9 m) using ε(Ni2+,Br–) = (0.268 ± 0.020) kg·mol–1 (see the solid lines in Figure V-25 and Figure V-26).

V.4.2.6 Aqueous Ni(II) – iodine complexes

V.4.2.6.1 Aqueous Ni(II) – iodo complexes

No quantitative data are available for the formation of 2NiI qq

− complexes. The formation of NiI+ and NiI2(aq) was postulated in [74KUT/GAL], to explain the polarographic pre-waves of Ni(II) reduction in 5 M Ca(NO3)2 solutions containing 0.2 – 1 M KI. The spec-trophotometric data reported in [68AND/KHA2] indicate formation of a considerable amount of 2

4NiI − in 9 – 11 M HI solutions (see Appendix A).

Figure V-25: Plot of 10log γ ± versus molality of aqueous NiBr2 solutions at 298.15 K. Solid line: SIT, ε(Ni2+, Br–) = 0.268; : Experimental data from [78LIB/MAC].

-0.4

-0.2

0

0.2

0.4

0.6

0.8

1

1.2

0 0.5 1 1.5 2 2.5 3 3.5 4 4.5

[NiBr2] / mol·kg−1

g 10

γ ±lo

g 10 γ ±


Figure V-26: Plot of osmotic coefficient versus molality of aqueous NiBr2 solutions at 298.15 K. Solid line: SIT, ε(Ni2+, Br–) = 0.268; : Experimental data from [78LIB/MAC], [79GOL/NUT].

V.4.3 Determination of the Ni2+ – −4ClO ion interaction coefficient

The ion interaction coefficients can be derived from the ionic strength dependence of the osmotic and mean activity coefficients of the corresponding electrolyte solutions. Two papers report relevant experimental data for Ni(ClO4)2 solutions. Libus and Sadowska used an isopiestic method to determine the osmotic coefficients in 0.1 – 3.5 m Ni(ClO4)2 solutions [69LIB/SAD], and the mean activity coefficients were calculated from the Gibbs-Duhem equation. Malatesta et al. determined the activity coefficients in 0.0001 – 0.9 m Ni(ClO4)2 solutions from the emf of liquid-membrane cells [99MAL/CAR]. Goldberg et al. in [79GOL/NUT] have also reported recommended values for activity and osmotic coefficients of Ni(ClO4)2 solutions, but their data are based on [69LIB/SAD].

The ion interaction coefficient ε(Ni2+, 4ClO− ) was previously derived in [89IUL/POR] (ε(Ni2+, 4ClO− ) = 0.375 kg·mol–1) from the experimental data reported in [69LIB/SAD]. Since no uncertainty assignment is given in [89IUL/POR], and in the meantime, another relevant data set has been published [99MAL/CAR], the value of ε(Ni2+, 4ClO− ) has been re-determined in the present review.

0.8

1.2

1.6

2

2.4

2.8

3.2

0 0.5 1 1.5 2 2.5 3 3.5 4 4.5

[NiBr2] / mol·kg−1

φ


According to the specific ion-interaction theory the mean activity coefficient γ ± of Ni(ClO4)2 can be expressed as (V.74):

10log γ ±2

4

2 4 (Ni ,ClO )91 1.5

mm

m

A II

I+ −= − + ε

+, (V.74)

while the osmotic coefficient φ reads:

( ) 243

2 ln(10) 1 2ln(10)1 1 1.5 2ln 1 1.5 ε (Ni ,ClO )91.5 1 1.5m m m

m m

A I I II I

+ −

= − + − − + + +

φ (V.75)

with A denoting the Debye-Hückel parameter.

In [69LIB/SAD] the osmotic coefficients have been determined by the isopi-estic method.

The value of ε(Ni2+, 4ClO− ) was recalculated based on the smoothed data of Ta-ble II in [69LIB/SAD] and using Equation (V.75). The limited validity of Equation (V.75) was taken into account by weighting schemes which reduce the influence of the data measured at higher ionic strength. So each data point was either not weighted for comparison or weighted by the inverse of its x value (I) and the inverse of its x2 value (I 2), respectively. In Figure V-27 it is shown that the φ data of [69LIB/SAD] can be reasonably reproduced by Equation (V.75) up to m(Ni2+, 4ClO− ) ≈ 2 mol·kg–1.

In [99MAL/CAR] the activity coefficients have been obtained by measuring the emf of liquid-membrane cells. In Table III of [99MAL/CAR] the estimated uncer-tainties are given along with the activity coefficients of Ni(ClO4)2. The value of ε(Ni2+, 4ClO− ) was recalculated using Equation (V.74). The statistically estimated uncer-tainty (± 0.004) seems to be rather low. In Figure V-28 experimental data 10log ±γ are plotted versus m(Ni(ClO4)2. They fit nicely into the SIT approach when an uncertainty of ± 0.03 for ε is provided for.

The mean value of the entirely independent investigations (emf measurements [99MAL/CAR] and isopiestic measurements [69LIB/SAD]) is selected:

ε(Ni2+, 4ClO− ) = (0.37 ± 0.03) kg·mol–1.

The uncertainty is rather pessimistic but takes into account that the weighting scheme influences the numerical value of ε(Ni2+, 4ClO− ) considerably.


Figure V-27: Osmotic coefficient plotted as a function of molality of aqueous Ni(ClO4)2 solutions at 298.15 K. Figure V-28: Plot of 10log ±γ vs. molality of aqueous Ni(ClO4)2 solutions at 298.15 K.

0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.00.5

1.0

1.5

2.0

2.5

3.0

3.5

weighting scheme ε (Ni2+, ClO4−)

none 0.393 x 0.375 x2 0.360

φ

[Ni(ClO4)

2] / mol·kg−1

smoothed data of [69LIB/SAD]

0.0 0.2 0.4 0.6 0.8 1.0-0.30

-0.25

-0.20

-0.15

-0.10

-0.05

0.00

0.05

log 10

γ ±

[Ni(ClO4)2] / mol·kg−1

data from [99MAL/CAR] ε = (0.36 ± 0.03) / kg·mol−1

V.5 Group 16 (chalcogen) compounds and complexes 161

V.5 Group 16 (chalcogen) compounds and complexes V.5.1 Sulphur compounds and complexes V.5.1.1 Nickel sulphides

V.5.1.1.1 Solid nickel sulphides

V.5.1.1.1.1 Crystallography and mineralogy of nickel sulphides

V.5.1.1.1.1.1 Heazlewoodite, Ni3S2(cr)

Ni3S2(cr) can be synthesised by fusing nickel and sulphur (3 Ni : 2 S) in an evacuated silica glass tube. The crystal structure of this synthetic product was determined by Pea-cock [47PEA]. It is trigonal, space group: R32, with unit cell dimensions a0 = 5.741 Å, c0 = 7.139 Å, Z = 1 (for rhombohedral cell a0 = 4.080 Å, α = 89°26’), ρ(calc.) = 5.87 g·cm–3; ρ(exp.) = 5.82 g·cm–3 (according to JCPDS-ICDD card No.8-126). Fleet refined the crystal structure of artificial α-Ni3S2 using X-ray powder intensity data, and essen-tially confirmed Peacock’s results a0 = 5.7465 Å, c0 = 7.1349 Å, Z = 1, ρ(calc.) = 5.86 g·cm–3 (according to JCPDS-ICDD card No.30-863) [77FLE].

Heazlewoodite is the natural form of the compound Ni3S2(cr) found intergrown with magnetite, Fe3O4, in narrow bands in the serpentine at Heazlewood, Tasmania. It occurs massive and granular, without crystal form or cleavage; it is characterised by the following properties: fracture even; brittle, hardness 4, lustre metallic, colour pale yel-lowish bronze, streak light bronze, non-magnetic, opaque.

V.5.1.1.1.1.2 α-Ni7S6

This Ni-sulphide phase is orthorhombic with a0 = 3.274 Å, b0 = 16.157 Å, c0 = 11.359 Å, space group: Bmmb. Based on ρ(exp.) = 5.36 g·cm–3, the unit-cell content is 22.5 Ni and 19.3 S [72FLE]. There are three non-equivalent S sites and five non-equivalent Ni sites per unit cell. Four of the Ni sites are in square pyramidal coordina-tion with S and one is in tetrahedral coordination. Each Ni site is related to at least one neighbouring Ni site by an interatomic distance similar to that of metallic Ni (2.492 Å). It is apparent that metallic (Ni-Ni) bonding has been significant in stabilising this struc-ture.

V.5.1.1.1.1.3 Pentlandite, (Fe, Ni)9S8(cr)

Pentlandite (JPCRD card 8-90, [52ROB/BRO]) is an important ore of nickel. It is com-monly associated with other sulphides such as pyrite (FeS2), chalcopyrite, (CuFeS2) and pyrrhotite (Fe1–xS) in basic igneous rock intrusions. The largest deposit of this important ore of nickel is at Sudbury, Ontario. Important deposits are mined in Manitoba, Russia, and Western Australia. However, the pure end-member, Ni9S8, of pentlandite does not seem to occur in nature and has not been synthesised [87FLE], thus no thermodynamic data are proposed for this phase in the present review.


V.5.1.1.1.1.4 Godlevskite, Ni9S8(cr)

Godlevskite forms light yellow microscopic crystals with metallic lustre and gray streak. The hardness is 4 – 5. Fleet showed that godlevskite has a crystal structure based on a distorted cubic close-packed array of 32 S atoms per unit cell, with 20 Ni atoms in tetrahedral coordination and 16 in square-pyramidal coordination [87FLE]. The ideal stoichiometry was established to be Ni9S8(cr). The empirical formula of the specimen was (Ni8.7Fe0.3)S8, the unit cell is orthorhombic, characterised by a0 = 9.3359 Å, b0 = 11.2185 Å, c0 = 9.4300 Å, Z = 4, space group: C222, ρ(calc.) = 5.273 g·cm–3.

Originally the X-ray single-crystal and powder diffraction patterns of godlevskite were reported to be consistent with the space groups C222, Cmm2, Amm2, and Cmmm, with a0 = 9.18 Å, b0 = 11.29 Å, c0 = 9.47 Å, and the stoichiometric formula was assumed to agree with β-Ni7S6 (according to JCPDS-ICDD card No. 8-107) [69KUL/ERS].

V.5.1.1.1.1.5 Millerite, β-NiS

Millerite is a low temperature hydrothermal mineral found in cavities in carbonate rocks and as an alteration product of other nickel minerals. The high temperature form α-NiS has not been found in nature. The β-NiS phase is rhombohedral with hexagonal axes a0 = 9.620 Å; c0 = 3.149 Å, space group: R3m, Z = 9, ρ(calc.) = 5.374 g·cm–3 (according to JCPDS-ICDD card No.12-41). The crystal structure of β-NiS has been refined by Grice and Ferguson [74GRI/FER] as well as Rajamani and Prewitt [74RAJ/PRE]. The colour of the mineral is brassy yellow, its luster is metallic, crystals are opaque, fracture is un-even, hardness is 3 – 3.5, ρ(exp.) = 5.3 – 5.5 g·cm–3, the streak is dark green to almost black. Associated minerals are calcite, CaCO3, quartz, SiO2, dolomite, CaMg(CO3)2, bravoite (basically a nickel-rich pyrite), Fe0.7Ni0.2Co0.1S2, chalcopyrite, CuFeS2, grossu-lar, Ca3Al2(SiO4)3, fluorite, CaF2, and pyrrhotite, Fe1–xS. Notable occurrences include south central Indiana, Keokuk, Iowa, Gap Mine, Pennsylvania and Sterling Mine, New York, USA; Freiberg, Germany; Glamorgan, Wales, England and Sherbrooke and Planet Mines, Quebec.

V.5.1.1.1.1.6 Polydymite, Ni3S4(cr)

Gricaenko et al. [53GRI/SLU] investigated a sample from the Ural mountains, Russia, with respect to its X-ray powder patterns and synthesised polydymite, Ni3S4, from aqueous solution. The d values of the natural mineral and the artificial product agreed with each other. Polydymite has a spinel structure, the dimension of the cubic unit cell is a0 = 9.48 Å, Z = 8, space group: Fd3m, ρ(calc.) = 4.83 (4.75) g·cm–3 (the numerical value in parenthesis agrees with Ni3S4), ρ(exp.) = 4.81 g·cm–3 (according to JCPDS-ICDD card No. 8-107).

The colour of the mineral is violet grey, copper red, light grey, or steel grey, lustre is metallic, its hardness is 4.5 – 5.5, its cleavage [001] is indistinct, the streak is black grey, the crystals are opaque. Polydymite occurs in Heazlewood, Tasmania, Aus-


tralia, Port Radium, Great Bear Lake, Canada, and in various sites in Germany and USA. Examples for the latter localities are the Grüne Au Mine, Siegen District, Rhine-land-Palatinate and the Gold Hill district, Boulder, Colorado.

V.5.1.1.1.1.7 Vaesite, NiS2(cr)

The synthesis and X-ray diffraction pattern of NiS2(cr) has been reported by de Jong and Willems [27JON/WIL]. This nickel sulphide has a pyrite, FeS2, type structure. Ma-terial approaching NiS2(cr) in composition and having a pyrite like lattice has been found in the Kasompi mine of Zaire (formerly the Belgian Congo) by Johannes Vaes of the Union Minière du Haut Katanga, whom it was named after. Vaes was an important contributor to the mineral development of Zaire.

Two specimens close to the composition NiS2(cr), but containing minor amounts of Fe or Fe and Co, were examined microscopically, by an X-ray study, and analysed chemically [45KER]. The crystal structure is cubic, space group PA3, with lattice constant a0 = 5.670 Å, Z = 4, ρ(calc.) = 4.45 g·cm–3 (according to JCPDS-ICDD card No. 11-99). The density listed in this card has been calculated for the natural min-eral. For pure NiS2(cr) a slightly different value, ρ(calc.) = 4.476 g·cm–3, would be ob-tained with a0 = 5.670 Å. The colour of the mineral is grey, its lustre is metallic, the crystals are opaque, its cleavage [001] is perfect. Vaesite occurs with pyrite randomly disseminated in dolomite. The type locality is the Kasompi mine, Katanga, Zaire.

V.5.1.1.1.2 Ni – S phase diagram

The phase diagram of the binary system Ni – S, given by Elliott [65ELL], is based on a systematic investigation of the phase relationships in this system performed by Kullerud and Yund [62KUL/YUN] as well as Sokolova [56SOK]. Craig and Scott [76CRA/SCO] presented a phase diagram which is in accordance with experimental studies of Kullerud and Yund [62KUL/YUN] as well as Arnold and Malik [75ARN/MAL]. Apart from the elements the following mixture phases and stoichiometric compounds were reported: liquid phase (miscibility gap at high mole fractions of sulphur), Ni3±xS2 (non-stoichiometric high-temperature modification), Ni1–xS(α) (hexagonal, non-stoichiometric high-temperature form), Ni7±xS6 (non-stoichiometric high-temperature modification), Ni7S6 (stoichiometric low-temperature modification), Ni3S2 (heazle-woodite), NiS(β) (millerite), Ni3S4 (polydymite), and NiS2 (vaesite). Whereas a monotectic reaction between NiS2 and the two liquid phases of the miscibility gap was proposed in [62KUL/YUN], Arnold and Malik [75ARN/MAL] re-investigated the phase relations for the sulphur-rich part of the phase diagram and found a syntectic equilibrium between NiS2 and two liquid phases, NiS2 L1 + L2. In 1980 Sharma and Chang [80SHA/CHA] optimised the Gibbs energy functions of all phases in the system Ni – S in order to calculate the phase diagram for the first time. Based on experimental work of Rau [76RAU] as well as Lin et al. [78LIN/HU] the high-temperature modifica-tion Ni3±xS2 was shown to consist of two non-stoichiometric phases, viz. β1-Ni3S2 and β2-Ni4S3, where both phases exhibit a broad homogeneity range.


In contrast to earlier investigations [54ROS], [62KUL/YUN] the composition of the stoichiometric low-temperature phase, coexisting with Ni3S2 and NiS(β), has been found to be Ni9S8 [94STO/FJE], [96SEI/FJE]. Moreover, the thermodynamic properties of Ni7±xS6, Ni9S8, NiS (high- and low-temperature forms), and Ni3S2 were re-investigated experimentally by Stølen et al. [91STO/GRO], [94STO/FJE], [95GRO/STO]. In particular, in the case of NiS the enthalpy of transition estimated by Biltz et al. [36BIL/VOI] was revised and a new value was obtained by [95GRO/STO] based on careful heat capacity measurements employing an adiabatic-shield calorimeter. Thus, a new optimisation of the Gibbs energy functions of some phases in the system Ni – S has been necessary in order to recalculate the phase diagram and to recommend reliable thermodynamic data for the solid nickel sulphides.

The phase diagram of the binary system Ni – S is depicted in Figure V-29 and Figure V-30. The calculation of the phase diagram is based on Gibbs energy equations listed in Table V-15. It is worth mentioning that the small homogeneity range of Ni7S6 has been neglected, i.e., this phase is treated as a stoichiometric compound.

Figure V-29: Phase diagram for the binary system Ni – S with T plotted versus mole fraction of sulphur. ○ [10BOR], ■ [54ROS], ▲ [70NAG/ING], [75MEY/WAR], × [75RAU2].

0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0

600

800

1000

1200

1400

1600

1800

T / K

xS


Figure V-30: Section of the calculated phase diagram for the composition range 0.30 < xS < 0.55. Dashed lines have been extrapolated, indicating regions where phase equilibria cannot be calculated by using the Gibbs energy functions of Table V-15.

Table V-15: Gibbs energy functions for the binary system Ni – S relative to metallic nickel, Ni(s), and gaseous sulphur, S2(g); f mGο∆ = a + bT + cT ln(T), Li = a + bT + cT ln(T).

Range of validity a (J·mol–1) b (J·K–1·mol–1) c (J·K–1·mol–1)

Liquid phase (three-suffix Margules model) Ni 900 < T < 1900 17472.000 – 10.120000 0.00000000 NiS 900 < T < 1900 – 123776.00 152.91400 – 13.513000 S 900 < T < 1900 – 65357.000 165.39600 – 13.513000 LNi–NiS 900 < T < 1900;

0 < xS < 0.45 – 13691.322 2.8470168 0.00000000

LNiS–Ni 900 < T < 1900; 0 < xS < 0.45

– 104089.64 18.798027 0.00000000

LNiS–S 900 < T < 1900; 0 < xS < 0.45

28684.680 – 29.748923 0.00000000

LS–NiS 900 < T < 1900; 0 < xS < 0.45

35296.291 4.2627929 0.00000000

(continued on next page)

0.30 0.35 0.40 0.45 0.50 0.55500

600

700

800

900

1000

1100α-NiS

β-N

iS

Ni 7S 6

Ni 9S 8

Ni 3S 2

β2-Ni4S3β1-Ni3S2

liquid

T / K

xS


Table V-15: (continued)

Range of validity a (J·mol–1) b (J·K–1·mol–1) c (J·K–1·mol–1)

β1-Ni3S2 (subregular model) Ni 845 < T < 1040 – 102153.21 216.88700 0.00000000 S 845 < T < 1040 0.00000000 0.00000000 0.00000000 L1 845 < T < 1040;

0.35 < xS < 0.50 – 93628.460 – 256.37000 0.00000000

L2 845 < T < 1040; 0.35 < xS < 0.50

595976.19 – 908.72000 0.00000000

β2-Ni4S3 (subregular model) Ni 845 < T < 1040 307930.46 – 177.62880 0.00000000 S 845 < T < 1040 0.00000000 0.00000000 0.00000000 L1 845 < T < 1040;

0.35 < xS < 0.50 – 896420.15 515.25100 0.00000000

L2 845 < T < 1040; 0.35 < xS < 0.50

– 500648.48 150.70800 0.00000000

NiS(α) (sublattice model) Ni:S 500 < T < 1250 – 161700.00 216.85000 – 18.707460 Va:S 780 < T < 1250 – 108112.14 106.48130 0.00000000 LVa-Va 780 < T < 1250 – 296408.36 3313.7040 – 419.61114 Stoichiometric phases Ni3S2 500 < T < 850 – 331689.70 163.98400 0.00000000 Ni7S6 675 < T < 850 – 9005.7621 – 7709.455 1064.58400 NiS(β) 500 < T < 700 – 155944.00 86.675500 0.00000000 Ni9S8 500 < T < 700 – 1259000.0 667.67880 0.00000000 Ni3S4 500 < T < 630 – 560503.30 338.15692 0.00000000 NiS2 600 < T < 1250 – 244721.70 160.65761 0.00000000

V.5.1.1.1.2.1 Ni – S melt

Sharma and Chang [80SHA/CHA] proposed a three-suffix Margules model for the ex-cess Gibbs energy function of the Ni – S melt. In the present review this concept is likewise adopted. The molar excess Gibbs energy is given by the sum of binary interac-tions between the three species (Ni, S and NiS) of the liquid phase

m,

( )exi j i i j j

i j

G x x L x L x= +∑ , (V.76)

where the temperature dependence of the binary interaction parameters is defined as

Li = a + bT + cT ln(T). (V.77)

The liquidus line in the Ni-rich part of the phase diagram agrees excellently with various experimental studies, as can be seen in Figure V-29. Hence, the Gibbs en-


ergy function of the liquid phase, listed in Table V-15, is consistent with [08BOR], [10BOR], [54ROS], [70NAG/ING], [73VEN/GEI], [75MEY/WAR].

V.5.1.1.1.2.2 β1-Ni3S2 and β2-Ni4S3

According to Sharma and Chang [80SHA/CHA] the excess Gibbs energy function of the mixture phases β1-Ni3S2 and β2-Ni4S3, containing the components Ni and S, are modelled by application of a subregular model

m S S 1 2(1 )[ (1 2 )]exSG x x L L x= − + − (V.78)

with the temperature dependence of the parameters L1 and L2 being likewise given by Equation (V.77).

The predicted homogeneity range of the high-temperature phases β1-Ni3S2 and β2-Ni4S3 is in accordance with phase equilibrium studies performed by Bornemann [10BOR] as well as sulphur partial pressure measurements reported by Rau [76RAU] and Lin et al. [78LIN/HU].

V.5.1.1.1.2.3 Ni3S2(cr) (heazlewoodite)

The standard enthalpy of formation of heazlewoodite, Ni3S2(cr), was determined di-rectly by drop calorimetry [86CEM/KLE], resulting in,

f mH ο∆ (Ni3S2, cr, 298.15 K) = – (217.2 ± 1.6) kJ·mol–1,

which is selected in the present evaluation. In another calorimetry investigation the reaction between elemental nickel and sulphur did not go to completion [92VID/KOR]. The result disagreed with [86CEM/KLE], and was not considered further in this review.

Heat capacity measurements at low temperatures were performed by Weller and Kelley [64WEL/KEL] and Stølen et al. [91STO/GRO]. From these data the stan-dard entropy at 298.15 K can be deduced, leading to mS ο (Ni3S2, cr, 298.15 K) = (133.89 ± 0.84) J·K–1·mol–1 [64WEL/KEL] and mS ο (Ni3S2, cr, 298.15 K) = 133.2 J·K−1·mol–1 [91STO/GRO], respectively. Since these values agree well with each other,

mS ο (Ni3S2, cr, 298.15 K) = (133.5 ± 0.7) J·K–1·mol–1

is selected for the standard entropy of Ni3S2(cr). High temperature heat capacity meas-urements were carried out by Ferrante and Gokcen [82FER/GOK] and Stølen et al. [91STO/GRO]. From these two data sets a heat capacity function valid from 298.15 to 835 K has been obtained in the present assessment, see Appendix A for details. At 298.15 K the heat capacity is reported to be 118.2 J·K–1·mol–1 [91STO/GRO], 117.7 J·K–1·mol–1 [82FER/GOK], and 117.65 J·K–1·mol–1 [64WEL/KEL]. The heat capacity at 298.15 K, selected in the present assessment, corresponds to:

,mpCο (Ni3S2, cr, 298.15 K) = (118.2 ± 0.8) J·K–1·mol–1.

The error limits (± 0.8 J·K–1·mol–1) are chosen such that the selected value for ,mpCο is consistent with all experimental data. The above selections lead to:


f mGο∆ (Ni3S2, cr, 298.15 K) = – (211.2 ± 1.6) kJ·mol–1.

Figure V-31 shows a plot of the Gibbs energy of reaction for:

3Ni(cr) + 2H2S(g) 2H2(g) + Ni3S2(cr) (V.79)

as a function of temperature. The experimental data reported by Rosenqvist [54ROS] and Liné and Lafitte [63LIN/LAF] are in close agreement with calculated values using the present selection of thermodynamic properties of Ni3S2(cr).

Figure V-31: Gibbs energy of reaction for (V.79) plotted versus temperature. ● [54ROS], solid line: calculation using the thermodynamic quantities selected in this review.

V.5.1.1.1.2.4 Ni7S6(cr) and Ni9S8(cr)

The standard entropy of Ni9S8(cr) was determined from low temperature heat capacity measurements by Stølen et al. [94STO/FJE]. Unfortunately, the thermodynamic proper-ties were measured on a mixture of (⅓Ni3S2 + ⅔Ni9S8) rather than on pure Ni9S8(cr). Taking into account the low-temperature heat capacity function of Ni3S2(cr) [91STO/GRO] the standard entropy at 298.15 K is calculated to be:

mS ο (Ni9S8, cr, 298.15 K) = (481 ± 7) J·K–1·mol–1.

660 680 700 720 740 760 780 800 820

-110

-105

-100

-95

-90

3 Ni + 2 H2S(g) Ni3S2 + 2 H2(g)

T / K

∆ rG

° m / k

J·mol

–1


This quantity is selected with fairly large error bounds because the entropy was not determined directly on pure Ni9S8(cr). The heat capacity function valid from 298.15 to 640 K is likewise obtained from measurements on the mixture of (⅓Ni3S2 + ⅔Ni9S8) carried out by the same authors [94STO/FJE], see Appendix A for details. No additional investigations on Ni9S8(cr) are reported in literature so far. The selected heat capacity of Ni9S8 at 298.15 K is:

,mpCο (Ni9S8, cr, 298.15 K) = (407 ± 5) J·K–1·mol–1.

According to Cemic and Kleppa [86CEM/KLE] the standard enthalpy of for-mation of the high-temperature form of Ni7S6(cr) at 298.15 K can be determined by drop calorimetry, where the reaction temperature was 833 K, yielding f mH ο∆ = − (582.8 ± 5.7) kJ·mol–1. However, Seim et al. [96SEI/FJE] state that the high-temperature form of Ni7S6(cr) cannot be obtained by quenching samples from tempera-tures above 675 K to room temperature. The formation of several metastable phases of different structures was reported. If, in turn, the quenched samples of Cemic and Kleppa consisted of a mixture of ⅓Ni3S2 + ⅔Ni9S8 (which should form from Ni7S6(cr) below 675 K according to the phase diagram), the standard enthalpy of formation of this mix-ture at 298.15 K would amount to – 579 kJ·mol–1 which coincides remarkably well with the experimental result of Cemic and Kleppa. The value – 579 kJ·mol–1 is based on a combination of the standard enthalpy of formation for Ni3S2(cr) with that for Ni9S8(cr) ( f mH ο∆ = – 760 kJ·mol–1) which is consistent with the Gibbs energy equation (Table V-15) used for the calculation of the phase diagram (Figure V-29). Thus, for the standard enthalpy of formation of Ni9S8(cr) at 298.15 K,

f mH ο∆ (Ni9S8, cr, 298.15 K) = – (760 ± 9) kJ·mol–1

can be selected, introducing large error bounds because no direct experimental determi-nation of this quantity has been performed so far. The above selections lead to:

f mGο∆ (Ni9S8, cr, 298.15 K) = – (746.8 ± 9.3) kJ·mol–1.

In Figure V-32 the temperature dependence of the partial pressure of sulphur is depicted for phase equilibria involving Ni9S8(cr), Ni7S6(cr), Ni3S2(cr), and β2-Ni4S3. Usually, the partial pressure of sulphur is obtained from measurements of the ratio p(H2S)/p(H2) after a hydrogen containing gas phase has been equilibrated with the solid phases. Figure V-32 reveals that calculations of phase equilibria, using the present thermodynamic model, coincide well with experimental data taken from various sources [39SCH/FOR], [52SUD], [54ROS], [78LIN/HU]. The present model disagrees, how-ever, with the experimental data of Jellinek and Zakowski [25JEL/ZAK], probably be-cause they investigated inadequately defined nickel phases.


Figure V-32: Partial pressure of sulphur for various phase equilibria in the binary sys-tem Ni – S plotted against temperature. [39SCH/FOR], ▲ [52SUD], ● [54ROS], □ [78LIN/HU], solid lines: calculations using the thermodynamic data of the present assessment.

V.5.1.1.1.2.5 α−NiS and β−NiS

The high-temperature modification of NiS crystallises in a hexagonal NiAs-type struc-ture and exhibits a wide range of homogeneity which was studied extensively as a func-tion of temperature by Rau [75RAU2]. Based on the experimental data of Rau a statisti-cal thermodynamic model [69LIG/LIB], [77LIN/IPS] was introduced by Sharma and Chang [80SHA/CHA], describing the defect interactions in this non-stoichiometric phase. It can be shown that the statistical thermodynamic model applied in [80SHA/CHA] is equivalent to a sublattice model which is used in the present assess-ment. α-NiS is envisaged to be composed of two sublattices, one of which contains only sulphur atoms and the other which consists of both nickel atoms and vacancies. This takes into account that deviations from the ideal stoichiometry (NiS) lead exclusively to nickel deficient compositions. The Gibbs energy function for this phase can be written as:

Ni S Ni:S Va S Va:S Ni Ni Va Va S S mR ( ln ln ln ) exmG y y G y y G T y y y y y y Gο ο= + + + + + (V.80)

600 650 700 750 800 850-14

-13

-12

-11

-10

-9

-8

-7

-6

Ni7S6

Ni9S8Ni3S2

β2-Ni4S3

log 10

[p(S

2) / b

ar]

T / K


and the non-ideal interaction between vacancies on the nickel sublattice can be ex-pressed as :

2m Va Va VaexG y L −= (V.81)

The symbols yNi and yVa denote the site fractions of nickel atoms and vacancies, respectively, on the nickel sublattice and the site fraction of sulphur atoms on the sul-phur sublattice is always equal to yS = 1. The temperature dependence of the defect in-teraction parameter LVa–Va is given by Equation (V.77). The Gibbs energies of the refer-ence states Ni:SGο and Va:SGο refer to the nickel sublattice completely occupied by Ni at-oms (stoichiometric composition NiS) and vacancies (empty nickel sublattice), respec-tively.

The homogeneity range of α-NiS, predicted by the thermodynamic model pre-sented in Table V-15, agrees well with experimental data reported by Rau [75RAU2], see Figure V-29. The standard enthalpy of formation of NiAs-type α-NiS (stoichiomet-ric composition) was determined by Cemic and Kleppa [86CEM/KLE] employing drop calorimetry,

f mH ο∆ (NiS, α, 298.15 K) = – (88.1 ± 1.0) kJ·mol–1

and this value is selected in the present data assessment. The enthalpy of α-NiS forma-tion determined by [71ARI/MOR] agrees with the selected value within the limits of uncertainty, but is less well documented. As the high-temperature modification of Ni(II) sulphide can be quenched easily to room temperature, the heat capacity function of stoichiometric α-NiS can be derived from calorimetric measurements [95GRO/STO] over the range from 298.15 to 900 K, see Appendix A. Furthermore, the heat capacity function of millerite, β-NiS, was determined from 298.15 to 660 K [95GRO/STO], see Appendix A for details. Several investigations dealt with the enthalpy of transition be-tween millerite and NiAs-type nickel monosulphide which are summarised in Table V-16. The most reliable one seems to be the study of Grønvold and Stølen [95GRO/STO]. Hence, combining this value for the enthalpy of transition with the enthalpy of forma-tion of α−NiS and the pertinent heat capacity functions leads to the selected value for the standard enthalpy of formation of β-NiS:

f mH ο∆ (NiS, β, 298.15 K) = – (94.0 ± 1.0) kJ·mol–1.

The standard entropy of millerite was determined at 298.15 K by application of low temperature heat capacity measurements [64WEL/KEL],

mS ο (NiS, β, 298.15 K) = (52.97 ± 0.33) J·K–1·mol–1

which is likewise selected in the present evaluation. Combining this value with the en-tropy of transition, ( trs mH ο∆ /Ttrs) = (9.99 ± 0.08) J·K–1·mol–1 [95GRO/STO], and the respective heat capacity functions results in the standard entropy for stoichiometric α−NiS,

mS ο (NiS, α, 298.15 K) = (60.89 ± 0.34) J·K–1·mol–1.


According to Weller and Kelley [64WEL/KEL] the heat capacity of β-NiS at 298.15 K amounts to 47.11 J·K–1·mol–1 which is in excellent agreement with the result of Grønvold and Stølen [95GRO/STO],

,mpCο (NiS, β, 298.15 K) = (47.08 ± 0.30) J·K–1·mol–1,

selected in the present review. In addition, it should be mentioned that the heat capacity of millerite was likewise measured in the past by Tilden [02TIL], [02TIL2], ,mpCο = 47.35 J·K–1·mol–1, and Regnault [1841REG], ,mpCο = 48.61 J·K–1·mol–1, respectively. The error limits of the selected value (± 0.30 J·K–1·mol–1) are chosen such that the selected value for ,mpCο is consistent with the experimental data of [02TIL], [02TIL2], [64WEL/KEL] and [95GRO/STO]. The selected value for the heat capacity of stoichiometric α-NiS at 298.15 K is:

,mpCο (NiS, α, 298.15 K) = (49.76 ± 0.30) J·K–1·mol–1.

Table V-16: Temperatures and enthalpies of transition for β-NiS α-NiS.

Ttrs / K trs mH ο∆ ( kJ·mol–1) Reference

669 2.6 [36BIL/VOI] 652 6.4 [77CON/SRI] 667 5.9 [80DUB/CLA]660 6.6 [95GRO/STO]

The temperature dependence of the Gibbs energy function for stoichiometric α−NiS, i.e., the nickel sublattice is completely occupied, is shown in Figure V-33. Tak-ing the enthalpy of formation, determined by Cemic and Kleppa [86CEM/KLE], and the

,mpC function as well as the standard entropy of α-NiS, measured by Grønvold and Stølen [95GRO/STO], the f mG∆ function of stoichiometric α-NiS does not match the values proposed by Sharma and Chang [80SHA/CHA], as can be deduced from Figure V-33. The Gibbs energy function for stoichiometric α-NiS, retained in the present re-view (cf. Table V-15), is consistent with both the high temperature data of Sharma and Chang and the experimental results obtained at temperatures below 850 K according to Cemic and Kleppa as well as Grønvold and Stølen. The Gibbs energy functions for the species S (empty nickel sublattice) and the parameter for the non-ideal interaction be-tween vacancies on the nickel sublattice are adopted from [80SHA/CHA].

The above selections lead to:

f mGο∆ (NiS, α, 298.15 K) = – (87.8 ± 1.0) kJ·mol–1

and f mGο∆ (NiS, β, 298.15 K) = – (91.3 ± 1.0) kJ·mol–1.


Figure V-33: Gibbs energy function for the stoichiometric composition of NiAs-type Ni(II) sulphide, α-NiS, plotted versus temperature.

V.5.1.1.1.2.6 NiS2(cr), vaesite

Cemic and Kleppa [86CEM/KLE] measured the standard enthalpy of formation of vae-site, NiS2(cr), at 298.15 K by drop calorimetry and obtained f mH ο∆ (NiS2, cr, 298.15 K) = – (124.9 ± 1.0) kJ·mol–1. The optimised Gibbs energy function of NiS2(cr) listed in Table V-15 is based on decomposition pressure measurements [64LEE/ROS] and phase equilibrium studies [75ARN/MAL]. An experimental determination of the heat capacity function at high temperatures has not yet been reported in literature. Applying the esti-mated heat capacity function for NiS2(cr) published in the JANAF-tables [98CHA],

,mpCο (NiS2, cr) = (64.4378 + 0.020753 (T/K)) J·K–1·mol–1 (V.82)

the Gibbs energy function given in Table V-15 yields for the standard enthalpy of for-mation and the standard entropy at 298.15 K: f mH ο∆ (NiS2, cr, 298.15 K) = – 121.1 kJ·mol–1 and mS ο (NiS2, cr, 298.15 K) = 88 J·K–1·mol–1, respectively. This value for

f mH ο∆ agrees reasonably well with that given by [86CEM/KLE].

Ogawa [77OGA] has published low-temperature measurements of the heat ca-pacity of NiS2(cr) from 17 to 347 K. Integration of these results between 0 and 298.15 K leads to the standard entropy of NiS2(cr), mS ο (NiS2, cr, 298.15 K) = (79.6 ± 2.4)

300 400 500 600 700 800 900 1000 1100 1200 1300 1400

-130

-120

-110

-100

-90

-80

-70

-60

-50

[86CEM/KLE],[95GRO/STO] [80SHA/CHA] optimized line

T / K

∆ fG° m

/ kJ

·mol

−1

Ni + 1/2 S2(g) NiS

∆ fG

° m / k

J·mol

–1


J·K−1·mol–1, see Appendix A for details. Taking this value for mS ο and the heat capacity function estimated in the JANAF tables (Equation (V.82)), the re-evaluation of decom-position pressure measurements, reported by Leegaard and Rosenqvist [64LEE/ROS], results in f mH ο∆ (NiS2, cr, 298.15 K) = – 128.3 kJ·mol–1. This value, however, is not consistent with the phase diagram data [75ARN/MAL] and the Gibbs energy function listed in Table V-15. As the quantity mS ο (NiS2, cr, 298.15 K) = 79.6 J·K–1·mol–1 is based on low-temperature heat capacities,

mS ο (NiS2, cr, 298.15 K) = (80 ± 8) J·K–1·mol–1

and

f mH ο∆ (NiS2, cr, 298.15 K) = – (128 ± 7) kJ·mol–1

are selected. The discrepancy between the Gibbs energy function, necessary to predict the high-temperature phase equilibria between NiS2(cr), NiS(α) and the liquid phases, and the selected thermodynamic properties of vaesite is taken into account by the fairly large error bounds.

From the above selections, the selected Gibbs energy of formation is:

f mGο∆ (NiS2, cr, 298.15 K) = – (123.8 ± 7.4) kJ·mol–1.

V.5.1.1.1.2.7 Ni3S4(cr), polydymite

Experimental investigations on the thermodynamic properties of polydymite, Ni3S4(cr), have not yet been published. Hence, thermodynamic data for this compound cannot be recommended in the present assessment.

V.5.1.1.1.2.8 Ni4.5Fe4.5S8(cr), pentlandite

Cemic and Kleppa [87CEM/KLE] determined the enthalpy of formation of synthetic pentlandite, Ni4.5Fe4.5S8(cr), by drop calorimetry and obtained f mH ο∆ (Ni4.5Fe4.5S8, cr, 298.15 K) = – (837 ± 15) kJ·mol–1.

V.5.1.1.1.3 Discussion of selected thermodynamic properties of nickel sulphides

A comparison of predicted and experimentally determined temperatures for invariant three-phase equilibria is given in Table V-17. Except for the peritectoid decomposition of Ni9S8(cr) into NiS(α) and Ni7S6(cr), the predicted temperatures, employing the ther-modynamic quantities of the present review (cf. Table V-15), agree well with experi-mental data [62KUL/YUN], [75ARN/MAL], [94STO/FJE]. The peritectoid decomposi-tion of Ni9S8(cr) was determined by means of DTA [94STO/FJE]. The discrepancy be-tween the experimental temperature and the predicted value may be caused by (i) an inconsistency of the present thermodynamic model or (ii) the equilibrium temperature was overestimated in the DTA experiment, as onset temperatures usually depend on the heating rate.


Table V-17: Temperatures for invariant three-phase equilibria in the binary system Ni − S.

T / K (calc.) T / K (exp.) Reference

β2-Ni4S3 + Ni7S6(cr) + Ni3S2(cr) 797.1 797 [62KUL/YUN]

Ni7S6(cr) + β2-Ni4S3 + NiS(α) 850.1 850.5 [94STO/FJE]

Ni7S6(cr) + Ni3S2(cr) + Ni9S8(cr) 675.3 675.2 [94STO/FJE]

Ni9S8(cr) + NiS(α) + Ni7S6(cr) 677.7 709 [94STO/FJE]

Ni3S4(cr) + NiS(α) + NiS2(cr) 628.8 629 [62KUL/YUN]

liquid + NiS2(cr) + NiS(α) 1259.8 1266 1258

[75ARN/MAL] [62KUL/YUN]

NiS2(cr) + L1 + L2 1305.7 1295 [75ARN/MAL]

The partial pressure of sulphur for phase equilibria, involving NiS(α), β2-Ni4S3 and NiS2(cr), is plotted versus temperature in Figure V-34. Experimental data from various sources [36BIL/VOI], [39SCH/FOR], [54ROS], [59LAF], [64LEE/ROS], [78LIN/HU], [90OSA/ROS] compare well with thermodynamic calculations of the pre-sent review (Gibbs energy functions of Table V-15, solid lines) as well as with [80SHA/CHA] (dashed lines). In the case of NiS2 the selected thermodynamic data (dotted line) coincides even better with experimental decomposition pressures than that proposed by [80SHA/CHA].

Finally, thermodynamic data can be selected for the following phases of the binary system Ni – S: Ni3S2(cr), Ni9S8(cr), NiS(β), NiS(α) (stoichiometric composition), and NiS2(cr). The heat capacity functions for these compounds are summarised in Table V-18. The standard enthalpies of formation, the standard entropies and the heat capaci-ties at 298.15 K are listed in Table V-19 to Table V-21. The selected quantities of this assessment compare reasonably well with those of other data compilations [74MIL], [95ROB/HEM], [98CHA]. The thermodynamic functions reported by [78SCH/MIL] are inconsistent with any of these data sets, and thus were disregarded.


Figure V-34: Plot of equilibrium partial pressure of sulphur versus temperature. ▼ [36BIL/VOI], [39SCH/FOR], ▲ [54ROS], ■ [59LAF], ● [64LEE/ROS], ○ [78LIN/HU], [90OSA/ROS], solid lines: calculations using Gibbs energy functions of Table V-15, dashed lines: calculations using the thermodynamic model of [80SHA/CHA], dotted line: calculation using the selected thermodynamic data for NiS2.

Table V-18: Heat capacity functions of nickel sulphides; ,mpCο = a + bT + cT 2 + eT –2. (These are selected values except for NiS2(cr), vaesite).

Compound Range of validity a (J·K–1·mol–1)

b (J·K–2·mol–1)

c (J·K–3·mol–1)

e (J·K3·mol–1)

Ni3S2(cr), heazlewoodite 298.15 < T/K < 835 165.79352 – 0.10316 0.00012 – 2444624 Ni9S8(cr) 298.15 < T/K < 640 1079.2452 – 2.0599400 0.00198 – 20839639 NiS(β), millerite 298.15 < T/K < 660 54.008 – 0.01437 0.000032282 – 490363 NiS(α) 298.15 < T/K < 660 46.676 0.0199807 0.00000000 – 255499.4 660 < T/K < 1000 50.07833 0.0149 0.00000000 – 336916.92 NiS2(cr), vaesite estimated 64.4378 0.020753 0.00000000 0.00000000

0.90 1.00 1.10 1.20 1.30 1.40 1.50 1.60-8

-7

-6

-5

-4

-3

-2

-1

0

β2-Ni4S3

α-NiS

NiS2

log 10

[p(S

2) / b

ar]

1000 K / T


Table V-19: Standard enthalpy of formation of nickel sulphides at 298.15 K.

Compound f mH ο∆ (kJ·mol–1)

[74MIL] [95ROB/HEM] [98CHA] this review

Ni3S2(cr), heazlewoodite – (215.9 ± 12.6) – (216.3 ± 3.0) – (216.3 ± 5.0) – (217.2 ± 1.6) Ni9S8(cr) – (742 ± 38) – (760 ± 9) NiS(β), millerite – (94.1 ± 4.2) – (91.0 ± 3.0) – (87.9 ± 6.3) – (94.0 ± 1.0) NiS(α) – (88.1 ± 1.0) NiS2(cr), vaesite – (133.9 ± 8.4) – (131 ± 17) – (128 ± 7)

Table V-20: Standard entropy of nickel sulphides at 298.15 K.

Compound mS ο (J·K–1·mol–1)


Ni3S2(cr), heazlewoodite (133.9 ± 2.5) (133.2 ± 0.3) (133.9 ± 0.4) (133.5 ± 0.7) Ni9S8(cr) (440 ± 38) (481 ± 7) NiS(β), millerite (53.0 ± 0.8) (53.0 ± 0.4) (53.0 ± 0.4) (52.97 ± 0.33) NiS(α) (60.89 ± 0.34) NiS2(cr), vaesite (67.8 ± 8.4) (72.0 ± 8.4) (80 ± 8)

Table V-21: Heat capacity of nickel sulphides at 298.15 K.

Compound ,mpCο (J·K–1·mol–1)


Ni3S2(cr), heazlewoodite 117.6 118.16 117.7 (118.2 ± 0.8) Ni9S8(cr) (407 ± 5) NiS(β), millerite 47.11 47.11 47.11 (47.08 ± 0.30) NiS(α) (49.76 ± 0.30)

V.5.1.1.2 Solubilty of NiS(s) and aqueous nickel hydrogen sulphido species

Apart from theoretical estimations of the stability constants of aqueous nickel sulphide complexes by Dyrssen [85DYR], [88DYR], [90DYR/KRE], only two experimental studies on the thermodynamic properties of nickel sulphide complexes are available so far [94ZHA/MIL], [96LUT/RIC]. These experimental data have been recalculated by applying the SIT model [97GRE/PLY2], see Table V-22.


Table V-22: Recalculated stability constants and SIT parameters for nickel sulphide complexes at T = 298.15 K and I = 0.

Reaction 10log iοb ∆ε Reference

(V.83): Ni2+ + HS– NiHS+ (5.18 ± 0.20) – (0.97 ± 0.39) [94ZHA/MIL], [96LUT/RIC]

(V.84): 2Ni2+ + HS– Ni2HS3+ (9.92 ± 0.10) – (0.05 ± 0.22) [96LUT/RIC]

(V.85): 3Ni2+ + HS– Ni3HS5+ (14.008 ± 0.099) (0.59 ± 0.22) [96LUT/RIC]

Based on these stability constants the distribution of sulphide-bearing aqueous Ni(II) species has been calculated as a function of total Ni(II) molality in aqueous solu-tions saturated with H2S at T = 298.15 K and a constant ionic strength of 1.00 mol·kg–1 NaCl, see Figure V-35. At fairly low pH values high Ni(II) molalities are stable in solu-tions saturated with H2S. At these conditions ([Ni(II)]tot > 10–3.5 mol·kg–1) the uncom-mon complexes Ni2HS3+ and Ni3HS5+ become the most dominant species in aqueous solution. As this situation seems to be unrealistic and no studies on the structure of these complexes are reported in literature, we select thermodynamic data only for the aqueous species NiHS+:

10 1,1log οb ((V.83), 298.15 K) = (5.18 ± 0.20).

This selection, with NEA-TDB selected auxiliary data, yields:

f mGο∆ (NiHS+, 298.15 K) = – (63.1 ± 2.5) kJ·mol–1.

Owing to the high uncertainty of the second dissociation constant of H2S the following dissolution reaction of nickel sulphide:

NiS(cr) + 2H+ Ni2+ + H2S(aq) ,0*

sK ο (V.86)

is thermodynamically better defined than

NiS(cr) Ni2+ + S2– ,0sK ο . (V.87)

To the best of our knowledge no reliable studies on the solubility of nickel sul-phides in aqueous media have been published yet. Some investigations refer to amor-phous solid phases, resulting in too high values for the solubility constants [14THI/GES], [70CAR/LAI]. In many cases reaction times were far too short to attain solid-solute equilibrium, yielding values for the solubility constants that are too low [14THI/GES], [24MOS/BEH], [47DON]. Clearly thermodynamic quantities based es-sentially on experimental data of [14THI/GES] are unreliable [93MAK/VOE]. Recalcu-lated values for both the solubility constant, 10 ,0

*log sK ο , corresponding to the dissolu-tion reaction (V.86) and the solubility product, 10 ,0log sK ο , using the currently accepted value for the second dissociation constant of hydrogen sulphide (H2S(aq) 2H+ + S2–

10 2log K ο = – 25.99 [92GRE/FUG]), are listed in Table V-23.


Figure V-35: Distribution of nickel sulphide complexes as a function of total molality of nickel (II) in aqueous solutions saturated with H2S, T = 298.15 K, I = 1.00 mol·kg–1 NaCl. Solid line: Ni2+; dashed line: NiHS+; dotted line: Ni2HS3+; dash-dotted line: Ni3HS5+.

10-6 10-5 10-4 10-3 10-2 10-110-12

10-10

10-8

10-6

10-4

10-2

NiHS+

Ni3HS5+

Ni2HS3+

Ni2+

p(H2S) = 1 barI = 1.00 mol·kg−1 NaCl

[Ni xH

S2x−1

] / m

ol·k

g−1

[Ni(II)]tot

/ mol·kg−1


Table V-23: Solubility constants for Ni(II) sulphide at T = 298.15 K and I = 0.

10 ,0*log sK ο 10 ,0log sK ο Reference

– 0.5 – 27.5 [1883THO], [09BRU/ZAW]

– 3.98 – 30.98 [24MOS/BEH], [31KOL]

9.20 – 17.8 [70CAR/LAI]

– (1.53 ± 0.30)(a) – 28.5(a) this review

– (2.10 ± 0.30)(b) – 29.1(b) this review

(a) NiS(α). (b) NiS(β), millerite.

The solubility data scatter remarkably and none of these is consistent with the solubility constants selected in this data assessment. The solubility constant given by Bruner and Zawadzki [09BRU/ZAW] is based on a calorimetric value for the enthalpy of reaction for the precipitation of NiS from aqueous solution:

Ni2+ + S2– NiS,

reported by Thomson [1883THO]. Thiel and Gessner [14THI/GES] carried out a com-prehensive study of the dissolution behaviour of NiS prepared by precipitation from aqueous solutions at various conditions. They found three fractions, viz. I-NiS, II-NiS and III-NiS, with different solubilities in hydrochloric acid. The variations of the solu-bility of these fractions, observed by the authors, can be explained in terms of different dissolution kinetics rather than in terms of different thermodynamic properties, since reaction times were certainly too short to attain equilibrium between the solid phase and the aqueous solution. Licht [88LIC2] has recalculated the results of [14THI/GES] and obtained Gibbs energies of formation for I-NiS, II-NiS and III-NiS, respectively. The solubilities of Ni(II) sulphides in solutions of hydrochloric acid according to [88LIC2], [14THI/GES] as well as the present assessment are depicted in Figure V-36. Whereas the solubility of I-NiS, which seems to correspond to an amorphous material, is too high, the solubilities of II-NiS and III-NiS are by far too low because of too short reac-tion times. The fractions II-NiS and III-NiS are apparently crystalline products. This assumption is likewise confirmed by Dönges [47DON] who has pointed out that, be-sides amorphous products, the metastable high-temperature modification, NiS(α), is precipitated from aqueous solutions. Furthermore, it should be mentioned that millerite, NiS(β), is the only nickel sulphide identified so far in natural low-temperature anoxic-sulphidic environments according to Thoenen [99THO].


Figure V-36: Solubility of Ni(II) sulphides in aqueous solutions saturated with H2S as a function of initial concentration of hydrochloric acid at 298.15 K. Dashed lines refer to I-NiS, II-NiS and III-NiS as experimentally investigated by Thiel and Gessner [14THI/GES] and thermodynamically interpreted by Licht [88LIC2]. Solid lines corre-spond to NiS(α) and NiS(β) using the thermodynamic data selected in the present com-pilation. Dotted lines are related to calculations taking into account the uncommon aqueous species Ni2HS3+ and Ni3HS5+, respectively.

V.5.1.2 Nickel sulphates

V.5.1.2.1 Aqueous nickel (II) sulphato complexes

The complexation reactions of Ni2+ with 24SO − have been the subject of a large number

of investigations (Table V-24). These include studies by cryoscopy [55BRO/PRU], [56KEN], [58KEN], [71ISO], conductivity [1895FRA], [56FIA/SHE], [73KAT], [74MAK/MAS], [76SHI/TSU], [79FIS/FOX], [93SHE], [2000TSI/MOL], [2001MOL/TSI], potentiometry and other electrochemical techniques [59NAI/NAN], [61TAN/OGI], [63TAN/SAI], [65POT], [68PRA], [70TUR/RUV], solubility [67GES/NEU], [70MIR/MAK], spectroscopy [64LAR], [70MIR/MAK], [77ASH/HAN], [78BEC/LIU], [80LIB/SAD] calorimetry [69IZA/EAT], [74BLO/RAZ], [85DRA/MAD] and isopiestic measurements [80LIB/SAD], [83HOL/MES]. There is a broad consensus, primarily based on the spectroscopic evi-dence, that, for temperatures below 100°C, association of these ions in dilute solutions

-2.0 -1.5 -1.0 -0.5 0.0 0.5-12

-10

-8

-6

-4

-2

0

III-NiS

II-NiS

β-NiS

α-NiS

I-NiS

log 10

([Ni(I

I)]to

t / m

ol·k

g−1)

log10

([HCl]initial

/ mol·kg−1)


is predominantly an outer-sphere interaction, with sulphate gradually displacing water molecules in the inner sphere around Ni2+ at higher temperatures and concentrations [77ASH/HAN], [78BEC/LIU]. Indeed, for 25°C, it is not clearly defined what does and what does not constitute a nickel sulphate complex. The nature of a specific type of ex-periment and the interpretation of experimental results affect any proposed value for the formation constant of NiSO4(aq) according to reaction (V.88):

Ni2+ + 24SO − NiSO4(aq). (V.88)

In particular, the reported value of the formation constant extrapolated to zero ionic strength is usually dependent on the particular method used to assign values to the ionic activity coefficients. In the present review, the objective is to assign a value that is consistent with the use of a particular formulation of the specific ion interaction equa-tions (the SIT approach, cf. Appendix B). Among the assumptions inherent in the SIT treatment is that although the association constant depends on the total ionic strength of the solution, on the ionic interactions between Ni2+ and anions of the supporting electro-lyte, and on the ionic interactions between 2

4SO − and cations of the supporting electro-lyte, all specific interactions between Ni2+ and 2

4SO − are accounted for by the NiSO4 formation constant. Nevertheless, Holmes and Mesmer [83HOL/MES] have shown that a Pitzer-type ion-interaction model can be used to provide a satisfactory interpretation of isopiestic data without assuming formation of a discrete NiSO4 complex, even at 140°C.

Several studies on ionic strength effects have used nickel sulphate as a model 2:2 electrolyte [55BRO/PRU], [73KAT], [75TAM]. In such cases, for the sake of con-sistency, it has been necessary to abandon the more sophisticated treatments, and to re-evaluate experimental data for complexation using the SIT. Perhaps because of the in-terest in nickel sulphate as a “typical” 2:2 electrolyte in aqueous solution, few of the association studies have been done as a function of concentration of a supporting elec-trolyte. Exceptions include the cryoscopic study of Kenttämaa [56KEN], [58KEN] and the polarographic studies of Tanaka and Ogino, [61TAN/OGI], and Tanaka, Saito and Ogino [63TAN/SAI]. Some spectroscopic studies have also been done using supporting electrolytes, often at very high concentrations [70TUR/RUV], [77ASH/HAN].

In the temperature range (0 to 50°C) in which an outer-sphere complex pre-dominates over inner sphere complexes, several studies [59NAI/NAN], [73KAT], [76SHI/TSU] have been used to determine the enthalpy of the complexation reaction (V.88). In addition, the enthalpy of reaction has been determined calorimetrically by Izatt et al. [69IZA/EAT], though a useful value can be obtained from their enthalpy titrations only if an independent method is used to determine the association con-stant [73POW] (cf. Appendix A). The comparison of derived values for the enthalpy of association also provides a tool for evaluating the association constant data.


Table V-24: Experimental values for the association constants and other thermodynamic quantities for nickel sulphate.

Method t/°C Medium (aq) Reported value Recalculated value(a) Reference

cond 25 0 (extrap.) K1 = 250 (220.7 ±13.4) [32MON/DAV], recalc.of [1895FRA]

fpt 0 0.005–0.09 self K1 = 125–244 (161.3 ± 20.0) [55BRO/PRU] fpt – 0.2 0.0484 m KClO4 K1 = 51 [56KEN] fpt – 0.8 0.258 m KClO3 K1 = 19 [58KEN] – 2.8 1.15 m KNO3 K1 = 4.90 pot 0 < 0.05 m HCl K1 = 121 (98.0 ± 23.0) [59NAI/NAN] 10 < 0.05 m HCl K1 = 151 (126.8 ±16.0) 15 < 0.05 m HCl K1 = 174 (166.2 ± 21.8) 25 < 0.05 m HCl K1 = 211 (172.8 ± 31.9) 35 < 0.05 m HCl K1 = 247 (210.2 ± 31.2) 45 < 0.05 m HCl K1 = 289 (236.9 ± 30.0) r mH ο∆ = 13.85 kJ·mol–1 see text r mS ο∆ = 90.8 J·K–1·mol–1 fpt – 3.06 1.15 m KNO3 K1 = 5.01

K2 = 33.6 [59ROS/ROS],

recalc. of [58KEN]

pol 25 0.2 M KNO3 K1 = 11.6 [61TAN/OGI] pol 25 1.0 M NaClO4 K1 = 3.7 [63TAN/SAI] K2 = 26 ir RT 1.0 M self β inner sphere / β outer sphere = 0.1 [64LAR] pot 35 0 (extrap.) K1 = 111 [68PRA] cal 25 0 (extrap.) K1 = 646 [69IZA/EAT] r mH ο∆ = 1.71 kJ·mol–1 see text r mS ο∆ = 59.0 J·K–1·mol–1 spec/sol 25.0 3.0 M LiClO4 K1 = 1.8 [70MIR/MAK] pol 25.0 5.0 M NaClO4 K1 = 15.6 [70TUR/RUV] fpt 0 0.005–0.03 self K1 = 118–250 (153.8 ± 20.0) [71ISO] cond 0.0 0 K1 = 156 (213.8 ± 4.4) [73KAT] 5.0 0 K1 = 162 (217.9 ± 4.9) 10.0 0 K1 = 168 (223.9 ± 4.9) 15.0 0 K1 = 176 (230.3 ± 4.7) 20.0 0 K1 = 183 (238.3 ± 5.3) 25.0 0 K1 = 187 (247.3 ± 5.8) 30.0 0 K1 = 191 (257.1 ± 7.1) 35.0 0 K1 = 201 (267.8 ± 7.8) 40.0 0 K1 = 209 (279.9 ± 8.8) 45.0 0 K1 = 217 (295.6 ± 9.4) 25.0 0 r mH ο∆ = 5.31 kJ·mol–1 (5.3 ± 2.0)



Table V-24: (continued)

Method t/°C Medium (aq) Reported value Recalculated value(a) Reference

r mS ο∆ = 61.5 J·K–1·mol–1 cal 25.0 0 r mH ο∆ =5.72–5.93 kJ·mol–1 [73POW],

recalc of [69IZA/EAT]

cal 25.0 1.5 M LiClO4/ 0.75M Li2SO4

r mH ο∆ = 2.59 kJ·mol–1 [74BLO/RAZ]

var 25. 0.018 self +? K1 = 280 [74MAK/MAS] 0.048 self +? K1 = 170 0.069 self +? K1 = 110 cond 0.0 0. K1 = 270 [75TAM],

recalc. of [73KAT]

cond 15.0 0. K1=457 (0.1 MPa) (187.7 ± 2.1) [76SHI/TSU] K1=227 (160 MPa) 25.0 0. K1=490 (0.1 MPa) (193.8 ± 1.9) K1=269 (160 MPa) 40.0 0. K1=575 (0.1 MPa) (226.7 ± 2.8) K1=313 (160 MPa) 25.0 0 r mH ο∆ = 6.90 kJ·mol–1 (6.0 ± 2.5) kJ·mol–1

(0.101 MPa)

25.0 0 r mS ο∆ = 109 J·K–1·mol–1

(0.101 MPa)

25.0 0 ∆1V° = 8.6 cm3·mol–1 uv 25.0(b) 5 M Na2SO4/NaClO4 K1= 5.89 [77ASH/HAN] 47 K1=6.8 60 K1=5.5 70 K1=5.4 cond 25.0 0 r mH ο∆ = 5.31 kJ·mol–1 r mS ο∆ = 61.5 J·K–1·mol–1

[77KAT] recalc.of [73KAT]

cond 15.0 0. K1=200 (0.1 MPa) (183.3 ± 30.4) [79FIS/FOX] K1=105 (203 MPa) 25.0 0. K1=196 (0.1 MPa) (199.4 ± 28.5) K1=93.5 (203 MPa) 25.0 0 ∆1V° = 11.6 cm3·mol–1 cond 25.0 0. K1=169 25.0 0 r mH ο∆ = 5.24 kJ·mol–1

[81YOK/YAM] recalc. of [73KAT]

r mS ο∆ = 60.2 J·K–1·mol–1 cond 25.0 < 0.06 M K1=230.1 [93SHE] cond. 20.0 0 K1=139 [2000TSI/MOL],

[2001MOL/TSI] (a) Uncertainties are the statistical uncertainties from the re-analysis of the data. The overall estimated uncer-

tainties are discussed in the text. (b) This value was reported by the authors based on an extrapolation to 25°C.


Table V-25 provides a summary of the results from re-analysis of the conduc-tance data for NiSO4(aq) using the series form [80PET/TAB]1 of the Lee-Wheaton equation [79LEE/WHE], modified so that the distance parameter is forced to the appro-priate value for the SIT model used in the TDB. No value could be recalculated from the work of Shehata [93SHE] because values of the concentration-dependent conduc-tances were not reported. Results from the other conductance studies at 25°C are com-pared in Figure V-37. The molar conductances agree well at low concentrations at 25°C. At higher concentrations (≥ 10–3 M), values are more divergent, especially those from the early study of Frank [1895FRA]. The values of Fisher and Fox [79FIS/FOX] and Shimizu, Tsuchihashi and Furumi [76SHI/TSU] are in very good agreement. As can be seen from Table V-25, the agreement is slightly poorer for 15°C. Values of Λo from analysis of the molar conductances from these two studies are substantially smaller than the values calculated from the molar conductances reported by Katayama [73KAT]. Nevertheless, all four studies lead to values of 10log K ο ((V.88), 298.15 K) between 2.29 and 2.39.

Table V-25: Results of reanalysis of NiSO4(aq) conductance data using a form of the Lee-Wheaton equation that is compatible with the SIT procedure. t/°C KA

(a) ± σKA (statistical)

Assessed (2σ)

log10KA Λo Reference

0 213.8 ± 4.4 2.33 68.82 [73KAT] 5 217.9 ± 4.9 2.34 80.48 [73KAT] 10 223.9 ± 4.9 2.35 92.96 [73KAT] 15 230.3 ± 4.7 2.36 106.01 [73KAT] 183.3 ± 30.4 2.26 105.28 [79FIS/FOX] 187.7 ± 2.1 2.27 103.63 [76SHI/TSU] 20 238.3 ± 5.3 2.38 119.82 [73KAT] 118.9 ± 7.2 2.08 123.34 [2000TSI/MOL] 25 247.3 ± 5.8 20 2.39 134.35 [73KAT] 220.7 ± 13.4 60 2.34 136.57 [1895FRA] 193.8 ± 1.9 30 2.29 132.99 [76SHI/TSU] 199.4 ± 28.5 60 2.30 132.73 [79FIS/FOX] 30 257.1 ± 7.1 2.41 149.50 [73KAT] 35 267.8 ± 7.8 2.43 165.16 [73KAT] 40 279.9 ± 8.8 2.45 181.47 [73KAT] 226.7 ± 2.8 2.36 179.12 [76SHI/TSU] 45 295.6 ± 9.4 2.47 198.58 [73KAT]

(a) The constant KA from the Lee-Wheaton treatment is assumed to be equivalent to K ο (V.88).

1 In 1979, Dr. A.D. Pethybridge kindly provided one of the reviewers (R.J. L.) with a version of the com-

puter program that has been used here (with modifications) for the analysis of conductance data.


As discussed in Appendix A, it appears that the reported molar conductance values from at least two of the studies [73KAT], [76SHI/TSU] are for “rounded” con-centration values. If the values reported are “smoothed” molar conductances, the true experimental scatter may have been greater. The “statistical” uncertainties listed in Table V-25 are the values from the data re-analysis [80PET/TAB]. Fitting of the (unavailable) primary data from those studies probably would have led to larger calculated statistical uncertainties in KA (assumed equivalent to K ο (V.88)). Estimated (2σ) uncertainties for all four values of K ο ((V.88), 298.15 K) are listed in Table V-25. Based solely on the conductance data at 25°C, at I = 0 the value of K ο ((V.88), 298.15 K) is (228.1 ± 15.5), or 10log K ο ((V.88), 298.15 K) = (2.36 ± 0.03).

Figure V-37: The variation of molar conductances of aqueous nickel sulphate solutions with concentration at 25°C. : [1895FRA], : [73KAT], : [76SHI/TSU], : [79FIS/FOX]

Earlier reviews [82WAG/EVA], [99ARC] gave substantial weight to the emf study of Nair and Nancollas [59NAI/NAN]. As shown in Appendix A, recalculation of values from that study to be consistent with the SIT procedure described in Appendix B, requires values for the dissociation constant of 4HSO− , and several interaction coeffi-cients, ε(Ni2+,Cl–), ε(H+,Cl–), ε(H+, 4HSO− ) and ε(Ni2+, 4HSO− ). The values calculated for the formation constant of NiSO4(aq) are extremely sensitive to the activity coeffi-cient model, to the values used for the interaction coefficients and to the values assumed

0 5 10 15 20 25 30 35 40180

190

200

210

220

230

240

250

260

mol

ar c

ondu

ctan

ce

Λ

/ S·

cm2 ·m

ol−1

104·[NiSO4(aq)] / mol·dm−3


for the first protonation constant for the sulphate ion. As discussed in Appendix A, this work can be used to determine a value at I = 0 of K ο ((V.88), 298.15 K) = (173 ± 80), or 10log K ο ((V.88), 298.15 K) = (2.24 ± 0.20).

If the estimated uncertainty bounds are considered, the value from Nair and Nancollas overlaps with those from the conductance studies. Using the weighted aver-age of the conductance and emf data ((228.1 ± 15.5) and (173 ± 80), respectively, weighted according to their uncertainties), the selected value of K ο ((V.88), 298.15 K) is:

K ο ((V.88), 298.15 K) = (226.1 ± 15.2)

or 10log K ο ((V.88), 298.15 K) = (2.35 ± 0.03).

None of the studies at higher ionic strength at 25°C is a systematic study in which the concentration of a single supporting electrolyte was varied. The results are summarised in Table V-26.

Table V-26: Experimental values for the first association constant of nickel sulphate determined in the presence of supporting electrolytes (see Appendix A). Values of

10 1log K ο were determined using the SIT (Appendix B).

Medium t / °C 10 1log K 10 1log K ο Reference

0.0484 m KClO4 – 0.2 1.70(a) 2.34 [56KEN] 0.258 m KClO3 – 0.8 (1.27 ± 0.30) (2.35 ± 0.30) [58KEN] 1.15 m KNO3 – 2.8 (0.69 ± 0.20) (2.16 ± 0.20) 0.2 M KNO3 25 (1.06 ± 0.20) (2.13 ± 0.20) [61TAN/OGI] 1.0 M NaClO4 25 (0.57 ± 0.20) (1.95 ± 0.20) [63TAN/SAI] 3.0 M LiClO4 25 0.26 1.07(b) [70MIR/MAK] 5.0 M NaClO4 25 1.19 1.70(c) [70TUR/RUV]

(a) The added nickel sulphate has a substantial effect on the ionic strength and the derived value of log10K1. (b) The value is anomalously low compared to all other reported values. (c) The ionic strength is beyond the range for use of the SIT (Appendix B).

Except for the spectroscopic value of Mironov et al. [70MIR/MAK], the results are qualitatively in agreement with the results of the conductivity and emf studies at low ionic strengths. The association constant values from the polarographic studies [61TAN/OGI], [63TAN/SAI], [70TUR/RUV] do tend to be somewhat smaller than the values selected above. The estimated uncertainties in the values from [61TAN/OGI] and [63TAN/SAI] are large, and the raw data are unavailable for re-analysis. Therefore, those results are not incorporated in the weighted-average selected value. The cryos-copic results from experiments in different saturated solutions [56KEN], [58KEN] are in excellent agreement with the value (2.33 ± 0.06) for 0°C from the conductance data of Katayama and the average value of (2.20 ± 0.06) from the two cryoscopic studies in


water [71ISO], [55BRO/PRU]. They are also in marginal agreement with the value (1.99 ± 0.10) recalculated from the emf data of Nair and Nancollas [59NAI/NAN].

At high sulphate concentrations, there is some evidence for formation of 2

4 2Ni(SO ) − in several studies (e.g., [58KEN], [59ROS/ROS], [63TAN/SAI]). However, the evidence is reasonably ambiguous, and may only reflect systematic errors in the experiments [58KEN]. No value is selected for 2K ο in the present review. There is con-siderable spectroscopic evidence that an inner-sphere sulphato complex exists at high sulphate concentrations and may become more important as the temperature is increased [77ASH/HAN], [78BEC/LIU]. Near 25°C, inner-sphere association is only about 10% of the outer-sphere association. No values are selected in this review for the formation constants of the inner sphere sulphato complexes [78BEC/LIU]. The polymeric species

42 4 4Ni (SO ) − proposed by Brintzinger and Osswald [34BRI/OSS] also is not considered

credible [41KIS/ACS].

V.5.1.2.1.1 Enthalpy of formation of NiSO4(aq)

There have been several studies [59NAI/NAN], [73KAT], [76SHI/TSU] of the tempera-ture dependence of the formation constant of the outer-sphere complex NiSO4(aq), and, as discussed above, several other measurements at temperatures between – 2.8 and 45°C. The recalculated values of 10 1log K ο

are plotted in Figure V-38.

The sets of values from most of the conductance and cryoscopic measurements are in good agreement. The temperature-dependencies of the values from these studies follow similar parallel curves. The higher temperature results of Nair and Nancollas [59NAI/NAN] follow the same type of curve for temperatures ≥ 15°C (288.15 K), but the values for the two lowest temperatures are inconsistent with the cryoscopic and con-ductivity results.

The recalculated values of 10 1log K ο from Katayama [73KAT] are used to cal-culate (5.3 ± 2.0) kJ·mol–1 for the average value of r mH ο∆ (V.88) between 0 and 45°C. Similarly, r mH ο∆ = (6.0 ± 2.5) kJ·mol–1 (average for 15 to 40°C) from the sparse data of Shimizu, Tsuchihashi and Furumi [76SHI/TSU] and r mH ο∆ = (12.6 ± 3.0) kJ·mol–1 (av-erage for 0 to 45°C, 2σ uncertainty) from the work Nair and Nancollas [59NAI/NAN].

Izatt et al. [69IZA/EAT] determined the enthalpy of reaction at 25°C by a titra-tion calorimetry method. The intent was to determine both the association constant and the enthalpy of association in a single experiment. Unfortunately, the enthalpy and asso-ciation constant values, as derived iteratively [69IZA/EAT] or by the method of least squares [73POW], are highly correlated. Powell showed that useful enthalpy values could be obtained only if a value was assumed for 1K ο . Starting from the 25°C value from Nair and Nancollas [59NAI/NAN], Powell used the data of Izatt et al. [69IZA/EAT] to obtain r mH ο∆ = (5.8 ± 0.2) kJ·mol–1. The calculation requires values for the enthalpy of protonation of sulphate and the enthalpy of dilution of the tetramethyl-ammonium sulphate titrant. Izatt et al. [69IZA/EAT] and Powell [73POW] also incor-


porated values for an association constant and enthalpy of association for the Me4NClO4(aq) ion pair. Determination of the values for Me4NClO4(aq) required using an explicit value for the association constant between Na+ and 2

4SO − . As discussed in Appendix A, a reasonable value from this work is r mH ο∆ = (5.8 ± 2.0) kJ·mol–1, based on Powell’s recalculation and an increased estimated uncertainty. In the present review, we exclude the value from Nair and Nancollas [59NAI/NAN], and accept the weighted average of the results from the other three studies [69IZA/EAT], [73KAT], [76SHI/TSU]:

r mH ο∆ ((V.88), 298.15 K) = (5.66 ± 0.81) kJ·mol–1.

From the selected values of 10 1log K ο and r mH ο∆ , the following quantities are calculated:

f mGο∆ (NiSO4, aq, 298.15 K) = – (803.2 ± 0.9) kJ·mol–1

f mH ο∆ (NiSO4, aq, 298.15 K) = – (958.7 ± 1.3) kJ·mol–1

and mS ο (NiSO4, aq, 298.15 K) = – (49.3 ± 3.1) J·K–1·mol–1.

Figure V-38: Values of 10 1log K ο for nickel sulphate association at different temperatures. : [1895FRA], : [55BRO/PRU], : [58KEN], : [59NAI/NAN], : [61TAN/OGI], : [63TAN/SAI], : [71ISO], : [73KAT], : [76SHI/TSU], : [79FIS/FOX], : [2000TSI/MOL].

270 280 290 300 310 320

2.0

2.1

2.2

2.3

2.4

2.5

log 10

K1o

T / K


V.5.1.2.1.2 Heat capacity values

Drakin, Madanat and Karlina [85DRA/MAD] reported measurements of the apparent molar heat capacities for nickel sulphate solutions at 50°C, 70°C and 90°C. The data are too sparse to be used to estimate the partial molar heat capacity of the ion pair.

V.5.1.2.2 Solid nickel sulphates

Several hydrated nickel sulphate solids from NiSO4·7H2O to NiSO4·H2O, including two forms of NiSO4·6H2O, a tetrahydrate, and a dihydrate have been reported [04STE/JOH], [39SIM/KNA], [39ROH], [64KOH/ZAS], though the domains of stability for and kinet-ics of interconversion between some of the intermediate hydrates remain contentious. In saturated solutions at 1 bar, the heptahydrate converts to the α-hexahydrate just above 300 K. Therefore, samples of the heptahydrate often contain amounts of the hexahy-drate, and samples of the α-hexahydrate can have the heptahydrate as a contaminant. This has complicated attempts to measure the physical properties for the individual sol-ids. Dehydration of the monohydrate to anhydrous NiSO4(s) occurs in a stream of wa-ter-saturated N2 only above 630 K [78SIR/SEM]. Decomposition of NiSO4(s) to NiO(s) and sulphur oxides becomes important at temperatures above 900 K [83GAU/BAL].

V.5.1.2.2.1 NiSO4·7H2O(cr)1

The heptahydrate is the stable nickel sulphate hydrate at 298.15 K. Stout et al. [66STO/ARC] reported low-temperature heat capacity measurements (from 16.1 K to temperatures above 279 K) for samples that were slightly hypostoichiometric and hy-perstoichiometric in water (NiSO4·6.867H2O(cr) and NiSO4·7.301H2O(cr)) with respect to the heptahydrate, and combined the results with those from an earlier study [41STO/GIA] to obtain mS ο (NiSO4·7.00H2O, cr, 298.15 K). There do not appear to be any other comparable measurements for this solid, and the entropy and heat capacity values for NiSO4·7.00H2O(cr) at 298.15 K from Stout et al. are selected in the present review. Though some of the auxiliary data used by Stout et al. differ slightly from those used in the present review, such differences are much less than the rather small esti-mated experimental uncertainties (cf. Appendix A for [66STO/ARC]). The selected values are:

mS ο (NiSO4·7.00H2O, cr, 298.15 K) = (378.95 ± 1.00) J·K–1·mol–1

,mpCο (NiSO4·7.00H2O, cr, 298.15 K) = (364.6 ± 4.0) J·K–1·mol–1.

The value of the difference between the enthalpy of solution of α-NiSO4·6.00H2O(cr) and NiSO4·7.00H2O(cr) in water ((7.68 ± 0.10) kJ·mol–1) has been reported by Stout et al. [66STO/ARC]. In the present review,

sol mH ο∆ (NiSO4·6.00H2O, α, 298.15 K) = (4.485 ± 0.200) kJ·mol–1 in water. Using this and the enthalpy of hydration, the value of sol mH ο∆ (NiSO4·7.00H2O, cr, 298.15 K) =

1 The dissolution of this compound is discussed in Section V.2.1.3.1.


(12.17 ± 0.23) kJ·mol–1 in water is obtained. The value 15.644 kJ·mol–1 was reported by Przepiera et al. [93PRZ/WIS] as the heat of solution of the heptahydrate in water to produce a solution with a 5 × 10–4 mass ratio of salt to water. If the value is assumed to be for the enthalpy of solution to this stated dilution (approximately 0.002 m), the appli-cation of a heat of dilution correction to I = 0 gives a value (12.5 kJ·mol–1) in good agreement with that based on Goldberg et al. and Stout et al.

Using the selected values of f mGο∆ (Ni2+) = – (45.77 ± 0.77) kJ·mol–1 (Section V.2.1.1), and sol mGο∆ = (12.94 ± 0.11) kJ·mol–1 (Section V.2.1.3) with standard auxiliary data,

f mGο∆ (NiSO4·7.00H2O, cr, 298.15 K) = – (2462.7 ± 0.9) kJ·mol–1.

With the selected entropy value,

mS ο (NiSO4·7.00H2O, cr, 298.15 K) = (378.9 ± 1.0) J·K–1·mol–1

is calculated and

f mH ο∆ (NiSO4·7.00H2O, cr, 298.15 K) = – (2977.3 ± 1.0) kJ·mol–1.

Bell [40BEL] reported vapour pressure measurements that showed that at 298.15 K removal of water from NiSO4·7H2O(cr) is less favoured than removal of D2O from NiSO4·7D2O (see Appendix A). In this review, no values are selected for the ther-modynamic quantities of NiSO4·7D2O.

V.5.1.2.2.1.1 α-NiSO4·6H2O

In contact with saturated solutions, the α-hexahydrate becomes the stable nickel sul-phate hydrate near 302 K, and is transformed to the β-hexahydrate at approximately 327 K. Stout et al. [66STO/ARC] reported low-temperature heat capacity measurements (from 13.7 to 321.0 K) for samples that were slightly hyperstoichiometric (NiSO4·6.010H2O) in water with respect to the hexahydrate, and combined the results with those for 1.1 to 21.8 K from an earlier study [64STO/HAD] and with results for the heptahydrate to obtain mS ο (NiSO4·6.00H2O, 298.15 K). There do not appear to be any other comparable measurements for this solid, and the entropy and heat capacity values for NiSO4·6.00H2O at 298.15 K from Stout et al. are selected in the present review. Though some of the auxiliary data used by Stout et al. differ slightly from those used in the present review, such differences are much less than the rather small estimated ex-perimental uncertainties (cf. Appendix A for [66STO/ARC]). The Reviewers select:

mS ο (NiSO4·6.00H2O, α, 298.15 K) = (334.45 ± 1.00) J·K–1·mol–1

,mpCο (NiSO4·6.00H2O, α, 298.15 K) = (327.9 ± 1.0) J·K–1·mol–1.

Values of the enthalpy of solution of α-NiSO4·6.00H2O in water have been re-ported by Stout et al. [66STO/ARC] and Goldberg et al. [66GOL/RID] and, allowing for the enthalpies of dilution, the values from the two sources are in good agreement.


The more extensive set of measurements by Goldberg et al. resulted in more dilute final solutions of nickel sulphate, and are preferred here. Goldberg et al. used enthalpy of dilution values of Lange [59LAN] and Lange and Miederer [56LAN/MIE] to extrapo-late the enthalpy of solution to I = 0, and obtained sol mH ο∆ = 4.81 kJ·mol–1 (as is normal for 2:2 electrolytes, the experimental heat of dilution for NiSO4 at low ionic strength changes much more strongly with concentration than would be predicted from simple Debye-Hückel theory, cf. Appendix A for [66GOL/RID]). The largest uncertainty in this heat of solution value comes from the extrapolation to I = 0. In the present review the selected value is:

sol mH ο∆ ((V.16), NiSO4·6.00H2O, α, 298.15 K) = (4.485 ± 0.200) kJ·mol–1.

This is one of the key quantities in the derivation of thermodynamic quantities of Ni2+ (cf. Section V.2.1.3).

Based on f mH ο∆ (NiSO4·7.00H2O, 298.15 K) = – (2976.94 ± 0.95) kJ·mol–1, and the difference between the enthalpies of solution of α-NiSO4·6.00H2O and NiSO4·7.00H2O in water ((7.68 ± 0.10) kJ·mol–1) reported by Stout et al. [66STO/ARC],

f mH ο∆ (NiSO4·6.00H2O, α, 298.15 K) = – (2683.8 ± 1.0) kJ·mol–1.

The value of f mGο∆ (NiSO4·6.00H2O, α, 298.15 K) is calculated from the se-lected values of f mH ο∆ (NiSO4·6.00H2O, 298.15 K) and mS ο (NiSO4·6.00H2O, α, 298.15 K) and the TDB auxiliary data in Table IV-1:

f mGο∆ (NiSO4·6.00H2O, α, 298.15 K) = – (2225.5 ± 1.1).

V.5.1.2.2.1.2 Other hydrated nickel sulphate solids

Solubility data [34BEN/THI], [39ROH] indicate that the β-hexahydrate is the stable solid in contact with saturated solutions of nickel sulphate in water for temperatures between 325 and 358 K, where lower hydrates begin to predominate. As might be ex-pected, addition of sulphuric acid results in the appearance of the lower hydrates at lower temperatures [39ROH].

Based on their differential scanning calorimetry results, Rabbering et al. [75RAB/WAN] reported r mH∆ = (6.53 ± 0.17) kJ·mol–1 for the reaction:

α-NiSO4·6.00H2O β-NiSO4·6.00H2O. (V.89)

Even a difference in the heat capacities of the two solids of 3 J·K–1·mol–1 would change this value by only 0.1 kJ·mol–1 (approximately) between 298.15 K and the temperature of the experiment. Therefore, a value of the transition enthalpy

r mH ο∆ ((V.89), α → β, 298.15 K) = (6.53 ± 0.20) kJ·mol–1 is assumed, and

f mH ο∆ (NiSO4·6.00H2O, β, 298.15 K) = – (2677.3 ± 1.0) kJ·mol–1


is retained for selection. If the transition temperature is 325 K, and the heat capacities of the two solids are assumed to be temperature independent and equal between 298.15 and 325 K, the selected value for molar entropy is:

mS ο (NiSO4·6.00H2O, β, 298.15 K) = (354.5 ± 5.0) J·K–1·mol–1

where the uncertainty has been increased to allow for the approximations in the calcula-tion and the uncertainty in the transition temperature. These selections yield:

f mGο∆ (NiSO4·6.00H2O, β, 298.15 K) = – (2224.9 ± 1.8) kJ·mol–1.

Rabbering et al. [75RAB/WAN] and Siracusa et al. [78SIR/SEM] carried out thermal decomposition studies by a variety of techniques, and found evidence for de-composition of the β-hexahydrate to a tetrahydrate at temperatures near 400 K and to the monohydrate near 440 K. Nandi et al. [79NAN/DES] found that the dehydration pattern for thermogravimetric measurements carried out at 5 K·min–1 were different for well-crystallised and powdered samples of the heptahydrate and hexahydrate. However, the intermediate hydrates were not well characterised.

Several studies [62CAI/POB], [64KOH/ZAS], [75RAB/WAN], [78SIR/SEM] have shown that the monohydrate does not readily lose water below 500 K. Goldberg et al. [66GOL/RID] determined the heat of solution of a nickel sulphate hydrate having the composition NiSO4·5.059H2O, and a value for the enthalpy of hydration of NiSO4·H2O(s) to the hexahydrate was calculated by assuming that the mixed solid was a mixture of NiSO4·6.000H2O(s) and NiSO4·H2O(s). It seems more likely that the solid was a mixture of NiSO4·6.000H2O(s) and NiSO4·4H2O(s) [65JAM/BRO]. Kohler and Zäske [64KOH/ZAS] reported equations for water vapour pressures as a function of temperature for mixtures of the hexahydrate and tetrahydrate, the tetrahydrate and a dihydrate, and the dihydrate and a monohydrate. Simon and Knauer [39SIM/KNA] also identified the tetrahydrate as a probable intermediate product in the isobaric decomposi-tion of the heptahydrate. The results from these studies are in rough agreement, but in the present review, no thermodynamic values are selected for these lower hydrates.

V.5.1.2.2.1.3 NiSO4(cr)

Although Goldberg et al. [66GOL/RID] indicated that samples of NiSO4(cr) dissolved too slowly in water to allow determination of its heat of solution, Adami and King [65ADA/KIN] reported the heat of solution in 4.360 m HCl(aq) at 303.15 K. The ther-modynamic cycle involved comparison with the heat of dissolution of metallic nickel in a similar medium. The results of [65ADA/KIN] have been recalculated in the present review (Appendix A) to obtain the selected value:

f mH ο∆ (NiSO4, cr, 298.15 K) = – (873.28 ± 1.57) kJ·mol–1.

Weller [65WEL] measured the heat capacity of anhydrous nickel sulphate be-tween 52.0 and 296.3 K. Stuve et al. [78STU/FER] also measured the heat capacity by adiabatic calorimetry, but primarily at lower temperatures, between 9.7 and 69.6 K (a


sharp thermal transition was observed at 35.5 K), and used their results with those of Weller [65WEL] to determine:

,mpCο (NiSO4, cr, 298.15 K) = (97.6 ± 0.1) J·K–1·mol–1

mS ο (NiSO4, cr, 298.15 K) = (101.3 ± 0.2) J·K–1·mol–1.

In the absence of other values, these values of mS ο (NiSO4, 298.15 K), ,mpCο (NiSO4, 298.15 K) are selected in the present review (with an estimated uncertainty

of ± 0.2 J·K−1·mol–1 for mS ο ). These selections yield:

f mGο∆ (NiSO4, cr, 298.15 K) = – (762.7 ± 1.6) kJ·mol–1.

Stuve et al. [78STU/FER] also reported the results of a limited set of drop calorimetry experiments (402.9 to 1001.5 K). The authors fitted an equation to the ex-perimental enthalpy differences such that the heat capacity values meshed smoothly with the value obtained from adiabatic calorimetry for 298.15 K. The authors indicated that their equation1 was applicable, within 0.6 per cent for temperatures between 298 and 1200 K. As discussed below, NiSO4 decomposes towards the upper end of this tem-perature range, and extrapolation of the heat capacities to 1200 K does not seem justi-fied. Conversion of the equation from calorie to joule units leads to:

1000K,m 298.15K[ ]pCο (NiSO4, cr) = (118.934 + 3.6761 × 10–2 (T/K) – 2.8681 × 106 (T/K)–2)

J·K–1·mol–1

and this equation for ,mpCο (T) from Stuve et al. [78STU/FER] is accepted in the present review.

There have been several studies of the equilibrium conversion of NiSO4 to NiO from 880 K to 1200 K according to:

NiSO4 NiO + SO3

with possible concurrent decomposition of SO3:

SO3 SO2 + ½O2.

Both pressure measurements [25MAR], [78KAR/MAL] and electrochemical measurements [66ING], [64ALC/SUD], [73SKE/ESP], [77SAD/KAW], [83GAU/BAL] have been reported. The measurements lead to values of f mGο∆ that agree within 5 – 6 kJ·mol–1 near 900 K, and 7 – 8 kJ·mol–1 at 1150 K. The values from Alcock et al. [64ALC/SUD] show better agreement with the vapour-pressure study of Karwan et al. [78KAR/MAL] than do the other electrochemical studies. In the present review, no at-tempt has been made to calculate heat-capacity values from the second derivatives of equations fitted to these measurements.

1 The rather similar equation for mH (T) – mH ο (298 K), probably based on prepublication results from

[78STU/FER], is incorrect as published in the review of Mah and Pankratz [76MAH/PAN].


V.5.2 Selenium compounds and complexes Nickel-selenium compounds have recently been reviewed by the NEA TDB review team on Se, (see [2005OLI/NOL] for the corresponding NEA TDB critical review).

V.5.3 Tellurium compounds and complexes V.5.3.1 Nickel tellurides

The binary phase diagram nickel/tellurium as summarised by Lee and Nash [90LEE/NAS] shows a complex behaviour with six intermetallic solid phases. This dia-gram is essentially based on a seminal paper authored by Klepp and Komarek [72KLE/KOM]. The phases β1, β2, 1′β and NiTe2–x(cr) have variable stoichiometry.

In an early stage of this review it was decided that non-stoichiometric alloy systems will not be covered. Thus only the two γ-phase line components, NiTe0.775(γ1) and NiTe0.85(γ2) among the six solid phases, need to be discussed here. The thermody-namic data of the Ni – Te system have been critically assessed by Ball et al. [92BAL/DIC].

V.5.3.1.1 NiTe0.775(γ1)

NiTe0.775(cr) is orthorhombic and has the following unit cell parameters: a = 3.916 Å, b = 6.860 Å, c = 12.32 Å [72KLE/KOM]. The γ1-phase appears to be stable up to 1015.5 K, when it decomposes peritectoidally into γ2 and β2. There have not been any heat capacity measurements reported for γ1. Ball et al. [92BAL/DIC] obtained estimates of ,mpCο (NiTe0.775, cr, 298.15 K) = 45.8 (48.7) J·K–1·mol–1 and mS ο (NiTe0.775, cr, 298.15 K) = 71.0 (67.6) J·K–1·mol–1 by taking a weighted average of heat capacities of the β and NiTe2–x(cr) phases. The values in brackets have been calculated with the cationic and anionic contributions to the heat capacity at 298 K and the revised ‘Latimer’ en-tropy contributions [93KUB/ALC]. If reasonable uncertainties of ± 5.0 J·K−1·mol–1 are chosen for ,mpCο∆ and mS ο∆ , both methods of estimation lead to results that are consis-tent within these limits.

The enthalpy of formation of NiTe0.775(cr) has been estimated from emf meas-urements in the γ + NiTe2–x(cr) two-phase field [92BAL/DIC]. Neither of these standard entropy or standard enthalpy estimations ( f mH ο∆ (NiTe0.775, cr, 298.15 K) = − (41.7 ± 4.4) kJ·mol–1) is based directly on measurements of the NiTe0.775(cr) phase, so no recommendation can be given in this review. This is not of any particular importance here as the γ1 phase is of peripheral interest only.

V.5.3.1.2 NiTe0.85(γ2)

Based on thermoanalytical studies Klepp and Komarek [72KLE/KOM] assigned some-what arbitrarily a composition of NiTe0.85(cr) to the high temperature γ2 phase. An at-tempt to characterise this phase by its X-ray powder diffraction pattern failed, because below 963 K it decomposes quickly into NiTe0.775(cr) + NiTe2–x(cr) and the phase equi-


librium could not be frozen in by quenching [72KLE/KOM]. Therefore, estimated thermodynamic quantities for the γ2 phase are even less reliable than for the γ1 phase. According to Ball et al. [92BAL/DIC], the same values of the standard heat capacity and entropy were estimated for both γ phases (based on x(Ni) + x (Te) = 1), resulting in

,mpCο (NiTe0.85, cr, 298.15 K) = 47.7 (50.7) J·K–1·mol–1 and mS ο (NiTe0.85, cr, 298.15 K) = 74.0 (70.7) J·K–1·mol–1. The values in brackets have been calculated with the cationic and anionic contributions to the heat capacity at 298 K and the revised ‘Latimer’ en-tropy contributions [93KUB/ALC]. The latter estimation method clearly leads also to identical heat capacity and entropy values for both γ phases based on x(Ni) + x (Te) = 1.

The enthalpy of formation of NiTe0.85(cr) has again been estimated from emf measurements in the γ + NiTe2–x(cr) two-phase field [92BAL/DIC]. Without confirma-tion by experimental data the estimated standard entropy and standard enthalpy ( f mH ο∆ (NiTe0.85, cr, 298.15 K) = – (50.1 ± 5.6) kJ·mol–1) of this phase cannot be rec-ommended in this review.

V.6 Group 15 compounds and complexes V.6.1 Nitrogen compounds and complexes V.6.1.1 Solid nickel nitrates

The hydrated nickel nitrate solids have been the subject of sporadic thermodynamic studies over the last 150 years, but the basic thermodynamic quantities for these materi-als are not well defined. Though careful synthesis yields solids with various Ni(NO3)2:H2O ratios, the regions of stability (if any) for the lower hydrates are still not well established [1899FUN], [34SIE/SCH], [98ELM/GAB], [2001MIK/MIG]. When hydrated nickel nitrate is heated at a pressure of approximately 0.1 MPa, the last water of hydration is lost between 450 and 500 K. Release of nitrogen oxides occurs at tem-peratures below 600 K, and continued heating leads to formation of basic nickel nitrates and then NiO [98ELM/GAB], [2001MIK/MIG].

V.6.1.1.1 Ni(NO3)2·9H2O(cr)

This solid has been reported as a stable solid in the Ni(NO3)2 – H2O system for tempera-tures below 270 K [34SIE/SCH]. The temperature for the equilibrium conversion of the nonahydrate to the hexahydrate in contact with saturated aqueous solution has not been well established. No values for Ni(NO3)2·9H2O(cr) are selected in the present review.

V.6.1.1.2 Ni(NO3)2·6H2O(cr)

The stable hydrate in equilibrium with a solution saturated in nickel nitrate at 298.15 K is Ni(NO3)2·6H2O(cr) [34SIE/SCH]. The water content of the commercially available solid tends to vary slightly. The solid can easily lose water on exposure to dry air [72AUF/CAR], [98ELM/GAB], but has also been reported to be slightly deliquescent in

V.6 Group 15 compounds and complexes 197

moist air [72AUF/CAR]. In a closed system, the hydrate begins to melt (or partially dissolve in its water of hydration) at 328 K [2001MIK/MIG]. Auffredic, Carel and Weigel [72AUF/CAR] measured the enthalpy of solution of the hexahydrate in water at 298.15 K. Calorimetric measurements were done for ratios of H2O:Ni(NO3)2·6H2O be-tween 500 (0.11 m) and 10000 (0.0056 m), and a value of (31.23 ± 1.00) kJ·mol–1 for

sol mH ο∆ was reported. For the more dilute solutions, there is considerable scatter in the measurements as a function of molality, and the reported value of sol mH ο∆ was simply an extrapolated value from almost random values for concentrations between 0.01 m and 0.0056 m. As discussed below, and in the Appendix A discussion for [72AUF/CAR], the lack of concentration dependence over this concentration range is not unexpected. The average value of the enthalpies of solution for final concentrations below 0.01 kg·mol–1 is (31.28 ± 0.87) kJ·mol–1. A reasonable estimate of the integral heat of dilution of a nickel nitrate solution from a solution with a solvent:salt ratio of 8000:1 is – (0.52 ± 0.10) kJ·mol–1 [59LAN] (see the Appendix A discussion of [72AUF/CAR]). In the present review, this estimated heat of dilution is applied to the average value of the enthalpies of solution for final concentrations below 0.01 kJ·mol–1, and for

Ni(NO3)2·6.00H2O(cr) Ni2+ + 2 3NO− + 6H2O(l) (V.90)

r mH ο∆ ((V.90), 298.15 K) = (30.76 ± 1.00) kJ·mol–1

is selected. Values based on measurements of the vapour pressure of water over the hexahydrate [72AUF/CAR] were not incorporated in the calculation of the assessed value because, as discussed below, the stoichiometry of the dehydrated salt in equilib-rium with the hexahydrate and H2O(g) is not firmly established. This selection yields, using NEA-TDB auxiliary data:

f mH ο∆ (Ni(NO3)2·6.00H2O, cr, 298.15 K) = – (2214.5 ± 1.6) kJ·mol–1.

Goldberg et al. [79GOL/NUT], reviewed the available literature data, and pro-vided tables of recommended values for the activity coefficients and osmotic coeffi-cients for solutions of Ni(NO3)2 in water at 298.15 K. However, there were no studies that allowed the tables to be extended to saturated solutions of Ni(NO3)2·6.00H2O(cr) at 298.15 K (5.47 m [34SIE/SCH]), and no major studies seem to have been published since that review. No values of sol mG∆ at saturation are estimated here.

Heat capacity measurements for Ni(NO3)2·6H2O(cr) were reported by Vasileff and Grayson-Smith [50VAS/GRA] for temperatures from 65 to 300 K. The authors es-timated the entropy change between 65 and 273 K to be 364.4 J·K–1·mol–1 (87.1 cal·K−1·mol–1). From the set of smoothed specific heats, the selected value of the molar heat capacity of the solid at 298.15 K is calculated to be

,mpCο (Ni(NO3)2·6H2O, cr, 298.15 K) = (463.0 ± 2.0) J·K–1·mol–1,


where the uncertainty has been estimated in the present review. The temperature range of the measurements is not adequate for the calculation of a value for

mS ο (Ni(NO3)2·6H2O, cr, 298.15 K).

V.6.1.1.3 Ni(NO3)2·4H2O(cr), Ni(NO3)2·3H2O(cr), Ni(NO3)2·2H2O(cr)

Dehydration of the hexahydrate leads to several lower hydrates, but the hydrate formed seems to depend markedly on the method used to carry out the dehydration. In few of the experiments were the dehydrated solids thoroughly characterised, nor was it estab-lished that the solids were stable over long periods.

Chukurov et al. [73CHU/DRA] and Auffredic, Carel and Weigel [72AUF/CAR] have reported measurements of the heat of solution of a compound with the stoichiometry Ni(NO3)2·4H2O(cr) in water (8.49 kJ·mol–1 and 8.66 kJ·mol–1, respec-tively). The latter group prepared their solid by dehydration of the hexahydrate over calcium chloride at room temperature. They also reported [72AUF/CAR2] a value of the heat of solution (– 27.41 kJ·mol–1) of Ni(NO3)2·2H2O(cr) (prepared by dehydration of the tetrahydrate over calcium chloride at 50°C). Recalculation of the heats of dilution (cf. Appendix A for [72AUF/CAR]) gives heats of solution, corrected to “infinite dilu-tion”, for the dihydrate of – (27.64 ± 1.00) kJ·mol–1 and for the tetrahydrate (8.23 ± 1.00) kJ·mol–1 (based on [72AUF/CAR]) and (8.49 ± 1.00) kJ·mol–1 (based on [73CHU/DRA]).

In contrast to Funk [1899FUN] who reported a stable trihydrate phase (55 to 95°C), Sieverts and Schreiner [34SIE/SCH] found that both the tetrahydrate (55 to 80°C) and dihydrate (85 to 120°C) were stable phases in the Ni(NO3)2 – H2O system, and could also be formed at equilibrium at 25°C in the Ni(NO3)2 – H2O – HNO3 system at high nitric acid mole fractions. Studies of the dehydration of the hexahydrate using differential thermal analysis [74WEN], [98ELM/GAB], differential scanning calo-rimetry [99MIK/MIG] and thermogravimetric analysis [66KAL/PUR], [2001MIK/MIG] suggest that although a tetrahydrate can be obtained, it is easily converted to the tri-hydrate at slightly higher temperatures. For example, Kalinichenko and Purtov [66KAL/PUR] reported that dehydration of the tetrahydrate to the trihydrate occurred between 164°C and 170°C.

Sano [37SAN] and Auffredic, Carel and Weigel [72AUF/CAR], reported measurements of the vapour pressure of water over Ni(NO3)2·6H2O(cr) as a function of temperature (Figure V-39). The measured vapour pressures differ systematically by a factor of 1.5 at similar temperatures. Sano interpreted his results in terms of the reac-tion:

Ni(NO3)2·6.00H2O(cr) Ni(NO3)2·3.00H2O(cr) + 3H2O(g) (V.91)

whereas Auffredic, Carel and Weigel interpreted their results in terms of the reaction:

Ni(NO3)2·6.00H2O(cr) Ni(NO3)2·4.00H2O(cr) + 2H2O(g). (V.92)


As can be seen from the figure, the slope of the plot of – log10 2H Op against 1/T is similar for both sets of results. This may indicate that whichever lower hydrate was involved in the equilibrium, it was the same solid for both sets of experiments. If the vaporisation corresponds to Reaction (V.92), the calculated enthalpy of vaporisation (110 kJ·mol–1) corresponds well with 116 kJ·mol–1, derived from the enthalpies of solu-tion of the hexahydrate [72AUF/CAR] and the tetrahydrate [72AUF/CAR], [73CHU/DRA]. This suggests that the vaporisation follows Reaction (V.92), rather than (V.91).

Similarly, the enthalpy for Reaction (V.93),

Ni(NO3)2·4.00H2O(cr) Ni(NO3)2·2.00H2O(cr) + 2H2O(g) (V.93)

r mH ο∆ = 124 kJ·mol–1, as derived from the reported calorimetric data [72AUF/CAR], [72AUF/CAR2], is similar to the value obtained from measurements of the vapour pres-sure of water over the tetrahydrate (122 kJ·mol–1) [72AUF/CAR2]. Nevertheless no value is selected in the present review for the Gibbs energy of Reaction (V.92) or (V.93) because there was no proper characterisation of the lower hydrate at equilibrium during any of these measurements.

Figure V-39: Measured water vapour pressures over hydrated Ni(NO3)2·6H2O(cr) [37SAN], [72AUF/CAR] and Ni(NO3)2·4H2O(cr) [72AUF/CAR2].

2.9 3.0 3.1 3.2 3.3 3.4 3.5-3.0

-2.8

-2.6

-2.4

-2.2

-2.0

-1.8

-1.6

-1.4

[37SAN] [72AUF/CAR] [72AUF/CAR2]

log 10

(p(H

2O)/b

ar)

103·K/T


In the present review, the average calorimetry-based value is selected for the enthalpy of dissolution of the tetrahydrate:

Ni(NO3)2·4.00H2O(cr) Ni2+ + 2 3NO− + 4H2O(l) (V.94)

r mH ο∆ ((V.94), 298.15 K) = (8.36 ± 0.80) kJ·mol–1.

The selected uncertainty is slightly larger than the statistical uncertainty be-cause the uncertainty in the heat of dilution is common to both results. The correspond-ing value for the dihydrate [72AUF/CAR2] has also been selected:

Ni(NO3)2·2.00H2O(cr) Ni2+ +2 3NO− + 2H2O(l) (V.95)

r mH ο∆ ((V.95), 298.15 K) = – (27.64 ± 1.00) kJ·mol–1.

In neither case should the selection of a value for the enthalpy of solution be taken to mean that it is certain that the solid is a stable phase in the Ni(NO3)2-H2O sys-tem. These selections yield, using NEA-TDB auxiliary data:

f mH ο∆ (Ni(NO3)2·4.00H2O, cr, 298.15 K) = – (1620.4 ± 1.4) kJ·mol–1

and

f mH ο∆ (Ni(NO3)2·2.00H2O, cr, 298.15 K) = – (1012.7 ± 1.6) kJ·mol–1.

V.6.1.1.4 Ni(NO3)2(anhydrous)

Except for salts containing singly charged cations, synthesis of anhydrous nitrate salts is very difficult, as the final steps of dehydration tend to cause loss of nitrogen oxides or nitric acid. Although some authors [82MU/PER], [2001MIK/MIG] claim anhydrous nickel nitrate forms as an intermediate decomposition product during heating of the hydrate, only Chukurov et al. [73CHU/DRA] claim to have carried out any chemical thermodynamic measurements ( f mH ο∆ = – (402.1 ± 2.9) kJ·mol–1, from determination of the heat of solution in dimethyl sulphoxide). The summary of their paper, all that was available to the reviewers, contains no details on the preparation, handling or analysis of the anhydrous salt. No values for the thermodynamic quantities for anhydrous Ni(NO3)2 are selected in the present review.

V.6.1.2 Aqueous Ni(II)-nitrato complexes

There are conflicting opinions concerning the interaction of Ni(II) with nitrate ion. Mi-ronov et al. have failed to detect nitrato complexes of Ni(II) [70MIR/MAK]. On the other hand, neutron diffraction of 2 and 4.3 m Ni(NO3)2 solutions revealed ion pair for-mation between Ni2+ and nitrate ions [97HOW/NEI]. Quantitative information on nitrate complexation (ion pair formation) of Ni(II) has been published in [73FED/SHM] and [73HUT/HIG]. The Russian authors reported formation of mono-, bis- and tris-complexes [73FED/SHM], while only 3NiNO+ was reported in [73HUT/HIG]. In light of the results collected for similar weak complexes of Ni(II) (e.g., for halogeno com-plexes), the formation of higher complexes ( 2

3Ni(NO ) nn

− , n > 1) is unlikely, except in


highly concentrated nitrate solutions. Therefore, the formation of such complexes is not considered in this review. The available association constant values for Reaction (V.96):

Ni2+ + 3NO− 3NiNO+ (V.96)

are reported in Table V-27.

The experimental data reported in [73FED/SHM] were re-evaluated taking into account only formation of the 3NiNO+ species (see Appendix A). To minimise the me-dium effect, only half of the experimental data, for which [ 3NO− ] ≤ [ 4ClO− ], were taken into account. The recalculated constants are also listed in Table V-27.

Table V-27: Formation constants (logarithmic values) of the 3NiNO+ complex (ion pair).

Method Medium t (°C) 10 1log b (a) 10 1log b (b) 10 1log οb (c) Reference

kin 1 M Na(ClO4, NO3) 25 45

(0.04 ± 0.08) (0.03 ± 0.12)

(0.02 ± 0.40) (0.01 ± 0.40)

(0.97 ± 0.60)

[73HUT/HIG]

sol 4 M Li(ClO4, NO3)

3 M Li(ClO4, NO3)

2 M Li(ClO4, NO3)

1 M Li(ClO4, NO3)

0 corr

25 – (0.30 ± 0.04)

– (0.55 ± 0.09)

– (0.44 ± 0.06)

– (0.22 ± 0.03)

(0.45 ± 0.08)(e)

– (0.01 ± 0.40)(d)

– (0.30 ± 0.40)(d)

– (0.21 ± 0.40)(d)

– (0.20 ± 0.40)(d)

(0.49 ± 0.45)

[73FED/SHM]

(a) Reported values. (b) Accepted values corrected to molal scale. (c) Corrected to I = 0 by means of the SIT. (d) Recalculated value, taking into account the solely formation of 3NiNO+ species (see Appendix A). (e) Extrapolated value, using the Vasil’ev equation.

Extrapolation of 10 1log b values to Im = 0 in lithium perchlorate media are shown in Figure V-40. The weighted linear regression resulted in the values of

10 1log οb (V.96) = (0.49 ± 0.45) and ∆ε((V.96), LiClO4) = – (0.08 ± 0.14) kg·mol–1. From the latter value ε( 3NiNO+ , 4ClO− ) = (0.44 ± 0.14) kg·mol–1 can be derived. Calcu-lation of this value also requires the value of ε(Ni2+, 3NO− ) and, in the absence of data to allow calculation of ε(Ni2+, 3NO− ) while also considering association, the value derived in Section V.6.1.2.1 is used. Based on this result, ∆ε((V.96), NaClO4) = (0.04 ± 0.15) kg·mol–1 can be estimated, and this is used to determine 10 1log οb from the value of

10 1log b reported in [73HUT/HIG]. Both values of 10 1log οb in Table V-27 are derived from constants at higher ionic strength that have large uncertainties because of medium effects. Therefore, in this review the lower value, calculated from [73FED/SHM], is selected with an increased uncertainty:

10 1log οb ((V.96), 298.15 K) = (0.5 ± 1.0).


Only a single value is reported for the enthalpy of Reaction (V.96) [75ARU]. Recalculating the experimental data using the tentative values of 10 1log οb ((V.96), 298.15 K) and ∆ε((V.96), Ni(ClO4)2) = – (0.11 ± 0.15) kg·mol–1, r mH∆ ((V.96), 298.15 K) = (6.0 ± 4.0) kJ·mol–1 can be estimated for I = 1 M (Ni(ClO4)2), from the data reported in [75ARU] (see Appendix A). The selected value yields, using NEA-TDB auxiliary data:

f mGο∆ ( 3NiNO+ , 298.15 K) = – (159.4 ± 5.8) kJ·mol–1.

Figure V-40: Extrapolation to I = 0 of the constants for Reaction (V.96) in lithium per-chlorate media based on the data in [73FED/SHM]. The recalculated and original values are marked with full and open squares, respectively (see Appendix A).

-0.5

0

0.5

1

1.5

2

0 1 2 3 4 5

I / mol·kg−1

log 1

0 β 1

+ 4

D

Ni2+ + 3NO− 3NiNO+


V.6.1.2.1 Determination of the Ni2+ – 3NO− ion interaction coefficient

Due to the low and rather uncertain value of 10 1log οb ((V.96), 298.15 K), the descrip-tion of the Ni(II)-nitrate system in terms of ion-ion interactions may be in many cases more straightforward. Therefore, the ion interaction coefficient ε(Ni2+, 3NO− ) is also derived in this review from the osmotic and mean activity coefficients of Ni(NO3)2 solu-tions, using Equations (V.74) and (V.75).

The most reliable values for the mean activity and osmotic coefficient of aque-ous Ni(NO3)2 solutions at 298.15 K are given in [82SAD/LIB] and [82SAR/COV]. Plots of 10log γ ± and φ as a function of the molality, m, of Ni(NO3)2 (m = Im/3) are depicted in Figure V-41 and Figure V-42, respectively. The two data sets of activity coefficients values are somewhat different, but the disagreement is acceptable, taking into account the similarity of osmotic coefficients. A fit of Equations (V.74) and (V.75) to the ex-perimental data up to an ionic strength of 15 mol·kg–1 yields ε(Ni2+, 3NO− ) = (0.182 ± 0.010) kg·mol–1, see solid lines in Figure V-41 and Figure V-42.

Figure V-41: Plot of 10log γ ± versus molality of aqueous Ni(NO3)2 solutions at 298.15 K. Experimental data from: [82SAD/LIB] ( ); [82SAR/COV] ( ). Solid line: SIT, ε(Ni2+, 3NO− ) = 0.182.

-0.4

-0.2

0

0.2

0.4

0.6

0.8

-0.2 0.4 1.0 1.6 2.2 2.8 3.4 4.0 4.6 5.2

[Ni(NO3)2] / mol·kg−1

log 1

0 γ±


Figure V-42: Osmotic coefficient plotted as a function of molality of aqueous Ni(NO3)2 solutions at 298.15 K. Experimental data from: [82SAD/LIB] ( ); [82SAR/COV] ( ). Solid line: SIT, ε(Ni2+, 3NO− ) = 0.182.

V.6.2 Phosphorus compounds and complexes V.6.2.1 Solid nickel phosphorus compounds

V.6.2.1.1 Nickel phosphides

The thermodynamic measurements for nickel phosphides were recently summarised by Schlesinger [2002SCH] (in that reference, red phosphorus (V) was used as the standard state for phosphorus). Schlesinger concluded that the details of the Ni-P phase diagram remain controversial, and that the available data are insufficient to allow calculation of assessed thermodynamic values for any nickel phosphides. No values are selected in the present review.

V.6.2.1.2 Nickel phosphates

A number of nickel phosphate solids have been reported, such as Ni3(PO4)2·8H2O [66GME], Ni3(PO4)2·7H2O [66GME], Ni3(PO4)2·1.25H2O [86TOD/HAS], NiHPO4·3H2O [87CUD/LEC], (NiHPO4)2·3H2O [87CUD/LEC], Ni(H2PO4)2·2H2O

0.8

1

1.2

1.4

1.6

1.8

2

2.2

2.4

2.6

-0.2 0.4 1 1.6 2.2 2.8 3.4 4 4.6 5.2

[Ni(NO3)2] / mol·kg−1

φ


[90TRO] and Ni3(PO4)2 [94HAS/TOD]. Nevertheless, chemical thermodynamic data for these solids are almost non-existent.

The most thoroughly studied solid is Ni3(PO4)2·8H2O, which is reported [96DOJ/NOV] to undergo three phase transitions below 300 K (at 174, 247 and 289 K). Though solubility measurements have been reported for Ni3(PO4)2·8H2O [61CHU/ALY2], [82CHI/SAB], Ni3(PO4)2·7H2O [76PAN/PAT] and NiHPO4 [76PAN/PAT], none are of adequate quality to allow chemical thermodynamic quanti-ties to be calculated for these solids (cf. Appendix A).

V.6.2.2 Aqueous nickel phosphorus species

V.6.2.2.1 Simple nickel phosphato complexes

The literature on complex formation between phosphate and Ni(II) ions in solution is not extensive (see Table V-28). Most measurements have been based on pH titrations [67SIG/BEC], [72FRE/STU], [76TAY/DIE], [96SAH/SAH], though there have also been polarographic [91SWI/PAW] and solubility [89ZIE/JON] studies. Caminiti [82CAM] showed that in acidic highly concentrated solutions oxygen atoms from phos-phate groups could replace water molecules in the inner sphere of hydrated Ni2+ ions. However, under most conditions the complexes are weak, and difficult to identify un-ambiguously because of protonation equilibria involving both the ligand and the com-plexes. Based on a curve-fitting analysis of potentiometric titration curves, Childs [70CHI] suggested a wide range of species: +

2 4MH PO , MHPO4(aq), 3 4 2MH (PO )− , M2H2(PO4)2(aq) and 4MPO− . Few of these have been reported in the Ni(II) – phosphate system.

Most of the studies [67SIG/BEC], [72FRE/STU], [76TAY/DIE], [91SWI/PAW], [96SAH/SAH] were inspired by possible parallels between phosphate complexation and biochemical interactions between phosphate esters and metal ions. These studies have been carried out over a fairly limited pH range (usually between 4 and 6) at low ionic strength ( ≤ 0.2 M). In none of these publications was the primary data published. Most papers did report experimental values for one or more protonation constants of the phosphate ion itself. This provides some basis for qualitative evaluation of the experiments. Except in the work of Taylor and Diebler [76TAY/DIE], the authors interpreted their results in terms of a single complex, NiHPO4(aq),

Ni2+ + 24HPO − NiHPO4(aq), (V.97)

and except for the polarographic result of Swiatek and Pawlowski [91SWI/PAW], the reported values are consistent within the experimental uncertainties. For reasons dis-cussed in Appendix A, the result from the polarographic study is not used here. The other values have been corrected to I = 0, and the weighted average of the three results at 25°C provides the selected values:

10log K ο ((V.97), 298.15 K) = (3.05 ± 0.09)


and r mGο∆ ((V.97), 298.15 K) = – (17.4 ± 0.5) kJ·mol–1.

The value 10log K ο ((V.97), 298.15 K) = (2.88 ± 0.30) obtained from the work of Frey and Stuehr [72FRE/STU] is compatible with this selection, which yields, using NEA-TDB auxiliary data:

f mGο∆ (NiHPO4, aq, 298.15 K) = – (1159.2 ± 1.8) kJ·mol–1.

Ziemniak, Jones and Combs [89ZIE/JON] reported solubilities of Ni(II) oxide in aqueous sodium phosphate solutions (pH298 > 10) for temperatures from 290 to 560 K. As discussed in Appendix A and elsewhere [95LEM], it is difficult to interpret these solubility results quantitatively because of the limited range in hydroxide concen-tration and the wide variation in the ionic strength. The values for the formation con-stants of NiHPO4(aq) from the pH studies [67SIG/BEC], [96SAH/SAH] and for the species, 2

2 4Ni(OH) HPO − , as proposed by Ziemniak, Jones and Combs [89ZIE/JON], [95LEM] are not inconsistent for the pH values at which the experimental measure-ments were obtained.

Table V-28: Experimental values for the association constants and other thermodynamic quantities for nickel phosphate complexes.

Method t (°C) Medium (aq) Reported value Recalculated value (a)

10log K ο Reference

ligand: 24HPO −

Ni2+ + 24HPO − NiHPO4(aq)

pot 25.0 0.1 M NaClO4 10 1log K = 2.08 (2.97 ± 0.30) [67SIG/BEC]

pot 15.0 0.1 M KNO3 10 1log K = 2.00 (2.88 ± 0.30) [72FRE/STU]

pot 25.0 0.1 M NaClO4 10 1log K = 2.11 (3.00 ± 0.20) [76TAY/DIE]

pol 25.0 0.2 M NaClO4 10 1log K = 3.26 [91SWI/PAW]

pot 25.0 0.1 M NaNO3 10 1log K = 2.20 (3.07 ± 0.10) [96SAH/SAH]

β-Ni(OH)2(cr) + 24HPO − 2

2 4Ni(OH) HPO −

sol 25.0 0.1 M NaClO4 10log K = – (5.0 ± 1.0) [89ZIE/JON]

ligand: 2 4H PO−

Ni2+ + 2 4H PO− 2 4H NiPO+

pot 25.0 0.1 M Li/NaClO4 10 1log K = 0.5 [76TAY/DIE] (a) Uncertainties are the statistical uncertainties from the re-analysis of the data.

The overall estimated uncertainties are discussed in the text.


V.6.2.2.2 Aqueous nickel diphosphato complexes

As in the NEA-TDB uranium review [92GRE/FUG], the only polyphosphate(V) species considered in the present review are the pyrophosphates (diphosphato complexes). Other polyphosphoric acid species have negligible equilibrium concentrations at total phosphate concentrations < 0.045 mol·dm–3 and at temperatures below 200°C [74MES/BAE]. There have been potentiometric (pH measurement) [56YAT/VAS], [64HAM/MOR], [73PER/SEC], calorimetric [56YAT/VAS2], spectrophotometric [58VAI/RAM] and amperometric [78KAR/HUB] studies of the formation of pyrophos-phate complexes of Ni(II) (Table V-29).

The complexes of the highly charged pyrophosphate ion with nickel are gener-ally stronger than the phosphate complexes, but interpretation of the experiments is be-set by the same difficulties as the interpretation of the phosphate studies with respect to unambiguous identification of the species. In the literature, results of studies of associa-tion of Ni2+ with pyrophosphate have been discussed in terms of the following equilib-ria:

Ni2+ + 42 7P O − 2

2 7NiP O − (V.98)

Ni2+ + 32 7HP O − 2 7NiHP O− (V.99)

Ni2+ + 2 42 7P O − 6

2 7 2Ni(P O ) − (V.100)

The results of Vaid and Rama Char [58VAI/RAM] were obtained under ill-defined conditions, and are in poor agreement with results from the other studies. The equilibrium constant of Yatsimirskii and Vasilev for Reaction (V.98) is consistent with other values, but was obtained at moderately high and variable ionic strength. Values from the measurements of Hammes and Morrell [64HAM/MOR] in 0.1 M Me4NCl(aq) for Reactions (V.98) and (V.99), after correction to I = 0, are 10log K ο (V.98) = (8.73 ± 0.25) and 10log K ο (V.99) = (5.14 ± 0.25), at 298.15 K.

Perlmutter-Hayman and Secco [73PER/SEC] did careful measurements of nickel pyrophosphate complexation at four different ionic strengths (KNO3(aq)), though for reasons discussed in Appendix A, only their values obtained in 0.1 M and 0.2 M KNO3(aq) are considered to be consistent. As discussed in Appendix A, the authors neglected association of pyrophosphate with K+, though the values they used for pyro-phosphate protonation were determined in the corresponding aqueous KNO3 solutions.

The apparent 298.15 K values (corrected to I = 0) of 10log K (V.98) ((8.73 ± 0.25) [64HAM/MOR], (7.69 ± 0.25) and (7.78 ± 0.30) [73PER/SEC]) are in poor agreement; the value of Frey and Stuehr [72FRE/STU] for 288.15 K (7.94 ± 0.25) is in only marginal agreement with that of Perlmutter-Hayman and Secco (7.53 ± 0.25), and the value at 303.15 K of Karwelk and Huber [78KAR/HUB] is lower than would be expected from the results of Perlmutter-Hayman and Secco. The disagreement is pri-marily the result of neglect of complexation of the singly-charged cations, K+ and

4NH+ , with pyrophosphate. As indicated in Appendix A, use of reasonable potassium


pyrophosphate complexation constant values brings the results of Perlmutter-Hayman and Secco into agreement with those of Hammes and Morrell. The raw data of Perlmut-ter-Hayman and Secco [73PER/SEC] are unavailable for re-analysis, and in this review, the selected value for 10log K ο (V.98) is the value based on the work of Hammes and Morrell [64HAM/MOR] at 298.15 K:

10log K ο ((V.98), 298.15 K) = (8.73 ± 0.25).

Hence,

r mGο∆ ((V.98), 298.15 K) = – (49.83 ± 1.43) kJ·mol–1.

The apparent formation constants from Perlmutter-Hayman and Secco [73PER/SEC] and Frey and Stuehr [72FRE/STU] for Reaction (V.99) at 288.15 K agree well within the uncertainty bounds. After extrapolation to I = 0, the formation constants from Perlmutter-Hayman and Secco ((5.02 ± 0.25) and (5.02 ± 0.30), based on their measurements in 0.1 and 0.2 M KNO3(aq), respectively) and from Hammes and Morrell (5.14 ± 0.25) also seem to be in good agreement. Unfortunately, the neglect of K+ asso-ciation with pyrophosphate [73PER/SEC] (see Appendix A) and the unavailability of the raw data of Perlmutter-Hayman and Secco [73PER/SEC] preclude a proper re-analysis of the data from that study. The value of 10log K (V.99) of Hammes and Morrell [64HAM/MOR] is selected in the present review for 298.15 K:

10log K ο ((V.99), 298.15 K) = (5.14 ± 0.25).

Hence,

r mGο∆ ((V.99), 298.15 K) = – (29.34 ± 1.43) kJ·mol–1.

There appears to be only one reported value [56YAT/VAS] for the cumulative second association constant (Reaction (V.100)), 2b = (1.54 ± 0.50) × 107. As discussed in Appendix A, this value was calculated from results obtained for solutions of ionic strength that ranged from 0.2 to 1.0 M. This value suggests that significant quantities of

62 7 2Ni(P O ) − should form only in solutions of moderate to high ionic strength where

formation of 22 7NiP O − is less strongly favoured. Furthermore, the effect of association

of Na+ with pyrophosphate [94STE/FOT] on the reported value has not been estab-lished. No value for the second association constant is selected in the present review.

The values for the enthalpies of Reactions (V.98) and (V.99) from the forma-tion constant measurements for 0.1 M KNO3 solutions [73PER/SEC] ( r mH ο∆ ((V.98), (278 – 308) K) = (30.6 ± 7.4) kJ·mol–1, ( r mH ο∆ ((V.99), (278 – 308) K) = (47.9 ± 10.0) kJ·mol–1) are accepted to be the same as the values at I = 0 without correction for either ionic strength or association of K+ with pyrophosphate. The uncertainties are estimates,

r mH ο∆ ((V.98), 298.15 K) = (30.6 ± 10.0) kJ·mol–1

r mH ο∆ ((V.99), 298.15 K) = (47.9 ± 15.0) kJ·mol–1.


The value for Reaction (V.98) is not inconsistent with the rough calorimetric value reported by Yatsimirskii and Vasilev (see discussion of [56YAT/VAS] in Appen-dix A). No value is selected for the enthalpy of Reaction (V.100).

It must be emphasised that the values selected in this review for formation of 2 7NiHP O− and 2

2 7NiP O − should not be used for solutions more than 0.01 M in alkali metal ions unless explicit values are introduced for the pyrophosphate-alkali metal ion association constants (the values of De Stefano et al. [94STE/FOT] may be useful, cf. Appendix A for that reference).

The corresponding Gibbs energy of formation are:

f mGο∆ ( 22 7NiP O − , 298.15 K) = – (2031.1 ± 4.8) kJ·mol–1

f mGο∆ ( 2 7NiHP O− , 298.15 K) = – (2064.3 ± 4.8) kJ·mol–1

Table V-29: Experimental values for the association constants and other thermodynamic quantities for nickel pyrophosphate complexes.

Method t /°C Medium (aq) Reported value

Recalculated value (a, b)


ligand: 42 7P O − (pyrophosphate)

Ni2+ + n 42 7P O − 2 4

2 7Ni(P O ) nn

−

sol 25.0 var 10 1log K = 5.82 [56YAT/VAS] 10 2log b = 7.19 10 ,0log sK (Ni2P2O7) = – 12.77

cal 25.0 var r 1H∆ = (17.61 ± 0.17) kJ·mol–1 [56YAT/VAS2] 10 2log K = 1.60 r 2H∆ = – (9.25 ± 0.25) kJ·mol–1

sp 25.0 var 10 1log K = 3.62 [58VAI/RAM] pot 25.0 0.1 M Me4NCl 10 1log K = 6.98 (8.73 ± 0.25) [64HAM/MOR] pot 15.0 0.1 M KNO3 10 1log K = 6.22 (7.94 ± 0.25) [72FRE/STU] pot 5.0 0.1 M KNO3 10 1log K = 5.63 [73PER/SEC]

15.0 0.1 M KNO3 10 1log K = 5.81 (7.53 ± 0.25) 25.0 0.02 M KNO3 10 1log K = 6.39 25.0 0.05 M KNO3 10 1log K = 6.04 25.0 0.1 M KNO3 10 1log K = 5.94 (7.69 ± 0.25) 25.0 0.2 M KNO3 10 1log K = 5.60 (7.78 ± 0.30) 35.0 0.1 M KNO3 10 1log K = 6.21 0.1 M KNO3 r m ((278 308) K)H∆ − = (30.6 ± 7.4) kJ·mol–1

amp 30.0 0.1 M NH4NO3 10 1log K = 5.35 (7.11 ± 0.30) [78KAR/HUB]



Table V-29 (continued) Method t /°C Medium (aq) Reported value

Recalculated value (a, b)


ligand: 32 7HP O − (hydrogen pyrophosphate)

Ni2+ + 32 7HP O − 2 7HNiP O−

pot 25.0 0.1 M Me4NCl 10 1log K = 3.83 (5.14 ± 0.25) [64HAM/MOR] pot 15.0 0.1 M KNO3 10 1log K = 3.50 (4.79 ± 0.25) [72FRE/STU] pot 5.0 0.1 M KNO3 10 1log K = 3.15 [73PER/SEC]

15.0 0.1 M KNO3 10 1log K = 3.55 (4.84 ± 0.25) 25.0 0.02 M KNO3 10 1log K = 4.01 25.0 0.05 M KNO3 10 1log K = 3.78 25.0 0.1 M KNO3 10 1log K = 3.71 (5.02 ± 0.25) 25.0 0.2 M KNO3 10 1log K = 3.39 (5.02 ± 0.30) 35.0 0.1 M KNO3 10 1log K = 4.07 0.1 M KNO3 r m ((278 308) K)H∆ − = (47.9 ± 10.0) kJ·mol–1

(a) Uncertainties are the estimated uncertainties as discussed in the text. (b) Values from most studies were re-calculated neglecting explicit association of K+ or +

4NH with pyro-phosphate, and cannot be averaged directly with the values obtained in solutions of tetraalkylammonium salts.

V.6.3 Arsenic compounds V.6.3.1 Nickel arsenides

The nickel-arsenic phase diagram [61YUN], [87SIN/NAS] is moderately complex, with the compounds Ni5As2(cr), Ni11As8(cr) (nickel-deficient Ni3As2(cr)), NiAs(cr) and NiAs2(cr) (with a β to α conversion near 870 K).

V.6.3.1.1 NiAs(cr)

There are at least two sources available for the enthalpy of formation of NiAs(cr). Pre-del and Ruge [72PRE/RUG] reported – 36.0 kJ·g-at–1 (– 72.0 kJ·mol–1) for the formation enthalpy based on the heat of dissolution in liquid tin at an unstated temperature. As discussed in Appendix A, a third law analysis results in the value f mH ο∆ (NiAs, cr, 298.15 K) = – (73.60 ± 8.00) kJ·mol–1. Skeaff et al. [85SKE/MAI] used an electro-chemical technique to determine the Gibbs energies of reaction of NiAs(cr) with oxygen to form NiO (bunsenite) and As4O6(g) at temperatures from 711 to 833 K. As discussed in Appendix A, a third law analysis with the heat capacity equation of Muldagalieva et al. [95MUL/CHU], and auxiliary data consistent with those used in the present review, leads to f mH ο∆ (NiAs, cr, 298.15 K) = – (69.73 ± 5.00) kJ·mol–1. The values are consis-tent within the lower uncertainties. The selected value of f mH ο∆ is the weighted average of the two experimental values, and the selected value of mS ο is the recalculated value from Skeaff et al. [85SKE/MAI]:


f mH ο∆ (NiAs, cr, 298.15 K) = – (70.82 ± 4.24) kJ·mol–1

mS ο (NiAs, cr, 298.15 K) = (50.76 ± 6.00) J·K–1·mol–1.

The heat capacity equation of Muldagalieva et al. [95MUL/CHU] is selected in the present review,

800K,m 298.15K[ ]pCο (NiAs, cr) = (36.54 + 41.87 × 10–3 (T/K) + 1.58×105 (T/K)–2) J·K–1·mol–1

as is the value of

,mpCο (NiAs, cr, 298.15 K) = (50.8 ± 4.0) J·K–1·mol–1.

The heat capacity equation was claimed by the authors to be applicable over the temperature range 298.15 to 675 K, and in the present review it is accepted over a slightly expanded range to 800 K to accommodate the recalculation of the equilibrium data of Skeaff et al. [85SKE/MAI]. Nozue et al. [97NOZ/KOB] measured the specific heat of NiAs(cr) for temperatures between 1.7 K and 30 K, but data linking these values to those near 300 K are lacking. The above selections yield:

f mGο∆ (NiAs, cr, 298.15 K) = – (66.6 ± 4.6) kJ·mol–1.

V.6.3.1.2 Ni11As8(cr)

The heat capacity equation of Muldagalieva et al. [93MUL/ISA] is selected in the pre-sent review,

800K,m 298.15K[ ]pCο (Ni11As8, cr) = 380.12 + 367.27×10–3 (T/K) + 0.018×105 (T/K)–2 J·K–1·mol–1

as is the value of:

,mpCο (Ni11As8, cr, 298.15 K) = (490 ± 40) J·K–1·mol–1.

The heat capacity equation was claimed by the authors to be applicable over the temperature range 298.15 to 700 K, and in the present review it is accepted over a slightly expanded range to 800 K to accommodate the recalculation of the equilibrium data of Skeaff et al. [85SKE/MAI].

Skeaff et al. [85SKE/MAI] used an electrochemical technique to determine the Gibbs energies of reaction of Ni11As8(cr) with oxygen to form NiO (bunsenite) and As4O6(g) at temperatures from 805 to 991 K. As discussed in Appendix A, a third law analysis with the heat capacity equation of Muldagalieva et al. [93MUL/ISA] and auxil-iary data consistent with those used in the present review leads to the selected values:

f mH ο∆ (Ni11As8, cr, 298.15 K) = – (743.0 ± 35.0) kJ·mol–1

mS ο (Ni11As8, cr, 298.15 K) = (518 ± 60) J·K–1·mol–1.

The above selections yield:

f mGο∆ ( Ni11As8, cr, 298.15 K) = – (715.8 ± 39.3) kJ·mol–1.


V.6.3.1.3 Ni5As2(cr)

Skeaff et al. [85SKE/MAI] also determined the Gibbs energies of reaction of Ni5As2(cr) with oxygen to form NiO (bunsenite) and As4O6(g) at temperatures from 937 to 1139 K. No experimental heat capacity values appear to have been reported for Ni5As2(cr). Based on the approximation r ,mpC∆ = 0 used by the original authors, and auxiliary data from the present review, f mH ο∆ (Ni5As2, cr, 298.15 K) = – (244.6 ± 20.0) kJ·mol–1,

mS ο (Ni5As2, cr, 298.15 K) = (196.2 ± 30.0) J·K–1·mol–1. Stolyrova [81STO] has reported a much less negative enthalpy of formation, – (204.0 ± 5.8) kJ·mol–1, for the same solid, and Skeaff et al. [85SKE/MAI] noted that her reported value is inconsistent with the observed stability of Ni5As2(cr) solid with respect other phases [87SIN/NAS]. For that reason, the values of Skeaff et al. are selected in the present review:

f mH ο∆ (Ni5As2, cr, 298.15 K) = – (244.6 ± 20.0) kJ·mol–1

mS ο (Ni5As2, cr, 298.15 K) = (196.2 ± 30.0) J·K–1·mol–1.


f mGο∆ (Ni5As2, cr, 298.15 K) = – (237.6 ± 21.9) kJ·mol–1.

V.6.3.1.4 NiAs2(cr)

An experimental value for f mH ο∆ of – (90.06 ± 2.30) kJ·mol–1 (1σ uncertainty) has been reported by Stolyrova [82STO] for this compound. Skeaff et al. [85SKE/MAI] were unable to establish a proper equilibrium with this compound in their apparatus, although some experiments suggested approximately the same value. In the present review, the value is selected, but with an increased uncertainty:

f mH ο∆ (NiAs2, cr, 298.15 K) = – (90.1 ± 8.0) kJ·mol–1.

V.6.3.2 Nickel arsenates

V.6.3.2.1 Solid nickel arsenates

Several anhydrous [58TAY/HEY] and hydrated [66GME] nickel arsenates have been identified, but quantitative chemical thermodynamic data for these solids are sparse. Chukhlantsev [56CHU3] published values for the solubility of Ni3(AsO4)2·8H2O (annabergite) and its apparent solubility product [56CHU4] (K = (3.1 ± 2.4)×10–26, at 20°C for:

Ni3(AsO4)2·8H2O(cr) 3Ni2+ + 2 34AsO − + 8 H2O(l). (V.101)

Nishimura et al. [88NIS/ITO] measured the solubility at 25°C of what was probably the same solid, based on measurements over a wider range of pH (4 to 8), and reported 10log K (V.102) = 9.9 ( 10log K (V.101) = – 26.76) for:

Ni3(AsO4)2·8H2O(cr) + 4H+ 3Ni2+ + 2 2 4H AsO− + 8H2O(l). (V.102)


Langmuir et al. [99LAN/MAH] re-analysed these solubility data using a model that included formation of the complex NiHAsO4(aq), and reported 10log K (V.101) = − 28.38. The experiments of Nishimura et al. [88NIS/ITO] which involved equilibration for periods between one week and four months, indicate greater stability for the solid than was determined by Chukhlantsev [56CHU4]. The value of 10log K ο (V.101) se-lected in the present review:

10log K ο (V.101) = – (28.1 ± 0.5)

is based on the work of Nishimura et al. [88NIS/ITO] (see Appendix A), the auxiliary data for the present review, and the effect of formation of NiHAsO4(aq) as estimated by Langmuir et al. [99LAN/MAH].

From this, r mG∆ (V.101) can be calculated and

f mGο∆ (Ni3(AsO4)2·8H2O, cr, 298.15 K) = – (3491.6 ± 8.8) kJ·mol–1.

Omarova and Sharipov measured the heats of solution of NiSO4·7H2O, Na3AsO4, Na2SO4, H2O and Ni3(AsO4)2·8H2O in aqueous HCl, and reported

f mH ο∆ (Ni3(AsO4)2·8H2O, s, 298.15 K) = – (4179.0 ± 9.2) kJ·mol–1 [80OMA/SHA]. Ar-tamonova and Kasenov [89ART/KAS] reported f mH ο∆ (Ni3(AsO4)2·8H2O, s, 298.15 K) = – (4065.4 ± 21.4) kJ·mol–1, based on the heat of reaction of Na3AsO4(cr) with an aqueous nickel chloride solution to precipitate hydrated Ni3(AsO4)2. In the latter ex-periment, the less negative value found for the enthalpy of formation suggests that the initially precipitated compound might not have been well crystallised, and that the value of Omarova and Sharipov [80OMA/SHA] is to be preferred. In the present review, in the absence of information concerning the auxiliary data used by the authors, the value of Omarova and Sharipov [80OMA/SHA] is selected with an estimated uncertainty of 20 kJ·mol–1:

f mH ο∆ (Ni3(AsO4)2·8H2O, 298.15 K) = – (4179.0 ± 20.0) kJ·mol–1,

which leads to:

mS ο (Ni3(AsO4)2·8H2O, 298.15 K) = (540.8 ± 73.3) J·K–1·mol–1.

V.6.3.2.2 Aqueous nickel arsenato complexes

Langmuir et al. [99LAN/MAH] reported an estimated value of 10log K = 2.90 at 298.15 K for

Ni2+ + 24HAsO − NiHAsO4(aq) (V.103)

based on work described in a thesis by Whiting [92WHI].

The value is similar to the value (3.05 ± 0.09) for the corresponding phosphate complex (cf. Section V.6.2.2.1), and is an acceptable analogue value. In the present re-view, the value from Langmuir et al. [99LAN/MAH] is selected, but because of the


unavailability of experimental values for comparison, an uncertainty of ± 0.3 is as-signed,

10log K ο ((V.103), 298.15 K) = (2.9 ± 0.3).

This selection implies:

f mGο∆ ( NiHAsO4, aq, 298.15 K) = – (776.9 ± 4.4) kJ·mol–1.

V.6.3.3 Nickel arsenites

V.6.3.3.1 Solid nickel arsenites

Several nickel arsenites [66GME] have been identified, but there are very limited chemical thermodynamic data for these solids. Chukhlantsev [57CHU] dissolved sam-ples of nickel orthoarsenite in dilute nitric acid solutions at 20°C over 12 hours to form solutions with aqueous nickel concentrations of 8.7 × 10–3 mol·dm–3 at a “pH” value of 6.75 and 3.1×10–3 at a “pH” value of 7.10 mol·dm–3. The author did not report a solubil-ity product based on their measurements but, as discussed in Appendix A, if the dissolu-tion of the solid is assumed to correspond to the reaction:

Ni3(AsO3)2·xH2O(cr, hydr.) + 6H+ 3Ni2+ + 2HAsO2(aq) + (2 + x)H2O(l) (V.104)

the selected equilibrium constant can be calculated:

10log K ο ((V.104), 298.15 K) = (28.7 ± 0.7).

From this, r mG∆ (V.104) can be calculated and, if the water of hydration is not explicitly included in the chemical formula,

f mGο∆ (Ni3(AsO3)2, cr, hydr, 298.15 K) = – (1253.6 ± 9.3) kJ·mol–1.

V.7 Group 14 compounds and complexes V.7.1 Carbon compounds and complexes V.7.1.1 Nickel carbonates

V.7.1.1.1 Solid nickel carbonates

V.7.1.1.1.1 NiCO3(cr)

V.7.1.1.1.1.1 Crystallography and mineralogy of nickel carbonate

V.7.1.1.1.1.1.1 NiCO3(cr), gaspéite The crystal structure of NiCO3(cr) was definitively determined by Isaacs [63ISA]. It is of the calcite type, hexagonal, space group: R 3 c, Z = 6, with unit cell dimensions a0 = 4.609 Å, c0 = 14.737 Å (according to JCPDS-ICDD card No. 12-771).

V.7 Group 14 compounds and complexes

215

The natural occurrence of a new mineral (Ni0.49Mg0.43Fe0.08)CO3 as a vein en-closed in siliceous dolomite in the Gaspé Peninsula, Quebec, was reported by Kohls and Rodda [66KOH/ROD]. The name gaspéite was selected for the nickel carbonate end member (approved by IMA 1966). The vein consists of essentially pure magnesian gaspéite with minor amounts of annabergite (Ni3(AsO4)2·8H2O), magnesite (MgCO3) and dolomite (CaMg(CO3)2). Enclosed in the magnesian gaspéite are a small amount of serpentine ((Mg, Fe)3Si2O5(OH)4) and a (Cr, Al) spinel. Well-formed crystals of mil-lerite (NiS), niccolite (NiAs), annabergite, gersdorffite (NiAsS) and magnesite are found outside the vein in the buff siliceous dolomite.

The X-ray powder diffraction patterns of magnesian gaspéite differ slightly from synthetic NiCO3(cr), see JCPDS-ICDD card No. 23-437. Gaspéite from the Gaspé Peninsula is light green, its luster is vitreous to dull and it has a yellow-green streak. The Mohs hardness is 4.5 to 5, the specific gravity is ρ(obs.) = (3.71 ± 0.01) and ρ(calc.) = 3.748 g·cm–3.

V.7.1.1.1.1.1.2 Ni2(CO3)(OH)2, nullaginite Nullaginite occurs as a secondary mineral in serpentinised peridotites in the Nullagine district (Western Australia). Mineral and name were approved by the IMA Commission in 1978. It is bright green with a luster varying from dull to silky. Composition and X-ray powder diffraction pattern of nullaginite were investigated by Nickel and Berry [81NIC/BER]. Its crystal system is monoclinic, space group: P21/m, cell dimension: a0 = 9.236 Å, b0 = 12.001 Å, c0 = 3.091 Å, Z = 4, β = 90.48°, Vcell = 342.60 Å3, ρ(obs.) = 3.56 and ρ(calc.) = 4.099 g·cm–3 (according to JCPDS-ICDD card No 35-501). The large discrepancy between ρ(calc.) and ρ(obs.) is attributed to the presence of admixed water in the specimen. The ideal composition of Nullaginite is Ni2(CO3)(OH)2, but it occurs also together with Cr, Mg, Fe, and Cu and belongs to the rosasite group. This group is composed of minerals with either a monoclinic or triclinic symmetry and a general formula of M2(CO3)(OH)2. The M ion can be Co2+, Cu2+, Zn2+, Mg2+ and/or Ni2+. Minerals of this group composed of colour stimulating metals such as cobalt, cop-per and nickel are extremely colourful, usually green to blue. Most of these minerals range from rare to extremely rare. No synthesis of nullaginite has been reported and no thermodynamic data have been determined so far.

V.7.1.1.1.1.1.3 Ni2(CO3)(OH)2·H2O, otwayite The crystal structure of otwayite was investigated by Nickel et al. [77NIC/ROB]. Its fibre-rotation and X-ray powder diffraction patterns can be indexed on an orthorhombic unit cell with a = 10.18 Å, b = 27.4 Å, c = 3.22 Å, Vcell = 898.16 Å3, Z = 8, ρ(obs.) = 3.41 and ρ(calc.) = 3.346 g·cm–3 (according to JCPDS-ICDD card No. 29-868).

Otwayite occurs in Heazlewood (Tasmania), in the Nullagine region, but also together with Widgiemoolthalite in Widgiemooltha (both in Western Australia). Min-eral and formula were approved by the IMA in 1977. Otwayite was named after Charles Albert Otway (1922 –), miner and prospector of Cosnells in Australia [77NIC/ROB].


216

No synthesis of otwayite has been reported and no thermodynamic data have been de-termined so far.

V.7.1.1.1.1.1.4 Ni3(CO3)(OH)4·4H2O, zaratite Isaacs studied the structure of zaratite [63ISA]. The cell dimensions are: a = 6.16 Å, Z = 1; Vcell = 233.74 Å3, ρ(calc.) = 2.31 g·cm–3, crystal system: isometric, space group: un-known (in part amorphous).

Zaratite occurs in Cape Ortegal, Galicia, (Spain); in Broken Hill, New South Wales, (Australia); and in Pennsylvania, (USA). The mineral was named after Antonio Gil y Zárate (1793 – 1861). Data for physical and chemical properties vary significantly [63ISA]. The average chemical composition is close to Ni3(CO3)(OH)4·4H2O, however, analytical results scatter in the following range: NiO 56.9 – 61.2%, CO2 13.5 – 15.7%, H2O 23.2 – 27.1%. Zaratite is rare, forming emerald green secondary crusts on other nickel minerals. Its hardness is 3.5. From X-ray investigation of artificially synthesised as well as natural specimens it was concluded that zaratite is not a single mineral and should not be called one. Consequently no thermodynamic data can be ascribed to this poorly defined mixture of nickel hydroxide carbonates.

V.7.1.1.1.1.1.5 Ni5(CO3)4(OH)2·4–5H2O, widgiemoolthalite The crystal structure of widgiemoolthalite was investigated by Nickel et al. [93NIC/ROB]. It has an ideal composition of Ni5(CO3)4(OH)2·4–5H2O. The monoclinic unit cell is similar to that of hydromagnesite, space group: P21/c. The cell dimensions are: a = 10.06 Å, b = 8.75 Å, c = 8.32 Å, Z = 2; β = 114.3°, Vcell = 667.48 Å3, ρ(obs.) = 3.13 and ρ(calc.) = 2.97 g·cm–3 (according to JCPDS-ICDD card No. 46-1398).

Widgiemoolthalite is found together with a number of other secondary nickel minerals in the weathered zone in the Widgie 132N mine in Widgiemooltha [93NIC/ROB]. It is bluish green and occurs mainly as a fibre. The mineral was discov-ered in 1992, named after its location and approved by IMA 1993. In reality the compo-sition of widgiemoolthalite is close to (Ni, Mg)5(CO3)4.15(OH)1.7·5.12 H2O. No thermodynamic data of this mineral are known.

V.7.1.1.1.1.2 Heat capacity and entropy of NiCO3(cr)

A high temperature heat capacity function claimed to be valid at 298 ≤ (T/K) ≤ 700 for NiCO3(cr) is given in thermodynamic data compilations:

,mpCο (NiCO3, cr) = (a + b (T/K) + e (T/K)–2) J·K–1·mol–1 700K

,m 298K[ ]pCο (NiCO3, cr) = (88.701 + 38.91×10–3 (T/K) – 12.34×105(T/K)–2) J·K–1·mol–1

(V.105)

Originally the following coefficients were given: a = 92.05, b = 38.91×10–3 and e = – 12.34×105 [77BAR/KNA]. In Knacke et al. [91KNA/KUB] and Kubaschewski et al. [93KUB/ALC] the coefficients b and e have been retained, but a


217

has been slightly modified (a = 88.701). The only source of original ,mpCο measurements up to 300 K found so far is [64KOS/KAL]. The latter authors estimated the maximum error of the standard entropy to be ± 0.4 J·K–1·mol–1, however, they presented their re-sults only graphically:

mS ο (NiCO3, cr, 298.15 K) = (85.4 ± 2.0) J·K–1·mol–1.

Robie and Hemingway assigned an uncertainty of ± 2.0 J·K–1·mol–1 to this quantity, which appears reasonable because, as mentioned above, the ,mpCο function can only be evaluated graphically [95ROB/HEM]. Both the numerical value for this quan-tity and its error limits have been selected for this compilation. The so-called ‘Latimer’ contributions would result in mS ο (NiCO3, cr, 298.15 K) = (81.7 ± 5.0) J·K–1·mol–1 whereby the actually selected value falls within these error limits [93KUB/ALC].

The following value can be read off directly from Fig. 1 of [64KOS/KAL]:

,mpCο (NiCO3, cr, 298.15 K) = (90.3 ± 4.1) J·K–1·mol–1.

The uncertainty is assigned in such a way that the ,mpCο values calculated ac-cording to Equation (V.105) as well as estimated according to the cationic and anionic contributions, given by Kubaschewski et al. [93KUB/ALC], fall within the error limits. Although neither the experimental nor the semi-empirical basis for Equation (V.105) could unambiguously be retraced, it has been retained in the current review.

V.7.1.1.1.1.3 Thermodynamic analysis of solubility data on gaspéite

Solubility measurements are especially valuable for the determination of an internally consistent set of thermodynamic data for nickel carbonates. Other methods turned out to be less convenient. Thus the calorimetric determination of the dissolution enthalpy ac-cording to Reaction (V.106) is bound to be inaccurate, because the result depends to a high extent on the state of carbon dioxide. It is well known that keeping CO2 quantita-tively either in solution or in the gas phase is quite difficult.

NiCO3(cr) + 2H+ Ni2+ + CO2(g) + H2O(l) (V.106)

V.7.1.1.1.1.3.1 Analysis of hydrothermal decomposition Another method to determine thermodynamic data of nickel carbonate consists in the investigation of the decomposition equilibrium (V.107). Tareen et al. [91TAR/FAZ] determined the standard molar quantities of formation of gaspéite, NiCO3(cr) from hydrothermal T, p decomposition data of this reaction, but probably overestimated the precision achieved: f mGο∆ = – (628.36 ± 1.24) kJ·mol–1 (– (632.3 ± 6.0) kJ·mol–1),

f mH ο∆ = – (703.38 ± 1.24) kJ·mol–1 (– (709.2 ± 6.0) kJ·mol–1).

NiCO3(s) NiO(s) + CO2(g) (V.107)

The numerical values in brackets have been found when the experimental re-sults of [91TAR/FAZ] were re-evaluated according to the present state of the art, see Appendix A.


218

V.7.1.1.1.1.3.2 Analysis of solubility studies It was shown recently that solubility data obtained from NiCO3·5.5H2O(cr), hellyerite, were erroneously ascribed to NiCO3(cr), gaspéite [2001GAM/PRE]. Thus, in fact it was Reiterer [80REI] who determined the solubility constant of gaspéite defined by Reac-tion (V.106) for the first time. He employed the pH variation method to investigate the temperature dependence of 10 p, ,0

*log sK (V.106) at constant ionic strength I = 1.0 mol·kg−1 (Na)ClO4. Kloger [82KLO] carried out additional experiments at I = 0.2 mol·kg–1 (Na)ClO4 and 363.15 K.

The solubility data of synthetic NiCO3 were thermodynamically analysed to obtain the respective standard molar quantities of formation f mGο∆ and f mH ο∆ . An analogous set of equations was used as described in Section V.3.2.2.3 (Equations (V.41) – (V.44)). Solubility constants were extrapolated to infinite dilution using the Specific Ion-interaction Theory [97GRE/PLY2]. All calculations were based on auxiliary quanti-ties which were either recommended by CODATA (see NEA TDB auxiliary data set) or recently critically evaluated (see Section V.2.1).

The following set of reliable thermodynamic quantities has been derived [82GAM/REI], [98GAM/KON], [2002WAL/PRE]:

10 p, ,0*log sK ((V.106), 298.15 K) = (7.16 ± 0.18)

f mGο∆ (NiCO3, cr, 298.15 K) = – (636.4 ± 1.3) kJ·mol–1

f mH ο∆ (NiCO3, cr, 298.15 K) = – (713.3 ± 1.4) kJ·mol–1.

These values obtained from solubility measurements of pure synthetic NiCO3(cr) fall well within the error limits of those re-evaluated from decomposition studies, but are clearly more precise.

V.7.1.1.1.2 NiCO3·5.5H2O(cr)

V.7.1.1.1.2.1 Crystallography and mineralogy of nickel carbonate hydrate

The preparation and X-ray powder diffraction pattern of nickel carbonate hexahydrate, NiCO3·6H2O(cr), were first described by Rossetti-François [52ROS]. The crystal system is monoclinic-prismatic, space group: C2/c, cell dimensions: a0 = 10.770 Å, b0 = 7.299 Å, c0 = 18.681 Å, Z = 8; β = 94.00°, Vcell = 1464.94 Å3, ρ(obs.) = 1.97 g·cm–3, ρ(calc.) = 1.975 g·cm–3 (according to JCPDS-ICDD cards 12-276 and 24-523, where the empirical formulae are NiCO3·6H2O(cr) in the former and NiCO3·5.5H2O(cr) in the latter).

The natural occurrence of hellyerite, NiCO3·5.5H2O(cr), in the Lord Brassey nickel mine near Heazlewood (Tasmania), was first reported by Williams et al. [59WIL/THR]. The Mohs hardness is 2.5, its luster is vitreous (glassy) and its streak is white. The X-ray powder diffraction pattern was found to be similar to that of synthetic nickel carbonate hexahydrate, prepared by the method of Rossetti-François.


219

The single crystal structure determination by Threadgold [63THR] of the natu-ral product may be not perfect due to extensive twinning of the material. Hellyerite forms predominantly twin crystals to enhance the low crystal symmetry. Therefore the crystal structure needs to be refined to resolve its subtleties.

Hellyerite has a prominent layered structure parallel to (001) and is made up of two distinct layers, a NiCO3(H2O)4 layer and an almost planar sheet of H2O molecules [63THR]. This layer structure with its unknown arrangement of H2O molecules is pre-sumably responsible for chemical and thermogravimetric analyses of hellyerite resulting in a stoichiometric formula of NiCO3·5.5H2O(cr) i.e., NiCO3(H2O)4·1.5H2O(cr) instead of NiCO3·6H2O(cr) [2001GAM/PRE]. Crystallographic investigations to solve the hel-lyerite structure definitively are in progress.

V.7.1.1.1.2.2 Heat capacity and entropy of NiCO3·5.5H2O(cr)

Apparently no experimental low-temperature heat capacity data of NiCO3(H2O)4·1.5H2O(cr) have been reported so far. A crude estimate was obtained by analogy to magnesium carbonate and its hydrates. When lansfordite (MgCO3·5H2O(cr)), nesquehonite (MgCO3·3H2O(cr)) and magnesite (MgCO3(cr)) are compared the heat capacity can roughly be approximated by Equation (V.108) [99KON/KON],

,mpCο (MgCO3·nH2O, cr, 298.15 K) = (76.09 + n × 58.0) J·K–1·mol–1. (V.108)

With Equation (V.105) the following value can tentatively be used for the heat capacity of NiCO3·5.5H2O(cr):

,mpCο (NiCO3·5.5H2O, cr, 298.15 K) = (405.4 ± 50.0) J·K–1·mol–1.

The entropy of NiCO3·5.5H2O(cr) was determined by solubility measurements,

mS ο (NiCO3·5.5H2O, cr, 298.15 K) = (311.1 ± 10.0) J·K–1·mol–1,

see Section V.7.1.1.1.2.3, and is selected in this review.

The standard entropy obtained from the thermodynamic analysis of solubility data agrees reasonably well with mS ο (NiCO3·5.5H2O, cr, 298.15 K) = 307.9 J·K−1·mol–1 estimated according to Latimer [52LAT].

V.7.1.1.1.2.3 Thermodynamic analysis of solubility data on hellyerite

V.7.1.1.1.2.3.1 Solubility measurements of hellyerite, NiCO3·5.5H2O(cr) As pointed out in Section V.7.1.1.1.1.3 solubility measurements of metal carbonates provide an important method to determine the thermodynamic quantities. For hellyerite, however, these data can be unambiguously assigned only when the stoichiometric com-position is known. As formula and crystal structure are strongly interdependent, they should be determined simultaneously. So far neither has been established definitively and the thermodynamic quantities derived from solubilities must be used with caution.


220

With hellyerite prepared by a new method, solubility measurements were car-ried out at different temperatures and at I = 1.0 m (Na)ClO4 [2001GAM/PRE]. In addi-tion, the solubility was also determined at 25°C and different ionic strengths [2002WAL/PRE]. In either case the pH variation method was used to study the dissolu-tion reaction according to Reaction (V.109):

NiCO3·5.5H2O(cr) + 2H+ Ni2+ + CO2(g) + 6.5H2O(l) (V.109)

The thermodynamic calculations of this compilation are tentatively based on the formula NiCO3·5.5H2O(cr) which agrees with Threadgold’s crystal structure analy-sis [63THR] and the composition of the synthetic product [2001GAM/PRE]. From the weak temperature dependence of the solubility constant, 10 p, ,0

*log sK , of hellyerite (see Figure V-43) f mH ο∆ (NiCO3·5.5H2O, cr) and mS ο (NiCO3(H2O)4·1.5H2O, cr) were calcu-lated using a non-linear least squares optimising routine.

Figure V-43: Temperature dependence of hellyerite solubility. I = 1.0 mol·kg–1 NaClO4, ● 278.15, × 288.15, ■ 298.15, ▲ 308.15, 313.15 K, solid line: specific ion-interaction formalism [97GRE/PLY2], dashed line: Pitzer model [91PIT].

270 280 290 300 310 32010.4

10.6

10.8

11.0

11.2

11.4

11.6

log 10

* K p,s,0

T / K

NiCO3·5.5H2O(cr) +2 H+ Ni2+ + CO2(g) + 6.5 H2O


221

Thermodynamic auxiliary data for H2O(l) and CO2(g) were taken from Table IV-1 (T = 298.15 K) and from CODATA [89COX/WAG] (T ≠ 298.15 K). Data for Ni2+ were used as selected in this review (see Section V.2.1). Nickel hydroxo and carbonato complexes are negligible [76BAE/MES], [2003BAE/BRA] in the pH ranges where the experiments of [2001GAM/PRE] and [2002WAL/PRE] were carried out.

Figure V-44 shows the solubility constant of NiCO3·5.5H2O(cr) plotted as a function of ionic strength at 298.15 K. Both the Pitzer equations [91PAP/PIT], [91PIT], [99KON/KON] and the SIT, [97GRE/PLY2] were applied to model the ionic strength dependence of 10 p, ,0

*log sK . Whereas the Pitzer equations seem to predict the ionic strength dependence of the solubility constant more accurately than the SIT approach, the experimental data as a function of temperature coincide best with the latter, see Figure V-43.

Figure V-44: Ionic strength dependence of hellyerite solubility. t = 25°C, Ionic medium: I = 0.5 to 3.0 mol·kg–1 NaClO4, [2002WAL/PRE], solid line: specific ion-interaction formalism [97GRE/PLY2], dashed line: Pitzer model [91PIT].

A plot of 2+ +2

6.5 210 H ONi H

log ( )a a a−⋅ ⋅ versus log10 2COp shown in Figure V-45 de-scribes the thermodynamic analysis of the data presented in [11AGE/VAL], [30MUL/LUB] and [2002WAL/PRE] in comparison with the optimised line (slope = − 1.00). The activities of Ni2+ and H+ were calculated by using the SIT model. This re-sult reveals a ratio of n[Ni]/n[CO2] close to unity in the solid phase, as is expected for NiCO3·5.5H2O(cr). A re-evaluation of the experimental data of Müller and Luber

0.0 0.2 0.4 0.6 0.8 1.0 1.2 1.4 1.6 1.8 2.010.5

10.6

10.7

10.8

10.9

11.0

11.1

11.2

11.3

11.4

11.5

lo

g 10* K p,

s, 0

(I / mol·kg−1) 1/2



222

[30MUL/LUB] as well as Ageno and Valla [11AGE/VAL] resulted in 10 p, ,0*log sK ο =

(10.56 ± 0.10), which compares favourably with

10 p, ,0*log sK ο ((V.109), 298.15 K) = (10.63 ± 0.10)

calculated from experimental data of [2001GAM/PRE], [2002WAL/PRE]. Thus it can be concluded that Ageno and Valla [11AGE/VAL] studied NiCO3·5.5H2O(cr) and not NiCO3 ( 10 p, ,0

*log sK ο = (7.13 ± 0.18)), see Appendix A. Obviously the solubility con-stants of hellyerite (Equation (V.109)) and gaspéite (Equation (V.106)) differ by 3.5 orders of magnitude (see Section V.7.1.1.1.1.3). This presents a further example that solubility data measured at different temperatures and ionic strengths provide a useful basis for the calculation of thermodynamic quantities when and only when the solid samples investigated are analytically and structurally unambiguously characterised. Moreover, a careful thermodynamic analysis showed that the experimental information of rather old solubility studies either falls in line with the optimised set of thermody-namic properties obtained [11AGE/VAL], [30MUL/LUB] or reveals that presumably a basic nickel carbonate was investigated [38SMU].

Figure V-45: Thermodynamic analysis of hellyerite solubility. Data of [11AGE/VAL]: ■, datum of [30MUL/LUB]: ▲, data of [2002WAL/PRE]: ◊ I = 0.5, ○ I = 1.0, + I = 2.0, × I = 3.0 mol·kg–1 NaClO4. Solid straight line: calculated with 10 p, ,0

*log sK ο = 10.63 se-lected in this review.

-1.5 -1.0 -0.5 0.0 0.5 1.0 1.5 2.08.0

8.5

9.0

9.5

10.0

10.5

11.0

11.5

12.0

12.5

log 10

{a

Ni2+ ·

a H2O

6.5 ·

a H+−2

}

log10 {p(CO2)/bar}


T = 298.15 K


223

Thus the selected standard Gibbs energy and standard enthalpy of formation have been derived from solubility measurements at various temperatures and ionic strengths.

f mGο∆ (NiCO3·5.5H2O, cr, 298.15 K) = – (1920.9 ± 1.0) kJ·mol–1

f mH ο∆ (NiCO3·5.5H2O, cr, 298.15 K) = – (2313.0 ± 3.1) kJ·mol–1.

A convenient overview of thermodynamic data on nickel carbonates and hy-droxide is provided by a predominance diagram as depicted in Figure V-46.

Figure V-46: Three-dimensional predominance diagram for the system Ni2+–CO2–H2O. 2+Ni

a = 10–4; black lines, solid phases: NiCO3, β-Ni(OH)2; gray lines, solid phases: NiCO3·5.5H2O(cr), β-Ni(OH)2; aqueous phase: Ni2+.

V.7.1.1.2 Aqueous nickel carbonato complexes

Recently the formation of carbonato complexes in the system Ni2+ – H2O – CO2 has been critically discussed and re-evaluated in a seminal review of Hummel and Curti [2003HUM/CUR]. So far only one paper has been published describing an attempt to determine the equilibrium constant, β1, of Reaction (V.110) experimentally [87EMA/FAR].

Ni2+ + 3HCO− 3NiHCO+ (V.110)

290310

330350

370

-4

-3

-2

-1

0

1

4.0 5.0 6.0 7.0 8.0 9.0 10.0

NiCO3·5.5H2

O

NiCO3β - Ni(OH)2

β - Ni(OH)2

Ni2+log 10

[pCO

2 / p0 ]

pHT / K


224

Emara et al. clearly misinterpreted their data and did not provide enough in-formation to allow recalculation. Consequently, the stability constant of 3NiHCO+ re-ported in [87EMA/FAR] cannot be included in this review.

Values of equilibrium constants for Reaction (V.110) estimated by various procedures differ considerably 0.96 ≤ 10 1log οb ≤ 3.08 (T = 298.15 K) [76ZHO/BEZ], [77MAT/SPO], [79MAT/SPO], [84FOU/CRI]. The stability constant of the carbonato complex according to Reaction (V.111) has also been estimated leading to an even lar-ger discrepancy: 2.56 ≤ 10 1log K ο ≤ 6.87 (T = 298.15 K) [76ZHO/BEZ], [77MAT/SPO], [79MAT/SPO], [81TUR/WHI], [84FOU/CRI].

Ni2+ + 23CO − NiCO3(aq) K1 (V.111)

One estimate exists for the equilibrium constant ( 10 2log K ο = 3.24) of Reaction (V.112) [77MAT/SPO].

NiCO3(aq) + 23CO − 2

3 2Ni(CO ) − K2 (V.112)

As the basis of the individual estimation procedures is rather dubious, varia-tions of up to more than four log-units in these stability constants are to be expected [2003HUM/CUR]. Again neither of these values appeared suitable to be included in this review.

Hummel and Curti proposed estimating K1 for Equation (V.111) using either the good correlation between the equilibrium constants of Ni(II) and Co(II) complexes and the poor data available for K1 of CoCO3(aq) or the rather poor correlation between Ni(II) and Zn(II) complexes and the excellent data for K1 for ZnCO3(aq) [2003HUM/CUR]. Both methods result in similar lower and upper bounds 4 < 10 1log K ο ((V.111), 298.15 K) < 5.5.

A comparison of the stabilities of transition metal hydrogen carbonato as well as carbonato complexes led to 1 < 10 1log οb ((V.110), 298.15 K) < 2 and 10 2log K ο < ( 10 1log K ο – 2) [2003HUM/CUR]. Even the careful and competent guesswork of Hummel and Curti resulted in rough estimates only. Fortunately, in a recent paper, Baeyens et al. [2003BAE/BRA] investigated Ni-carbonato and -oxalato complexes by an ion exchange method. Ni-carbonato complexes were investigated at constant ionic strength I = 0.5 mol·dm–3 NaClO4 / NaHCO3 and (22 ± 1)°C. The experimentally ob-tained complexation constant ( 10 1log K (295.15 K) = (2.9 ± 0.3)) was extrapolated to I = 0 with the SIT approach to give [97GRE/PLY2]:

10 1log K ο ((V.111), 298.15 K) = (4.2 ± 0.4).

This result was finally selected for the present review. The somewhat higher uncertainty was assigned, because Baeyens et al. carried out their measurements at 22°C instead of 25°C, and used a relatively simple approximation to extrapolate

10 1log K ο to I = 0. This selection yields:


225

f mGο∆ (NiCO3, aq, 298.15 K) = – (597.6 ± 2.4) kJ·mol–1.

Only upper bounds can be given for the stabilities of 3NiHCO+ and 2

3 2Ni(CO ) − : 10 1log οb ((V.110), 298.15 K) < 1.4 and 10 2log K ο ((V.112), 298.15 K) < 2 [2003BAE/BRA]. Both upper bounds compare well with the lower limits of the range predicted by Hummel and Curti [2003HUM/CUR], but do not qualify for being in-cluded in the list of selected values in this review. Figure V-47 shows the nickel speci-ation when NiCO3(cr) dissolves in dilute sodium carbonate solutions under atmospheric conditions. The formation of NiCO3(aq) becomes evident only at pH > 8.5. At pH > 9 NiCO3(aq) predominates.

Figure V-47: Calculated solubility of NiCO3(cr) in dilute sodium carbonate solutions under atmospheric conditions.

7.4 7.6 7.8 8.0 8.2 8.4 8.6 8.8 9.0 9.2 9.4

-8.0

-7.5

-7.0

-6.5

-6.0

-5.5

-5.0

-4.5

-4.0

-3.5

log 10

([Ni(I

I)] /

mol

·kg−1

)

pH

25°C, I = 1.0 mol·kg−1 NaClO4log10(p (CO2) / bar) = − 3.5

[Ni2+] [NiOH+] [NiCO3] [Ni(II)]tot


226

V.7.1.2 Nickel cyanides

V.7.1.2.1 Aqueous Ni(II) cyano complexes

The first quantitative data on the aqueous Ni(II) cyanide complexes were generated at the end of XIXth century, when Varet reported the heat of formation of the 2

4Ni(CN) − complex [1896VAR]. Varet's r mH∆ value is in good agreement with more recent de-terminations, but there has been an increase of nearly twenty orders of magnitude in the reported 10 4log b value of 2

4Ni(CN) − , since it was first measured by Masaki in 1932 [32MAS]. Besides the numerical value of 10 4log b , conflicting results also have been obtained concerning the composition and protonation state of the species formed. There-fore, the discussion is divided into two parts.

V.7.1.2.1.1 Complexes in neutral and alkaline solution

Most authors have agreed that the formation of NiCN+, Ni(CN)2(aq) and 3Ni(CN)− can not be detected in the equilibrated solutions. Although, Shibata et al. claimed to detect Ni(CN)2(aq) and possibly Ni(CN)3(H2O)– by electrophoresis of γ-ray irradiated K2Ni(CN)4 solutions [66SHI/FUJ], the assignment of the measured peaks was rather speculative. The absence of the above-mentioned species indicates that the stepwise formation constants K1, K2 and K3 are much less than 4

4b , and thus K4 is probably greater than 1014 [68KOL/MAR]. This very large value is likely related to the transfor-mation of an octahedral to a square planar complex in the fourth stepwise complex for-mation process.

The first reported value for 10 4log b [32MAS] was not accurate, since the elec-trode reactions involving Ni(II) are not reversible [50HUM/KOL].

Ni2+ + 4CN– 24Ni(CN) − (V.113)

Reliable values were reported from spectrophotometric [59FRE/SCH], kinetic [61MAR/BYD], [68KOL/MAR], [76PER], and pH-metric [63CHR/IZA], [71IZA/JOH], [74PER] measurements. Freund and Schneider used an incorrect HCNpK value, therefore the experimental data reported in [59FRE/SCH] were re-evaluated (see discussion in Appendix A). The experimentally determined 10 4log b values are col-lected in Table V-30. Although only a limited number of data are available in NaClO4 media, the precision of the constants is assumed to be good enough to perform an SIT analysis (Figure V-48). The weighted linear regression using five data points yielded the selected value of:

10 4log οb ((V.113), 298.15 K) = (30.20 ± 0.12).

From the slope in Figure V-48, ∆ε((V.113), NaClO4) = – (0.465 ± 0.045) kg·mol–1 can be calculated. Using the selected value for ε(Ni2+, 4ClO− ), and ε(Na+, CN−) = (0.07 ± 0.03) kg·mol–1 reported in [92BAN/BLI], ε(Na+, 2

4Ni(CN) − ) = (0.185 ± 0.081) kg·mol–1 can be derived. In KClO4 media, too narrow an ionic strength range is covered to perform the SIT analysis [59FRE/SCH], but after correction for the Debye-Hückel


227

contributions, the formation constants determined in these low ionic strength solutions are almost constant, and are nearly identical to the selected 10 4log οb ((V.113), 298.15 K) value (see insert in Figure V-48).

At ratios [CN–]/[Ni2+] > 4, the pale yellow solution of diamagnetic 24Ni(CN) −

becomes successively orange, red and deep red upon addition of cyanide ion. The rea-son for this change is evidently the formation of higher complexes, but the composition and stability of these have been disputed for a long time. Several early reports [23JOB/SAM], [43SAM], [59KIS/CUP], [59BLA/GOL], [59BLA/GOL2] interpreted the colour change as evidence for the formation of hexacyanonickelate(II) ion,

46Ni(CN) − . Some of them reported surprisingly high equilibrium constant values for the

reaction: 24Ni(CN) − + 2CN– 4

6Ni(CN) −

[59KIS/CUP], [59BLA/GOL2]. Later, the formation of a weak pentacyano complex and the non-existence of the earlier-suggested 4

6Ni(CN) − ion were definitely proven [60MCC/JON], [62BEC/BJE], [64GEE/HUM], [65COL/PET], [71PIE/HUG]. The re-ported equilibrium constants for the reaction:

24Ni(CN) − + CN– 3

5Ni(CN) − (V.114)

are collected in Table V-30. Since the pentacyano complex is rather unstable, high cya-nide concentrations were used to achieve its formation, which resulted in considerable replacement of the original background electrolyte by NaCN. Due to this medium ef-fect, considerably higher uncertainties have been assigned to the equilibrium constants than originally reported. The SIT analysis of the constants determined in NaClO4 media (Figure V-49), resulted in 10 5log K ο ((V.114), 298.15 K) = – (1.70 ± 0.36), and ∆ε((V.114), NaClO4) = (0.00 ± 0.11) kg·mol–1. From the latter value, ε(Na+, 3

5Ni(CN) − ) = (0.25 ± 0.14) kg·mol–1 can be derived. Using equilibrium constant from above, the overall formation constant of the 3

5Ni(CN) − species is

Ni2+ + 5 CN– 35Ni(CN) − (V.115)

10 5log οb ((V.115), 298.15 K) = (28.5 ± 0.5).


228

Table V-30: Experimental equilibrium data for the Ni(II) cyanide system.

Method Medium Im t (°C) 10log xb reported

10log xb

retained(a) Reference

Ni2+ + 4CN– 24Ni(CN) −

pot NaCN 0.78 M 25 (11.83 ± 0.17) [32MAS] pol KCl 0.10 ? > 24 [50HUM/KOL] sp K(ClO4) 0.1

0.048 0.014 0.0028 0 corr

25

(31.0 ± 0.08)

(29.37 ± 0.20)(b) (29.45 ± 0.20)(b) (29.79 ± 0.20)(b) (30.12 ± 0.20)(b)

[59FRE/SCH]

kin NaCl 1.03 25 31.5 [61MAR/BYD] gl (self) 0 corr 25 (30.1 ± 0.2) (30.1 ± 0.2) [63CHR/IZA] kin NaClO4 0.10 25 (30.5 ± 0.3) (30.48 ± 0.6) [68KOL/MAR] gl (self) 0 corr 10

25 40

(32.20 ± 0.20) (30.22 ± 0.05) (27.43 ± 0.09)

(32.20 ± 0.30) (30.22 ± 0.15) (27.43 ± 0.20)

[71IZA/JOH]

gl NaClO4 3.50 25 (31.12 ± 0.08) (30.85 ± 0.15) [74PER] kin NaClO4 3.50 25 (31.08 ± 0.05) (30.81 ± 0.15) [76PER]

24Ni(CN) − + CN– 3

5Ni(CN) −

sp Na(ClO4) 1.43 15.4 25.2 33.6

– (0.66 ± 0.02) – (0.72 ± 0.02) – (0.77 ± 0.02)

– (0.69 ± 0.30) – (0.75 ± 0.30) – (0.80 ± 0.30)

[60MCC/JON]

ir Na(ClO4) 4.95 25 – (0.55 ± 0.02) – (0.65 ± 0.30) [63PEN/BAI] sp Na(ClO4) 2.84 23

25 – (0.70 ± 0.02)

– (0.77 ± 0.30) [64GEE/HUM]

sp KF-KCN 4.00 25 (0.03 ± 0.01)(c) [65COL/PET] sp Na(ClO4) 2.21 20

25 – (0.77 ± 0.01)

– (0.84 ± 0.30)(d) [71PIE/HUG]

(a) Accepted values corrected to molal scale. (b) Re-evaluated value, see Appendix A. (c) In molal units. (d) Corrected to 25°C, using the selected enthalpy values.


229

Figure V-48: Extrapolation to I = 0 of experimental formation constants of 24Ni(CN) −

(V.113) determined in NaClO4 medium (insert shows the corresponding plot of data reported for KClO4 medium).

Figure V-49: Extrapolation to I = 0 of experimental equilibrium constants for Reaction (V.114) determined in NaClO4 medium.

0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5

30

31

32

33

34

log 10

β4 +

4D

I / mol·kg−1

0.00 0.02 0.04 0.06 0.08 0.10 0.12

29.5

30.0

30.5

Ni2+ + 4CN– 24Ni(CN) −

0 1 2 3 4 5-3.0

-2.5

-2.0

-1.5

-1.0

-0.5

0.0

log 1

0 K5 −

4D

I / mol·kg−1

24Ni(CN) − + CN– 3

5Ni(CN) −


230

The selected thermodynamic formation constants correspond to:

r mGο∆ ((V.113), 298.15 K) = – (172.4 ± 0.7) kJ·mol–1,

r mGο∆ ((V.115), 298.15 K) = – (162.7 ± 2.9) kJ·mol–1

and thus the standard molar Gibbs energies of formation are:

f mGο∆ ( 24Ni(CN) − , 298.15 K) = (449.6 ± 10.1) kJ·mol–1,

f mGο∆ ( 35Ni(CN) − , 298.15 K) = (626.6 ± 12.9) kJ·mol–1.

The reaction enthalpy of formation of the tetracyano complex has been studied calorimetrically by Varet [1896VAR] and by Izatt and his co-workers [63CHR/IZA], [71IZA/JOH] (see Table V-31). The value determined in [63CHR/IZA] is selected in this review:

r mH ο∆ ((V.113), 298.15 K) = – (180.7 ± 4.0) kJ·mol–1.

Izatt et al. also determined the standard enthalpy of Reaction (V.113) at two other temperatures (10 and 40°C), which allowed calculation of the standard molar heat capacity of the reaction r ,mpCο∆ ((V.113), 298.15 K) = (150.6 ± 42.0) J·K–1·mol–1 [71IZA/JOH].

Neglecting the ionic strength dependence, r mH ο∆ ((V.114), 298.15 K) = − (10.4 ± 4.0) kJ·mol–1 can be estimated from the temperature variation of 10 5log K [60MCC/JON]. The combination of the above two enthalpy values yielded:

r mH ο∆ ((V.115), 298.15 K) = – (191.1 ± 8.0) kJ·mol–1.

Table V-31: Experimental enthalpy values for Reactions (V.113) and (V.114).

Method Medium Im t (°C) r mH∆ (kJ·mol–1) reported

r mH∆ (kJ·mol–1) retained

Reference

Ni2+ + 4CN– 24Ni(CN) −

cal ? ? 14 – (180 ± 4.2) [1896VAR] cal (self) 0 corr 25 – (180.7 ± 2.0) – (180.7 ± 4.0) [63CHR/IZA] cal (self) 0 corr 10

40 – (189.1 ± 1.2) – (183.7 ± 0.8)

– (189.1 ± 2.0) – (183.7 ± 2.0)

[71IZA/JOH]

24Ni(CN) − + CN– 3

5Ni(CN) −

sp Na(ClO4) 1.43 15 – 34 – 12.5 – (10.4 ± 4.0)(a) [60MCC/JON]

(a) Calculated from the temperature dependence of the corresponding formation constant, see Appendix A.

From these data

f mH ο∆ ( 24Ni(CN) − , 298.15 K) = (353.7 ± 14.7) kJ·mol–1,

and


231

f mH ο∆ ( 35Ni(CN) − , 298.15 K) = (490.6 ± 19.4) kJ·mol–1

are selected in this review.

These above selections yield:

mS ο ( 24Ni(CN) − , 298.15 K) = (245.0 ± 36.6) J·K–1·mol–1,

mS ο ( 35Ni(CN) − , 298.15 K) = (278.8 ± 51.1) J·K–1·mol–1.

V.7.1.2.1.2 Protonated complexes

The cyano complexes of several metal ions are reported to undergo protonation in the acidic pH-range [87BEC], but rather contradictory results are available for Ni(II). The protonation constants of 2

4Ni(CN) − ( 10 1log K = (5.4 ± 0.2), 10 2log K = (4.5 ± 0.2), 10 3log K = (2.6 ± 0.4) at t = 25°C and I = 0.1 M NaClO4) were first reported by Kolski

and Margerum, based on a kinetic study of the acid dissociation of 24Ni(CN) −

[68KOL/MAR]. Their equilibrium measurements between pH = 4 – 5, using a spectrophotometric method, resulted in systematic changes in the calculated 10 4log b values. This was taken as further proof for the presence of protonated species, although the UV-VIS spectrum of 2

4Ni(CN) − was found to be unaffected by the protonation. Two successive protonations of 2

4Ni(CN) − have been reported using a pH-metric method in [74KAB], but the results are uncertain, and thus this work was not considered in the present review (see Appendix A). Recently, Monlien et al. reported a kinetic study on the cyanide exchange with the 2

4Ni(CN) − complex [2002MON/HEL]. The kinetic data confirmed the presence of protonated species, although the protonation did not affect the 13C NMR spectrum of 2

4Ni(CN) − . On the other hand, spectrophotometric [59FRE/SCH] and careful potentiometric [63CHR/IZA], [71IZA/JOH], [74PER] studies performed between pH = 4 – 7 did not indicate the existence of protonated species in notable concentrations, although their kinetic role in the acid dissociation of 2

4Ni(CN) − was later corroborated [76PER].

To sum up, the kinetic role of the protonated species in some reactions of 24Ni(CN) − is well established, but the presence of such species in equilibrated solutions

is at least questionable. Additional experimental data would be needed for a definitive conclusion.

V.7.1.3 Nickel thiocyanates

V.7.1.3.1 Aqueous nickel thiocyanato complexes

The thiocyanate ion is a widely studied ligand, mainly due to its ambidentate nature. Aside from some general reviews [75NEW2], [79GOL/KOH], a critical survey of sta-bility constants of thiocyanate complexes has been published [97BAH/PAR]. The latter publication listed the majority of available thermodynamic data for the Ni(II) – thiocy-anate system, but recommended values are given only for 10 1log b (V.116) at I = 1.0 and 1.5 M.


232

Thiocyanic acid is a moderately strong acid. Only a few badly scattered data are available for estimation of its protonation constant [97BAH/PAR]. Among them,

10 1log K = – 0.87 at I = 2.0 M [82NEV/ANG] seems to be the most reliable. Thus, the protonation of the thiocyanate anion is negligible even at the most acidic conditions used [64TRI/CAL], [68MAL/TUR] for studying the Ni(II)-thiocyanate interaction.

The thiocyanate anion is N-coordinated in most of its Ni(II) compounds [77CHE/BRO]. The aqueous Ni(II)-thiocyanate complexes have octahedral geometry, although the [Ni(SCN)4(H2O)2]2– species is probably in equilibrium with a minimal concentration of tetrahedral [Ni(SCN)4]2– [88BJE2]. Some early papers reported qualita-tive information on the Ni(II) – thiocyanate system [38CSO], [42MAJ] based on spec-trophotometric studies. Later, several experimental techniques, namely ion exchange, spectrophotometric, polarographic, distribution, IR, potentiometric and kinetic methods, have been used to derive equilibrium data for the Ni(II)-thiocyanate complexes. De-pending on the ligand-to-metal ratio, the formation of four species is generally recog-nised:

Ni2+ + qSCN– 2Ni(SCN) qq

− (V.116)

with q = 1 – 4. The reported formation constants are listed in Table V-32. For reasons discussed in Appendix A, the data reported in [61DAV/SMI], [61MOH/DAS], [62FRO/LAR], [62TRI/CAL], [71DAS/DAS], [71LAN/CUM], [74NET/BAT] were not considered in the present review. In some cases [53FRO2], [64TRI/CAL], [74DIC/HOF], the formation constants reported for 20°C were corrected to 25°C, using the enthalpy values selected below. A majority of the accepted experimental data refers to the formation of the NiSCN+ species in perchlorate media. SIT analysis of these data shows acceptable consistency (cf. Figure V-50). The weighted linear regression using 12 data points yielded the selected value of:

10 1log οb ((V.116), q = 1, 298.15 K) = (1.81 ± 0.04).

Since no experimental data are available, in the above treatment ε(H+, SCN–) and ε(Li+, SCN–) ion interaction parameters were assumed to be equal to ε(Na+, SCN−). From the slope in Figure V-50, ∆ε((V.116), q = 1, 4ClO− ) = − (0.109 ± 0.025) kg·mol–1 can be calculated. Using the selected values for ε(Ni2+, 4ClO− ) and ε(Na+, SCN–), ∆ε((V.116), q = 1, 4ClO− ) leads to a value of ε(NiSCN+, 4ClO− ) = (0.31 ± 0.04) kg·mol–1.

Less data are available for the formation of the Ni(SCN)2(aq) and 3Ni(SCN)− species. The SIT treatment of the accepted data (cf. Figure V-51 and Figure V-52) re-sulted in the following selected thermodynamic formation constants:

10 2log οb ((V.116), q = 2, 298.15 K) = (2.69 ± 0.07),

10 3log οb ((V.116), q = 3, 298.15 K) = (3.02 ± 0.18).

From the slopes, ∆ε((V.116), q = 2, 4ClO− ) = – (0.091 ± 0.043) kg·mol–1 and ∆ε((V.116), q = 3, 4ClO− ) = (0.14 ± 0.12) kg·mol–1 can be derived. These parameters


233

lead to the values of ε(Ni(SCN)2(aq), Na+ + 4ClO− ) = (0.38 ± 0.06) kg·mol–1 (also see [97ALL/BAN]) and ε(Na+, 3Ni(SCN)− ) = (0.66 ± 0.13) kg·mol–1. The use of an interac-tion coefficient like ε(Ni(SCN)2(aq), Na+ + 4ClO− ) for a neutral species is not a standard procedure within the context of the SIT as described in Appendix B, but appears to be justified in the current case. Further experimental work could prove useful.

Several experimental values are published for 10 4log b (or 10 4log K ) [64TRI/CAL], [71LAN/CUM], (or [88BJE2]). In [64TRI/CAL] there was a rather small excess of ligand (

SCNc − ≤ 0.5 M), and for both ionic strengths K3 < K4 is reported. Con-

sidering that no geometrical change of Ni(II) occurs during the stepwise complex formation, the reported K4 values are probably overestimated. The method used in [71LAN/CUM] and [90BJE2] is not compatible with the SIT. Although there is good evidence for the existence of 2

4Ni(SCN) − [88BJE2], the available data cannot be used to derive a selected value.

The above listed selected 10log qb values correspond to:

r mGο∆ ((V.116), q = 1, 298.15 K) = – (10.33 ± 0.22) kJ·mol–1,

r mGο∆ ((V.116), q = 2, 298.15 K) = – (15.35 ± 0.40) kJ·mol–1,

r mGο∆ ((V.116), q = 3, 298.15 K) = – (17.24 ± 1.03) kJ·mol–1, and hence

f mGο∆ (NiSCN+, 298.15 K) = (36.60 ± 4.08) kJ·mol–1,

f mGο∆ (Ni(SCN)2, aq, 298.15 K) = (124.27 ± 8.04) kJ·mol–1,

f mGο∆ ( 3Ni(SCN)− , 298.15 K) = (215.09 ± 12.07) kJ·mol–1.

Table V-32: Experimental equilibrium data for the Ni(II) – thiocyanate system.


10log xb

retained(a) Reference

Ni2+ + SCN– NiSCN+ cix (Na)ClO4 1.05 20

25 (1.18 ± 0.02) –

– (1.12 ± 0.06)(b)

[53FRO2]

sp Ni(NO3)2 0 corr(c) 23 (1.67 ± 0.03) (1.69 ± 0.15) [58YAT/KOR] sp pot

Ni(ClO4)2 + KSCN 0 corr 25 35

(1.50 ± 0.04) (1.82 ± 0.02)

[61MOH/DAS]

sp ? 0.5 M 1.4 1.38 [61DAV/SMI] ir (NaClO4) 3.5 1.18 [62FRO/LAR] pol (Na, H)ClO4 1.05(c) ? (1.17 ± 0.02) [62TRI/CAL] sp (Na)ClO4 0 corr

0.43 25 (1.762 ± 0.005)

– (1.77 ± 0.10)(d) (1.20 ± 0.08)(d)

[62WIL2]



234



10log xb

accepted(a) Reference

Ni2+ + SCN– NiSCN+ dis (NaClO4) 1.62

3.50

20 25 20 25

(1.14 ± 0.02) – (1.19 ± 0.02) –

– (1.08 ± 0.10)(b) – (1.09 ± 0.10)(b)

[64TRI/CAL]

pol (H)ClO4 0.67 15 25 35

(1.34 ± 0.02) (1.24 ± 0.02) (1.17 ± 0.03)

(1.33 ± 0.10) (1.23 ± 0.10) (1.16 ± 0.10)

[68MAL/TUR]

dis (Na)ClO4 0.26 25 (1.3 ± 0.2) (1.29 ± 0.20) [69SUB/COR] sp (LiClO4) 3.48 25 (1.34 ± 0.01) (1.27 ± 0.10) [70MIR/MAK] pot KSCN 0 corr 25

35 45

(1.98 ± 0.02) (1.85 ± 0.02) (1.74 ± 0.02)

[71DAS/DAS]

aix KSCN 0 corr ? 2.22 [71LAN/CUM] pol (KNO3) 0.2

0.2 0.2(c) 0.2

5 15 25 35

1.55 1.46 1.37 1.29

[71TUR/MAL]

kin (NaClO4) 1.05 25 45

(1.09 ± 0.06) (1.00 ± 0.03)

(1.07 ± 0.10) (0.98 ± 0.10)

[73HUT/HIG]

kin 0 corr 20 25

1.83 –

– (1.79 ± 0.10)(b)

[74DIC/HOF]

cal (NaClO4) 1.05 25 (1.14 ± 0.02) (1.12 ± 0.02) [74KUL3] ir (LiClO4) 3.48 ? 0.3 – 1.44 [74NET/BAT] dis (NaClO4) 1.05 25 1.1 (1.08 ± 0.10) [76MUR/KUR]

Ni2+ + 2SCN– Ni(SCN)2(aq) cix (Na)ClO4 1.05 20

25 (1.64 ± 0.04) –

– (1.54 ± 0.06)(b)

[53FRO2]

dis (NaClO4) 1.62

3.50

20 25 20 25

(1.76 ± 0.06) – (1.68 ± 0.05) –

– (1.63 ± 0.10)(b) – (1.48 ± 0.10)(b)

[64TRI/CAL]

aix KSCN 0 corr ? 2.43 [71LAN/CUM] pol (KNO3) 0.2

0.2 0.2(c) 0.2

5 15 25 35

2.23 2.04 1.86 1.69

[71TUR/MAL]

cal (NaClO4) 1.05 25 (1.58 ± 0.05) (1.54 ± 0.05) [74KUL3] dis (NaClO4) 1.05 25 1.6 (1.56 ± 0.20) [76MUR/KUR]



235



10log xb

accepted(a) Reference

Ni2+ + 3SCN– 3Ni(SCN)− cix (Na)ClO4 1.05 20

25 (1.81 ± 0.07) –

– (1.66 ± 0.10)(b)

[53FRO2]

dis (NaClO4) 1.62

3.50

20 25 20 25

(1.70 ± 0.10) – (1.3 ± 0.3) –

– (1.52 ± 0.20)(b) – (1.01 ± 0.30)(b)

[64TRI/CAL]

aix KSCN 0 corr ? 2.62 [71LAN/CUM] cal (NaClO4) 1.05 25 (1.6 ± 0.2) (1.54 ± 0.20) [74KUL3]

Ni2+ + 4SCN– 24Ni(SCN) −

dis (NaClO4) 1.62 3.50

20 20

(2.04 ± 0.20) (1.54 ± 0.35)

[64TRI/CAL]

aix KSCN 0 corr ? 2.15 [71LAN/CUM] sp NaSCN 0 corr 25 10 4log K = – (0.92 ± 0.18)(e) [88BJE2] (a) Accepted values corrected to the molal scale. The accepted values reported in Appendix A are expressed

on the molar or molal scales, depending on which units were used originally by the authors. (b) Corrected to 25°C, using the accepted enthalpy values. (c) Data for several ionic strengths, see Appendix A. (d) Re-evaluated value, see Appendix A. (e) ‘Semi-thermodynamic’ constant, the activity of thiocyanate ion (using ±γ of NaSCN) was taken into

account.

Figure V-50: Extrapolation to Im = 0 of the accepted formation constants reported for the formation of NiSCN+ in perchlorate media.

0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0

1.6

1.8

2.0

2.2

2.4

log 10

β1 +

4D

I / mol·kg−1


236

Figure V-51: Extrapolation to Im = 0 of the accepted formation constants reported for the formation of Ni(SCN)2(aq) in perchlorate media.

Figure V-52: Extrapolation to Im = 0 of the accepted formation constants reported for the formation of 3Ni(SCN)− in perchlorate media.

0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.52.2

2.4

2.6

2.8

3.0

3.2

3.4

log 10

β2 +

6D

I / mol·kg−1

0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.51.6

2.0

2.4

2.8

3.2

3.6

log 10

β3 +

6D

I / mol·kg−1


237

Five reliable reports are available for the reaction enthalpies of the formation of 2Ni(SCN) xx

− (x = 1 – 3) complexes. The reported values are collected in Table V-33.

Table V-33: Experimental reaction enthalpy values for the formation of Ni(II)-thiocyanate complexes.

Method Medium Im t (°C) r mH∆ (kJ·mol–1) reported

r mH∆ (kJ·mol–1) retained for selection

Reference

Ni2+ + SCN– NiSCN–

cal NiClO4 + KSCN < 0.06 25 – (10.46 ± 0.40) – (9.25 ± 2.00)(a) [67NAN/TOR] pol (H)ClO4 0.67 15 – 35 – 14.43 – (14.5 ± 2.0) [68MAL/TUR] pol (KNO3) 0.2 5 – 35 – – (14.3 ± 2.0)(b) [71TUR/MAL] kin (NaClO4) 1.05 25 – 45 – – (8.3 ± 2.0)(b) [73HUT/HIG] cal (NaClO4) 1.05 25 – (12.02 ± 0.15) (12.0 ± 0.5) [74KUL3]

Ni2+ + 2SCN– Ni(SCN)2(aq)

pol (KNO3) 0.2 5 – 35 – – (29.5 ± 2.0)(b) [71TUR/MAL] cal (NaClO4) 1.05 25 – (20.92 ± 1.15) – (20.9 ± 1.2) [74KUL3]

Ni2+ + 3SCN– 3Ni(SCN)−

cal (NaClO4) 1.05 25 – (29.12 ± 5.15) – (29.1 ± 5.2) [74KUL3] (a) Re-evaluated value, see Appendix A. (b) Calculated from the temperature dependence of the corresponding formation constant, see Appendix A.

These data do not allow correct evaluation of the ionic strength dependence, therefore it was assumed that reaction enthalpies are independent of the ionic strength. Assigning higher weight to the most reliable calorimetric data published in [74KUL3], and considering the above assumption, the weighted averages of the retained values are as follows:

r mH ο∆ ((V.116), q = 1, 298.15 K) = – (11.8 ± 5.0) kJ·mol–1,

r mH ο∆ ((V.116), q = 2, 298.15 K) = – (21 ± 8) kJ·mol–1,

r mH ο∆ ((V.116), q = 3, 298.15 K) = – (29 ± 10) kJ·mol–1,

From the above selected r mH ο∆ values, the following standard enthalpies of formation can be derived:

f mH ο∆ (NiSCN+, 298.15 K) = (9.59 ± 6.46) kJ·mol–1,

f mH ο∆ (Ni(SCN)2, aq, 298.15 K) = (76.79 ± 11.35) kJ·mol–1,

f mH ο∆ ( 3Ni(SCN)− , 298.15 K) = (145.19 ± 15.65) kJ·mol–1.


mS ο (NiSCN+, 298.15 K) = (7.5 ± 25.4) J·K–1·mol–1,

mS ο (Ni(SCN)2, aq, 298.15 K) = (137.8 ± 46.5) J·K–1·mol–1,


238

mS ο ( 3Ni(SCN)− , 298.15 K) = (261.5 ± 66.2) J·K–1·mol–1.

V.7.2 Silicon compounds and complexes V.7.2.1 Solid nickel silicates

V.7.2.1.1 Nickel orthosilicate Ni2SiO4(cr)

V.7.2.1.1.1 Crystal structure and phase transitions

The only stable silicate in the system NiO – SiO2 is the orthosilicate 2NiO·SiO2 [54EIT], [66GME], [63PHI/HUT], [69LEV/ROB], [91KNA/KUB]. The existence of the pure metasilicate NiSiO3 has not been confirmed by any X-ray investigation [66GME]. However, an estimation of the thermodynamic data and the stability for the nickel metasilicate (pyroxene) was performed by Campbell and Roeder [68CAM/ROE] and by Navrotsky [71NAV]. In the present review, the thermodynamic data for the NiSiO3 are not evaluated because the pure nickel pyroxene is unstable relative to silica and the nickel olivine at a total pressure of one atmosphere [68CAM/ROE].

The 2NiO·SiO2, orthorhombic crystals, have the olivine structure, but at higher temperature and pressure a spinel transformation can occur. This transformation was the subject of numerous studies [62RIN], [65AKI/FUJ], [68NAV/KLE], [73NAV], [73NAV2], [76NAV/KAS], [74YAG/MAR], [74MA] because of great geophysical in-terest. The approximate equilibrium pressure for the olivine – spinel transition in Ni2SiO4 at 650°C was found to be 18 kb [62RIN]. This value is about 40% lower than that determined by [74MA] who found a linear pressure – temperature dependence of the transformation, expressed by the equation: p/bar = 23000 + 11.8 t /°C. The follow-ing thermodynamic data for the olivine (ol.) – spinel (sp.) transition in Ni2SiO4 were obtained by Navrotsky [73NAV] by measuring the heats of solution of both polymorphs in a molten oxide at 986 K: the enthalpy of the transition Ni2SiO4(ol.) → Ni2SiO4(sp.),

trs mH ο∆ (986 K) = (5.9 ± 2.9) kJ·mol–1, the heat content increments, ( mH (986 K) –mH ο (298 K)), olivine, (107.7 ± 1.8) kJ·mol–1, spinel, (106.2 ± 0.8) kJ·mol−1 and the en-

tropy of the olivine – spinel transition trs mS ο∆ = 12.6 to 14.6 J·K–1·mol–1. Although the spinel polymorph of Ni2SiO4 also exists in metastable form at atmospheric pressure, the thermodynamic data for this high pressure compound are not comprehensively treated in the present review.

Ma [74MA] found that Ni2SiO4 olivine melts incongruently at high pressures to NiO plus melt and that it is a stable phase until melting occurs. Since such melting behaviour persists at a pressure as low as 4.9 kbar, he expected that similar melting be-haviour also occurs at atmospheric pressure. The melting curve extrapolated to atmos-pheric pressure gave a melting temperature of 1848 K [74MA]. The phase diagram NiO-SiO2, drawn from data obtained by quenching and direct observational techniques shows, however, that the decomposition temperature of the nickel olivine is


239

(1818 ± 5) K [63PHI/HUT], [69LEV/ROB]. The dissociation reaction occurring prior to melting is

Ni2SiO4(cr) 2NiO(cr) + SiO2(cris.) (V.117)

where (cris.) means cristobalite.

The uncertainty of the melting and decomposition temperature of Ni2SiO4 oli-vine was minimised by refining the thermodynamic data concerning this silicate [87NEI]. The decomposition temperature of Ni2SiO4 may therefore be given as (1820 ± 5) K [87NEI]. This is in perfect agreement with Phillips et al. [63PHI/HUT].

V.7.2.1.1.2 Crystallography and mineralogy of nickel orthosilicate

Liebenbergite (Ni2SiO4) is an end-member mineral of olivines being a group of minerals of the general formula X2SiO4, where X is a divalent metal cation (Mg, Fe, Mn, Ni, Ca, and Co) [82BRO]. The pure nickel olivine (liebenbergite) does not occur naturally [73WAA/CAL], [2001HEN/RED]. The natural occurrence of liebenbergite (Ni, Mg)2SiO4 of an approximate empirical formula Ni1.5Mg0.5SiO4 has been reported from Scotia Talc mine, Bon Accord, Barbeton distr., Transvaal, South Africa [73WAA/CAL]. Liebenbergite was named after W. R. Liebenberg, director of the Na-tional Institute of Metallurgy of South Africa. The crystal system of liebenbergite is orthorombic (space group Pbnm; Z = 4) [87BOS]. Cell data for the end-member Ni2SiO4 are a = 4.727 Å, b = 10.120 Å, c = 5.911 Å, V = 284.98 Å3 [2001HEN/RED]. The X-ray powder diffraction pattern is listed on JCPDS-ICDD card No. 15 –388.

V.7.2.1.1.3 Heat capacity and standard entropy

Heat capacity measurements of the nickel silicate olivine were reported by [82WAT], [65EGO/SMI], and [84ROB/HEM]. However, the heat capacities reported by Watanabe in the temperature range 350 to 700 K [82WAT] are 4.6% greater than the heat capaci-ties obtained by [84ROB/HEM] with the more accurate low-temperature calorimeter. Therefore, the results of Watanabe are not incorporated in this review. Egorov and Smirnova [65EGO/SMI] have determined the specific heat capacities of the nickel or-thosilicate in the temperature range between 555 to 1570 K using a copper block calo-rimeter. The well-described, careful measurements allowed a determination of ,mpCο to within ± 0.5%. These high-temperature data were used in this review to extend the heat capacity equation up to 1570 K. Robie et al. [84ROB/HEM] have measured the heat capacities of Ni2SiO4-olivine between 5 and 387 K by cryogenic adiabatic-shield calo-rimetry and between 360 and 1000 K by differential scanning calorimetry. The standard molar heat capacity at 298.15 K, determined from these low-temperature measurements, was reported to be (123.20 ± 0.20) J·K–1·mol–1 [84ROB/HEM]. This value, the only one experimentally determined by accurate calorimetric measurements, is selected by the present review:

,mpCο (Ni2SiO4, olivine, 298.15 K) = (123.20 ± 0.50) J·K–1·mol–1.


240

An assumption of a higher uncertainty for the accepted value of ,mpCο (Ni2SiO4, olivine, 298.15 K) was necessary to achieve the coincidence of the standard heat capac-ity value at 298.15 K obtained from the low temperature measurements [84ROB/HEM] and from the high temperature thermal heat capacity function selected by this review (see below).

According to [84ROB/HEM] between 300 and 1300 K the heat capacities can be represented by the equation:

,mpCο (T) = (289.73 – 0.024015 T + 131045 T –2 – 2779.0 T –1/2) J·K–1·mol–1 to within ± 0.5%.

It should be mentioned that these authors carried out measurements only in the temperature range between 5 and 1000 K. Therefore, the extrapolation using the above equation up to 1300 K seems to be less reliable. In this review, we have used the calo-rimetric heat capacity measurements of [65EGO/SMI] and [84ROB/HEM] and fitted the thermal heat capacity function in the temperature range 270 < (T/K) < 1570. The se-lected experimental data and the fitted curve are shown in Figure V-53. The coincidence of both data sets is satisfactory. Thus, we can select the following thermal function of the heat capacity for Ni2SiO4-olivine between 270 and 1570 K:

,mpCο (T) = (147.61612 + 0.0486860 T – 1.50143·10–5- T 2 – 3.34214·106 T –2) J·K–1·mol–1 with an uncertainty of ± 0.5%.

Figure V-53: The standard molar heat capacity of Ni2SiO4-olivine as a function of tem-perature. The data are taken from [84ROB/HEM] ( ) and [65EGO/SMI] ( ). The fit-ted curve (solid line) corresponds to the thermal heat capacity function selected in this review.

0 200 400 600 800 1000 1200 1400 1600110

120

130

140

150

160

170

180

190

T / K

C0 p, m

/ J·m

ol−1

·K−1

C° p,m

/ J·K

–1·m

ol–1


241

The standard molar entropy at 298.15 K for Ni2SiO4-olivine as determined by calorimetric measurements by Robie et al. [84ROB/HEM] is:

mS ο (Ni2SiO4, olivine, 298.15 K) = (128.10 ± 0.20) J·K–1·mol–1.

This value was also selected by the present review. The entropy value calcu-lated from the temperature dependence of the Gibbs energy of the reaction of nickel with silica and oxygen e.g., as estimated by [68CAM/ROE], [76MAH/PAN], was not accepted by this review.

V.7.2.1.1.4 Enthalpy of formation

The enthalpy of formation of Ni2SiO4 olivine from the component oxides was measured by solution calorimetry in a molten oxide solvent at (965 ± 2) K by Navrotsky [71NAV].

The value of r mH∆ ((V.118), 965 K) for the reaction:

2NiO(cr) + SiO2(q) Ni2SiO4(ol.) (V.118)

was reported to be – (13.9 ± 1.9) kJ·mol–1. Using the heat capacities of the components, the standard molar enthalpy of this reaction at 298.15 K was calculated: r mH ο∆ ((V.118), 298.15 K) = – (6.3 ± 6.7) kJ·mol–1 [78KOT/MUL]. The main interest, for this review, is the determination of the standard enthalpy of formation of Ni2SiO4-olivine from the elements in their standard state i.e., f mH ο∆ for the reaction:

2Ni(cr) + Si(cr) + 2O2(g) Ni2SiO4(cr) (V.119)

The enthalpy of Reaction (V.118), together with the enthalpies of the forma-tion of SiO2(q), where (q) means quartz, and NiO(cr) and the thermal functions of heat capacity of all components can of course be used for the calculation of

f mH ο∆ (298.15 K). Generally, the enthalpy of the reaction based solely on calorimetric data is the least reliable quantity, and contributes the major part of the uncertainty in the determination of the Gibbs energy of reaction. On the other hand, the standard entropy and the high-temperature heat capacities as well as the Gibbs energy of a reaction do provide more accurate data [87NEI]. Thus, using the third-law method, i.e., the equa-tion:

r mH ο∆ (298.15 K) = r mGο∆ (T) – r ,m298.15

T

pCο∆∫ dT +

T ( r mS ο∆ (298.15 K) + r ,m298.15( / )

T

pC Tο∆∫ dT), (V.120)

the standard molar enthalpy of a reaction at 298.15 K can be calculated. Based on his emf-measurements and the heat capacity values, among others those for the Ni2SiO4-olivine determined by [84ROB/HEM], O´Neill [87NEI] used this method to calculate

r mH ο∆ (298.15 K) for the reaction:

2Ni(cr) + SiO2(cr) + O2(g) Ni2SiO4(cr) (V.121)


242

to be – 482.42 kJ·mol–1. Using this value it was possible to calculate the Gibbs energy of Reaction (V.118) and to achieve very good agreement between the calculated and experimentally observed decomposition temperature of Ni2SiO4-olivine, i.e., (1820 ± 5) K [87NEI].

Robie and Hemingway [95ROB/HEM] used their accurate heat capacity meas-urements, combined with results from molten salt calorimetry, thermal decomposition of the Ni2SiO4-olivine into its constituent oxides, and equilibrium studies, both by CO reduction and solid state electrochemical cell measurement for Reaction (V.121) [87NEI], and calculated the following standard molar enthalpy of formation of Ni2SiO4-olivine (liebenbergite): f mH ο∆ (298.15 K) = – (1396.5 ± 3.0) kJ·mol–1.

In the present review, the standard molar enthalpy of formation of Ni2SiO4−olivine at 298.15 K was also determined by the third-law method using the thermal heat capacity function of Ni2SiO4-olivine, the equation representing the Gibbs energy of Reaction (V.121) in the temperature range 970 – 1770 K, both derived by this review as described in Sections V.7.2.1.1.3 and V.7.2.1.1.5, as well as thermodynamic data of elemental nickel selected by this review. Elemental silicon, oxygen and silica (quartz) data are taken from the NEA-TDB auxiliary dataset. The value obtained and selected is:

f mH ο∆ (Ni2SiO4, olivine, 298.15 K) = – (1396.0 ± 5.0) kJ·mol–1

and is in an excellent agreement with the value recommended by [95ROB/HEM]. A higher uncertainty than calculated was ascribed to f mH ο∆ (298.15 K) to overcome the problem of inconsistency arising during calculation of the decomposition temperature of Ni2SiO4-olivine into NiO and SiO2-cristobalite. Using the selected value of f mH ο∆ it is possible to calculate the decomposition temperature to be (1820 ± 5) K. It should be noted that our calculations of f mH ο∆ were carried out using the thermodynamic data for SiO2-quartz and not for SiO2-cristobalite. However, no thermodynamic data of SiO2(cristobalite) are available in the CODATA compilation [89COX/WAG] or in the TDB auxiliary data appendices used till now. The reason for it is probably the fact, that the thermodynamic data of SiO2(cristobalite) are less reliable. This issue was also dis-cussed briefly by [87NEI]. Using the thermodynamic data of SiO2-cristobalite given by [95ROB/HEM] and the same procedure for the calculation of f mH ο∆ (298.15 K) of Ni2SiO4-olivine as described above for SiO2-quartz, we have calculated:

f mH ο∆ (298.15 K) = – (1401.15 ± 3.40) kJ·mol–1 and the proper decomposition tempera-ture of Ni2SiO4-olivine (1820 ± 5)°C. Different sets of values for low and high tempera-ture SiO2(cristobalite) have been proposed by [98CHA]. The other value of the standard molar enthalpy of formation of Ni2SiO4-olivine proposed in the literature

f mH ο∆ (298.15 K) = – 1378.2 kJ·mol–1 [61LEB/LEV] was not used by this review (see Appendix A).


243

V.7.2.1.1.5 Gibbs energy of formation

The Gibbs energy of formation of Ni2SiO4-olivine from metallic nickel, oxygen and silica according to Reaction (V.121) has been determined by equilibration of this sili-cate or metallic nickel and silica with a CO/CO2 atmosphere [61LEB/LEV], [63BUR/ABB], by determination of the oxygen fugacity over the mixture of nickel ox-ide-silica and the gas atmosphere of known CO2/H2-ratio [68CAM/ROE], or electro-chemically by use of cells with a solid electrolyte [64TAY/SCH], [75LEV/GOL], [84ROG/BOR], [98ROG/KOZ], [2000ROG/KOZ], [87NEI] and [86JAC/KAL]. The results of the paper [86JAC/KAL] were not used for further evaluations, because they were presented only in graphical form. Thus, additional uncertainties in recovering the values from the original report were expected. A comparison of all results obtained by the determination of the Gibbs energy of formation of Ni2SiO4-olivine from metallic nickel, oxygen and silica, reported in the literature, is shown on Figure V-54. Although, different reactions were investigated, comparison was possible by recalculation, using the original data reported in the papers under consideration, to obtain the Gibbs energy of Reaction (V.121). Reaction (V.121) is the overall reaction in the solid electrolyte cells used for the determination of r mGο∆ . Thus, it has been directly investigated in pa-pers [64TAY/SCH], [75LEV/GOL], [84ROG/BOR], [98ROG/KOZ], [2000ROG/KOZ], [87NEI]. The study of Campbell and Roeder [68CAM/ROE] also directly refers to the equilibrium of Reaction (V.121). On the other hand, the studies of nickel orthosilicate reported by Lebedev and Levitskii [61LEB/LEV] and by Burdese et al. [63BUR/ABB] are concerned with the equilibrium of the reaction:

2Ni(cr) + SiO2(cr) + 2CO2(g) Ni2SiO4(cr) + 2CO(g). (V.122)

The calculations of the Gibbs energy of Reaction (V.121) based on the Gibbs energy determinations of Reaction (V.122) requires the r mGο∆ of Reaction (V.123)

2CO(g) + O2(g) 2CO2(g). (V.123)

Using the r mGο∆ values for Reactions (V.122) and (V.123) reported by [61LEB/LEV] and [63BUR/ABB], the Gibbs energies of Reaction (V.121) were calcu-lated and are also shown on the diagram r mGο∆ versus T, cf. Figure V-54, together with results of the solid electrolyte cell measurements.

A comparison of the Gibbs energy of Reaction (V.121) as obtained by different authors, Figure V-54, shows very good agreement between all solid electrolyte emf-studies and the oxygen fugacity study of Campbell and Roeder [68CAM/ROE]. The results of Lebedev and Levitskii [61LEB/LEV] and of Burdese et al. [63BUR/ABB] are significantly different. Therefore, they were disregarded by the present review and were not used for the non-linear curve fit. It should be mentioned that no significant differ-ence between the value of the Gibbs energy of Reaction (V.121) was found when SiO2 cristobalite or quartz was used in solid electrolyte emf-studies reported by O´Neill


244

[87NEI]. The equation describing the fitted curve shown in Figure V-54 represents the Gibbs energy of Reaction (V.121) in the temperature range 970 – 1770 K:

r mG∆ (T) = (– (522.89819 ± 8.94994) + (0.39674 ± 0.05714) T – (0.02629 ± 0.00700) T lnT) kJ·mol–1.

As already discussed in Section V.7.2.1.1.4, this equation was used in the pre-sent review for the calculation of the molar enthalpy of formation of Ni2SiO4-olivine at 298.15 K. The standard entropy change corresponding to Reaction (V.119)

r mS ο∆ (298.15 K) was calculated using mS ο (Ni2SiO4, olivine, 298.15 K) and the standard entropy of elemental nickel selected in this review. The other data i.e., the standard en-tropies of elemental silicon, oxygen and silica (quartz) at 298.15 K were taken from the NEA TDB auxiliary dataset. The Gibbs energy of formation of Ni2SiO4-olivine at 298.15 K was selected to be:

f mGο∆ (Ni2SiO4, olivine, 298.15 K) = – (1288.44 ± 5.00) kJ·mol–1.

The selected value is in good agreement with f mGο∆ = – (1289.0 ± 3.1) kJ·mol–1 reported by [84ROB/HEM].

Figure V-54: Standard Gibbs energy of Reaction (V.121) as a function of temperature. The data are taken from [68CAM/ROE] , [64TAY/SCH] , [75LEV/GOL] , [84ROG/BOR] , [98ROG/KOZ] , [2000ROG/KOZ] , [87NEI] +, [87NEI] SiO2(cr) , [61LEB/LEV] and [63BUR/ABB] . The solid line refers to the non-linear curve fit used by this review. The data of [61LEB/LEV] and [63BUR/ABB] were disregarded.

1000 1100 1200 1300 1400 1500 1600 1700 1800

-320

-300

-280

-260

-240

-220

-200

-180

-160

∆rG

0 m /

kJ·m

ol−1

T / K

∆rG

° m / k

J·mol

–1

2Ni(cr) + SiO2(cr) + O2(g) Ni2SiO4(cr)


245

V.7.3 Germanium compounds and complexes V.7.3.1 Solid nickel germanates

V.7.3.1.1 Nickel orthogermanate Ni2GeO4(cr)

V.7.3.1.1.1 Crystal structure of nickel orthogermanate

Ni – orthogermanate has a spinel structure and can easily be synthesised from the com-ponent oxides [66GME]. Germanates have long been regarded as suitable analogues for silicates, based on crystal chemistry systematics and, therefore, used for investigations to predict the strength of silicate spinel in the transition zone of the Earth [2001LAW/ZHA]. Even at high temperatures and pressures Ni2GeO4-spinel is a stable phase [66GME]. The cubic face-centred cell of the nickel orthogermanate, space group: Fd3m, has dimensions a = 8.221 Å, Z = 8 (according to JCPDS-ICDD card No. 10 – 266).

V.7.3.1.1.2 Enthalpy of formation from NiO(cr) and GeO2(cr)

Despite it being of geophysical interest, there are few thermodynamic data regarding nickel orthogermanate reported in the literature so far. Moreover, the reported data are incomplete. The direct determination of thermodynamic data for Ni2GeO4(cr) was per-formed by Navrotsky and Kleppa [68NAV/KLE] and by Navrotsky [71NAV]. These authors used oxide-melt solution calorimetry and determined the enthalpy of formation of Ni2GeO4-spinel from NiO(cr) and GeO2(quartz) at 965 K to be – (39.8 ± 1.7) kJ·mol−1. The enthalpy of formation values of Ni2GeO4(cr) from the constituent oxides given by other authors [78KOT/MUL] and [72KUB] were taken from the original work of [68NAV/KLE] or [71NAV]. However, the value reported by [78KOT/MUL], − (18.1 ± 1.7) kJ·mol–1, seems to be erroneously quoted. Lack of the thermal heat capac-ity function and the Gibbs energy of formation or entropy values for Ni2GeO4(cr) do not allow accurate calculation of thermodynamic data for this germanate at 298.15 K. Thus, the present review does not recommend any standard thermodynamic data for Ni2GeO4(cr).

V.8 Group 13 compounds and complexes V.8.1 Boron compounds and complexes Intermetallics are not treated in this review, therefore, thermodynamic data regarding borides are not included in this chapter.

V.8.1.1 Solid nickel borates

V.8.1.1.1 Crystal structure

Depending on the temperature, nickel borates can form crystals and glassy phases with different degrees of polymerisation, which are very often difficult to identify from the crystallographic point of view [66GME]. The numerous water-free and water-


246

containing nickel borate phases were characterised only by their stoichiometric compo-sition and not by X-ray diffraction methods [66GME]. The X-ray diffraction pattern is known for Ni3B2O6 (JCPDS card No. 22-745, 26-1284 and 75-1809).

V.8.1.1.2 Thermodynamic data for NiO·2B2O3(s)

So far no direct determination of thermodynamic data for water-free solid nickel borates has been reported in the literature. One of the reasons for this is probably the difficulty in obtaining well-defined crystalline phases. The only available data, published in a compilation of thermodynamic data for borate systems [74SLO/JON], are estimated values of the standard enthalpy of formation and the entropy for NiO·2B2O3 at 298.15 K. The value of r mH ο∆ (NiO·2B2O3, s, 298.15 K) = – (62.8 ± 20.0) kJ·mol–1 re-ported by [74SLO/JON] was obtained by a linear relationship between the enthalpy of reaction of the crystalline oxides and the ratio of cationic charge z to radius r. In view of the uncertainties with respect to the ionic radii selected, for (i) the z·r–1 versus r mH ο∆ correlation and for (ii) the Ni2+ ion, this value has not been accepted by this review.

According to Latimer´s method the standard entropies of predominantly ionic compounds can be obtained additively from values determined for the cationic and ani-onic constituents [51LAT]. Using the revised contribution of Ni2+ [93KUB/ALC] and the value for 2

4 7B O − derived by Richter and Vreuls [79RIC/VRE], the standard entropy was estimated to be mS ο (NiO·2B2O3, s, 298.15 K) = (129.9 ± 6.0) J·K–1·mol–1, which only slightly differs from the value given by [74SLO/JON] (137.6 ± 6.0). Generally this method is quite reliable, although the uncertainty would have to be increased to the 2σ level (± 12.0). The problem, however, is that no well defined solid compound is known to which this entropy could be assigned. Thus it was not accepted by this review.

V.8.1.1.3 Solubility of Ni(BO2)2·4H2O(s)

Shchigol [61SHC] studied the solubility of nickel orthoborate, Ni(BO2)2·4H2O(s). An average value of the solubility product of Ni(BO2)2·4H2O ,0sK ο = 2.16×10–9 mol3·dm–9 at 22°C was obtained. The recalculation (see Appendix A) resulted in

10 ,0log sK ο (Ni(BO2)2·4H2O, s) = – (8.80 ± 0.19). The nickel orthoborate used in the study [61SHC] was not structurally characterised by X-ray diffraction. It was synthesised by reacting nickel hydroxide or carbonate with orthoboric acid and only analytically inves-tigated with respect to its stoichiometric composition. Reliable thermodynamic quanti-ties can be determined from solubility measurements only when the solid phase is un-ambiguously characterised with respect to its structure. The results of solubility meas-urements of the nickel orthoborate given by [61SHC] do not fulfil this criterion, there-fore, no thermodynamic data for Ni(BO2)2·4H2O(s) were selected in this review.

No evidence of Ni(BO2)2·4H2O formation was found in a solubility study on the system NiCl2-ZnCl2-H3BO3-H2O [89BAL].


247

V.8.1.2 Aqueous nickel borato complexes

The existence of at least one soluble nickel borate complex was postulated by Shchigol [61SHC]. Nickel orthoborate Ni(BO2)2·4H2O(s) dissolves better in orthoboric acid solu-tions than in pure water. The formation of the nickel borate complex in solution was believed to occur according to the reaction:

Ni(BO2)2·4H2O(s) + H3BO3(aq) HNi(BO2)3(aq) + 5 H2O(l) (V.124)

The formula assigned to the proposed complex species was 2 3Ni(BO )− .

The dissociation constant of the aqueous nickel borate complex, according to Reaction (V.125), was calculated by Shchigol [61SHC], 3,1b (V.125) = 2.82×10–9 mol3·dm–9.

2 3Ni(BO )− Ni2+ + 3 2BO− (V.125)

An inspection of the experimental data revealed that they were misinterpreted, although Equation (V.124) may be correct, see Appendix A. Consequently the reported value of the dissociation constant was not selected by the present reviewers.


248

249

Chapter VI

VI Discussion of auxiliary data selection

VI.1 Group 15 auxiliary species VI.1.1 Additional consistent auxiliary data for As compounds As was the case during preparation of an earlier TDB volume [92GRE/FUG], auxiliary data for arsenic compounds are not being critically assessed at this time, and several self consistent values from the U.S. National Bureau of Standards compilation have been accepted as an interim measure. The values of ( mH (T) – mH ο (298 K)) for As(cr) are from Herrick and Feber [68HER/FEB], although it is recognised that the value of

mS ο (As, cr, 298.15 K) from the same paper is likely a better value than the TDB value, and differs from it by 0.6 J·K–1·mol–1. The high-temperature thermodynamic quantities for As4O6(g) are based on the work of Behrens and Rosenblatt, but are adjusted to ac-count for the differences in the values of mS ο (As4O6, cubic, 298.15 K). The consistent values: f mH ο∆ = (As4O6, g, 298.15 K) = – (1196.25 ± 16.00) kJ·mol–1 and mS ο = (As4O6, g, 298.15 K) = (408.6 ± 6.0) J·K–1·mol–1, and the enthalpy and entropy differences of Behrens and Rosenblatt [72BEH/ROS] are used in the present review. The similar measurements of Murray et al. [74MUR/POT] on the equilibrium between gas-phase As4O6 and the monoclinic solid are less well documented, and have not been taken into account here. The main contributions to uncertainties in the values for the values for As4O6(g) are from the uncertainties in the corresponding values for As4O6(cubic).

VI.2 Other auxiliary species VI.2.1 Auxiliary data for KCl(cr), KBr(cr) and KI(cr) The enthalpy of formation values for these solids in Glushko et al. [82GLU/GUR] are good, well-evaluated values that are very close to CODATA consistent. The value for

f mH ο∆ (KCl, cr, 298.15 K) used in the first book in the CODATA series [87GAR/PAR] (but not included in the CODATA selected data [89COX/WAG]) is only 0.02 kJ·mol–1 more positive than the value from [82GLU/GUR].

VI Discussion of auxiliary data selection

250

Detailed assessments of the aqueous enthalpy of solution values at 298.15 K for KCl(cr), KBr(cr) and KI(cr) were given in [82GLU/GUR]. In the present review, to obtain TDB-consistent values for the enthalpies of formation, these enthalpy of solution values are combined with the TDB enthalpy of formation values for K+ and Cl–, Br– and I– for 298.15 K.

Table VI-1: Enthalpy of formation values for KCl(cr), KBr(cr), and KI(cr) at 298.15 K

solid sln mH ο∆ / kJ·mol–1 f mH ο∆ (KX, cr) / kJ·mol–1 [82GLU/GUR] selected

KCl(cr) (17.241 ± 0.018) – (436.461 ± 0.129) KBr(cr) (19.780 ± 0.080) – (393.330 ± 0.188) KI(cr) (20.230 ± 0.100) – (329.150 ± 0.137)

The selected values of the enthalpies of formation differ by ≤ 0.15 kJ·mol–1 from those tabulated by Glushko et al. [82GLU/GUR]. A TDB reassessment of the en-thalpies of solution is not attempted, as it would involve analysis of a very large body of experimental data, and it is unlikely that the final result would be more accurate than the values from [82GLU/GUR]. It is possible that the enthalpy of solution values selected by Glushko et al. [82GLU/GUR] for these solids influenced the CODATA selections for the enthalpies of formation of the aqueous ionic species – the discussion in the CODATA volume [89COX/WAG] appears to lack the details that would be necessary to check this point. However, any extra uncertainties and inconsistencies from this source are expected to be small.

251

Part IV

Appendices

253

Appendix A Equation Section 1

This appendix comprises discussions relating to a number of key publications, which contain experimental information cited in this review. These discussions are fundamen-tal in explaining the accuracy of the data concerned and the interpretation of the ex-periments, but they are too lengthy or are related to too many different Sections to be included in the main text. The notation used in this appendix is consistent with that used in the present book, and not necessarily consistent with that used in the publication un-der discussion.

[1883THO]

Thomsen measured the heats of solution of NiCl2(cr) and NiCl2·6H2O(cr) in water. The final solution had a water to salt ratio of 400:1 (0.139 m). The reported heat of solution of the hexahydrate was 1.16 kcal·mol–1 (4.85 kJ·mol–1) at 291.65 K. In the experiment, dilution was to 400 moles of H2O per mole of salt. Archer [99ARC] indicated that the NBS tables [82WAG/EVA] corrected the heat of solution value to I = 0 and 298.15 K, and obtained 0.138 kJ·mol–1 using – 188 J·K–1·mol–1 for the heat capacity of reaction, and – 3.502 kJ·mol–1 for the integral heat of dilution. Based on the enthalpy of solution data of Manin and Korolev [95MAN/KOR] and the enthalpy of dilution (see Section V.4.1.3.2, Reference [59HAM] and the Appendix A discussion of [72AUF/CAR]), the enthalpy of dilution to I = 0 of the 1:400 solution of nickel chloride solution is here es-timated as – (3.0 ± 1.0) kJ·mol–1. The total uncertainty from the heat of solution meas-urements themselves and the heat capacity correction is estimated as 1.5 kJ·mol–1, and therefore the heat of solution of NiCl2·6H2O to infinite dilution in water at 298.15 K can be estimated as (0.64 ± 1.80) kJ·mol–1. This leads to f mH ο∆ (NiCl2·6H2O, cr) = − (2103.33 ± 2.01) kJ·mol–1.

The reported heat of solution of the anhydrous salt was – 19.17 kcal·mol–1 (− 80.21 kJ·mol–1), again at 291.65 K and dilution to 400 moles H2O per mole salt. If it is assumed that the heat of dilution is reasonably independent of the change in tempera-ture, and that the heat capacity changes for the hydration reaction are not strong func-tions of temperature, the heat of hydration of NiCl2 to NiCl2·6H2O at 298.15 K can be estimated as – 85.06 kJ·mol–1. Thomsen and other early experimentalists found that NiCl2(cr) dissolved only with difficulty [66GME]. This might cause the reliability of the direct enthalpy of solution measurements of NiCl2(cr) to be cast in some in doubt. How-ever, the heat of solution of the anhydrous solid is not inconsistent with the results of Manin and Korolev [95MAN/KOR].

A. Discussion of selected references

254

Based on the NEA TDB auxiliary value for f mH ο∆ (H2O, l, 298.15 K), and the selected enthalpy of formation of NiCl2(cr) (– (304.90 ± 0.11) kJ·mol–1), Thomsen’s values would lead to an enthalpy of hydration of – (2104.9 ± 1.8) kJ·mol–1, where the uncertainty (estimated here) is slightly lower than if the two heat of solution measure-ments were independent, because it is probable there were compensating errors.

[1895FRA]

In this early paper, results are reported for the conductance of nickel sulphate solutions for six concentrations between 9.8 × 10–4 M and 3.1 × 10–2 M. The work appears to have been done carefully. Re-analysis is problematic because only two of the solutions were < 0.0025 M in nickel sulphate [79LEE/WHE]. Reanalysis using the solutions at the three lowest concentrations led to K1 = (221 ± 13). The statistical deviation is un-doubtedly smaller than the true uncertainty because of the limited data and the difficul-ties in properly calibrating the apparatus as discussed by Robinson and Stokes [59ROB/STO]. In the present review, we estimate the uncertainty in K1 at 25°C to be ± 60.

[1899FUN]

Solubility measurements are reported for Ni(NO3)2·6H2O(cr) (– 21°C to 41°C) and Ni(NO3)2·3H2O(cr) (58°C to 95°C) in water. The equilibrium mixtures below 0°C may also have contained the nonahydrate. The analyses on the solids indicated that good crystals of a tetrahydrate salt could not be obtained.

[01MEU]

Meusser measured the solubility of several hydrated forms of Ni(IO3)2, and of anhy-drous Ni(IO3)2, between 0 and 90°C. Near 25°C, the least soluble compound was identi-fied by the author as the β dihydrate. Later authors [73NAS/SHI], [97PRA/LAN] have concluded that the α-dihydrate of Meusser was actually an ill-defined amorphous solid, and Meusser’s β form is what is now normally referred to as the dihydrate, Ni(IO3)2·2H2O. The solubility of this solid was measured at 8, 18, 50, 75 and 80°C. Equilibration times range from 24 h at the lowest temperature to 10 minutes at 80°C.

The reported experimental (percent weight as Ni(IO3)2) values were converted to molalities, and the apparent solubility products were corrected to I = 0. These were plotted against (T/K)–1 to obtain values for 10 ,0log sK ο at 298.15 K and an average value of r mH ο∆ (A.1) over the temperature range of the measurements.

Ni(IO3)2·2H2O Ni2+ + 2 3IO− + 2H2O(l) (A.1)

If only the values to 50°C are included, 10 ,0log sK ο ((A.1), Ni(IO3)2·2H2O, 298.15 K) = – (5.17 ± 0.11) and r mH ο∆ (A.1) = (21.4 ± 6.1) kJ·mol–1, whereas, if the points at 75 and 80°C are included, 10 ,0log sK ο ((A.1), Ni(IO3)2·2H2O, 298.15 K) = − (5.17 ± 0.08) and r mH ο∆ (A.1) = (22.5 ± 2.2) kJ·mol–1.


255

A similar analysis based on the solubility measurements (30°C to 90°C) for a yellow anhydrous nickel iodate, probably the β–Ni(IO3)2 [73NAS/SHI], leads to

10 ,0log sK ο ((A.1), Ni(IO3)2·2H2O, 298.15 K) = – (4.427 ± 0.015) and r mH ο∆ (A.1) = − (7.31 ± 0.69) kJ·mol–1. This appears to be the origin of the values listed for Ni(IO3)2 in the US NBS tables [82WAG/EVA]. In the present review, greater uncertainties (± 0.1 in

10 ,0log sK ο and 4.0 in r mH∆ ) are estimated because of the temperature extrapolation and the difficulties in establishing the purity of the solid at pseudo-equilibrium.

The tetrahydrates are difficult to purify, and convert easily to the dihydrate in the presence of liquid water [73NAS/SHI]. This agrees with observation that the meas-ured solubilities in water are greater than those of the dihydrate at all temperatures [01MEU]. However, the solubilities reported for the tetrahydrate are not used further in the present review.

[06THO]

See the entry for [1883THO].

[07WEI]

The solubility of crystalline (millerite) as well as precipitated NiS in pure water at 291.15 K was determined by means of conductivity measurements. The solubility of millerite was found to be 1.629 × 10–5 mol·dm–3 whereas that of a precipitated product amounts to 3.987 × 10–5 mol·dm–3. No experimental details are given.

[09BRU/ZAW]

Based on the enthalpy of formation of NiS(s) given by Thomsen [1883THO] ( f mH ο∆ (NiS, cr, 298.15 K) = – 72.76 kJ·mol–1) the solubility product was calculated, leading to 10 ,0log sK ο = – 24.15 at T = 298.15 K. The entropy of formation for NiS(s) was assumed to be zero, which is likewise confirmed by the thermodynamic data of the pre-sent assessment. The authors used values for the second dissociation constant of H2S(aq),

H2S(aq) 2H+ + S2–

10 2log K ο = – 23.26 and the standard potential of Ni2+, 2+Ni/NiEο = 0.219 V, which are no

longer considered reliable. Employing the currently accepted values, 10 2log K ο = – 25.99 [92GRE/FUG] and 2+Ni/Ni

Eο = 0.2372 V (this assessment), the recalculated solubility product amounts to 10 ,0log sK ο = – 27.5.

[10BOR]

The phase relations in the system Ni – S in the composition range from 0 to 31 wt% sulphur were investigated by thermal (cooling curves) and microscopic techniques. The eutectic between metallic nickel and Ni3S2(cr) was found to lie at 918 K and 21.5


256

wt% S. Moreover, it was reported that the high-temperature modification of Ni3S2(cr) melts incongruently at a temperature around 1083 K.

[11AGE/VAL]

Seemingly the solubility of NiCO3 was studied by Ageno and Valla for the first time, however, in their preliminary communication the authors described no attempt to char-acterise the chemical and physical state of the samples investigated. Regardless of this uncertainty the result was ascribed to anhydrous NiCO3 and accepted by thermody-namic compilations for calculation of f mGο∆ (NiCO3, cr, 298.15 K) [35KEL/AND], [52LAT].

Ageno and Valla equilibrated their solid phase with water under a carbon diox-ide containing atmosphere in a closed vessel at 25°C. The Ni(II) concentration and the partial pressure of CO2 were determined experimentally. Column 2 of their table on carbonate solubilities is labelled incorrectly. Under the conditions prevailing in their study an elementary stoichiometric consideration predicts:

2[Ni2+] ≈ [ 3HCO− ].

Consequently it has been assumed that the numerical values listed in the solu-bility table under NiCO3/2 have to be ascribed to 2 [Ni(II)] instead.

[14THI/GES]

Nickel sulphide was prepared by adding Na2S or (NH4)2S to aqueous solutions of nickel chloride (typically 0.04 mol·dm–3) at ambient (around 298 K) and elevated tempera-tures, respectively. In addition, NiS was precipitated by reaction of H2S with NiCl2 solu-tions containing sodium acetate (typically 0.1 mol·dm–3) at ambient as well as elevated temperatures. In some cases the precipitated product was aged in the mother liquor for 10 weeks.

The solubility of NiS at room temperature was investigated by treating both freshly prepared and aged solid phases with HCl solutions (2 mol·dm–3) saturated with H2S. The amount of dissolved nickel(II) was determined by means of electrolysis. Three different fractions of the solid sulphide could be discerned with respect to their solubil-ity behaviour. Fraction I dissolved rapidly in hydrochloric acid at room temperature. The solubility was estimated to be 0.1 mol·dm–3 at pH = 2. This high value indicates that fraction I corresponds to an amorphous product. Fraction II showed a solubility in H2S saturated 2 M HCl of approximately 1 mmol·dm–3 at room temperature. As the re-action time was in the range from 0.5 to 42 hours, equilibrium was certainly not at-tained. Finally, the solubility of fraction III in H2S saturated 2 M HCl was reported to be approximately 0.02 mmol·dm–3 at room temperature. At temperatures around the boiling point of the hydrochloric acid solution the solubility amounted to 0.22 mmol·dm–3 with the reaction time ranging from 0.5 to 2 hours. Again, this time is not expected to be suf-ficient to attain equilibrium between the aqueous and the solid phase. According to


257

Dönges [47DON] the hexagonal NiAs-type modification of nickel sulphide is obtained when the precipitation is performed slowly by using H2S. Hence, it can be concluded that fractions II and III correspond to crystalline NiS of different particle sizes depend-ing on the conditions of precipitation and ageing. The different nickel(II) concentrations found by dissolving these fractions may be caused by the variation of the dissolution rate with the particle size. However, thermodynamic solubility constants cannot be de-rived from these experimental studies. As none of these nickel sulphide phases was characterised by X-ray diffraction analysis the original notation I, II, and III-NiS has been retained in this review.

[16DER/YNG]

Using a static method, water vapour pressures were determined as a function of tem-perature above mixtures of the tetrahydrate and hexahydrate (1.8°C to 36.25°C) and the dihydrate and tetrahydrate (25.15°C to 79.06°C). The work established the existence of the tetrahydrate as a separate phase, and the temperature of the hexahydrate/tetrahydrate conversion at equilibrium with water vapour was reported as 36.25°C. The reported water vapour pressure above the tetrahydrate/hexahydrate mixtures at 25°C was (1.40 ± 0.05) kPa, where the uncertainty is estimated in the present review.

[18ALM]

Almqvist suspended freshly precipitated nickel hydroxide in water for 24h, filtered off the residue, and determined the Ni(II) content of the clear solution. The value of the Ni(OH)2 solubility thus obtained ([Ni(II)] = 1.37 × 10–4 mol·dm–3) implies that, owing to the mass and charge balance of Equation (A.2), either the saturated solution is strongly alkaline or the uncharged hydroxo complex Ni(OH)2 is rather stable. As neither phenomenon has been observed, Almqvist’s solubility is too high, as clearly docu-mented in Figure 5 of [97MAT/RAI]. Therefore, it has not been taken into account in the calculation of a value of 10 ,0log sK . Probably the high solubility observed by Alm-qvist was due to colloidal Ni(OH)2 particles which passed through the filter.

[Ni(II)] = [Ni2+] + [NiOH+] + [Ni(OH)2]

[OH–] = 2 [Ni2+] + [NiOH+] (A.2)

[21BER/CRU]

Berger and Crut investigated the equilibrium of the reaction,

NiCl2(cr)+H2(g) Ni(cr) + 2 HCl(g) (A.3)

using a static method. When the experimental data depicted in Figure V-16 and listed in Table A-1 were compared with those of [53BUS/GIA], [53SAN], [54SHC/TOL] they did not seem to be of adequate quality for use in a third law analysis.


258

Table A-1: Equilibrium constants of Reaction (A.3) [21BER/CRU], [24CRU].

T / K log10(Kp / atm) log10(Kp / bar) – RT ln(Kp / bar) / kJ·mol–1

583 – 2.214 – 2.2083 24.6475 613 – 1.627 – 1.6213 19.0269 668 – 0.784 – 0.7783 9.9532 718 – 0.0556 – 0.0499 0.6857

[23VIL/BEN]

The solubility of nickel sulphate was determined by the floating equilibrium method. The following results were obtained at 25°C: 40.469 g NiSO4 per 100 g H2O and 40.720 g NiSO4 per 100 g H2O (floating equilibrium), 40.417 g NiSO4 per 100 g H2O (gravim-etry).

[24CRU]

Crut described the method applied and the experiments carried out previously [21BER/CRU] in more detail. The reported constants are the same as those given in [21BER/CRU]. Only one chloride data set at 668 K is not identical. This was consid-ered to be a misprint, as the other Kp values were exactly the same in [24CRU] as in [21BER/CRU]. The reported measurements were primarily for reduction of the chlo-ride. Only very limited measurements for the bromide system were carried out (at 718 and 848 K). Also, see the discussion of [24CRU2].

[24CRU2]

Crut [24CRU2] reported the heat of solution of anhydrous and hydrated NiBr2 as – 19.9 kcal·mol–1 (– 83.26 kJ·mol–1) and 0 kcal·mol–1, respectively. The paper also reports a somewhat different derived value for the heat of hydration of the anhydrous salt as − 19 kcal·mol–1 (– 79.50 kJ·mol–1). This value is consistent with the heats of solution of anhydrous and hydrated NiBr2, – 18.9 kcal·mol–1 (– 79.08 kJ·mol–1) and 0 kcal·mol–1, probably from the same experiments, reported in an earlier paper by the same author [24CRU]. The extent of hydration of the hydrated salt is not reported, nor is the salt purity, the final solution concentrations nor the measurement temperature. Because nickel bromide prepared at high temperatures dissolved slowly in water at room tem-perature, a sample of anhydrous NiBr2(cr) was prepared in a manner to enhance its rate of dissolution in water. The dissolution rate difference may reflect a structural differ-ence that could have been affected the value determined in the heat of solution meas-urements.

Bichowsky and Rossini [36BIC/ROS], in the forerunner to the US NBS Tables, estimated f mH ο∆ (NiBr2·3H2O, cr) using the value – 18.8 kcal·mol–1 (– 78.62 kJ·mol–1) for the enthalpy of hydration of NiBr2(cr) to NiBr2·3H2O(cr) from Crut’s 1923


259

dissertation (unavailable to the current reviewers). This value corresponds approxi-mately to the values reported in [24CRU]. The authors appear to have based their identi-fication of the solid as NiBr2·3H2O on work reported in the thesis or elsewhere. There is a similar, but smaller discrepancy in the values for the cobalt bromides from the same two papers [24CRU], [24CRU2]. The analysis and value for f mH ο∆ (NiBr2·3H2O, cr) from Bichowsky and Rossini [36BIC/ROS] is the origin of the value for the same solid (at 298.15 K) in the 1952 version of the US NBS tables [52ROS/WAG] (and probably of the slightly different value in the 1982 tables [82WAG/EVA]).

[24MOS/BEH]

The solubility of NiS at 293.15 K in H2S saturated 1 M H2SO4 was reported to be 1.04 × 10–4 mol·dm–3. After an equilibration time of 10 hours the dissolved amount of nickel(II) was determined gravimetrically. The solid phase was precipitated at elevated temperatures (333 – 343 K) by adding H2S to an aqueous solution of NiSO4·7H2O. No characterisation of the solid product was performed. The reaction time of 10 hours seems to be too short for equilibrium between the solid phase and the aqueous solution to be attained.

[24TAN]

The solubilities of the crystalline nickel sulphate hydrates were determined to calculate the respective temperature coefficients. For 25°C 40.55 and 40.26 g NiSO4 per 100 g H2O were given.

[25BRI]

Britton titrated a nickel chloride solution with sodium hydroxide and employed a hy-drogen electrode to follow the pH change. As only 1.66 moles of NaOH were needed to complete the reaction it is quite clear that a basic chloride had at least been co-precipitated with the pure nickel hydroxide. Thus, it is doubtful whether Britton’s ex-perimental data refer to Ni(OH)2 at all. In addition Jena and Prasad have previously re-ported that Britton miscalculated the solubility product [56JEN/PRA]. Re-evaluation of the buffer region of Britton’s Figure 2 has shown that 10 ,0log sK ο = – (16.3 ± 0.3) is con-sistent with the experimental data, and has been used instead of the value originally given (– 18.06) for the respective entry in Figure V-11 and Table V-6.

[25JEL/ZAK]

The partial pressure of sulphur in equilibrium with nickel sulphide was determined at 788.15 and 903 K. The ratio p(H2S)/p(H2) in the gas phase was measured by means of chemical analysis, resulting in sulphur partial pressures by taking into account the well-known equilibrium:

H2S(g) H2(g) + ½ S2(g).


260

The solid phase, NiS, was prepared by reaction of aqueous solutions of NiCl2 with (NH4)2S. The product was characterised by chemical analysis, but no X-ray diffraction analysis was performed. As the results (log10[p(S2)/bar] = – 9.0 for the proposed equilibrium between NiS and Ni7S6 at 788.15 K and log10[p(S2)/bar] = – 10.0 for the proposed equilibrium between Ni and β1−Ni3S2 at 903 K) do not coincide with any other literature data, this investigation is not considered further in the present assessment.

[25WIJ]

In the course of a study of nickel ammine complexes, in two experiments turbidity oc-curred, which was interpreted as an indication of formation of Ni(OH)2. When the solu-bility product ( 10 ,0log sK = − 13.8, 25°C) is extrapolated to I = 0 by means of the Davies equation the following constants were obtained ( 10 ,0log sK ο = – 14.64, 10 ,0

*log sK ο = 13.36) [62DAV]. These values are in line with others which refer to freshly precipitated probably amorphous Ni(OH)2, see Table V-6 [38OKA], [54FEI/HAR].

[26JEL/ULO]

The reduction equilibrium of nickel(II) chloride was studied by a flow method. The results were probably flawed by vaporisation of NiCl2, thus they were not subjected to a third law analysis. The experimental data depicted in Figure V-16 are listed in Table A-2.

Table A-2: Equilibrium constants of NiCl2(cr) + H2(g) Ni(cr) + 2 HCl(g).


573 – 2.574 – 2.5683 28.1739 613 – 1.398 – 1.3923 16.6060 723 – 0.136 – 0.1303 1.8033

[26JEL/ULO2]

The flow measurements for hydrogen reduction of NiBr2(cr) (683 K to 833 K) and NiI2(cr) (783 K to 983 K) suffer from problems similar to those discussed for the meas-urements on NiCl2(cr) [26JEL/ULO]. Furthermore, above 700 K, NiI2 has been reported to be subject to thermal decomposition [74WOR/COW]. The results of Jellinek and Uloth are not used in the present assessment.

[27JON/WIL]

Powdered NiS was heated with S at 170°C for 30 to 40 h. The empirical formula of the product was NiS2.18. The X-ray powder pattern was quite similar to the pattern for FeS2.


261

A cubic lattice constant a0 = 5.74 Å was determined.

[28JEL/RUD]

The equilibrium of the reaction:

NiF2(cr) + H2(g) Ni(cr) + 2HF(g) (A.4)

was studied from 573 K to 773 K. The measured equilibrium constants varied with flow rate. Therefore, although the derived value of f mH ο∆ (NiF2, cr, 298.15 K) is consistent with those from other types of experiments [67RUD/DEV], the values of Jellinek and Rudat are not used in the assessment of f mH ο∆ (NiF2, cr, 298.15 K) in the present re-view.

[28MUR]

Murata reviewed older measurements with – 0.597 V ≤ E°(NiSO4(aq) | Ni) ≤ – 0.183 V and concluded that the state of the nickel electrode was responsible for this conspicuous disagreement [1894NEU], [04EUL], [04MUT/FRA], [04SIE], [09SCH], [09SCH2], [10SCH], [18SMI/LOB]. In his own work he used powdered nickel which was reduced with hydrogen directly in the following cell

Ni | NiSO4 | KCl(sat.) | KCl(0.1 M), Hg2Cl2 | Hg.

Murata’s data have been recalculated with the standard potential of the Hg2Cl2 | Hg electrode ((0.2681 ± 0.0026) V, [89COX/WAG]) and the chloride activity of a 0.1 M KCl solution (aKCl = 0.077) [59ROB/STO]. The activity of Ni2+ in the solu-tion was calculated using the SIT model. The corrected values are E°(Ni2+ | Ni) = − (0.246 ± 0.003) V at 25°C and – (0.245 ± 0.005) V at 18°C. As the measurements were carried out under hydrogen, which according to other authors interferes, and no correction for the liquid junction was applied, Haring and Vanden Bosche’s [29HAR/BOS] result has been preferred to Murata’s.

[29HAR/BOS]

Haring and Vanden Bosche reviewed the literature on the standard potential of Ni2+ | Ni, discussed Murata’s work and included the following papers into the discus-sion [00KUS], [01PFA], [05COF/FOR], [08THO/SAG], [26GLA]. They pointed out that in many of the numerous earlier studies, which led to results with E°(Ni2+ | Ni) ranging from – 0.621 to – 0.178 V, appropriate corrections for activity coefficients and liquid junction potentials were lacking. Moreover, the absence of interfering gases such as oxygen or hydrogen and a constant temperature were not always ensured.

These authors carried out most careful measurements using the cell:

Ni | NiSO4(m) | Hg2SO4 | Hg

which involves no liquid junction. Finely divided, strain-free, electroplated nickel was used and precautions were taken to exclude oxygen and hydrogen completely. The


262

measured potentials were constant within 1-5 mV for 18 days. The cell behaved com-pletely reversibly. The reversibility of the electrode was demonstrated by the Nernst behaviour of the measured potentials with the activity of NiSO4, see also Figure V-2. Moreover, after anodic and cathodic polarisation and short circuiting, which produced a temporary disturbance, the equilibrium value was quickly regained.

The results are listed in Table A-3. It was necessary to recalculate the standard potential (E°(Ni2+| Ni, 298.15 K) = − (0.231 ± 0.002) V [29HAR/BOS]), because the SIT model leads to slightly different activity coefficients and the standard potential of the Hg|Hg2SO4 electrode had to be changed to the value consistent with the NEA-TDB auxiliary data. Using the function ln γ± = F (m (NiSO4)) given in the caption of Figure A-5 (page 295 under the discussion of [59ROB/STO]) for the recalculation of E°(Ni2+| Ni, 298.15 K) results in a value which agrees within 0.1 mV with the one derived using the SIT. f mGο∆ (Ni2+) selected in this review has been based on the latter value.

An attempt by these authors to measure the standard potential of the nickel electrode with the similar cell:

Ni | NiCl2(0.05 mol·kg–1) | Hg2Cl2 | Hg

was unsuccessful. Obviously nickel corroded in the chloride medium.

Table A-3: Electrode potential of the cell: Ni | NiSO4 (m) | Hg2SO4 | Hg [29HAR/BOS]

m (NiSO4)

(mol·kg–1)

E/ V

E° / V

[29HAR/BOS]

E° / V

SIT

E° / V

[59ROB/STO]

0.0504 0.967 0.852 0.84848 0.84794

0.0504 0.967 0.851 0.84848 0.84794

0.0504 0.969 0.853 0.85048 0.84994

0.0506 0.968 0.852 0.84953 0.84900

0.0506 0.967 0.851 0.84853 0.84800

0.0507 0.971 0.855 0.85256 0.85204

0.0508 0.968 0.852 0.84959 0.84907

0.0509 0.967 0.851 0.84862 0.84811

0.0511 0.968 0.852 0.84968 0.84917

0.0511 0.967 0.851 0.84868 0.84817

0.0512 0.967 0.851 0.84871 0.84821

0.0515 0.970 0.854 0.85180 0.85131



263

Table A-3 (continued)

m (NiSO4)

(mol·kg–1)

E/ V

E° / V

[29HAR/BOS]

E° / V

SIT

E° / V

[59ROB/STO]

0.0515 0.968 0.852 0.84980 0.84931

0.0516 0.968 0.852 0.84982 0.84934

0.0525 0.971 0.854 0.85308 0.85263

0.1017 0.958 0.852 0.84967 0.85007

0.1018 0.958 0.852 0.84969 0.85008

0.1020 0.958 0.852 0.84971 0.85011

0.1024 0.958 0.852 0.84977 0.85017

0.1052 0.958 0.852 0.85016 0.85056

0.1503 0.953 0.852 0.85023 0.85046

0.1504 0.954 0.853 0.85124 0.85147

0.1519 0.953 0.852 0.85038 0.85060

0.1543 0.953 0.852 0.85060 0.85080

0.1543 0.952 0.851 0.84960 0.84980

0.1617 0.954 0.852 0.85226 0.85241

[30MUL/LUB]

Müller and Luber determined a single value of the solubility of nickel carbonate hy-drate, which they believed to be NiCO3·6H2O, at p(CO2) = 50 atm, without giving the temperature of equilibration explicitly [30MUL/LUB]. When it is tentatively assumed that the experiment was carried out at 25°C, Müller and Luber’s result agrees fairly well with results from [11AGE/VAL] and [2001GAM/PRE], [2002WAL/PRE] indicating that the same nickel carbonate phases were used in these studies.

[31KOL]

The solubility product of NiS was calculated by taking into account the solubility of nickel(II) sulphide in 1 M H2SO4 (T = 293.15 K), determined by Moser and Behr [24MOS/BEH], leading to 10 ,0log sK ο = – 26.96. The author used a value for the second dissociation constant of H2S(aq),

H2S(aq) 2H+ + S2–,

10 2log K ο = – 21.97 which is no longer considered to be correct. Employing the currently accepted value, 10 2log K ο = – 25.99 [92GRE/FUG], the solubility product can be recalcu-lated, resulting in 10 ,0log sK ο = − 30.98.


264

[32MAS]

Potentiometric measurements of the nickel(II) – cyanide system are described. The ad-dition of Ni(CN)2, NiSO4 or Ni(NO3)2 to sodium cyanide solution resulted in the forma-tion of a single nickel(II) containing species, 2

4Ni(CN) − . Its formation constant at 25°C and in 0.78 M NaCN solution was found to be 10 4log b = (11.83 ± 0.17). The author used nickel metal as the working electrode. Since the electrode reactions involving nickel(II) are not reversible [50HUM/KOL], the reported formation constant is not reli-able.

[32MON/DAV]

This paper provides a calculated value of 250 for K1 for nickel sulphate complexation based on the earlier measurements of Franke [1895FRA] (at the two lowest molarities for which results were reported, 9.768 × 104 and 19.53 × 104 M) at 25°C. The value of Λ

° was based on estimates of the molar ionic conductivities of sulphate (158

S·cm2·mol−1), and for Ni2+ (112.2 S·cm2·mol–1). There are no new experimental data.

[34BRI/OSS]

This paper reports the results of a dialysis study using aqueous solutions 1 M in ammo-nium sulphate and 0.05 M in various other metal sulphates. The sparse results for the nickel system are interpreted in terms of formation of 4

2 4 4Ni (SO ) − . There are insuffi-cient details to derive useful complexation constants from the data. Also see [41KIS/ACS].

[34COL2]

Colombier measured the potential of NiSO4(0.5 M) | Ni against a calomel reference electrode at 20°C. When the electrolyte was appropriately degassed and kept air-free the same equilibrium potentials were found for massive or powdered nickel, the latter hav-ing been electrolytically deposited or prepared by hydrogen reduction of NiO. As the author presented no details about the experimental data, the reference potential selected, how the liquid junction potential and the activity coefficients were accounted for, the final result, E°(Ni2+ | Ni, 293.15 K) = – (0.227 ± 0.002) V, cannot be recalculated and was not being considered any further.

[34SIE/SCH]

Sieverts and Schreiner [34SIE/SCH] carried out solubility measurements from – 27.8°C to 119.8°C. They found Ni(NO3)2·9H2O below – 11°C and Ni(NO3)2·6H2O to 50°C. They also reported that both the tetrahydrate (55 to 80°C) and dihydrate (85 to 120°C) were stable phases in the Ni(NO3)2 – H2O system, and could be formed at equilibrium at 25°C in the Ni(NO3)2 – H2O – HNO3 system at high nitric acid mole fractions. The so-lution compositions at saturation do not differ markedly from those of Funk [1899FUN] except at the lowest temperatures.


265

[35KEE/CLA]

This paper reports the measurements of the specific heat capacity of nickel in the tem-perature range from 1.1 to 19.0 K. The hypothesis was advanced that the deviation from Debye’s T 3 law is connected with the energy of electron conduction.

[36BIL/VOI]

The decomposition pressure of NiS2(cr) was measured as a function of temperature by means of a tensimetric method. The solid phase, NiS2(cr), was prepared by solid state reaction between NiS(cr) and sulphur. The composition of the product was checked by chemical analysis (no X-ray diffraction analysis was performed).

Furthermore, the phase transformation of millerite into its high-temperature modification was studied by an early DTA-like method. The transition temperature was reported to be 669 K and the heat of transition was estimated to be in the range between 2.1 and 2.9 kJ·mol–1. Millerite was precipitated at room temperature by bubbling H2S gas through an ammonia-containing aqueous solution of NiCl2. The solid compound was characterised by chemical analysis.

[36BRO/WIL]

The mean heat capacities from – 80 to 120°C have been determined for nickel and a few other metals. The errors were estimated to amount to ± 0.1%. A discussion of the theo-retical basis for the equations describing the temperature dependence of the heat capac-ity is given together with some calculations of heat capacities using data which involve no measurements of heat e.g., the expansion coefficient, compressibility etc.

[36JOB]

This is a spectrophotometric study on (among others) nickel(II) bromide complexes formed in aqueous HBr solution. The author identified two bromo complexes, NiBr2(aq) and 2

4NiBr − . The former species is reported to be dominant in 6 – 7 M, the latter above 10 M HBr. Since the ionic strength could not be maintained constant, the activity of bromide, instead of its concentration, was taken into account in the mass action expres-sions. These “semi-thermodynamic” constants include the ratios

( 2)NiBr NiBr/n n−

g g (where n = 2 or 4), which are assumed to be nearly constant with increasing HBr concentra-tions. The reported constants are 10 2log ο′b = – 3.24 and 10 4log ο′b = – 8.12. The graphi-cal data presented in [36JOB] were later re-evaluated in [72RET/HUM], and were found consistent with the formation of NiBr+ and NiBr2(aq) complexes, if the activity of water is also taken into account in the mass action expressions. In [72RET/HUM] the recalculated value of 10 2log ο′b is also reported ( 10 2log ο′b = – 3.7). Since negatively charged halogeno complexes of nickel(II) have not been identified in aqueous solutions neither with fluoride nor with chloride, only the latter value is considered in this review.


266

[36LOT/FEI]

Lotmar and Feitknecht investigated variations in ionic distances of bivalent basic metal halogenides with layered structures of the CdI2 type. In this context they needed the unit cell dimensions of nickel hydroxide as well as cobalt hydroxide.

[36TRA/SCH]

Trapeznikova et al. studied the heat capacity of NiCl2 at temperatures between 13 and 130 K. They used a glass calorimeter and a platinum thermometer heater. The discrep-ancy between their results and those of Busey and Giauque probably indicates that their glass calorimeter and their thermometer did not attain thermal equilibrium [52BUS/GIA].

[37BEL]

Measurements are reported for dissociation pressures of nickel chloride hexadeuterate to tetradeuterate. The transition temperatures for dehydration of the hexahydrate and the hexadeuterate at 1 atm are reported as 36.25 and 35.9°C, respectively. It was noted that the solvent pressure ratio over the Ni–Cl solvates (

2 2D O H O:p p ) rises with increasing tem-perature from 25 to 35°C. There are no experimental numbers reported for the hexahy-drate/tetrahydrate equilibrium.

[37DOM]


NiF2(cr) + H2O(g) NiO(cr) + 2HF(g) (A.5)

was studied from 773 K to 973 K. The measured equilibrium constants showed only a small variation with flow (at higher flow rates). In the present review, values of

r mH ο∆ ((A.5), 298.15 K) are calculated using the reported equilibrium constants, se-lected values of mS ο and ,mpC (T) from Tables III-1 and III-3 for NiF2(cr) and NiO(cr), and the tabulated values for H2O(g) and HF(g) in the CODATA compilation [89COX/WAG]. Then, using f mH ο∆ values from Tables III-1 and IV-1 for NiO(cr), H2O(g), and HF(g), an average value, f mH ο∆ (NiF2, cr, 298.15 K) = – (666.13 ± 0.64) kJ·mol–1, is calculated. The drift in the calculated value depends only slightly on the measurement temperature (< 1 kJ·mol–1 for the 200 K range). The value calculated in the present review is 18 kJ·mol–1 more negative than the value recalculated by Rudzitis et al. from the same data, though the values of r mH∆ (298 K) are in good agreement. Only part of the discrepancy can be explained as resulting from use of dif-ferent auxiliary data (primarily the value for f mH ο∆ (HF, g, 298.15 K)).

[37SAN]

Sano reported the values for the vapour pressure of water over (initial) Ni(NO3)2·6H2O from 300 K to 326 K using a Jackson spring manometer. The product of the dehydration


267

reaction was assumed to be Ni(NO3)2·3H2O, and the plot of – log10 2H Op against 1/T is linear, but no estimate of the uncertainties is given. In this paper, the thermodynamic parameters for the dehydration reaction were calculated incorrectly from the data (the equilibrium constant depends on

2

3H Op , not

2H Op as assumed by the author). No evi-dence was presented that the dehydration product was the trihydrate rather than the tet-rahydrate, the dihydrate or a non-equilibrium mixture of hydrated salts.

[37SAN2]

See [53SAN].

[38OKA]

Three titrations of pure as well as NaCl- or Na2SO4-containing nickel nitrate solutions with sodium hydroxide were carried out at 25°C. From the pH measured at each point taken on the titration curve (0.07 ≤ n(NaOH) / n(Ni(NO3)2) ≤ 1.62) the solubility prod-uct was calculated and extrapolated to zero ionic strength. A mean value of 10 ,0log sK ο = – 14.50 was obtained and has been taken as a basis for the value listed in Figure V-11 and Table V-6, respectively.

[38SMU]

Smurov precipitated nickel carbonate by mixing aqueous sodium carbonate and nickel chloride solutions. The solid phase obtained was washed free of chloride and used for the solubility study, but no attempt was made to characterise its stoichiometry and/or its crystal structure. The precipitate was equilibrated with water in a temperature range between 278.15 and 353.15 K. For each temperature the partial pressure of carbon diox-ide was varied from 0.0005 atm to

2 2CO H Op p+ = 1 atm. Smurov’s data have been ex-trapolated to zero ionic strength under the assumption that the solid phase had been hel-lyerite. As shown in Figure A-1 the solubility constants thus derived disagree not only with the values of hellyerite but also with those of gaspéite. The temperature depend-ence is more pronounced than is usually observed with neutral carbonates. This proba-bly indicates that a basic nickel carbonate was investigated instead. However, without stoichiometric and structural evidence, no thermodynamic quantities can be assigned to it.


268

Figure A-1: Solubility of precipitated nickel carbonate. ■ Experimental data [38SMU]; solid curve: experimental data fitted to 10 , ,0

*log p sK = A + B/T + C·lnT; dotted curve: solubility constant of hellyerite, NiCO3·5.5H2O; dash curve: solubility constant of gaspéite, NiCO3.

[38SYK/WIL]

The specific heat capacity – temperature dependence for four different types of nickel have been determined and the effects of the method of manufacture, the heat treatment and the chemical composition have been ascertained. The results, together with those obtained by previous investigators were reviewed and an attempt was made to evaluate the most probable ,mpCο versus T curve for nickel in the temperature range 100 to 600°C.

[39ROH]

After an equilibration period of four days at 25°C the solubility of NiSO4·7H2O was found to be 41.2 g NiSO4 per 100 g H2O.

The 6 values taken from [23VIL/BEN], [24TAN], [39ROH], and cited in the respective Appendix A entries, were used to calculate mean and standard deviation (2σ) of the NiSO4·7H2O solubility in water (see Section V.2.1.3.1).

260 280 300 320 340 360

6.5

7.0

7.5

8.0

8.5

9.0

9.5

10.0

10.5

11.0

11.5

log 10

* K0 p,s,0

T / K


269

[39SCH/FOR]

The high-temperature form of nickel(II) sulphide, NiS, was treated with hydrogen gas at three different temperatures, viz. 673, 773 and 873 K. The H2S(g) content in the gas phase was measured at constant temperature as a function of the sulphur content of the solid phases. Equilibrium compositions of the gas phase for univariant equilibria, in-volving NiS, Ni3S2 and Ni6S5 (according to Stølen et al. [94STO/FJE] the composition of this phase is rather Ni7S6 (high-temperature modification) and Ni9S8 (low-temperature modification), respectively), are given. No experimental details are re-ported.

[39SIM/KNA]

Results of study of the decomposition of nickel sulphate heptahydrate through nickel sulphate hexahydrate and nickel sulphate tetrahydrate at 0.015 bar (11 torr) are pre-sented. X-ray diffraction patterns are presented for some of the different solids. Detailed measurement results were not tabulated, but the graphical presentation and the reported thermodynamic quantities for the dehydration reactions are in reasonable agreement with results reported in the later study of Kohler and Zäske [64KOH/ZAS].

[40BEL]

This paper contains measurements of the water vapour pressures above samples of NiSO4·7H2O (at 20, 25 and 30°C) and NiSO4·7D2O (at 20 and 25°C). The values for the NiSO4·7H2O are very similar to those obtained by Schumb [23SCH], Bonnell and Bur-ridge [35BON/BUR] and [66STO/ARC]. The vapour pressures above the deuterate were found to be about 80% of the vapour pressures above the ordinary hydrate. The enthalpy of removal of D2O from NiSO4·7D2O was found to be approximately 11 kJ·mol–1 less favourable than dehydration of NiSO4·7H2O.

Equations for dissociation pressures of nickel chloride hexahydrate to tetrahy-drate and the nickel chloride hexadeuterate to tetradeuterate are presented. The equation for the deuterate reaction is based on the experimental vapour pressures reported in [37BEL]. The equation for the hexahydrate/tetrahydrate equilibrium matches the ex-perimental points from Derby and Yngve [16DER/YNG] so well that it raises the ques-tion of whether the reported equation is from a fit to the earlier data [16DER/YNG], rather than based on new experiments. The transition temperature reported in [37BEL] matches that in [16DER/YNG] exactly (within 0.01 K).

[41KIS/ACS]

This paper discusses problems with the dialysis method as used by Brintzinger and co-workers (cf. [34BRI/OSS]).


270

[41STO/GIA]

Results were reported for careful heat capacity and magnetic susceptibility measure-ments carried out between 1 and 15 K on a sample of “NiSO4·7H2O”. The results indi-cated a maximum in the heat capacity for 1.8 K. In a later paper [66STO/ARC], the re-sults were reinterpreted and combined with results for higher temperatures to determine the third-law entropy of the solid at 298.15 K.

[42FRI/WEI]

Mixtures of metallic nickel and nickel oxide were treated with CO/CO2 gas mixtures at temperatures between 1044 and 1289 K. The composition of the gas phase after equili-bration with the solid phases was determined by chemical analysis, yielding equilibrium constants for the reaction:

Ni(cr) + CO2(g) NiO(cr) + CO(g)

as a function of temperature.

[43NAS]

Näsänen titrated an excess of potassium hydroxide drop-wise with NiCl2 solutions and determined the equivalence point potentiometrically with a hydrogen electrode. The solubility product was calculated from the minimum of the buffer capacity. A recalcula-tion taking Näsänen’s experimental conditions into account has led to 10 ,0log sK ο = − 15.09 and 10 ,0

*log sK ο = 12.92, respectively. The latter value has been listed in Fig-ure V-11 and Table V-6. It agrees reasonably well with 10 ,0log sK ο = – 15.21 ( 10 ,0

*log sK ο = 12.79) given in the original paper. These values are at best valid for amor-phous Ni(OH)2 immediately after its precipitation, but in the course of titration some Ni(OH)2–xClx may have formed as well and thus influenced the result.

[45KER]

Two sulphides have been identified as lying close to the Ni and Co end members of the Co-, Ni-, Fe-disulphide system. In the light of this study the pyrite group forms an isostructural series with pyrite, FeS2, vaesite NiS2 and cattierite CoS2 as end members, while bravoite (Co, Ni, Fe)S2 is intermediate.

[47DON]

Amorphous NiS was prepared by reaction of nickel chloride solutions with Na2S and (NH4)2S. If the nickel sulphide is precipitated by bubbling H2S(g) through a 1 M NiCl2 solution containing sodium acetate, crystalline NiS is obtained which shows the hex-agonal NiAs-type structure according to X-ray diffraction analysis. After stirring the solid phase in 1 M HCl for 10 minutes the nickel content of the aqueous phase was ana-lysed spectrophotometrically. The solubility of NiS was reported to be in the range from


271

0.006 to 0.142 mol·dm–3. As the reaction time of 10 minutes is by far too short to attain equilibrium, no further evaluation of these data has been performed.

[47PEA]

Peacock investigated two specimens of heazlewoodite from Heazlewood, Tasmania, by chemical and X-ray diffraction analyses. The empirical formula of one of the heazle-woodite samples was close to Ni3.04Fe0.02S2. The X-ray diffraction patterns of the natural and the artificial Ni3S2 phases agreed very well with each other and also with those re-ported by Westgren [38WES]. The cell dimensions given by the latter author, however, were incompatible with the respective d spacings. Consequently, the cell dimensions given by Peacock and not those of Westgren were accepted for JCPDS-ICDD card No.8-126.

[49GAY/GAR]

Nickel hydroxide was prepared by adding an equal volume of 0.02 M NaOH to a hot solution of 0.0034 M Ni(NO3)2.The precipitate was washed free of alkali with water. Two sets of samples were prepared for each concentration of HCl (0.0025 – 0.1 m) or NaOH (0.0016 – 15 m). One was equilibrated for 5 to 7 days at 35°C and then for 7 days at 25°C, the other was equilibrated for 5 to 7 days at 25°C only. In this way the authors intended to approach equilibrium from supersaturation as well as undersatura-tion. Inspection of Figure V-11 shows that in the acid range an aqueous suspension of β–Ni(OH)2 saturated at higher temperatures is undersaturated at lower temperatures. As a rule of thumb a temperature decrease of 10 K slows down chemical reactions, such as the dissolution reaction according to Equation (V.40), by at least a factor of two. Long ago Schindler pointed out that the method applied by Gayer and Garrett may not be ef-fective with slowly dissolving metal hydroxides [63SCH]. Thus, it is not quite sure that in their work the investigated solutions were indeed saturated. When it is assumed that Gayer and Garrett's result is valid at 308.15 K it agrees very well with the experimental data of Gamsjäger et al. [2002GAM/WAL].

Jena and Prasad [56JEN/PRA] as well as Mattigod et al. [97MAT/RAI] pre-sume that there were systematic errors in pH measurements by Gayer and Garrett, but this cannot be proven just by shifting the reported pH values until their own results are confirmed. Moreover, Gayer and Garrett determined the concentration of nickel(II) em-ploying the dimethylglyoxime method; in addition the data in alkaline solutions were checked by radioactive tracer experiments. They calibrated the glass electrode for pH measurements in a then conventional way. Thus, errors between 0.5 and 0.8 pH units as suggested by Mattigod et al. are improbable. The solubility of β–Ni(OH)2 certainly de-pends on the particle size, but whereas it is quite certain that Mattigod et al. equilibrated their solutions with the finest portion of their samples, no such information is available of Jena and Prasad's or Gayer and Garrett's experiments.


272

The re-evaluation of Gayer and Garrett's data resulted in 10 ,0*log sK = 10.78 (as

compared to 10.8 given) only when the second column of their Table II was considered to contain the molality, +H

m , and not the activity, +Ha , as stated in the original paper.

The ionic strength was not controlled, instead activity coefficients of Harned and Owen were used [58HAR/OWE]. The nickel(II) concentrations, over the pH range studied, were sufficiently low to neglect the formation of 4+

4 4Ni (OH) . The authors re-ported the values, 10 ,1,2

*log sK ≈ – 7 and 10 ,1,3log sK = – (4.2 ± 0.6), for the reactions:

Ni(OH)2 (cr) Ni(OH)2 (aq) (A.6)

Ni(OH)2(cr) + OH– 3Ni(OH)− (A.7)

The authors mentioned, that the dissolution of nickel hydroxide in alkaline so-lutions up to 15 M NaOH was so slight that only estimates of its solubility could be made. When the reported experimental data were re-evaluated it became evident that, at best, only one constant, ,1,3sK , can be determined. The solubility of Ni(OH)2(cr) in-creases with increasing NaOH molality, but the solubility curve flattens out at [OH–] > 1 m. This may be due to ion-ion interactions at the high ionic strengths, but there is no evidence whatever for formation of 2

4Ni(OH) − . Equation (A.7) can be fit reasonably to the data at lower NaOH molalities. When the SIT model is applied:

log10 [Ni(II)]tot = 10 ,1,3log sK – ∆ε·[OH–]ini + log10[OH–]ini

The logarithmic solubility curve, as obtained by regression analysis, turns sharply downwards at high hydroxide molalities. This demonstrates that the SIT ap-proach is no longer valid, because the mass law requires an increasing solubility with an increasing NaOH molality.

The regression analysis results in 10 ,1,3log sK = – (4.40 ± 0.12) and ∆ε = (0.66 ± 0.17) kg·mol–1. Where ∆ε = ε(Na+, 3Ni(OH)− ) – ε(Na+,OH–) and ε(Na+,OH–) = 0.04 kg·mol–1. The value of ε(Na+, 3Ni(OH)− ) = (0.70 ± 0.12) kg·mol–1 is an unusually high value that leads to the strange looking solubility curve of Figure V-7. This indi-cates that the present analysis of the sparse data is probably oversimplified, especially for high base concentrations.

[50AKS/FIA]

Aksel’rud and Fialkov precipitated Ni(OH)2, which was not characterised any further, and reacted it with nickel sulphate solutions at 18°C, determined the equilibrium pH with hydrogen, quinhydrone and antimony electrodes, plotted y = log10[Ni2+] + 2log10KW + 2pH versus x = [Ni2+], and extrapolated graphically to zero concentration. The authors used log10KW valid at 25°C instead of 18°C. In Figure V-11 and Table V-6 the similarly extrapolated but appropriately corrected value is given. It should be em-phasised that by equilibrating Ni(OH)2 with NiSO4 solutions only minute amounts of the solid phase dissolve and thus the solubility relates to its finest probably amorphous fraction.


273

[50GLE/EIN]

Glemser and Einerhand investigated β-NiOOH, γ-NiOOH and other nickel oxide hy-droxides of higher oxidation state by X-ray powder diffractometry and electron micro-graphy.

[50HUM/KOL]

Polarographic studies on the reduction of the 24Ni(CN) − ion have been performed. The

results confirmed the irreversibility of the reduction, and indicated a minimum forma-tion constant of about 1024. Polarograms of saturated solutions of Ni(CN)2 showed waves due to the aquonickel(II) ion and the 2

4Ni(CN) − complex, but no free cyanide ions could be detected. This indicates that the solid is actually Ni[Ni(CN)4] (see also [51LON]. The solubility product of Ni[Ni(CN)4] in 0.1 M NaCl solution was reported to be 10log sK = – 8.77.

[50VAS/GRA]

Heat capacity measurements were reported for temperatures from 65 to 300 K for sev-eral hydrated transition-metal salts, including Ni(NO3)2·6H2O. Results were not listed for each measurement, although the specific heat values are indicated as points on a graph. “Smoothed” values are reported at 5 to 20 K intervals over the temperature range. A rather broad second-order transition for Ni(NO3)2·6H2O was found between 140 K and 170 K, with a peak at 149 K. The authors estimated the entropy change be-tween 65 K and 273 K to be 87.1 cal·K–1·mol–1 (364.4 J·K–1·mol–1). No magnetic transi-tion was found in the temperature range of the transition. The authors stated that the random error “should not exceed 0.001 cal·K–1·g–1” (1.22 J·K–1·mol–1). From the set of smoothed specific heats, the value of the molar heat capacity of the solid at 298.15 K is calculated to be (463.0 ± 2.0) J·K–1·mol–1, where the uncertainty has been estimated in the present review. The temperature range of the measurements is not adequate for the calculation of a value for mS ο (Ni(NO3)2·6H2O, 298.15 K).

[51COU]

Coughlin measured the enthalpy of nickel chloride from 298.15 to 1336 K, see Figure A-2. The NiCl2 samples for this investigation were provided by Busey and Giauque [52BUS/GIA]. The adopted melting point was Tfus(NiCl2, cr) = 1303 K.


274

Figure A-2: High temperature enthalpy of anhydrous NiCl2. Solid curve: 2 2 1 1

m m( ) (298.15) · ( 298.15) ·( 298.15 ) ·( 298.15 )H T H a T b T c Tο − −− = − + − + − , with a = 7.37949·10–2, b = 6.11344·10–6, c = 512.775. : Experimental data from [51COU].

[51HOO/WOY]


NiCl2(cr) + 2 HF(g) NiF2(cr) + 2 HCl(g) (A.8)

was studied from 477 K to 830 K. The concentrations of HCl(g) and HF(g) were meas-ured in a gas stream that had been brought to equilibrium over the solids. At the tem-peratures used, the measured equilibrium constants were independent of the flow rates of the nitrogen carrier gas containing HCl(g) or HF(g). In the present review, values of

r mH ο∆ ((A.8), 298.15 K) are calculated using the reported equilibrium constants, se-lected values of mS ο and ,mpCο (T) from Tables III-1 and III-3 for NiF2(cr) and NiCl2(cr), and the tabulated values for HCl(g) and HF(g) in the CODATA compilation [89COX/WAG]. Then, using f mH ο∆ values from Tables III-1 and IV-1 for NiCl2(cr), HCl(g), and HF(g), an average value, f mH ο∆ (NiF2, cr, 298.15 K) = – (652.76 ± 0.93) kJ·mol–1, is calculated. Calculated values of f mH ο∆ (NiF2, cr, 298.15 K) appear to be systematically more negative (by 1–2 kJ·mol–1) when based on the higher temperature measurements.

200 400 600 800 1000 1200 1400-20

0

20

40

60

80

100

120

140

160

180

( Hm(T

) − H

° m(298

.15)

) / k

J·mol

−1

T / K


275

[51LAT]

In this paper the original “Latimer” entropy contributions of metals and negative ions in solid compounds are listed.

[51LON]

The exchange of cyanide ion between the 24Ni(CN) − ion and the bulk solution was

complete within 30 seconds, which was too fast to be followed properly by the method used. A study of the Ni2+ – 2

4Ni(CN) − exchange was complicated by formation of the precipitate (Ni(CN)2) in the mixture of nickel(II) and tetracyanonickelate ions. The re-sults indicated two non-equivalent metal ions in the solid, which confirmed that its structure is indeed Ni[Ni(CN)4]. The nickel(II) exchange between 2

4Ni(CN) − and four nickel(II) complexes indicated a direct bimolecular interaction.

[52BUS/GIA]

Busey and Giauque measured the heat capacity of nickel from 15 to 300 K. The en-tropy, heat content and free energy functions have been calculated. The authors used 99.98% nickel, in contrast to the numerous low temperature heat capacity studies quoted in this paper where rather low purity Ni-metal was investigated. Calculations of thermodynamic properties of nickel were extended to 800 K on the basis of available data. However, the results of such extrapolation seem to be less reliable. The standard molar entropy of Ni determined by Busey and Giauque was equal to 29.86 J·K–1·mol–1.

[52BUS/GIA2]

The heat capacity of anhydrous nickel chloride was studied between 15 to 300 K. The anomalous region associated with the transition from the antiferromagnetic to the para-magnetic state was investigated in detail. A maximum of ,mpCο (T) was found near 52 K. This temperature corresponds with the temperature of the magnetic susceptibility maximum within the accuracy of the magnetic measurements.

[52CAR/BON]

Carr and Bonilla reported an approximate value of the standard electrode potential of nickel from the following cell:

Ni | NiSO4 | KCl(sat.) | KCl(sat.), Hg2Cl2 | Hg.

Their data have been recalculated with the standard potential of the Hg2Cl2 | Hg electrode ((0.2681 ± 0.0026) V, [89COX/WAG]) and the chloride activity of a saturated KCl solution a (KCl) = 2.829 [89LOB]. The activity of Ni2+ in the solu-tion was calculated using the SIT model supported by ChemSage. The results are listed in Table A-4.


276

Table A-4: Standard electrode potential of Ni2+ | Ni at 25°C.

NiSO4/M (mol·dm–3)

log10([NiSO4]/m) 2+10 Nilog a E (vs. std. CE) (mV)

E (vs. SHE) (mV)

E° (vs. SHE) (mV)

0.0001 – 4.0014 – 4.10627 – 590.4 – 349.0 – 227.54 0.0010 – 2.999 – 3.15545 – 571.3 – 329.9 – 236.56 0.0100 – 1.9987 – 2.41766 – 548.6 – 307.2 – 235.69 0.1000 – 0.9989 – 1.81732 – 527.7 – 286.3 – 232.54 1.0000 0.0017 – 1.26378 – 499.2 – 257.8 – 220.42

The authors rejected the results in rows 1, 4, and 5 altogether. As Figure A-3 shows only the data in row 5 deviate considerably from the curve predicted with a value of E°(Ni2+ | Ni, 298.15 K) = − (0.2331 ± 0.0082) V. This value agrees within the error limits with the value in the NBS tables [82WAG/EVA] and the result of [29HAR/BOS] as well as the compilation of [89BRA]. The reversibility of the cell was tested thor-oughly and thus the value of [52CAR/BON] confirms the result of [29HAR/BOS] by an independent method. When only the results in row 2 and 3 are accepted, as recom-mended by the authors, [29HAR/BOS] and [52CAR/BON] agree within 1 mV. The uncertainty from the use of an uncorrected non-thermodynamic diffusion potential re-mains, and as a consequence the result of [29HAR/BOS] has been preferred.

Figure A-3: Plot of E (vs. SHE) versus log10([NiSO4]/m)

-4 -3 -2 -1 0

-360

-340

-320

-300

-280

-260 Ni|NiSO4|KCl (sat.)|KCl (sat.), Hg2Cl2|Hg

25oC

E / m

V (v

s. SH

E)

log10 (m (NiSO4) / mol·kg−1)

exp. data (vs. SHE) [52CAR/BON] E / mV = − 233.1 − (R T /2 F) ln (a (Ni2+)) error limits: ± 8.2 mV


277

[52CHA/COU]

Approximate hydrolysis constants for a number of cations, including Ni2+, were deter-mined by potentiometric titrations at 30°C and 0.1 M (K)Cl ionic strength. The forma-tion of NiOH+ as the sole hydrolytic species was assumed. However, the total metal concentration used (about 0.01 M) is too high to neglect formation of the tetranuclear hydroxo complex. The experimental data are not reported. Therefore, the results in this paper have not been considered further in the present review.

[52GAY/WOO]

The pH values of solutions prepared from purified NiCl2 ([Ni2+]t = 0.0156 to 0.5 M) were measured to obtain data on the magnitude of the hydrolysis reactions. Only the formation constant, 1,1

*b , of NiOH+:

Ni2+ + H2O NiOH+ + H+,

was taken into account to explain the variation of the measured pH values. This paper was a short communication, and thus, several experimental details (such as the method of calculation and the source of the activity coefficients) were not reported. Further-more, the [Ni2+]t used was far too high to neglect formation of polynuclear hydroxo complexes. In addition, the effect of formation of chloro complexes was not addressed. The experimental data are inadequate for proper re-evaluation; therefore the reported constant was not considered in this review.

[52ROS]

An aqueous solution of nickel nitrate saturated with carbon dioxide was treated with ammonium carbonate at 0°C. After one week at 5°C bluish green crystals were formed. The analysis of this precipitate agreed within the error limits with the formula NiCO3·6H2O. The X-ray powder diffraction pattern was taken and listed.

[52SUD]

The reduction equilibria of synthetic nickel sulphides with hydrogen gas were measured with the help of a flow method. Ni7S6 was prepared by precipitation (reaction of H2S with an ammonia containing nickel sulphate solution). The solid product was annealed at 661 K in an evacuated silica tube for one hour. The composition was shown to be close to Ni7S6 by gravimetric analysis; no X-ray diffraction analysis was carried out. As the reduction experiments were mainly performed at temperatures below 675 K, the gas phase, characterised by a certain ratio p(H2S)/p(H2), is in equilibrium with Ni3S2 and Ni9S8 in accordance with [94STO/FJE]. Thus, the ratios p(H2S)/p(H2) given as a func-tion of temperature should correspond to the equilibrium:

½Ni9S8 + H2(g) 32 Ni3S2 + H2S(g).

In addition, the author prepared Ni3S2 by reduction of Ni7S6 with H2. Again, the solid product was characterised by chemical analysis, but no X-ray diffraction


278

analysis was performed. Ni3S2 was partly reduced with hydrogen gas at temperatures ranging from 855 to 1000 K and the ratio p(H2S)/p(H2) in the gas phase was measured. The results were erroneously referred to the equilibrium between heazlewoodite (low-temperature form of Ni3S2) and metallic nickel. However, at temperatures above 834 K the high-temperature form (β1-Ni3S2) is stable [80SHA/CHA], [91STO/GRO]. More-over, the temperature for the eutectic between Ni, β1-Ni3S2 and the liquid phase is 905 K [80SHA/CHA], so that above this temperature some of the material equilibrated with the gas phase was molten. For these reasons the data of the experiments on Ni3S2 are discarded.

[53BUS/GIA]

Busey and Giauque investigated the equilibrium reaction:

NiCl2(cr) + H2(g) Ni(cr) + 2 HCl(g) (A.9)

in order to provide evidence from the application of the third law of thermodynamics that a ferromagnetic substance approaches zero entropy at 0 K. This plausible implica-tion was experimentally confirmed for the first time by this study.

A third law analysis of the reported equilibrium constants was carried out em-ploying the Gibbs energy functions of Ni(cr) and NiCl2(cr) selected in this review, whereas the Gibbs energy functions of H2(g) and HCl(g) were taken from [89COX/WAG]. The experimental data depicted in Figure V-18 are listed in Table A-5.

Table A-5: Equilibrium constants of Reaction (A.9)

T / K Kp / atm Kp / bar – RT ln(Kp / bar) / kJ·mol–1

661.67 0.2503 0.2536 7.5476 676.65 0.3949 0.4001 5.1532 693.00 0.6365 0.6449 2.5272 645.33 0.1488 0.1508 10.1516 676.18 0.3913 0.3965 5.2011 707.11 0.9568 0.9695 0.1822 738.02 2.135 2.1633 – 4.7349 707.02 0.954 0.9666 0.1994 646.19 0.1514 0.1534 10.0720 692.14 0.6223 0.6305 2.6539 692.12 0.6238 0.6321 2.6400 707.42 0.9565 0.9692 0.1842 722.73 1.423 1.4419 – 2.1989 737.89 2.095 2.1228 – 4.6180 630.18 0.0872 0.08836 12.7133


279

[53FRO2]

The formation constants of 2Ni(SCN) xx

− (x = 1 – 3) complexes have been determined at 20°C, at pH = 3, in 1 M (Na)ClO4 solution, using Amberlite IR-105 cation exchanger. The thiocyanate concentration ranged from 0 to 0.5 M, [Ni2+]tot varied between 1 and 5 mM. The equilibrium nickel(II) concentration was determined spectrophotometrically using dithiooxalate, and that of the thiocyanate ion by a Volhard titration. The results obtained were qualitatively confirmed by spectrophotometric measurements of the aqueous nickel(II) – thiocyanate system. Considering the accepted enthalpy values for the formation of 2Ni(SCN) x

x− (x = 1 – 3) species, the reported formation constants were

corrected to 298.15 K: 10 1log b = (1.14 ± 0.06), 10 2log b = (1.58 ± 0.06), 10 3log b = (1.72 ± 0.10).

[53GRI/SLU]

Ni3S4 was synthesised by precipitation of NiS from aqueous solution with H2S. The suspended NiS was oxidised with oxygen to NiS2 which eventually disproportionated into Ni3S4 and H2SO4.

[53SAN]

This paper seems to be a slightly updated English version of [37SAN2]. Sano studied the reduction equilibria of cobalt(II) and nickel(II) chloride by hydrogen. He used the method of observing the total pressure and the partial pressure of hydrogen by means of a palladium membrane. Busey and Giauque tried to account for the difference of their results and Sano's results by assuming that thermal equilibrium was not attained be-tween Sano’s thermocouple and the reaction vessel. If Sano's temperatures were lower by about 14 K, the equilibrium constants would agree quite well with those of [53BUS/GIA]. In Figure V-18 the original experimental data have been depicted. They are listed in Table A-6.

Table A-6: Equilibrium constants of Reaction (A.9) [53SAN]


792 0.7345 0.7402 – 11.2236 760 0.4141 0.4198 – 6.1083 738 0.1985 0.2042 – 2.8853 709 – 0.1783 – 0.1726 2.3426 661 – 0.8016 – 0.7959 10.0716

[53SCH/POL]

Schwab et al. determined the solubility product of freshly precipitated Ni(OH)2 by measuring the pH when one half of the Ni2+ originally present in NiCl2 solutions had been precipitated by adding NaOH. The solid phase was not investigated any further.


280

All experiments were carried out at “room temperature”. The authors reported 10 ,0log sK = – 15.5. It is difficult to recalculate this value from the experimental data

given; in any case, two different values of the pH of half precipitation can be found in [53SCH/POL]. If the first (pH = 7.75) is used, this results in 10 ,0log sK = – 15.5. Use of the second value (pH = 8.4) would lead to 10 ,0log sK = – 14.2. Thus, it was concluded that to extrapolate either of these results to I = 0 is not justified, and consequently nei-ther was included in the final SIT calculations. In addition the hydrolysis of nickel (II) was studied by potentiometric titrations of NiSO4, Ni(NO3)2 and NiCl2 solutions. The initial concentration of nickel(II) was 2.44 mM for all titrations. The ionic strength was not controlled. The experimental data were interpreted assuming that NiOH+ was the only hydrolysis product formed. For all three solutions, irrespective of the anion pre-sent, 10 1,1

*log b = – 9 was reported. This value was not considered further in this review.

[53SWA/TAT]

Swanson and Tatge recorded the pattern of a NiO sample obtained from Johnson, Mat-they & Co., Ltd. The estimated purity of 99.99% was corroborated by spectroscopic analysis, which showed only faint traces of Mg, Si, and Ca. After comparison with eight other published patterns the cell dimensions given by [53SWA/TAT] were accepted for JCPDS-ICDD card No.4-835.

[54DOB]

Nickel sulphate solutions were titrated with sodium hydroxide and vice versa at 75°C. Each titration was carried out within 2 to 3 hours which appears to be a rather short pe-riod for equilibration even at 75°C. Ni(OH)2 and NiSO4·3Ni(OH)2 were precipitated in dilute and concentrated solutions, respectively. Neither solid phase was investigated with respect to its structure, e.g., by X-ray analyses. It is difficult to recalculate the re-sults properly, because the ionic strengths are not given in Dobrokhotov’s Table 1. The value for Ni(OH)2 listed by the author ( 10 ,0log sK = – 16.2, 10 ,0

*log sK = 9.25 at 75°C) agrees surprisingly well with our recent results. The recalculation of Dobrokhotov’s digitalised data from Figures 3 and 5.1 with the SIT model resulted in 10 ,0

*log sK = 9.48 at 75°C, whereby complex formation between Ni2+ and 2

4SO − had been taken into ac-count (see Figure V-11 and Table V-6).

[54FEI/HAR]

This is an abstract of a conference paper containing no experimental details at all. The authors maintain that the solubility products of Ni(OH)2, freshly precipitated and aged at 110°C, differ by 2.5 orders of magnitude. As mentioned in Section V.3.2.2.3, the solid phases were not characterised with respect to their X-ray patterns.

[54ROS]

This important contribution deals with the investigation of phase relations in the system Ni—S by measurement of the partial pressure of sulphur in equilibrium with various


281

solid phases. The nickel sulphides were prepared by solid state reaction from the ele-ments. The solid sulphides were equilibrated with hydrogen gas at temperatures be-tween 673 and 1273 K. The partial pressure of sulphur is given by the ratio p(H2)/p(H2S) which is recorded in situ by density measurements employing a buoyancy balance. The compositions of the solid phases were determined by chemical analysis before and after equilibration with hydrogen. From these equilibrium studies the stabil-ity ranges of the different sulphides as well as their Gibbs energies of formation were derived. Gibbs energy functions are given for the following reactions:

32 Ni+H2S ½Ni3S2+H2 808K

r m 673K[ ]G∆ = (– 75479.4 + 32.217 (T/K)) kJ·mol–1,

2Ni3S2+H2S Ni6S5+H2 798Kr m 673K[ ]G∆ = (– 12761.2 – 7.7404 (T/K)) kJ·mol–1,

Ni6S5+H2S 6Ni1–xS+H2 833Kr m 673K[ ]G∆ = (– 21840.5 + 14.477 (T/K)) kJ·mol–1,

32 Ni+H2S ½Ni3+xS2+H2 923K

r m 808K[ ]G∆ = (– 49998.8 – 13.514 (T/K)) kJ·mol–1,

Ni3+xS2+H2S 3Ni1–xS+H2 1083Kr m 833K[ ]G∆ = (– 69454.4 + 70.291 (T/K)) kJ·mol–1,

Ni1–xS+H2S NiS2+H2 1073Kr m 673K[ ]G∆ = (– 19664.8 + 55.815 (T/K)) kJ·mol–1.

It is worth mentioning that according to Stølen et al. [94STO/FJE], the compo-sition of the phase Ni6S5 is actually Ni7S6.

[54SHC/TOL]

Shchukarev et al. measured the equilibrium concentrations of HBr(g) and H2(g) at equi-librium over solid samples of nickel and nickel bromide for temperatures from 623 K to 928 K. The primary reaction was:

NiBr2(cr) +H2(g) Ni(cr) + 2HBr(g).

The values of the logarithm (base 10) of the reported equilibrium constant val-ues (converted to the 1 bar standard state) are listed in Table A-7, where:

2

2HBr(g)

10 10H (g)

/ barlog log

/ barp

pK

p

=

.

The authors indicated that equilibrium may not have been established success-fully at the lowest temperatures (below 700 K), and that almost complete conversion of the NiBr2 to Ni at 928 K made the value for the highest temperature somewhat suspect. The reported equilibrium constant values were used here with the selected entropy (Sec-tion V.4.1.4.1) and equations and values for ,mpCο (NiBr2, cr, T), Section V.4.1.4.2,

,mpCο (Ni, cr, T), Section V.1.2, and ,mpCο (H2, g) and ,mpCο (HBr, g) [89COX/WAG], to calculate values for f mH ο∆ (NiBr2, cr). The authors suggested that there was evidence for a transition for NiBr2(cr) just below 800 K (and an enthalpy of transition of 24 kJ·mol–1). Although two crystalline modifications of NiBr2 are known [34KET], the evidence from these sparse measurements [54SHC/TOL] seems to be insufficient to conclude that there is a transition with such a large transition enthalpy. The authors ap-


282

parently did not confirm the crystalline form of the nickel bromide after measurements at each temperature. The calculated values of f mH ο∆ for 298.15 K become progressively more negative with increasing temperatures of the experiments on which the f mH ο∆ values are based (and the results for nickel chloride, as obtained by the same authors, also drift slightly with temperature). The average value of f mH ο∆ (NiBr2, cr) from the measurements done between 700 and 900 K is – (215.02 ± 0.91) kJ·mol–1.

Table A-7: Enthalpy of formation values calculated from equilibrium gas compositions.

T/K 10log pK f mH ο∆ (NiBr2, cr)

623 – 2.915 – 210.86 683 – 1.930 – 212.25 723 – 1.392 – 214.37 748 – 1.070 – 214.36 773 – 0.777 – 214.48 823 – 0.295 – 215.48 873 0.130 – 216.42 928 0.497 – 218.24

Shchukarev et al. investigated the decomposition reaction:

NiCl2 Ni + Cl2

by studying the equilibrium:

NiCl2(cr) + H2(g) Ni(cr) + 2 HCl(g) (A.10)

and combining the latter with the well known formation equilibrium of hydrogen chlo-ride:

H2 + Cl2 2 HCl.

A third law analysis of the reported equilibrium constants was carried out using the same auxiliary Gibbs energy functions described under [53BUS/GIA]. The experi-mental data are depicted in Figure V-18 and listed in Table A-8.

Table A-8: Equilibrium constants of Reaction (A.10) [54SHC/TOL]

T / K Kp / atm Kp / bar – RT ln(Kp / bar) / kJ·mol–1

573 0.013 0.01317 20.6272 623 0.078 0.07903 13.1460 693 0.59 0.5978 2.9643 723 1.23 1.2463 – 1.3236 743 2.00 2.0265 – 4.3633 773 4.00 4.0530 – 8.9944 823 9.8 9.9299 – 15.7079


283

[55BRO/PRU]

This paper reports cryoscopic measurements for NiSO4 aqueous solutions (0.00 > t /°C > – 0.21), and demonstrates the dependence of the calculated value of K1 on the as-sumptions in the complexation model. Based on the arguments of Brown and Prue, it is clear that use of the SIT as is done in the TDB project imposes constraints on the data analysis, and that these constraints should lead to a specific value for K1. However, re-calculations show that the association constant is also strongly dependent on the concentration at which the data set is truncated. For the purposes of the present review, recalculations were done using the SIT, but only results for molalities ≤ 0.03 were used in the selection of the value of 10 1log K . When data for higher concentrations were in-cluded, the calculated value for the association constant increased, and 10 1log K = 300 was obtained if the entire data set was used.

[55CAT/STO]

This adiabatic calorimetry study (12 K to 300 K) appears to be the only low-temperature heat capacity study for NiF2(cr). From their large series of results the au-thors calculated mS ο (NiF2, cr, 298.15 K) = (73.60 ± 0.17) J·K–1·mol–1. A λ transition was found at (73.22 ± 0.05) K. The experimental values seem to have a non-systematic un-certainty of less 0.3 J·K–1·mol–1. In a later paper, the authors indicated that the heat ca-pacity values above 270 K may have a slightly greater uncertainty than those for lower temperatures [55STO/CAT]. Slightly lower values between 270 K and 300 K mesh bet-ter with the available high-temperature drop calorimetry results [70BIN/HEB].

[55KRA/WAR]

Krauss and Warncke used two different calorimetric methods to determine the specific heat of high-purity nickel. The measurements in the temperature range between 180 and 600°C were carried out using a continuous calorimetric technique and between 500 and 1160°C by means of “reverse calorimetry”. The source of heat in the last method was a Pt-rod of known temperature and heat capacity which was placed into the calorimeter containing a nickel sample. The separation of the determined heat capacity into a lattice-vibration term, a magnetic term and a residual term was briefly discussed in this paper. The reported data were used by this review for fitting of the thermal heat capacity func-tion for nickel crystal.

[56AHR/ROS]

The stabilities of the fluoride complexes of nickel(II), copper(II) and zinc(II) have been determined by a pH-metric method using a quinhydrone electrode at I = 1 M Na(ClO4, F) and at 293.15 K. The pK value of 2HF− was reported in an earlier paper (pK = 2.93, [56AHR/LAR]). Under the conditions used ([F–]t ≤ 400 mM, [Ni2+]t ≤ 100 mM) only the monocomplexes (MF+) were found to exist in appreciable concentration.


284

[56CUT/KSA]

Potentiometric titrations of NiClO4 solutions with NaOH were performed at (20 ± 0.1)°C. To determine the solubility constant, the pH of incipient precipitation was measured for nickel concentrations of 0.0041 – 0.034 M at constant ionic strength, I = 1.0 M NaClO4. The value for 10 ,0log sK ο , obtained by extrapolating these experimental data to I = 0 with the SIT model is listed in Table V-6. The hydrolysis constant was de-termined at nickel concentrations ranging from 0.0025 to 0.184 M. The ionic strength was not controlled, and the activity coefficients were calculated using an extended De-bye-Hückel expression. Such a general expression is normally assumed to be valid at ionic strengths below 0.1 M. In the course of this study, the ionic strengths varied from 0.007 to 0.55 M, with most data pairs at I > 0.1 M. The formation of the single complex NiOH+ was assumed, and a value of 10 1,1

*log οb = – 8.94 was obtained from graphical extrapolation to I = 0. The experimental data used in the calculations referred to very low degrees of hydrolysis, thus, large unavoidable errors render the result dubious. For this reason, the reported hydrolysis constant was not included in the detailed evaluation.

[56FIA/SHE]

This was a scoping study. Conductance measurements were carried out for dilute solu-tions of nickel sulphate. However, the data were incorrectly analysed in terms of a 1:2 nickel:sulphate complex, and the data are not available for re-analysis.

[56JEN/PRA]

Nickel hydroxide was prepared by adding an excess of sodium hydroxide solution to a solution of nickel chloride. The precipitate was freed from alkali by washing with water and was then dried in an air-oven at 98 – 100°C. This solid phase was equilibrated ei-ther with HCl (0.0189 – 0.1892 M) at 25°C or buffer solutions of (1:1) acetic acid (0.04 – 0.20 M) and sodium acetate at 29°C. The authors preferred the result of the buffer so-lutions, however a recalculation shows that the mean values of either series agree within an experimental error of ± 0.1 log10K units and thus both have been recorded in Fig-ure V-11 and Table V-6.

[56KEN]

This paper reports cryoscopic measurements for NiSO4 aqueous solutions in saturated potassium perchlorate solutions (eutectic freezing point 272.99 K, freezing point de-pressions, ∆T, are between 0.0673 K and 0.181 K). Thus, the measurements were done with less variation of ionic strength than for measurements done in water. The results of this study are assessed in our discussion of a later paper [58KEN] by the same author.

[56LAN/MIE]

Lange and Miederer determined the relative apparent molar enthalpy, φL, of NiSO4 and demonstrated that:


285

0,

limm

T p

Lm

φ

→

∂

∂ = 28.9 kJ·kg0.5·mol–1.5

leads to the slope, which is over two times higher than the value of AL predicted by the Debye-Hückel limiting law. Lange and Miederer’s experimental data were re-evaluated using the chord method described in [58HAR/OWE]. A still higher limiting slope has been obtained:

0,

limm

T p

Lm

φ

→

∂

∂ = (37.5 ± 2.6) kJ·kg0.5·mol–1.5.

The φL function thus derived was employed to extrapolate the enthalpies of NiSO4·6H2O(cr) solution determined by Goldberg et al. to zero ionic strength [66GOL/RID].

[56SOK]

The application of thermal, metallographic, X-ray, conductivity and specific gravity measurements revealed the phase relations in the binary system Ni – S in the region from 30.0 to 50.0 at% S. The phase diagram given by Elliott [65ELL] is based on re-sults of this study as well as the work of Kullerud and Yund [62KUL/YUN].

[56YAT/VAS]

The solubility product at 25.0°C for hydrated Ni2P2O7, and 1:1 and 1:2 complexation constants between Ni2+ and 4

2 7P O − were determined from a small set of solubility meas-urements in aqueous solutions of sodium pyrophosphate ranging in concentration from 0.032 to 0.123 M. The experiments seem to have been done carefully. Hydrolysis of pyrophosphate was considered to be small, and was neglected in the data analysis. This was probably not a bad approximation, as it can be calculated that < 5% of the unasso-ciated pyrophosphate in the final solutions was protonated. The species 2 7NiHP O− [64HAM/MOR], [73PER/SEC] was even less important under the experimental condi-tions. Because of the presence of many highly charged species, the ionic strength of the experimental solutions was high (0.5 to 1.2 M), and varied from solution to solution. Estimation of interaction coefficients would be highly speculative for some of the spe-cies, and it does not seem useful to attempt to use the value in this paper for the first association constant to determine thermodynamic values for I = 0. The value of 5.82 for

10 1log K , is consistent with later studies in constant ionic strength media [64HAM/MOR], [73PER/SEC].

The value determined for the cumulative second association constant (Reaction (A.11)) was 2b = 1.54 × 107

Ni2+ + 2 42 7P O − 6

2 7 2Ni(P O ) − (A.11)

(actually the cumulative dissociation constant, (6.5 ± 2.0) × 10–8) was reported). Re-examination of the experimental data suggests that the 1:2 complex does form, but evi-dence for its formation is only marginal except in those solutions with higher initial


286

pyrophosphate concentrations (> 0.08 M), and 10 2log K is probably < 1. The neglect of pyrophosphate complexation with sodium ions (an inherent assumption required within the context of the SIT) [94STE/FOT], introduces a further uncertainty. Therefore, no value for K2 is selected in the present review.

In the absence of added sodium pyrophosphate, the solubility of the nickel py-rophosphate in 0.5 M to 1.0 M KNO3(aq) was reported as 3.4 × 10–4 M, which is not inconsistent with the solubility product proposed by the authors (for higher ionic strengths) if the lower hydroxide concentration, and consequent formation of 2 7HNiP O− , is considered.

[56YAT/VAS2]

Measurements of the enthalpy of solution of aliquots of a 1.92 M solution of nickel ni-trate into solutions of sodium pyrophosphate (0.025 M to 0.184 M) were used to deter-mine the enthalpy of the complexation reactions of Ni2+ with 4

2 7P O − . The final nickel: pyrophosphate ratios were also varied at the four different pyrophosphate concentra-tions. The values r 1H∆ = (17.61 ± 0.17) kJ·mol–1 and r 2H∆ = – (9.25 ± 0.25) kJ·mol–1 at 25°C were reported. A value of 1

2K − = 2.5 × 10–2 derived from the calorimetric results is also reported. It is not clear that the enthalpies of the two complexation reactions are as easily separable as suggested by the authors, especially at the higher pyrophosphate concentrations. The complexation constants cannot be accurately calculated for these solutions (see the discussion above for [56YAT/VAS]), and therefore there are large uncertainties in the enthalpies of the complexation reactions. Recalculations suggest that the value for r 1H∆ is more positive than the reported value, that r 2H∆ is more nega-tive than the reported value, and that in none of the solutions does a single complex ac-count for > 90% of the nickel in solution. For lower pyrophosphate concentrations, the measured enthalpies also must include a contribution for deprotonation of small amounts (< 5%) of monohydrogenpyrophosphate.

[57CHU]

The nickel orthoarsenite, not identified except by the Ni:As ratio, was found to dissolve in dilute nitric acid solutions at 20°C over 12 hours to form solutions with aqueous nickel concentrations of 8.7 × 10–3 mol·dm–3 at a “pH” value of 6.75 and 3.1 × 10–3 mol·dm–3 at a “pH” value of 7.10. The Ni:As ratio of the solid was found to be un-changed after the experiment. The author did not report a solubility product based on their measurements. If the dissolution of the solid is assumed to correspond to the reac-tion:

Ni3(AsO3)2·xH2O + 6H+ 3Ni2+ + 2HAsO2(aq) + (2 + x)H2O (A.12)

and the concentrations are assumed to be sufficiently low that the ∆ε terms in the activ-ity coefficient equations are negligible, the average value of 10log K (A.12) is (28.69 ± 0.16). In the near-neutral solutions, ionisation of HAsO2 can be assumed small,


287

as can the extent of any complex formation. If the initial “pH” values are compared to those calculated based on the final concentrations of Ni2+, there are discrepancies of 19% and 31%. This suggests that the measured “pH” values were not particularly accu-rate. In the present review, the average value for 20°C is accepted as the value for 25°C, but with an increased uncertainty, i.e., 10log K ((A.12), 298.15 K) is (28.7 ± 0.7).

[57KIU/WAG]

This pioneering work demonstrates the usefulness of emf measurements on galvanic cells, involving solid electrolytes for the determination of the standard molar Gibbs en-ergy of formation of oxides, sulphides, selenides, and tellurides at elevated tempera-tures. These measurements have been carried out by means of the “open cell stacked pellet” technique. The reference electrode was a metal – metal oxide system of known Gibbs energy of formation. The emf of the cell Fe, FexO (wüstite) (0.85 ZrO2 + 0.15 CaO) Ni, NiO and the standard Gibbs energy change of the reaction:

Ni(s) + ½O2(g) NiO(s)

were determined in the temperature range from 1023.15 to 1413.15 K. The reported reproducibility of emf was ± 2 mV or better, corresponding to an uncertainty in

f mG∆ (T) of ± 0.418 kJ·mol–1. The standard molar Gibbs energy of formation of nickel oxide was the following function of temperature: f mG∆ = (– 235.171 + 0.08567 (T/K)) kJ·mol–1.

[57KIV/LUO]

The formation constants of nickel(II) and zinc(II) chloro complexes were determined by a polarographic method at 25°C using lead(II) as the indicator ion. The ionic strength was held at 2 M using sodium perchlorate. The concentration of chloride ion was varied between 0.4 – 1.4 M. The authors reported the formation of two nickel(II) chloro spe-cies: NiCl+ and NiCl2(aq) and their formation constants β1 = (0.56 ± 0.20) and β2 = (0.9 ± 0.5). From these data, it follows that K2 (= β2/β1 = 1.6) is three times higher than K1 (= β1). This is very unlikely taking into account the low complex forming ability of chloride ion with nickel(II), and is probably the result of some unidentified experimen-tal error. Consequently, these data could not be used to formulate conclusions concern-ing the nature and stability of chloro complexes.

[57MOR/ZEL]

This report describes a carrier-gas method for measuring vapour pressure of liquid met-als. Vapour pressures of nickel were determined at temperatures 1813 to 1893 K. The results can be represented by the following equation: log10 pmm = – 21030/T + 9.689, where pmm is vapour pressure (in mm of Hg). The heat of vaporisation of nickel at 0 K was determined to be 427.898 kJ·mol–1.


288

[58KEN]

This paper reports cryoscopic measurements for NiSO4 aqueous solutions in saturated potassium chlorate solution (freezing point 272.55 K) and saturated potassium nitrate solution (freezing point 270.32 K). The eutectic freezing point depressions, ∆T, are be-tween 0.0296 K and 0.178 K. Thus, together with the measurements reported in [56KEN], the variation in the complexation constant, K1, with ionic strength was exam-ined.

The experimental association constant values, K1, were extrapolated to the saturation ionic strength of the supporting electrolyte at the melting point. Therefore, the reported values at different limiting ionic strengths are also values for slightly dif-ferent temperatures. For the cryoscopic measurements using potassium chlorate, the ionic strength contribution of the supporting electrolyte is generally less than that from the nickel sulphate. Only in the potassium nitrate solutions is the contribution of the nickel sulphate to the total ionic strength reasonably minor. Therefore, the extrapolation method used by the author is probably appropriate only for the potassium nitrate and (perhaps) the potassium chlorate systems.

The author’s extrapolations to determine K1 in the saturated solution of sup-porting electrolytes are not strictly compatible with the TDB use of the SIT, but the re-ported K1 values are probably adequate if the uncertainties are increased. The values log10K1 = (1.27 ± 0.30) (– 0.8°C, I = 0.26 mol·kg–1 KClO3(aq)) and log10K1 = (0.69 ± 0.20 (– 2.8°C, I = 1.15 mol·kg–1 KNO3(aq)) are accepted for use in the present assessment. In this review, the SIT is used to determine the values at I = 0 (with the effective interaction coefficient for Ni2+ with 3NO− (0.18 ± 0.01) kg·mol–1, neglecting nitrate association (Section V.6.1.2.1), and the estimated interaction coefficient for Ni2+ with 3ClO− (0.28 ± 0.10) kg·mol–1). The latter value was estimated as the average of the interaction coefficients of Ni2+ with perchlorate and nitrate, with an increased uncer-tainty. The results are 10 1log K = (2.35 ± 0.30) (– 0.8°C) and 10 1log K = (2.16 ± 0.20) (– 2.8°C).

The curve fits to Kenttämaa’s data could be used to determine values of the second cumulative association constant, but these values of β2 are strongly dependent on activity coefficient assumptions and on the values used for the heats of solution. Kenttämaa did not propose β2 values (and considered that the value corresponding to β2 from the fitting process also incorporated systematic errors from the experiment and data analysis). Rossotti and Rossotti [59ROS/ROS] did report β2 values based on the measurements done by Kenttämaa using potassium nitrate mixtures. We concur with the opinion of the original author, and do not accept these values of approximately 10 2log K = 1 at I = 0.3 m and 1.2 m, although they are roughly comparable to a value reported by Tanaka, Saito and Ogino [63TAN/SAI] ( 10 2log K = 1.4 in 1.2 M NaClO4(aq)-CH3COONa mixtures at 25°C).


289

[58MUL]

Muldrow synthesised anhydrous NiCl2 and determined sol mH∆ of reaction:

NiCl2(cr) Ni2+ + 2Cl– (A.13)

at 298.15 K. The mean value of sol mH ο∆ (A.13) has been obtained from column 4 of Table A-9 by Equation (A.14)

sol mH ο∆ (A.13) = sol mH∆ (A.13)– φL, (A.14)

see Section V.2.1.3.2 (NiCl2·6H2O(cr) cycle). For the calculation of φL the last term of Equation (IX.44) in [97GRE/PLY2] had to be neglected, because the temperature de-rivative of ε(Ni2+, Cl–) is unknown. Systematic errors may have arisen from the fact that the NiCl2 samples were not completely dry (n(H2O) / n(NiCl2) ≈ 0.04 – 0.08).

Table A-9: Enthalpy of solution of NiCl2(cr) in H2O(l) at 298.15 K [58MUL].

b (NiCl2) (mol·kg–1)

sol mH∆ (A.13) (kJ·mol–1)

φL (kJ·mol–1)

sol mH ο∆ (A.13) (kJ·mol–1)

0.00545 – 82.508 0.667 – 83.175 0.00971 – 83.596 0.854 – 84.451 0.00192 – 82.341 0.417 – 82.758 0.00200 – 82.718 0.424 – 83.142 0.00373 – 82.760 0.564 – 83.323 0.00555 – 83.011 0.672 – 83.683 0.00138 – 81.881 0.357 – 82.238 0.00279 – 82.508 0.494 – 83.003 0.00685 – 83.303 0.736 – 84.040(a) 0.01036 – 83.387 0.878 – 84.265

(a) Carried out in dilute acid.

[58TRE]

The method of frontal analysis on an ion-exchange column was used to study the forma-tion of weak chloro complexes of nickel(II), cobalt(II) and copper(II). Although, the reported dissociation constant for the complex NiCl+ (Kdiss = (4.6 ± 0.1), which can be converted to 10 1log K = – (0.66 ± 0.01)) is close to other reported values, some experi-mental data, such as the temperature of the measurements, were not published. There-fore, the reported constant is rejected in this review.

[58VAI/RAM]

The only part of this work that was related quantitatively to pyrophosphate complexa-tion of nickel was a sparse set of spectrophotometric measurements from which an “in-


290

stability constant” of 2.4 × 10–4 was derived. The pH of the solutions is not indicated, though apparently the sodium salt was used. The association quotient is smaller by sev-eral orders of magnitude than values from other studies, and there are insufficient ex-perimental details to determine if the value might refer to a protonated complex. The reported value is not used in the present review.

[58YAT/KOR]

The formation of MSCN+ complexes (M = Mn(II), Fe(II), Co(II) and Ni(II)) has been studied by a spectrophotometric method, by detecting the decrease of the absorbance of Fe(SCN)2+ complex at 465 and 495 nm with increasing concentrations of M(II). The concentration of Fe(III) and SCN– was 5 and 0.5 mM, respectively, while the concentra-tion of Ni(II) ranged from 0.02 to 0.5 M. Consequently, the ionic strength varied be-tween 0.13 and 1.56 M. The Davies equation was used to calculate the activity coeffi-cients. For the purposes of the present review, the reported data were re-evaluated using the SIT (cf. Figure A-4). From the plot of this figure, 10 1log οb = (1.69 ± 0.15) and ∆ε = – (0.075 ± 0.030) kg·mol–1 can be derived.

Figure A-4: Extrapolation to Im = 0 (using the SIT) of the formation constants for the species NiSCN+ as reported in [58YAT/KOR].

[59ACH]

The pH of slightly alkaline 0.01 M K2SO4 and KClO4 solutions was measured before and after addition of a Ni(NO3)2 solution under N2 atmosphere. The concentration of

0.0 0.5 1.0 1.51.2

1.4

1.6

1.8

2.0

2.2

log 10

β1 +

4D

I / mol·kg−1


291

nickel(II) in the titration vessel varied between 0.003 and 0.024 M. The ionic strength was varied during the measurements, and the Güntelberg equation was used to calculate the activity coefficients. Since the inertness of sulphate media is highly questionable, only the data in perchlorate solution will be discussed further. The solutions were equilibrated for 30 minutes. The experimental data were interpreted only in terms of the formation of NiOH+. Therefore, we re-evaluated the reported experimental data assum-ing the presence of both NiOH+ and 4

4 4Ni (OH) + . The following values were obtained: 10 1,1

*log b = − (10.85 ± 0.10) and 10 4,4*log b = – (26.8 ± 1.0). This calculation revealed,

however, that the degree of hydrolysis is very low. The highest concentration of NiOH+ and 4

4 4Ni (OH) + was 5.4 × 10–7 and 4.4 × 10–9 M, respectively. Therefore, above data were not considered in the present review.

[59BLA/GOL]

The spin-lattice relaxation times of water protons have been measured to study complex formation between nickel(II) and cyanide ion in concentrated aqueous ammonia solu-tion (ρ = 0.88 g·cm–3). The relaxation time increased up to the ratio [CN–]/[Ni2+] = 4/1, where a sharp breakpoint was observed, according to the transformation of the para-magnetic 2+

3 6Ni(NH ) to the diamagnetic 24Ni(CN) − . At higher cyanide concentration the

relaxation time decreased with increasing cyanide concentration, and after a second breakpoint at the ratio [CN–]/[Ni2+] = 6/1 the relaxation time was unaffected by further cyanide addition. From these observations the authors were led to the conclusion that the addition of cyanide ion to 2

4Ni(CN) − results in quantitative formation of 46Ni(CN) − .

This finding was disputed later by many authors [60MCC/JON], [62BEC/BJE], [63PEN/BAI], [64GEE/HUM], [65COL/PET], [71PIE/HUG], and now the non-existence of 4

6Ni(CN) − is well established.

[59BLA/GOL2]

In this work an effort has been made to determine the formation constants of the 46Ni(CN) − species, using relaxation times derived from 1H NMR measurements. As

already mentioned in the discussion of [59BLA/GOL], the existence of this species can be ruled out, even in the case of high excess of cyanide ion.

[59FRE/SCH]

Complex formation in the nickel(II) – cyanide system has been investigated by spectro-photometry (267.5 nm) at 24.92°C and at several ionic strengths (I = 0.0028 to 0.1 M) using potassium perchlorate as the ionic medium, in the pH range from 5.3 to 7.7. The nickel(II) concentration was ~ 4 × 10–5 M. The [Ni2+]/[CN–] ratio was varied from 0.05 to 0.8 to obtain data so that the Job's method could be applied, and otherwise kept at around ¼. The time required for the equilibration ranged from a few days for solutions of higher pH to several weeks for those of lower pH. Acetate and phosphate buffers were used to maintain the pH, but correction was made only to account for the forma-tion of the acetato complex. Only formation of the complex 2

4Ni(CN) − was detected,


292

with a formation constant of 10 4log οb = (31.0 ± 0.1). The authors used HCNpK = 9.398, which is the average value of some early determinations. Since the above value is not correct, and 10 4log b depends on HCNpK to the fourth power, Christensen et al. [63CHR/IZA] recalculated the thermodynamic formation constant reported by Freund and Schneider, using the more reliable HCNpK ο = 9.216, which resulted in 10 4log οb = (30.3 ± 0.1). The method of this recalculation is unknown; therefore, a re-evaluation was carried out in the present review using the reported experimental data. Since the concentration of the acetate and phosphate buffers were not reported, corrections taking into account the acetato- and phosphato- complexes of nickel(II) was not possible. The values of HCNpK at different ionic strengths were calculated using the data reported in [92BAN/BLI]. The re-evaluation indicated that only 2

4Ni(CN) − formed in the equili-brated solutions, with 10 4log b = (29.39 ± 0.14) (I = 0.1 M), (29.46 ± 0.13) (I = 0.048 M), (29.79 ± 0.10) (I = 0.014 M), (30.12 ± 0.12) (I = 0.0028 M). For the purposes of this review, slightly larger uncertainties, ± 0.20, were estimated for each of these con-stants. The ionic strength range is too narrow to perform an SIT analysis, but these val-ues for 10 4log οb determined at low ionic strengths are nearly identical to the value of 10 4log οb selected in the present review.

[59KEN]

This paper reports cryoscopic measurements of NiCl2 solutions in saturated potassium chlorate (0.258 m, f.p. = 272.35 K) and potassium perchlorate (0.0485 m, f.p. = 272.99 K) solutions. The author determined the freezing point depression at different NiCl2 concentrations (in KClO3 2NiClc = 0.0193 – 0.121 m and ∆T = 0.10 – 0.61°C, in KClO4

2NiClc = 0.0234 – 0.1202 m and ∆T = 0.12 – 0.59°C), then extrapolated to 2NiClc = 0 at

the saturating ionic strength of the supporting electrode. The ionic strength contribution of NiCl2 is dominant in KClO4 solutions, therefore the extrapolation used by the author can be accepted, with an increased uncertainty, only for the data measured in KClO3 solutions. The value of 10 1log b = (0.23 ± 0.30) (– 0.8°C, I = 0.258 m KClO3) is ac-cepted in this review. From this value, 10 1log οb = (0.79 ± 0.40) (– 0.8°C) can be calcu-lated using the SIT, assuming the interaction coefficients ε(Ni2+, 3ClO− ) to be (0.28 ± 0.10) kg·mol–1 (see also [58KEN]) and ε(NiCl+, 3ClO− ) = 0.5 × {ε(Ni2+, 3ClO− ) + ε(K+, Cl–)} = (0.135 ± 0.100) kg·mol–1.

[59KIS/CUP]

Spectrophotometric measurements have been performed in the binary nickel(II)-cyanide and copper(I)-cyanide, and ternary nickel(II)-copper(I)-cyanide systems. It was con-cluded that 2

4Ni(CN) − and 46Ni(CN) − form in the binary nickel system. The equilibrium

constant for the reaction: 24Ni(CN) − + 2 CN– 4

6Ni(CN) −

was determined in 5 M potassium acetate ( 10log K = (3.33 ± 0.15)) and 10 M potassium nitrite ( 10log K = (3.42 ± 0.15)) solutions (the temperature of the measurements was not


293

reported). Today, the non-existence of 46Ni(CN) − is well established [60MCC/JON],

[62BEC/BJE], [63PEN/BAI], [64GEE/HUM], [65COL/PET], [71PIE/HUG]. According to Beck and Bjerrum [62BEC/BJE], the results reported in [59KIS/CUP] do not support the formation of a stable hexacyano complex, but rather of an unstable pentacyano spe-cies.

[59LAF]

The phase relationships between nickel monosulphide and Ni3S2 or Ni7S6 were studied by equilibrating NiS with hydrogen gas at temperatures ranging from 677.15 to 1006.15 K. The high-temperature form of NiS was in equilibrium with Ni7S6 from 677 to 823 K and with Ni3S2 (high-temperature modification) between 823 and 1006 K, re-spectively. The H2S(g) content of the gas, formed due to reduction of the monosulphide with hydrogen, was measured iodometrically and the composition of the solid phases was determined by X-ray diffraction analysis, magnetic susceptibility measurements, and chemical analysis. From the ratio of the partial pressures between H2S(g) and H2(g) the partial pressure of S2 can be easily calculated, leading to a temperature dependence of p(S2) which is in close agreement with the data reported by Rosenqvist [54ROS].

[59NAI/NAN]

The paper reports on a series of potentiometric measurements (0 to 45°C). The meas-urements were carried out in acidic media (HCl(aq)), and a Davies-type equation was used to calculate the activity coefficients (I < 0.05 m). The reported formation constant (A.15) for NiSO4(aq) at 25°C is (211 ± 24),

Ni2++ 24SO − NiSO4(aq). (A.15)

Sufficient data are provided to permit recalculation using the SIT and the TDB selected values [92GRE/FUG] for the first protonation constant (and enthalpy of proto-nation) for the sulphate ion. As in the original analysis, the initial concentrations, meas-ured potential of the cell:

H2(g), Pt| HCl(aq, m1), MSO4(aq, m2)|AgCl(s)|Ag(s),

the selected value for the equilibrium constant (I = 0) for:

H+ + 24SO − 4HSO− , (A.16)

and the selected activity coefficient equations (in this case the SIT equations) were used. A model assuming formation of a single 1:1 nickel sulphate complex was used, and calculations were carried out so that the ionic strength values in the activity coeffi-cient equations were identical to those calculated from the final species concentrations.

The values 10log K ο ((A.16), 298.15 K) = (1.98 ± 0.05) and r mH ο∆ ((A.16), 298.15 K) = (22.4 ± 1.1) kJ·mol–1 were those selected in a previous TDB review [92GRE/FUG]. From these, the association constants for Reaction (A.16) were calcu-lated assuming constant values for ,mpCο over the temperature range 0 to 45°C. As dis-


294

cussed below, this is probably not adequate. The values ε(Ni2+, Cl–) = (0.17 ± 0.02) and ε(H+, Cl–) = (0.12 ± 0.01) were from the same source (see also Section V.4.2.4). The value ε(H+, 4HSO− ) = (0.01 ± 0.02) is the value used by Grenthe et al. [92GRE/FUG] in their Appendix A discussion of [76PAT/RAM]. No specific value was adopted for ε(Ni2+, 4HSO− ), as this interaction was assumed to be incorporated in the value of the formation constant for the NiSO4(aq) complex. Preliminary calculations indicated that the reported data for experiment 8 at 298.15 K [59NAI/NAN] could not be used in re-calculations. The concentrations reported in Table 1 of the paper are not compatible with the reported complexation constant value or the reported final ionic strength.

The calculations showed that the values calculated for the formation constant of NiSO4(aq) are extremely sensitive to the activity coefficient model and to the values assumed for the first protonation constant for the sulphate ion. For example, for 298.15 K, the formation constant was recalculated to be (173 ± 64). However, if the value for 10log K ο (A.16) was changed from 1.98 to 1.96, well within the assessed un-certainty bound, the value became (140 ± 72). The uncertainties here are merely the 2σ uncertainties in the set of seven values recalculated from [59NAI/NAN], and do not include uncertainties in the auxiliary data. Even though the solute concentrations are low (I < 0.05), and the overall changes to the activity coefficients from the SIT interac-tion terms are small, the calculated value of 10log K ο (A.15) increases from (173 ± 64) to (201 ± 35) if all the ε values are set to zero (but the maximum change in

Cl−g from the interaction terms is 0.004).

In the present review, the value 10log K ο ((A.15), 298.15 K) = 173 is accepted, and an uncertainty of ± 80 is estimated. It is not clear whether the values of K(A.16) at temperatures other than 298.15 K are sufficiently well defined to generate usable values of K(A.15). The calculated values of K(A.15) at the six different temperatures can be used to calculate r mH∆ = – (13.8 ± 1.0) kJ·mol–1 if the enthalpy of reaction is assumed to be constant from 273 to 318 K. However, Figure V-38 suggests that the values of K(A.15) at the two lowest temperatures diverge from those obtained from conductance and cryoscopic studies.

[59ROB/STO]

The data m, γ± and φ of NiSO4 required in the plots of Figure A-5 and Figure A-6 were taken from Table 16, Appendix 8.10 of this textbook. The activity coefficient γ±(sat) has been calculated by a third order polynomial of ln γ± vs. m , as depicted in Figure A-5. For the activity of water, aW(sat), φ was extrapolated to the saturated solution molality with a fourth order polynomial in m, as depicted in Figure A-6.


295

Figure A-5: Extrapolation of activity coefficients to saturated solution at 25°C. Solid curve: calculated with a third order polynomial in m : ln γ± = A + B1·m0.5 + B2·m + B3·m1.5 with A = – (0.90807 ± 0.02571), B1 = – (3.60463 ± 0.09626), B2 = (1.42317 ± 0.10959) and B3 = – (0.07339 ± 0.3843), with R2 = 0.9998, and σ = 0.00699. ● data from [59ROB/STO], extrapolated value.

Figure A-6: Extrapolation of osmotic coefficients to saturated solution at 25°C. Solid curve: calculated with a fourth order polynomial in m: φ = A + B1·m + B2·m2 + B3·m3 + B4·m4; with A = 0.61907, B1 = – 0.48613, B2 = 0.46844, B3 = – 0.16846, B4 = 0.02584, and R2 = 0.99831 and σ = 0.00326. ● data from [59ROB/STO], extrapolated value, with m(NiSO4) = 0.7488 ;φ = 0.7488.

0.0 0.2 0.4 0.6 0.8 1.0 1.2 1.4 1.6 1.8-3.6

-3.4

-3.2

-3.0

-2.8

-2.6

-2.4

-2.2

-2.0

-1.8

-1.6

-1.4

ln γ ±

(m / mol·kg−1)0.5 (NiSO4)

-0.5 0.0 0.5 1.0 1.5 2.0 2.5 3.0

0.45

0.50

0.55

0.60

0.65

0.70

0.75

0.80

φ

m (NiSO4) / mol·kg−1


296

[59WIL/THR]

Hellyerite forms a colourful association with bright green, amorphous zaratite, Ni3CO3(OH)4·4H2O, along shear planes in serpentine containing heazlewoodite (Ni3S2) and other nickel sulphides. The crystals are of a cobalt or pale blue. The empirical for-mula was found to be Ni(CO3)1.18·5.67H2O. The mineral was named after Henry Hel-lyer, first Surveyor General of the Van Dieman’s Land Company, who was responsible for exploring and surveying much of North Western Tasmania during 1826 – 1830 (ap-proved by IMA 1959). Its rarity is probably due to the fact that hellyerite is relatively unstable and, if not kept in a cool, air-tight environment (Heazlewood is typically rather cold and damp), it decomposes to powdery zaratite-like phases, plus otwayite, Ni2(CO3)(OH)2·H2O, and other minerals.

[60LIS/ROS]

See comments under [66KEN/LIS].

[60MCC/JON]

See the discussion for [65COL/PET].

[60WIL2]

Williams discovered a new mineral at Heazlewood, Tasmania, and identified it as nickel hydroxide, Ni(OH)2, by its X-ray powder pattern.

[61CHU/ALY2]

A value of 10 ,0log sK for hydrated Ni3(PO4)2 at 20°C is reported as – 30.3. The value is based on sparse measurements, and no account is taken of complex formation between Ni2+ and 2

4HPO − or other phosphate species. The results from this study were not used further in the present review.

[61DAV/SMI]

Spectrophotometric measurements were performed at 242.5 – 252.5 nm to determine the association constant for the NiSCN+ species. The concentration of the metal ion was in five- to twenty-fold excess over that of the thiocyanate ion so that the kinetic data could be interpreted. The ionic strength was 0.5 M, the background electrolyte is not specified. The association constants were measured between 1.4 and 45°C, but only the value for 1.4°C is reported in the paper (although two interpolated values for 5.2 and 9.3°C are also given). Due to the lack of experimental details, the reported constants were not considered further in this review.

[61LEB/LEV]

Lebedev and Levitskii investigated the equilibrium of the reaction:

Ni2SiO4 + 2CO 2Ni + SiO2 + 2CO2


297

in the range 800 – 1100°C using a circulation technique with an automatic and a volu-metric gas analyser. The equilibrium of this reaction was approached both by reduction of nickel orthosilicate and by oxidation of nickel. Combining the experimentally deter-mined Gibbs energy change for the above reaction with the Gibbs energy for the reac-tion:

2 CO(g) + O2(g) 2 CO2(g),

the authors [61LEB/LEV] calculated the Gibbs energy for the reaction:

2Ni + SiO2 + O2 Ni2SiO4

in the range 800 – 1100°C. These r mGο∆ values are significantly different from results obtained by another direct equilibration technique [68CAM/ROE] and the emf-studies shown in Figure V-54. The results of Lebedev and Levitskii were, therefore, disre-garded by this review. A possible explanation of the observed discrepancy seems to be difficulty in attaining equilibrium for the reduction of nickel orthosilicate by carbon monoxide in the temperature range chosen by [61LEB/LEV]. The values of the standard enthalpy of formation f mH ο∆ (Ni2SiO4, olivine, 298.15 K) = – 1378.21 kJ·mol–1 and the Gibbs energy of formation f mGο∆ (Ni2SiO4, olivine, 298.15 K) = − 1265.24 kJ·mol–1 proposed by Lebedev and Levitskii were not used by this review and are significantly different from the recommended values.

[61LIS/WIL]

The emf of the cell Ag,AgBr|NaBr+M(ClO4)2+NaClO4|NaBr+NaClO4|AgBr,Ag (where M = Ni(II) or Co(II)) has been measured at I = 2 M (Na, M)ClO4 and between 273 – 323 K in order to determine the composition and stability of the nickel(II) and cobalt(II) complexes formed with bromide ion. The concentration of nickel(II) ranged from 0.08 to 0.31 M, while the chloride concentrations were nearly constant ([Cl–] ~ 0.008 M). The complex formation induced 1 to 6.5 mV differences between the electrodes. The average value of 10 1log b at 298.15 K is – (0.12 ± 0.02). This value has been accepted with an increased uncertainty (± 0.40). From the temperature dependence of 10 1log b ,

r mH∆ appears to fall from about 13 to 2.7 kJ·mol–1 over the range 0 to 50°C (see Figure A-7). The authors could not explain this behaviour, which is probably due to some uni-dentified experimental error. Assuming constant r mH∆ between T = 273 – 323 K, and that the medium effect is additive in the logarithmic scale, from these data r mH ο∆ for the reaction,

Ni2+ + Br– NiBr+ (A.17)

can be obtained as a rough estimation, r mH ο∆ (A.17) = (8 ± 5) kJ·mol–1.


298

Figure A-7: Plot of 10 1log b (A.17) reported in [61LIS/WIL] vs. 1/T.

[61MAR/BYD]

The kinetics of the reaction:

Ni(edta)2– + 4CN– 24Ni(CN) − + edta4– (A.18)

have been studied by spectrophotometry (285 nm) at 25°C in I = 0.11 and 1.01 M NaCl solutions around pH 11. From the kinetic data, 10 4log b (A.18) = 31.52 was derived at I = 1.01 M, using HCNpK = 9.36. Although, the latter pK is rather far from the correct value [92BAN/BLI], considering the pH used, the error in HCNpK has little effect on the reported 10 4log b (A.18) value.

[61MOH/DAS]

The thermodynamic stability constant of the NiSCN+ complex was determined by spec-trophotometric (t = (25 ± 1)°C) and potentiometric (t = 35°C) methods. The Davies equation was used to calculate the activity coefficients, and this is not compatible with the SIT. Therefore, the potentiometric data were re-evaluated, using the SIT approach. The resulting 10 1log b – Im plot is rather scattered and the ionic strength range was rela-tively narrow (see Figure A-8). Therefore the reported data were not considered further in this review.

-0.4

-0.35

-0.3

-0.25

-0.2

-0.15

-0.1

-0.05

0

0.003 0.0031 0.0032 0.0033 0.0034 0.0035 0.0036 0.0037

K / T

log 1

0 β1


299

Figure A-8: 10 1log b – Im plot obtained from the potentiometric data (t = 35°C) reported for the Ni2+ – SCN– system in [61MOH/DAS].

[61SHC]

Shchigol synthesised nickel orthoborate by reacting nickel hydroxide or carbonate with orthoboric acid and determined its stoichiometric composition analytically. No attempt was made to characterise this substance structurally. First, the solubility of nickel ortho-borate, Ni(BO2)2·4H2O(s), was studied at 22°C by equilibrating the solid phase with HCl solutions. Second, solubility measurements in KOH solutions were interpreted by assuming simultaneous equilibrium between Ni(BO2)2·4H2O(s), Ni(OH)2(s) and the aqueous medium. A third method consisted of precipitating nickel nitrate with borax solutions. It should be mentioned that neither the solid nickel borate nor the precipitated nickel hydroxide were structurally characterised. An average value of the solubility product of Ni(BO2)2·4H2O(s) ,0sK = 2.16×10–9 mol3·dm–9 was obtained by these three methods. A recalculation of the solubility product was based on the experimental data obtained from undersaturation (method 1) and oversaturation (method 3). The acid dis-sociation constant of boric acid was taken to be 10log aK ο = – 9.623 at 22°C [58HAR/OWE] and the Davies equation [62DAV] was used for extrapolation to I = 0. A value of 10 ,0log sK ο = – (8.80 ± 0.19) was found in reasonable agreement with Shchi-gol’s result ( 10 ,0log sK ο = – 8.666), see Figure A-9.

Shchigol found that the solubility of Ni(BO2)2·4H2O in aqueous media in-creases with increasing concentration of orthoboric acid. In electrolysis experiments the

1.9

2

2.1

2.2

0 0.1 0.2 0.3 0.4 0.5

I / mol˙kg-1

log 1

0 β1 +

4D


300

solution of the anode compartment was enriched in nickel indicating its presence in the complex borate anions.

The formation of the nickel and cobalt borate complexes in solution was be-lieved to occur according to the reaction:

M(BO2)2(s) + nH3BO3(aq) HnM(BO2)2+n(aq) + nH2O(l) (A.19)

The value n calculated from the solubility measurements on Ni(BO2)2·4H2O(s) at different concentrations of H3BO3 was calculated to be 0.7. Thus the co-ordination number for Ni2+ relative to the metaborate ion was assumed to be 3 and the proposed complex species was 2 3Ni(BO )− . While Figure A-10 shows that for nickel n = 1 and for cobalt n = 2, Shchigol’s conclusion that the dissolved Ni2+ essentially consists of

2 3Ni(BO )− must be refuted. Where do the balancing cations come from ? When it was attempted to explain the data by the formation of a neutral species like NiH(BO2)3, the solubility product did not remain constant. Additional experiments (careful pH meas-urements, reagents free of protolytic impurities) are necessary to establish the formula and stability of nickel borate complexes. After standing 5 to 6 days, fine crystals of nickel hexaborate NiB6O10 were precipitated from the complex nickel borate solution containing an excess of orthoboric acid possibly according to the reaction:

HNi(BO2)3(aq) + 3H3BO3(aq) NiB6O10(s) + 5H2O(l) (A.20)

Figure A-9: Solubility product of Ni(BO2)2·4H2O(s) at 22°C from data reported in [61SHC].

oversaturation, exp. slope: – 1.39, undersaturation, exp. slope both series: – 1.06 ; theor. slope: – 2.00

-3.5 -3.4 -3.3 -3.2 -3.1 -3.0 -2.9

-2.8

-2.6

-2.4

-2.2

-2.0

-1.8

log

a (N

i2+)

log a (BO2-)


301

Figure A-10: Nickel and cobalt borate complexes from data in [61SHC].

M2+ : Co2+, Co(BO2)2·2H2O(s) + 2 H3BO3(aq) CoH2(BO2)4(aq) + 4 H2O(l), M2+ : Ni2+, Ni(BO2)2·4H2O(s) + H3BO3(aq) NiH(BO2)3(aq) + 6 H2O(l), ([Ni2+]tot – [Ni2+]) = f ([H3BO3].

[61TAN2]

The standard potentials of cobalt and nickel in methanol were determined using the fol-lowing cell without liquid junction:

Ni | NiCl2, solvent | Hg2Cl2 | Hg.

Whereas for the cobalt cell measurements were carried out successfully in pure water, methanol was the solvent for the nickel cell. Apparently the author experienced the same difficulties with nickel in the aqueous medium which were already described by [29HAR/BOS]. Thus, the results of this work provide no relevant information on thermodynamic data of Ni2+ in aqueous solution.

[61TAN/OGI]

Polarographic measurements were carried out in an acetate buffer (the assumed “pH” was 3.66), and concentrations of sulphate and acetate were varied in the presence of 1.00 × 10–3 M nickel nitrate and 1.09 × 10–3 M nitrilotriacetate at a constant ionic strength of 0.2 M (maintained with potassium nitrate). A dropping mercury electrode was used, and polyoxyethylene lauryl ether (1–2 × 10–6 M) was used as a maximum suppressor. Although results are reported for 15, 25 and 35°C, the assumed “pH” value for the buffer at 15 and 35°C is not stated. The acetate molarities were varied from 0.025 to 0.100 M at 0.03 M sulphate, and the potassium sulphate molarities were varied from 0.01 to 0.05 M at 0.05 M acetate. The results indicated that for the experimental

-0.1 0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8

0.000

0.005

0.010

0.015

0.020

0.025

0.030

0.035

[M2+

] tot /

mol

·dm

-3

[H3BO3] /mol·dm-3

t = 22°C


302

conditions only a monosulphato complex was formed for Ni(II), with no evidence for higher complexes or a mixed sulphato-acetato complex. Values of (14.3 ± 0.4), (11.6 ± 0.9) and (14.6 ± 0.7) dm3·mol–1 were reported for K1 at 15, 25 and 35°C, respec-tively. The reported statistical uncertainties are undoubtedly underestimates of the true uncertainties. For the purposes of the present assessment, the value 10 1log K = (1.06 ± 0.20) at 25°C is used; the measurements at 15 and 35°C are rejected because of the lack of experimental details. At this low ionic strength, assuming that the ε value for the Ni2+–acetate interaction is the same as the value for the Ni2+– 3NO− interaction has little effect on the value calculated for I = 0 using the SIT, and 10 1log K ο = (2.13 ± 0.20) is obtained.

[62BEC/BJE]

The effect of different salts on the visible spectra of tetracyanonickelate(II) has been investigated, in order to elucidate the formation of further complexes, and to solve the discrepancy between two earlier papers, [59BLA/GOL] and [60MCC/JON]. All salts (KCN, KSCN, KI, KBr, KCl, KF, KNO3) added to 2

4Ni(CN) − caused an increase in absorbance around 375 – 500 nm, but the effects were rather different. The data qualita-tively confirmed the formation of four very weak complexes: 3

5Ni(CN) − , Ni(CN)4SCN3–

Ni(CN)4I3–, and to a much lesser extent Ni(CN)4Br3–. Numerical estimates of their sta-bility were not provided.

The above-mentioned salts and some other compounds (NH3, pyridine, ethyl-enediamine, butylamine, ethanol, methanol, dioxane, dimethylformamide) were also added to solutions containing 0.04 2

4Ni(CN) − and 0.98 M KCN. Relatively small con-centrations of all these compounds result in a decrease of absorbance between 375 – 500 nm. Since ammonia has nearly the same effect as fluoride ion or dioxane, the au-thors rejected the possibility of further complex formation (such as Ni(CN)5F4–), and explained the observation as a strong decrease of activity coefficients corresponding to the mass action expression:

F = 2 54 3Ni(CN) CN Ni(CN)

/f f f− − − .

[62CON/GIL]

Conway and Gileadi carried out electrochemical kinetic studies at the nickel oxide (NiII–NiIII) electrode. A polarisation decay method was employed in which the reversi-ble potential was approached both from the anodic and the cathodic directions.

[62FAU/CRE]

Dissociation of some unstable chloro complexes has been studied by cryoscopic titra-tion in saturated KNO3 (I = 1.38 m, f.p. = 270.34 K). Using [Ni(NO3)2] = 0.02 M and [KCl] = 0.05 – 0.25 M; the freezing-point depressions were found to be between 0.065 and 0.393 K. The recalculation of the reported data using non-linear regression resulted in a value ( 10 1log b = (0.36 ± 0.06)) almost identical to that reported by the authors


303

( 10 1log b = (0.38 ± 0.08)). The value of 10 1log b = (0.36 ± 0.30) is accepted for use in the present assessment. Assuming ε(NiCl+, 3NO− ) equal to 0.5((ε(Ni2+, 3NO− ) + ε(K+, Cl–)) = 0.09 kg·mol–1 the extrapolation using the SIT resulted 10 1log οb = (1.17 ± 0.40) (– 2.0°C).

[62FRO/LAR]

The frequency and intensity of the C–N vibration have been measured for the SCN– ion (at 2066 cm–1) and some MSCN+ complexes of 3d metal ions. In the presence of metal ions, new peaks at ~ 2100 cm–1 appeared in all systems, and these peaks correspond to formation of the metal complexes. The intensities of the free and bound SCN– ion were used to determine stability constants of the monothiocyanato complexes which formed. The NaSCN concentration was 0.36 M, [Ni2+]tot varied from 0.44 to 1.32 M, and (2.64 – 2[Ni2+]tot) M NaClO4 was used as a background electrolyte. Therefore, during the meas-urements, the ionic strength ranged from 3.44 to 4.32 M. Moreover, the temperature of the measurements was not reported in the paper. Although, the reported stability con-stant for the NiSCN+ complex ( 10 1log b = 1.18) is close to the value in other papers, taking into account the uncertain experimental conditions, the above value was not con-sidered further in this review. Based on the comparison of different evaluation methods of the experimental data, the authors suggested the formation of 80% inner-, and 20% outer-sphere complexes.

[62KUL/YUN]

A systematic study of the phase relations in the binary system Ni – S was performed in the temperature range between 473 and 1303 K. The phase equilibria were investigated by means of quenching experiments, differential thermal analysis (DTA) and, to a lesser extent, high temperature X-ray powder diffraction. The phases present in the quenched samples were detected by application of X-ray diffraction analysis and reflecting mi-croscopy. In the latter case the samples were polished and etched with concentrated nitric acid. The following stoichiometric phases were found by the authors: Ni3S2 (heazlewoodite), Ni7S6 (low temperature modification), NiS (millerite, rhombohedral structure), Ni3S4 (polydymite), and NiS2 (vaesite). The high temperature modifications of Ni3S2, Ni7S6, and NiS (hexagonal NiAs-type structure) were reported to show a non-stoichiometric homogeneity range. Whereas the homogeneity range of Ni7S6 was fairly narrow, the high temperature phases Ni1–xS and Ni3+xS2 showed large stability fields in the respective phase diagram. In the sulphur-rich portion of the system a two-liquid phase field over a wide range of composition was revealed. The authors proposed a monotectic reaction between NiS2 and the two immiscible liquids at 1264 K.

[62LOP]

In this work a similar set-up to the one described previously by [52CAR/BON] was used for the measurement of the Ni2+ | Ni standard electrode potential. To avoid the formation of a Ni(OH)2 surface layer on the Ni electrode the NiSO4 solution was acidi-


304

fied with H2SO4. According to other authors the nickel electrode corrodes in acid solu-tion [52CAR/BON]. As a consequence the result, E°(Ni2+ | Ni, 295.15 K) = − (0.248 ± 0.004) V, (recalculated with the CODATA value for the saturated calomel electrode and corrected for an obvious misprint in the error limits) has been rejected.

[62RIN]

A direct search was made for the spinel polymorph of Ni2SiO4(cr) using high-pressure techniques. An olivine – spinel transition was discovered at 650°C and (18000 ± 5000) bars. This result is in excellent agreement with the theoretical prediction. This enables the free energy of transition and the density of the closer packed polymorph to be found. From these data the transition pressure can be calculated.

[62TRI/CAL]

The decrease of the kinetic wave heights of the Ti(IV)-SCN complex at a Hg dropping electrode, in the presence of a second cation, allowed the stability constants of Ni(II), Co(II), Cd(II) and Mn(II) with thiocyanate ions to be determined at different ionic strengths (I = 0.7 – 5 M (Na,H)ClO4). Due to the weak complex formation, this method could not be applied to the detection of formation of the NiCl+ complex. Therefore, the authors also determined +

1 (NiSCN )b in a solution of 1.5 M (Na,H)ClO4 by means of spectrophotometry, using the UV-band at 275 nm, attributed to the NiSCN+ species. Then the decrease of the absorbance at 275 nm with increasing chloride concentration allowed calculation of +

1 (NiCl )b . The following values were reported for the NiSCN+ complex.

Table A-10: 10 1log b versus Im.

Im 0.73 1.05 1.62 2.21 2.84 6.58

10 1log b (1.21 ± 0.02) (1.17 ± 0.02) (1.12 ± 0.02) (1.10 ± 0.04) (1.11 ± 0.02) (1.27 ± 0.05)

Several experimental details are not given in the publication: (i) the tempera-ture of the measurement, (ii) the chloride concentration used to determine +

1 (NiCl )b (Chemical Abstracts reported [Cl–] < 1.3 M). The composition of the background elec-trolyte ((Na, H)ClO4) is also dubious. Therefore, the reported data were not taken into account in this review.

[62VAG/UVA]

Vagramian and Uvarov carried out emf measurements in the range of 18 ≤ t ≤ 250°C with the cell:

Ni | NiSO4(0.501 mol·kg–1) | Hg2SO4 | Hg.


305

Between 200 and 250°C the potential became a linear function of the tempera-ture and was extrapolated to 25°C. The resultant E°(298.15 K) = 0.952 V combined with the standard potential of the Hg2SO4 | Hg electrode ((0.6127 ± 0.0030) V, NEA-TDB auxiliary data) would lead to E°(Ni2+ | Ni, 298.15 K) = – 0.339 V. The value originally given (– 0.270 V) is based on an obsolete standard potential value for the Hg2SO4 | Hg electrode (0.682 V). Finally, this result has been rejected due to the ques-tionable calculation of the NiSO4 activity at t > 200°C and the rather long extrapolation to 25°C, although the slope (dE°/dT) of the solid straight line in Figure 1 of [62VAG/UVA] can be predicted accurately from the quantities selected by NBS and CODATA [82WAG/EVA] or using NEA-TDB auxiliary data.

[62WIL2]

The dissociation constants of CoSCN+ and NiSCN+ have been determined from UV spectrophotometric measurements at 25°C. The first set of data was measured at con-stant ionic strength (I = 0.417 M (Na)ClO4) and was used to calculate the molar absorp-tion coefficients of the MSCN+ complexes. A second set of data, measured at varying ionic strength, yielded the thermodynamic dissociation constants of thiocyanate com-plexes. The necessary activity coefficients were calculated from the Davies equation. The reported dissK ο for the NiSCN+ complex is (0.0173 ± 0.0002), which is equivalent to

10 1log οb = (1.762 ± 0.005). Since the activity model used by the author is not equivalent to the SIT treatment, the above value was not selected. Instead, the reported experimen-tal data were re-evaluated for the purposes of this review. The recalculation of the first data set (measured at I = 0.417 M (Na)ClO4) resulted in the value of 10 1log b = (1.21 ± 0.08), and ε (NiSCN+, 273 nm) = (149 ± 5). Using latter value for ε (NiSCN+, 273 nm), 10 1log οb = (1.77 ± 0.10) can be calculated from SIT analysis of the second data set.

[63BOL/JAU]

The hydrolysis of nickel(II) perchlorate ([Ni2+]t = 2 – 20 mM) in 0.25 M and 1.0 M Na-ClO4 media has been investigated by potentiometric titration between – log10[H+] = 7 and 7.7 and at temperatures ranging from 25 – 50°C ( ± 0.05°C). Only the formation of NiOH+,

Ni2+ + H2O NiOH+ + H+ 1,1*b ,

was taken into account to explain the experimental data. The values of 10 1,1*log b =

− 9.76, r mG∆ (298.15 K) = (13.25 ± 0.10) kcal·mol–1 and r mH∆ (298.15 K) = (7.76 ± 1.00) kcal·mol–1 are reported, probably for both ionic strengths used, since no “sensible” dependence of 10 1,1

*log b with ionic strength is reported.

Under the conditions used, a substantial amount of 44 4Ni (OH) + would be ex-

pected to form; therefore, the reported experimental data have been re-evaluated for the purpose of this review, taking into account the formation of the mononuclear and


306

tetranuclear hydroxo species. Considering the formation constants reported by Baes and Mesmer [76BAE/MES], the presence of the dinuclear Ni2OH3+ species can be neglected under the conditions used. The recalculation resulted in the following constants (2σ uncertainties).

Table A-11: Recalculated constants.

t (°C) / I (M)

25 / 0.25 30 / 1.0 40 / 1.0 50 / 1.0

10 1,1*log b – (9.84 ± 0.02) – (9.74 ± 0.02) – (9.64 ± 0.06) – (9.68 ± 0.20)

10 4,4*log b – (27.34 ± 0.08) – (27.12 ± 0.10) – (26.41 ± 0.20) – (25.26 ± 0.20)

Taking into account the limited experimental data available, the uncertainty ± 0.30 is accepted for all the above values. From the temperature dependence of the formation constants (see Figure A-11 and Figure A-12) 10 1,1

*log b = – (9.73 ± 0.30) and 10 4,4

*log b = – (27.68 ± 0.40) can be extrapolated for t = 25°C and I = 1.0 M. The tem-perature dependence of 10 1,1

*log b is too small, compared to the experimental error, to derive a value for the enthalpy of reaction, but r mH∆ (298.15 K, 1.0 M) = (174 ± 15) kJ·mol–1 can be obtained for the formation of 4

4 4Ni (OH) + .

Figure A-11: Temperature dependence of 10 1,1*log b values derived from

[63BOL/JAU].

0.0031 0.0032 0.0033-10.5

-10.0

-9.5

-9.0

-8.5

log 10

*β 1,

1

K / T


307

Figure A-12: Temperature dependence of 10 4,4*log b values derived from [63BOL/JAU].

[63BUR/ABB]

This study is similar to [61LEB/LEV]. The equilibrium of the reaction:

Ni2SiO4 + 2CO 2Ni + SiO2 + 2CO2

was investigated in the temperature range 700 – 1200°C. Results obtained by Burdese et al. do not coincide with results of other studies discussed in Section V.7.2.1.1.5 and shown in Figure V-54. Thus, they were not used for the evaluation of thermodynamic quantities selected by this review.

[63CHR/IZA]

The thermodynamic formation constant and heat of formation of 24Ni(CN) − in aqueous

solution have been determined at 25°C potentiometrically and calorimetrically, respec-tively. No background electrolyte was used, and the Debye-Hückel equation was ap-plied to extrapolate the experimental thermodynamic data to zero ionic strength. Al-though, the activity model used is not compatible with the SIT, most of the experiments referred to I < 0.007 M; therefore, the reported value ( r mH ο∆ = – 180.7 kJ·mol–1 at 25ºC) is accepted as valid at I = 0, but the uncertainty is estimated in this review as ± 4.0 kJ·mol–1.

0.0031 0.0032 0.0033-28

-27

-26

-25

-24

log 10

*β 4,

4

K / T


308

[63GEO/FER]

The authors measured the heat capacity from 1120 to 1919°C and the heat of nickel fusion by a calorimetric method. The reported heat of fusion is fus mH ο∆ = (17.47 ± 0.23) kJ·mol–1.

[63ISA]

The X-ray diffraction patterns of hydrothermally synthesised anhydrous nickel carbon-ate were taken, indexed and the constants of the unit cell ascertained. The T, p phase diagram of NiCO3(cr) was determined semi-quantitatively. Solid solution formation between NiCO3 and MgCO3 was established. In addition hellyerite, NiCO3·5.5H2O, and zaratite were investigated. The latter turned out not to be a single mineral, but a com-posite of amorphous and fibrous components.

[63KO/HEP]

Ko and Hepler claimed that S. Friedberg (Department of Physics, Carnegie Inst. Tech-nol., 1961) gave them mS ο (NiCl2·6H2O, cr) = 344.3 J·K–1·mol–1 “from unpublished heat capacity data”. No subsequent article containing these heat capacity data was found. Hence, no documentation of this value can be provided for.

[63LIN/LAF]

The partial pressure ratio p(H2S)/p(H2) of the gas phase in equilibrium with the high-temperature modification of Ni3S2 was determined as a function of composition at 853.15 and 933.15 K, respectively. The same experimental procedure was performed as described by Laffitte [59LAF]. The ratios p(H2S)/p(H2) for the coexistence of Ni3+xS2 with metallic nickel at 853.15 and 933.15 K are listed in Table A-12. The experimental values coincide well with those predicted by the present compilation of thermodynamic properties of nickel sulphides.

Table A-12: Comparison between experimental data [63LIN/LAF] and calculated val-ues of log10[p(H2S)/p(H2)] for the coexistence of Ni3+xS2 with metallic nickel.

T/K log10[p(H2S)/p(H2)], exp. log10[p(H2S)/p(H2)], calc.

853.15 – 3.04 – 3.03

933.15 – 2.90 – 2.79

[63NET/DRO]

The stability constants of NiCl+ and NiBr+ complexes have been determined by a spectrophotometric method at an ionic strength of 5.7 M (H(X, ClO4), where X = Cl– or Br−). The concentration of the complex forming anion was varied from 0 to 5.44 M. The calculation of the formation constants was based on the slight changes of extinction


309

coefficients at 395 and 400 nm. The values of the formation constants accepted in this review at I = 5.7 M for the NiCl+ and NiBr+ complexes are 10 1log b = – (0.5 ± 0.4) and – (0.3 ± 0.5), respectively.

[63PEN/BAI]

See discussion for [65COL/PET].

[63PHI/HUT]

Phillips et al. investigated the phase equilibria in the system NiO – Al2O3 – SiO2 by quenching and direct observational techniques. The crystalline phases were identified by petrographic microscopy and X-ray diffraction methods. The equilibrium diagram of the system NiO – SiO2 drawn from the experimental data shows that the only intermedi-ate compound is Ni2SiO4, which has the olivine structure. The authors found, that unlike other olivines which melt congruently, nickel olivine has an upper temperature of stabil-ity 1545°C and at temperatures 1545 and 1650°C, NiO and SiO2(cristobalite) coexist in equilibrium.

[63SHA/DES]

Solutions of unspecified nickel(II) salts ([Ni]tot = 0.029 to 0.0855 M) were titrated with NaOH solution at (28.0 ± 0.5)°C at an ionic strength of 1.0 M using NaClO4 as the ionic medium. The pH values of the solutions were determined by unspecified “external elec-trodes”, which were calibrated using a single buffer solution (pH = 7). The added vol-ume of NaOH solution was measured in “drops”, assuming 20.6 drops per cm3. The uniform drop-size over the range of 0 – 25 cm3 is questionable. The results were ana-lysed assuming the formation of NiOH+ as the sole hydrolytic species ( 10 1,1

*log b = − (10.03 ± 0.08)), although under the conditions used, a considerable amount of

44 4Ni (OH) + would be expected. Therefore we re-evaluated the reported data, assuming

(i) only the presence of 44 4Ni (OH) + ( 10 4,4

*log b = – (27.85 ± 0.3)) and (ii) the formation of both NiOH+ and 4

4 4Ni (OH) + ( 10 1,1*log b = – (10.03 ± 0.04) and 10 4,4

*log b = − (30.2 ± 0.7)). Both calculations revealed substantial errors in the measured pH, proba-bly due to the relatively low precision of the experimental work. Therefore, these con-stants were not considered further in this review. The correction applied by Plyasunova et al. in [98PLY/ZHA] for the data in [63SHA/DES] resulted in 10 1,1

*log b = − (10.39 ± 0.17), 10 1,2

*log b = − (10.4 ± 0.7) and 10 4,4*log b = – (28.2 ± 1.0). Even these

constants do not provide a better description of the experimental data.

[63TAN/SAI]

This work is a continuation of an earlier study [61TAN/OGI]. Polarographic measure-ments were carried out at 25°C in an acetate buffer (the assumed “pH” was 3.72), and concentrations of sulphate and acetate were varied. The sulphate complexation constant was then obtained by accounting for acetate complex formation. At this higher ionic strength, the reported nickel sulphate association constant was K1 = 3.7 dm3·mol–1. This


310

was determined using an extrapolation procedure that should have minimised any effect of acetate. The SIT is used with ε(Ni2+, 4ClO− ) equal to (0.37 ± 0.03) kg·mol–1 and

10 1log K ο = 1.95 is calculated, with an estimated uncertainty of ± 0.20.

At the higher acetate concentrations (to 0.5 M) and sulphate concentrations (to 0.20 M) there was evidence for formation of 2

4 2Ni(SO ) − (K = 26 dm6·mol–2), Ni(OAc)(SO4)– (K = 1 dm6·mol–2) and 3

4 2Ni(OAc)(SO ) − (K = 3 dm9·mol–3). There is insufficient information to re-analyse the experimental data, and the values of the higher complexation constants are not accepted in the present review.

[63THR]

In the course of structure refinement Threadgold points out that the water content of natural hellyerite corresponds to the composition NiCO3(H2O)4·1.5H2O, whereas the nickel content indicates NiCO3(H2O)4·2H2O [59WIL/THR]. From the packing model of the structure it was concluded that it is not possible for hellyerite to contain two water molecules per formula unit in the interlayer water sheet. Consequently, the agreement of the diffraction pattern with Rossetti-François’ NiCO3(H2O)4·2H2O [52ROS] was con-sidered a problem. It should be emphasised, however, that a new synthesis also resulted in crystals with the empirical formula NiCO3(H2O)4·1.5H2O [2001GAM/PRE], [2002WAL/PRE].

[64ALC/BEL]

This work deals with the experimental determination of the oxygen activity in lead melts by performing emf measurements over the temperature range from 820 to 1120 K. Galvanic cells of the type Pb(l) | YSZ | NiO, Ni were applied, where the Ni / NiO – oxygen buffer served as the reference electrode. The Gibbs energy function for the ref-erence electrode was obtained from measurements of the solid-gas equilibrium:


provided by Tomlinson and Young (quoted as a private communication). Tomlinson and Young determined the CO/CO2 ratios in equilibrium with Ni/NiO by chemical analysis as a function of temperature. They obtained the following Gibbs energy func-tion for the reaction mentioned above: r mG∆ (T) = (– (47.15 ± 0.71) + (0.293 ± 0.711) × 10–3 (T/K)) kJ·mol–1. Based on this equation, Alcock and Belford derived for the Gibbs energy function of formation for NiO: f mG∆ (T) = (– 235.27 + 0.08682 (T/K)) kJ·mol–1, 820 < T/K < 1120, with a standard deviation of ± 1.1 kJ·mol–1.

[64GEE/HUM]

The authors repeated (and extended) the measurements reported by Blackie and Gold in [59BLA/GOL]. It was found that the solutions, for which the ratio of cyanide to nickel(II) was between 4 and 8, have the same relaxation time as the solvent (both in 17 M ammonia solution and in water), in sharp contrast with the findings of Blackie and


311

Gold [59BLA/GOL]. At a much higher excess of cyanide, a marked change in the UV-VIS spectra, between 300 – 520 nm, was observed. This spectral change was used to determine the stability constant for the reaction:

24Ni(CN) − + CN– 3

5Ni(CN) −

at 23°C and I = 2.5 M Na(ClO4). The concentration of the tetracyanonickelate(II) ion ranged between 0.4 and 70 mM, while the concentration of cyanide ion varied from 0.05 to 2.5 M. The reported stability constant ( 10 5log K = – (0.69 ± 0.02)) was not cor-rected to 25°C, as the difference is assumed to be negligible.

[64GRI/LIB]

Equilibria established in a solution containing a cation (Cu2+, Co2+, Ni2+ or Cd2+) and chloride have been studied by means of chromatography on ion-exchange papers. Com-plex formation constants were determined for Ni2+ at 25°C and I = 3 M H(ClO4, Cl). The concentration of nickel(II) was 0.4 M, while the chloride concentration varied be-tween 0 – 3.0 M. The reported constants are β1 = (0.23 ± 0.03) M–1; β2 = (0.06 ± 0.03) M–2; β3 = (0.001 ± 0.006) M–3. The authors stated that the results are less accurate then ones obtained by means other standard methods. Due to the considerable change in the ionic medium during the measurements, only the formation of the first complex is con-sidered in this review. After recalculation of the experimental data using non-linear re-gression, a value of 10 1log b = – (0.45 ± 0.50) is accepted.

[64HAM/MOR]

Complexation concentration quotients for 25°C and a medium of 0.1 M (CH3)4NCl(aq) were reported for the formation of 2

2 7NiP O − ( 10log K (A.21) = 6.98) and 2 7NiHP O− ( 10log K (A.22) = 3.83).

Ni2+ + 42 7P O − 2

2 7NiP O − (A.21)

Ni2+ + 32 7HP O − 2 7NiHP O− (A.22)

The ligand solutions were prepared by titration of the free acid with tetramethylammonium hydroxide (to a pH value of approximately 10). Aliquots containing the ligand, metal (as Ni(NO3)2), and tetramethylammonium chloride were titrated with 0.0685 M HCl(aq), with the primary measurements being carried out at pH values between 4.5 and 6.5. The electrodes were probably calibrated against pH buffers rather than concentration standards, and the hydrogen ion activity coefficient, +H

g was assigned the value of 0.796; the SIT procedure described in Appendix B would have led to essentially the same value, +H

γ = 0.80.

For Reaction (A.21) in 0.1 M (CH3)4NCl, 10log K (A.21) = 6.98 and 10log K ο (A.21) = (8.73 + 0.1 ∆ε).

For Reaction (A.22) in 0.1 M (CH3)4NCl, 10log K (A.22) = 3.83 and 10log K ο (A.22) = (5.14 + 0.1 ∆ε).


312

As the ionic strength is low, the effect of the ∆ε term on the final calculated values of 10log K ο is limited. Rather than use estimated ε values for interactions involv-ing the pyrophosphate moieties (some of which are highly charged), it is simply as-sumed that ∆ε is zero. As a consequence, the uncertainty in the 10log K ο values is in-creased to ± 0.25 for the values from these measurements in 0.1 M Me4NCl solutions.

[64KOH/ZAS]

Thermal decomposition of hydrated nickel sulphates was carried out as a function of temperature at 0.026 bar (“20 torr”). Samples were equilibrated at constant temperature (± 0.1 K) and pressure for at least 24 hours. Equations were reported for water vapour pressures as a function of temperature for mixtures of hydrates (NiSO4·7H2O/NiSO4·6H2O; NiSO4·6H2O/NiSO4·4H2O; NiSO4·4H2O/NiSO4·2H2O; NiSO4·2H2O/NiSO4·1H2O), and for a saturated solution (18 to 59°C). Rehydration of the monohydrate at 60°C with water vapour pressure of 0.078 bar was reported to gen-erate a mixture of the mono-, di- and tetrahydrates. Decomposition of the monohydrate led to an amorphous solid, and vapour pressure measurements over mixtures of the an-hydrous solid and the monohydrate were not reproducible.

The authors indicated that they did not feel it was useful to use their empirical equations to calculate dissociation enthalpies and entropies, and values calculated from the equations are not used in the present review.

[64KOS/KAL]

The temperature functions of the heat capacities of hydrothermally synthesised calcite type carbonates MnCO3, FeCO3, CoCO3, and NiCO3 were investigated between 70 and 300 K. The respective standard entropies were calculated by the integration of ,mpC /T versus T curves.

[64LAR]

The author used infrared spectroscopic results to conclude that only about 10% of the nickel sulphate ion pairs are inner-sphere complexes in 0.3 M and 1 M NiSO4. The de-rived association constants are useful for comparisons for the different di- and trivalent systems studied, but are not suitable for calculating a selected value in the present re-view.

[64LEE/ROS]

The sulphur pressures (sulphur activities) for the equilibrium between NiS2 and Ni1–xS were studied at temperatures ranging from 673 to 873 K. Pure NiS2 was prepared by solid state reaction between sulphur and NiS. The solid product was checked by X-ray diffraction analysis before and after the decomposition experiments. The equilibrium sulphur pressure for the decomposition of NiS2 into Ni1–xS and S2(g) was measured by means of a Rodebush manometer [30ROD/HEN].


313

[64PER]

The first hydrolysis constant ( 1*K ≡ 1,1

*b ) of nickel(II) was determined at 20°C by po-tentiometric titrations using various ionic strength and nickel(II) nitrate concentrations: 0.0016 M ≤ I ≤ 1.503 M, 0.00050 M ≤ [Ni(II)]tot ≤ 0.01000 M. In another set of titra-tions, the temperature was varied from 15 to 42°C, applying three different ionic strengths: 0.0016 ([Ni(II)]tot = 0.0005 M), 0.0130 and 0.0430 M ([Ni(II)]tot = 0.001 M). In all titrations, KNO3 was used to adjust the ionic strength. The author assumed that the only important hydrolysed species was NiOH+. For the experiments made at [Ni(II)]tot ≤ 0.001 M, the formation of 4

4 4Ni (OH) + (and Ni2(OH)3+) can, indeed, be ignored, in the pH range of the experiments. The formation of 3NiNO+ was not considered by the au-thor, although at 20°C, where a wide range of ionic strength (I = 0.0016 – 1.5) was used, its concentration can not be neglected. Therefore, for the purposes of this review, the reported constants were corrected for the formation of 3NiNO+ , and the SIT ap-proach was used to derive 10 1,1

*log οb (see Figure A-11, page 306). From the plot in Figure A-13, for the reaction:

Ni2+ + H2O NiOH+ + H+ (A.23)

10 1,1*log οb = – (10.07 ± 0.07) and ∆ε = – (0.38 ± 0.10) kg·mol–1 are obtained. From the

latter value and ε(Ni2+, 3NO− ) = (0.182 ± 0.010) kg·mol–1, ε(NiOH+, 3NO− ) = − (0.27 ± 0.12) kg·mol–1 can be derived.

Figure A-13: Extrapolation to I = 0 of the equilibrium constants 10 1,1*log b (A.23) re-

ported in [64PER] at 20°C, using the SIT.

-10.8

-10.4

-10

-9.6

-9.2

-8.8

0 0.5 1 1.5 2

I / mol·kg−1

log 1

0*β 1

,1 +

2D


314

At the other temperatures, only a few percent of 3NiNO+ can be expected (Imax = 0.046 M), thus no correction was applied, but the SIT was used to derive 10 1,1

*log οb . The reported 10 1,1

*log οb values and those recalculated using the SIT are collected in Table A-13.

Table A-13: Reported and recalculated 10 1,1*log οb .

t (°C) 15 20 25 30 36 42

10 1,1*log οb (reported) – 10.22 – 10.05 – 9.86 – 9.75 – 9.58 – 9.43

10 1,1*log οb (recalculated) – 10.22 – 10.07 – 9.80 – 9.74 – 9.59 – 9.38

From the plot of 10 1,1*log οb vs. 1/T (see Figure A-14) r mH ο∆ (A.23) =

(51.7 ± 3.4) kJ·mol–1 can be derived.

Figure A-14: Plot of 10 1,1*log οb (A.23) recalculated from [64PER] vs. 1/T.

-24

-23.5

-23

-22.5

-22

-21.5

-21

0.00315 0.0032 0.00325 0.0033 0.00335 0.0034 0.00345 0.0035

K / T

ln *

βo 1,1


315

[64TAY/SCH]

This is one of the pioneering studies of thermodynamic data by emf-measurements with a solid electrolyte galvanic cell. Zirconia, doped with CaO and MgO, was used as an electrolyte to determine the free energy of the reaction:

Ni + 0.5 SiO2 + 0.5O2 0.5Ni2SiO4,

and the free energy of formation of numerous silicates and titanates at temperatures from 750 to 1200°C. The data of Taylor and Schmalzried coincide well with the newest emf-study of the above reaction and were used by this study to determine the Gibbs energy of formation of Ni2SiO4.

[64TRI/CAL]

The formation constants of thiocyanate complexes of Ni(II), Cd(II) and Ni(II) were de-termined by an extraction method at 20°C. The aqueous media were 1.5 and 3.0 M (Na)ClO4 solutions. The authors used rather acidic conditions (pH = 0.5 to 1.0), as only the protonated form of the ligand (HSCN) can be extracted into the organic phase (Ni(SCN)2(aq) is reported to be weakly extractable). The method used to determine the concentration of the extracted HSCN was not reported. The concentration of the thiocy-anate ion ranged from 0.01 to 0.55 M, while [Ni2+]tot was ~ 0.1 M. The experimental data were described by the formation of four 2Ni(SCN) x

x− (x = 1 – 4) complexes. For

both ionic strengths used, K3 < K4 is reported. Considering that no geometrical change of nickel(II) occurs during the stepwise complex formation [90BJE2], the reported K4 values are probably overestimated. Using the accepted enthalpy values for the formation of 2Ni(SCN) x

x− (x = 1 – 3) species, the reported formation constants were corrected to

298.15 K and the uncertainties were increased: 10 1log b = (1.10 ± 0.10), 10 2log b = (1.69 ± 0.10), 10 3log b = (1.62 ± 0.20) (I = 1.5 M), and 10 1log b = (1.15 ± 0.10),

10 2log b = (1.61 ± 0.10), 10 3log b = (1.21 ± 0.30) (I = 3.0 M).

[64WEL/KEL]

Low-temperature heat capacities were determined for NiS (millerite) and Ni3S2 (heazlewoodite) between 50 and 298.15 K. The solid compounds were obtained by re-action of certain amounts of nickel oxide with sulphur. All products were characterised by X-ray diffraction analysis. Standard entropy values at 298.15 K were derived by in-tegration of the respective heat capacity functions. In the case of millerite, the standard entropy is equal to mS ο (NiS, β, 298.15 K) = (52.97 ± 0.33) J·K–1·mol–1 and the heat capacity at 298.15 K is ,mpCο (NiS, β, 298.15 K) = 47.11 J·K–1·mol–1, which agrees well with the result published by Grønvold and Stølen [95GRO/STO], ,mpCο (NiS, β, 298.15 K) = 47.08 J·K–1·mol–1. The standard entropy of Ni3S2 was found to be

mS ο (Ni3S2, cr, 298.15 K) = (133.89 ± 0.84) J·K–1·mol–1 which is in close agreement with the value given by Stølen et al. [91STO/GRO] ( mS ο (Ni3S2, cr, 298.15 K) = 133.2 J·K−1·mol–1). Weller and Kelley [64WEL/KEL] obtained ,mpCο (Ni3S2, cr, 298.15 K) =


316

117.65 J·K−1·mol–1, which is in excellent agreement with the values recommended by Ferrante and Gokcen [82FER/GOK], ,mpCο (Ni3S2, cr, 298.15 K) = 117.7 J·K–1·mol–1, and Stølen et al. [91STO/GRO], ,mpCο (Ni3S2, cr, 298.15 K) = 118.2 J·K–1·mol–1.

[65ADA/KIN]

Adami and King reported heats of solution in 4.360 m HCl(aq) at 303.15 K. The ther-modynamic cycle used heats for the following reactions:

H2SO4·7.068H2O(l) 2H+(sln) + 24SO − (sln) + 7.068H2O(sln)

Ni(cr) + 2H+(sln, H2(sat)) Ni2+(sln, H2(sat)) + H2(g)

7.068H2O(l) 7.068H2O(sln)

NiSO4(cr) Ni2+(sln) + 24SO − (sln)

to obtain a value of (3.724 ± 0.460) kJ·mol–1 for the heat of the reaction:

Ni(cr) + H2SO4·7.068H2O(l) NiSO4(cr) + 7.068H2O(l) +H2(g) (A.24)

at 303.15 K. The authors corrected the value for the heat of Reaction (A.24) to (3.598 ± 0.460) kJ·mol–1, at 298.15 K.

The reported experimental uncertainties are 1σ uncertainties. Also, the authors estimated the uncertainly in the reaction involving dissolution of nickel with evolution of hydrogen as ± 0.3 kJ·mol–1, after applying a correction to account for heat absorbed in vaporising the water and hydrogen chloride that accompanied the evolution of hydro-gen. Solution calorimetry involving reactions resulting in gas evolution is notoriously difficult, and in the present assessment the 1σ uncertainty has been doubled. Based on this, a value of (3.724 ± 1.400) kJ·mol–1 is calculated for the heat of the Reaction (A.24) at 303.15 K (2σ uncertainty).

The difference in the values of r ,mpCο∆ for 298.15 K and of r ,mpCο∆ for the av-erage reaction temperature, 300.65 K, are small enough that they can be ignored relative to the uncertainties in the values of ,mpCο . Using the values at 298.15 K, ,mpCο (NiSO4, cr, 298.15 K) = (97.63 ± 0.10) J·K–1·mol–1 (see Section V.5.1.2.2.1.3, and [78STU/FER]),

,mpCο (Ni, cr, 298.15 K) = (26.07 ± 0.10) J·K–1·mol–1 (see Section V.1.2), ,mpCο (H2, g, 298.15 K) = (28.836 ± 0.002) J·K–1·mol–1 (cf. Table IV-1), ,mpCο (H2O, l, 298.15 K) = (75.351 ± 0.080) J·K–1·mol–1 (cf. Table IV-1) and ,mpCο (H2SO4·7.068H2O, l, 298.15 K) = (614 ± 10) J·K–1·mol–1 [56HOR/BRA], [82WAG/EVA], we obtain r ,mpCο∆ = (20 ± 10) J·K–1·mol–1, ∆( r mH ο∆ ) = (0.10 ± 0.10) kJ·mol–1, and r mH ο∆ = (3.59 ± 1.40) kJ·mol–1. Then, using f mH ο∆ (H2O, l, 298.15 K) = – 285.83 kJ·mol–1 and f mH ο∆ (H2SO4·7.068H2O, l, 298.15 K) = – 2897.05 kJ·mol–1 [82WAG/EVA], f mH ο∆ (NiSO4, cr, 298.15 K) can be calculated. The value for f mH ο∆ (H2SO4·7.068H2O, l, 298.15 K) is adjusted to − (2897.12 ± 0.64) kJ·mol–1 to allow for the differences in the values for f mH ο∆ ( 2

4SO − , 298.15 K) from Table IV-1 and from the US NBS table [82WAG/EVA]; the uncertainty


317

is estimated to incorporate the uncertainty for f mH ο∆ for the sulphate ion and the “hy-dration” to H2SO4·7.068H2O.

From these values, f mH ο∆ (NiSO4, cr, 298.15 K) is calculated to be − (873.28 ± 1.57) kJ·mol–1. The difference between this number and the value originally reported by Adami and King is related primarily to the different values used for

f mH ο∆ (H2SO4, l, 298.15 K). Adami and King used values from the 1952 US NBS Cir-cular 500 [52ROS/WAG].

[65AKI/FUJ]

In this paper the olivine – spinel transition in Fe2SiO4 and Ni2SiO4 have been investi-gated over the temperature range 700 to 1500°C in the pressure range 20 to 40 kilobars for Ni2SiO4 using a tetrahedral-anvil type of high-pressure apparatus. For both silicates the transition has proved to be reversible. The equation for the transition curve for Ni2SiO4 was determined to be p = 19 + 0.012 t, where p is in kilobars and t is in degrees centigrade. On the basis of this equation, the standard Gibbs energy of formation of spinel from the olivine was calculated to be (1200 + 0.98 T) cal·mol–1 for Ni2SiO4, where T is absolute temperature.

[65BUR/LIL]

The hydrolysis of nickel(II) was studied at (25.0 ± 0.1)°C in 3 M (Na)ClO4 medium by potentiometric titrations using glass electrodes. The concentration of nickel(II) ranged from 0.1 to 0.8 M. Since 4[ClO ]− was kept at 3 M during the measurements, the ionic strength ranged from 3.1 to 3.8 M. This is a common problem of investigations using rather concentrated nickel(II) solutions. Nevertheless, this is one of the highest quality papers concerning the hydrolysis of nickel(II). Both graphical and computer evaluation of the experimental data indicated that the major species is 4

4 4Ni (OH) + , under the ex-perimental conditions used. The remaining small deviations between the observed and calculated Z values (number of OH– bound per nickel(II)) were explained by small ana-lytical errors and the presence of Ni2OH3+ as a minor complex. The optimum fit resulted

10 4,4*log b = – (27.37 ± 0.02 (3σ)) and 10 1,2

*log b = – (10.0 ± 0.5). The authors listed some additional minor species, remarking that the experimental data give no support for their existence.

Due to the relatively important change of ionic strength (and thus activity coef-ficients) between the measurements at different nickel(II) concentrations, we re-evaluated the reported experimental data referring to both [Ni2+]tot = 0.1 to 0.4 M and [Ni2+]tot = 0.1 to 0.8 M. Using the whole dataset, and the model proposed by the authors,

10 4,4*log b = – (27.37 ± 0.04) (3σ) and 10 1,2

*log b = – (9.42 ± 0.10) were obtained. The agreement between the recalculated and original values is excellent for the main hydro-lytic species, and reasonable for the minor complex. The difference is very likely due to the more sophisticated data analysis software now available (for all calculations in this chapter the PSEQUAD program was used [91ZEK/NAG]). Introducing the species


318

NiOH+ into the equilibrium model, a substantial (40%) decrease of the fitting parameter was obtained. From the data referring to [Ni2+]tot = 0.1 – 0.4 M, 10 4,4

*log b = − (27.43 ± 0.04) (3σ), 10 1,1

*log b = – (9.86 ± 0.12) and 10 1,2*log b = – (9.87 ± 0.50), us-

ing the whole dataset 10 4,4*log b = – (27.42 ± 0.03), 10 1,1

*log b = – (9.89 ± 0.06) and 10 1,2

*log b = − (9.78 ± 0.16) were calculated. It is important to note that the authors also tried to include NiOH+ in the equilibrium model, but their calculations gave no real support for its existence. Taking into account the above values, 10 4,4

*log b = − (27.43 ± 0.05) (3σ), 10 1,1

*log b = – (9.86 ± 0.12) and 10 1,2*log b = – (9.8 ± 0.5) are ac-

cepted for use in this review.

[65BUR/LIL2]

The hydrolysis of nickel(II) was studied at (25.0 ± 0.1)°C in 3 M (Na)Cl medium by a method identical to the one described in [65BUR/LIL]. The total nickel(II) concentra-tion was varied between 0.2 and 1.4 M, which resulted in a change of ionic strength from 3.2 to 4.4 M. The analysis of experimental data indicated the presence of

44 4Ni (OH) + as the dominant species ( 10 4,4

*log b = – (28.42 ± 0.05)), while the devia-tions at lower pH values suggested the presence of Ni2OH3+ ( 10 1,2

*log b = – (9.3 ± 0.2)). For 10 1,1

*log b a value of – 10.0 was obtained from the calculations, but was rejected as too uncertain. During these measurements, a large part of the Na+ ion concentration was replaced by Ni2+, consequently the constancy of activity coefficients at different nickel(II) concentrations is questionable. Furthermore, the formation of chloro com-plexes (both binary and ternary) was not considered.

For the purpose of this review, an attempt has been made to minimise the in-fluence of the medium effect (using only the data reported for [Ni2+]tot = 0.2 and 0.4 M) and to take into account the presence of NiCl+ complex. For the reaction:

Ni2+ + Cl– NiCl+ (A.25)

10 1log b = – 0.59 can be calculated at Im = 3.2 m in NaCl medium, using the recom-mended value for 10 1log οb (A.25) and ε(NiCl+, Cl–) = ½{ε(Ni2+, Cl–) + ε(Na+, Cl–)} = (0.1 ± 0.03) kg·mol–1.

Considering the whole data set and the same model as proposed, 10 4,4*log b =

− (28.32 ± 0.09) (3σ) and 10 1,2*log b = – (9.41 ± 0.13) can be calculated, which are in

reasonable agreement with the published values (see also the discussion of [65BUR/LIL]). Using only the data sets reported for [Ni2+]tot = 0.2 and 0.4 M, and tak-ing into account the formation of NiCl+, 10 4,4

*log b = – (27.52 ± 0.10 (3σ)) and 10 1,2

*log b = – (8.55 ± 0.06) can be derived. Introducing the species NiOH+ into the equilibrium model, a 50% decrease in the fitting parameter was found, and 10 4,4

*log b = – (27.55 ± 0.06) (3σ), 10 1,1

*log b = – (9.56 ± 0.10) and 10 1,2*log b = − (8.95 ± 0.20) were

obtained. These values are accepted for use in this review, with increased uncertainties: 10 4,4

*log b = – (27.55 ± 0.20), 10 1,1*log b = – (9.56 ± 0.20) and 10 1,2

*log b = − (8.95 ± 0.40).


319

[65COL/PET]

Penneman and his co-workers published three consecutive papers concerning formation of the 2

4Ni(CN) − ion [60MCC/JON], [63PEN/BAI], [65COL/PET]. In the first paper [60MCC/JON], the complexes formed in aqueous solutions of 2

4Ni(CN) − and cyanide ions were studied by IR and spectrophotometric methods. The results indicated the for-mation of only one new species ( 3

5Ni(CN) − ), even at a very high excess of cyanide con-centrations. The magnetic measurements indicated diamagnetic behaviour of the penta-cyano species, which suggests square pyramid geometry around the metal ion. From the spectrophotometric data, measured at 430, 450 and 470 nm, 10 5log K = − (0.72 ± 0.02) was determined at 25.2°C and I = 1.34 M (NaClO4). From the IR measurements,

10 5log K = – (0.52 ± 0.20) was obtained. Identical spectrophotometric studies at 15.4 and 33.6°C have been performed. From the temperature dependence of 10 5log K (see Figure A-15) r mH∆ = – (10.4 ± 0.2) kJ·mol–1 can be obtained. The uncertainty in this value does not include systematic uncertainties, nor the uncertainty in determining an enthalpy of reaction from equilibrium data over a small temperature range. In the pre-sent review, we assign an uncertainty of ± 4.0 kJ·mol–1 to this value.

While [60MCC/JON] was in press, two papers by Blackie and Gold [59BLA/GOL], [59BLA/GOL2] appeared, with substantially different conclusions, therefore Penneman et al. reinvestigated the 2

4Ni(CN) − – CN– system [63PEN/BAI], using IR spectroscopy at 25°C and I = 4 M NaClO4. The authors suggested the forma-tion of two new species 3

5Ni(CN) − and 46Ni(CN) − . The latter complex was taken into

account to explain some deviations at very high cyanide concentrations ([CN–] > 2.5 M). The values of 10 5log K = – (0.55 ± 0.02) and 10 6log K = – (1.0 ± 1.0) have been re-ported. Due to the substantial medium effect, only formation of the pentacyano complex is considered in this review, and its formation constant is accepted with an increased uncertainty (± 0.30). In a further experiment, the authors found that the addition of NaCl to the solution of 2 M NaCN and 0.075 M 2

4Ni(CN) − (the main species is 35Ni(CN) − )

resulted in a significant decrease in the peak height attributed to 35Ni(CN) − . This obser-

vation was explained as formation of Ni(CN)5Cl4– ( 10log K = – 0.66).

Two years later, Coleman et al. in [65COL/PET] tried to solve the discrepancy between their two earlier papers [60MCC/JON] and [63PEN/BAI]. In this third paper the authors again used IR and spectrophotometric methods to study the 2

4Ni(CN) − – CN– system in a 4 m KF – KCN mixture at 25°C (30°C for IR). They showed that, under the conditions used, the formation of a single species, 3

5Ni(CN) − , is sufficient to account for the experimental data. Based on the spectrophotometric data, a value of 10 5log K = (0.03 ± 0.01) (in molal units) has been reported. From the IR measurements at 30°C,

10 5log K = – (0.08 ± 0.06) was derived. In light of these findings, the authors explained their earlier results reported in [63PEN/BAI] (the formation of 4

6Ni(CN) − and Ni(CN)5Cl4–) by the neglected medium effect caused by the substantial change of NaCN/NaClO4 (NaCl/NaClO4) ratio during the measurements.


320

Figure A-15: Temperature dependence of the association constant for the reaction 24Ni(CN) − + CN– 3

5Ni(CN) − reported in [60MCC/JON].

[65EGO/SMI]

This is a very careful study of the specific heat capacity of Ni2SiO4 and Zn2SiO4, the only one for the nickel orthosilicate for temperatures as high as 1570 K. Egorov and Smirnova used a copper block calorimeter and 99 NiCr / Constantan thermocouples for the measurement of the temperature change of calorimeter. The average, least square, precision of measurements performed in the range between 555 – 293 K and 1570 – 293 K was ± 0.4. This high temperature data were used in this review to fit the heat capacity equation up to 1570 K.

[65LIN/SEI]

According to Linke and Seidell’s critical evaluation of the nickel sulphate solubility in water the stable phase from the eutectic temperature to 31°C is NiSO4·7H2O. The re-ported solubilities differ by 1 – 3%.

[65MOR/REE]

The formation of nickel(II) chloro complexes in aqueous solution has been studied us-ing a cation-exchange method at 20°C, in a solution of 0.691 M H(ClO4,Cl). The distri-

0.0032 0.0033 0.0034 0.0035

-0.9

-0.8

-0.7

-0.6

log 10

K5

K / T


321

bution of nickel at equilibrium between the resin (Amberlite IR-120) and the solution has been determined spectrophotometrically. The reported formation constants are β1 = (1.7 ± 0.1) M–1; β2 = (0.92 ± 0.10) M–2. Due to the almost complete change of the ionic medium during the measurements ([Cl–] = 0 to 0.659 M) the real formation of the bis-complex is doubtful. As the presented data do not allow an independent recalculation, they were not considered in this review.

[65PAO]

The heat of solution of NiBr2 in water was reported as – (82.09 ± 0.17) kJ·mol–1 based on limited measurements and analysis. The solid was obtained by dehydration over sul-phuric acid in a vacuum desiccator, and the Ni and Br analyses were within 0.2% of theoretical. The number of samples measured was not reported, but some values in the same paper were based on only two measurements. Neither the dilution of the final solution nor the sample size was given. The enthalpy value probably was not corrected for dilution, as there is no description of a correction procedure. The uncertainty reported was probably 1σ. The resulting value is approximately 4 kJ·mol–1 less exothermic than that of Efimov and Furkalyuk [90EFI/FUR] (although the final solutions probably were somewhat different).

[65PAO/SAB]

A value of – (70.46 ± 0.33) kJ·mol–1 was determined for the enthalpy of solution of NiI2(cr) in water containing (n–Pr)4NI (0.0005 m) and a “small quantity” of HI. The final solution concentrations were approximately 0.00025 m in nickel. The solid was obtained by dehydration over sulphuric acid in a vacuum desiccator, and the Ni and I analyses were within 0.2% of theoretical. The reported value was an average of “at least two determinations”. The authors concluded from similar measurements on the dissolu-tion of ZnI2(cr) that the effects of association were small; yet, their own measurements suggest that the uncertainties from this source could be greater than 1 kJ·mol–1. The re-sulting value is less exothermic than that of Efimov and Furkalyuk [90EFI/FUR], though the final solutions also are not identical.

[65PAW/STA]

The specific heat capacity of high-purity nickel and several nickel-rich alloys from room temperature to 600°C is reported. The data, obtained with an adiabatic calorime-ter, were reproducible to within ± 0.25 per cent and have an estimated absolute error of less than ± 0.5 per cent. This high quality data were credited by this review and used for the determination of the heat capacity function.

[65POT]

This reports a study of the complex dielectric constants of 1 M NiSO4 in aqueous solu-tion as a function of temperature.


322

[65WEL]

Heat capacity measurements, 52.0 to 296.3 K, were reported for anhydrous NiSO4(cr). A maximum entropy for 298.15 K was estimated by extrapolation to 0 K and incorpora-tion of an estimated magnetic contribution. These heat capacity results were incorpo-rated with results at lower temperatures in the later analysis of Stuve et al. [78STU/FER].

[66BOD/DEH]

Bode et al. investigated the reactions and reaction products on the nickel oxide elec-trode of the alkaline storage batteries. The following aspect of Bode et al.’s model is well established by now: the charged (NiOOH) as well as the discharged (Ni(OH)2) material of nickel oxide hydroxide electrodes can exist in two forms. In this paper the synthesis and structure of α-Ni(OH)2 has been described and its formula 3Ni(OH)2·2H2O was determined. This nickel hydroxide phase consists of brucite type layers of Ni(OH)2 with larger layer distances than those of β-Ni(OH)2. This is a result of interlayers of water. The definitive characterisation of its structure was not possible, because no single crystals could be grown.

[66BUR/IVA]

The hydrolysis of nickel(II) has been studied at (25.0 ± 0.1)°C in 3 M (Na)NO3 medium by a method identical to the procedure described in [65BUR/LIL]. The nickel(II) con-centration ranged from 0.2 to 1.0 M, which resulted in a change of ionic strength from 3.2 M to 4.0 M. The analysis of experimental data indicated the presence of 3

3 3Ni (OH) + ( 10 3,3

*log b = – (21.58 ± 0.03)) as major and Ni2OH3+ ( 10 1,2*log b = − (9.6 ± 0.2)) as

minor complexes. Since no spectroscopic evidence was found up to 4 M nitrate concen-tration, for the formation of the species 3NiNO+ , the formation of binary or ternary ni-trato complexes was not considered by the authors. In fact, a substantial amount of ni-trato complex is present at 3 M nitrate concentration, therefore the reported constants are underestimated. In any other medium (NaClO4, NaCl, KCl, NaBr), the tetramer is reported to dominate, and the formation of a trimer has never been confirmed by others; therefore, we re-evaluated the reported experimental data. Considering the suggested hydroxo complexes and neglecting the formation of 3NiNO+ , 10 3,3

*log b = − (21.61 ± 0.03) and 10 1,2

*log b = – (9.32 ± 0.11) can be calculated, in good agreement with the reported values. However, an equally good fit to the experimental data can be achieved by assuming the presence of 4

4 4Ni (OH) + , NiOH+ and Ni2OH3+ ( 10 4,4*log b =

− (28.00 ± 0.04), 10 1,1*log b = − (9.70 ± 0.11) and 10 1,2

*log b = − (9.28 ± 0.2)). An at-tempt was made to consider the formation of 3NiNO+ species. However, the tentatively accepted equilibrium constant for the reaction,

Ni2+ + 3NO− 3NiNO+ , (A.26)

10 1log οb ((A.26), 298.15 K) = (0.5 ± 1.0)) represents only the higher limit of the ‘true’ constant due to the neglected medium effect. Using the SIT approach, 10 1log b (A.26),


323

298.15 K) = − (0.3 ± 1.0) can be calculated for 3.31 m NaNO3 medium. Accepting this value, 10 4,4

*log b = – (26.52 ± 0.05), 10 1,1*log b = – (9.14 ± 0.12) and 10 1,2

*log b = − (8.77 ± 0.25) were obtained. These constants are much higher than the values of [65BUR/LIL] and [65BUR/LIL2], which is very likely the consequence of the overes-timated formation constants of 3NiNO+ . Fixing the value of 10 4,4

*log b at – 27.4, 10 1,1

*log b = – (9.49 ± 0.15), 10 1,2*log b = – (9.02 ± 0.15) and 10 1log b ((A.26), 298.15 K)

= – (0.84 ± 0.04) can be calculated. Considering the great uncertainty of 10 1log οb ((A.26), 298.15 K), none of the above values are accepted for use in this re-

view.

[66FLO]

The formation and stability of the 2 5NiCl(H O)+ complex have been studied by a spec-trophotometric method, in 10 M Li(NO3, Cl) medium at 303 K, to explain the polaro-graphic behaviour of nickel(II) in concentrated chloride media. The concentration of LiCl was varied between zero and 10.0 M. This fundamental change in the ionic me-dium results in a large uncertainty in the reported formation constant (β1 = (0.094 ± 0.009)), although the author noted that the absorption spectra of 0.03 M nickel(II) in 10 M LiCl, 10 M HCl, 5 M CaCl2 and 5 M MgCl2 are almost identical. This indicates a relatively minor impact of the medium effect. Therefore, 10 1log b = − (1.0 ± 0.5) is accepted in this review. The ionic strength applied is too high to extrapo-late the reported constant to infinite dilution using the SIT. In addition, LiNO3 can not be regarded as an ‘innocent’ background electrolyte, since nickel(II) forms an ion pair with the nitrate ion.

An attempt to determine the stability constants of the nickel bromo complexes was complicated by the slow formation of bromine from oxidation of bromide by the nitrate ion. However, the results indicated that the bromo and chloro complexes of nickel(II) have similar stability.

[66GOL/RID]

In this paper, heats of solution in water at 298.15 K were reported for α-NiSO4·6.000H2O(s) (final solution compositions 0.00009 m to 0.051 m) and for a mixed hydrate with a nominal composition of NiSO4·5.059H2O (final solution composi-tions 0.02 m). The mixed solid was assumed to be a mixture of NiSO4·6.000H2O(s) and NiSO4·H2O(s), and the enthalpy of hydration of NiSO4·H2O(s) to the hexahydrate was calculated. The standard enthalpies of solution for the solids were calculated using heat of dilution values reported by Lange [59LAN], and Lange and Miederer [56LAN/MIE].

The enthalpy of solution of NiSO4·6.000H2O(s) in water was reported as sol mH ο∆ ((A.27), 298.15 K) = 1.15 kcal·mol–1 (4.81 kJ·mol–1).

NiSO4·6H2O(cr) Ni2+ + 24SO − + 6H2O(l) (A.27)


324

The largest uncertainty in this value comes from the extrapolation to I = 0. Measurements by Lange and co-workers [56LAN/MIE] on CdSO4, CuSO4, ZnSO4, and NiSO4 lead to values for the integral heats of dilution that approach I = 0 with a slope almost twice the expected Debye-Hückel limiting slope, AL. If an extended Debye-Hückel treatment (such as the SIT [97ALL/BAN]) is used to describe the behaviour of a strong 2:2 electrolyte, the electrolyte will be found to behave as if it is associated (as we have noted for NiSO4). However, because higher order electrostatic terms have been omitted, the extended Debye-Hückel theory is inadequate for extrapolation of enthalpies of dilution and heat capacity values to I = 0 [97MAL/ZAM], [98ARC/RAR]. Thus, the large slope found by Lange and Miederer [56LAN/MIE] is not unexpected.

In Figure A-16 the original data are plotted versus the square root of the ionic strength. The experimental enthalpies of dilution for NiSO4 reported by Lange and Miederer [56LAN/MIE] were subtracted from sol mH∆ values of [66GOL/RID], result-ing in values of sol mH ο∆ . The mean value obtained from Equation (A.28)

sol mH ο∆ (A.27) = sol mH∆ (A.27) – φL (A.28)

is sol mH ο∆ (A.27) = (4.485 ± 0.200) kJ·mol–1, and this value is accepted in the present review. For comparison φL was calculated with Equation (IX.44) of [97ALL/BAN], where the last term had to be neglected because the temperature derivative of ε(Ni2+, 2

4SO − ) is unknown. Figure A-16 shows that relying on the Debye-Hückel limit-ing law leads to an error of ~ 1.7 kJ·mol–1 in the enthalpy of solution of NiSO4·6H2O(cr) at infinite dilution.

It seems likely that the NiSO4·5.059H2O was a mixture of NiSO4·6H2O(s) and NiSO4·4H2O(s) rather than a mixture of the hexahydrate and monohydrate. Jamieson et al. [65JAM/BRO], using samples dehydrated at 100°C, noted changes in dissolution kinetics and enthalpies for hydrated nickel sulphates with less than 29% water (the wa-ter content of the tetrahydrate would be 31%). They also showed that vacuum dehydra-tion at 40°C gave products with a different heat of hydration per mole of water of hydration. Goldberg et al. [66GOL/RID] reported similar difficulties in dissolving more highly dehydrated samples, and presented no proof that dehydration of the hexahydrate over P2O5 led to formation of the monohydrate.


325

Figure A-16: Molar enthalpy of solution of NiSO4·6H2O(cr) at 25°C. Experimental data from Goldberg et al. [66GOL/RID] ○; dotted curve: calculated according to Equation (IX.44) in [97ALL/BAN]; dash dotted curve: calculated according to the chord method, see Equations (8 – 2 – 19) and (8 – 2 – 20) of Harned and Owen [58HAR/OWE]; solid curve: calculated with an adjustable Debye-Hückel parameter, AL, and a term linear in I,

sol mH ο∆ ((A.27), 298.15 K) derived from dilution enthalpies [56LAN/MIE] ×; mean value of sol mH ο∆ ((A.27), 298.15 K) , dash line.

[66KEN/LIS]

This paper reports calorimetric results for the heats of formation of (among others) NiCl+ and NiBr+ complexes at an ionic strength of 2.0 at 298 and 313 K. In a typical experiment 350 cm3 of the nickel(II) perchlorate solution (0.67 M) was mixed with 50 cm3 2.0 M sodium halide solution. In the first step, values of the stability constants of the complexes formed were taken from two earlier papers from Lister's laboratory [60LIS/ROS] and [61LIS/WIL]. In [60LIS/ROS] the authors measured (among others) the emf of the cell Ag,AgCl | NaCl+M(ClO4)2+NaClO4 | NaCl+NaClO4 | AgCl,Ag (where M = Ni(II) or Co(II), [M] = 0.12 – 0.31 M and [Cl–] = 0.008 M) at I = 2 M (Na,M)ClO4 and at 285, 298 and 313 K, to determine the composition and stability of the nickel(II) complexes formed with chloride ion. The complex formation induced 1-6 mV differences between the electrodes. The experimental data were interpreted by tak-

0.0 0.1 0.2 0.3 0.4 0.5

4

5

6

7

8

9

∆ solH

m / kJ

·mol

−1

(I / mol·kg−1) 0.5


326

ing into account the formation of two complexes, NiCl+ (K1) and Ni2Cl3+(K2). The latter complex was considered to explain a small drift in the K1 value. The experimental data reported in this earlier paper were reinterpreted in this paper by Kennedy and Lister [66KEN/LIS], as the temperature dependence of K1 and K2 determined earlier was not consistent with their calorimetric results. Taking into consideration only the formation of NiCl+, a coherent model can be given. Therefore, the authors recalculated the earlier reported K1 values for 298 and 313 K neglecting the formation of Ni2Cl3+. For the pur-pose of this review we re-evaluated the experimental data reported in [60LIS/ROS] for all the three temperatures. This calculation resulted in only slightly different values for 298 and 313 K than were tabulated in [66KEN/LIS]. The following formation constants are accepted in this review: 10 1log b = – (0.21 ± 0.20) (285 K), – (0.16 ± 0.20) (298 K) and – (0.14 ± 0.20) (313 K). It is worth mentioning that the impact of the medium effect on these constants, induced by the relatively high concentrations of Ni(ClO4)2 (0.12 − 0.31 M) used, is probably considerably lower than for most of the studies re-ported in Table V-12.

From the equilibrium constants at 285 and 298 K, a considerably higher reac-tion enthalpy ( r mH∆ = (5.6 ± 1.4) kJ·mol–1) can be calculated using the integrated van’t Hoff equation than from the data of 298 and 313 K ( r mH∆ = (3.1 ± 1.2) kJ·mol–1). The latter value is closer to the calorimetrically determined reaction enthalpies ( r mH∆ = (2.1 ± 0.3) kJ·mol–1 and (1.8 ± 0.2) kJ·mol–1 at 298 and 313 K, respectively). In the pre-sent review we re-evaluated the calorimetric data at 298 K of the equilibrium:

Ni2+ + Cl– NiCl+ (A.29)

using the recommended value for 10 1log οb and ∆ε((A.29), NaClO4). At I = 2 M, β1 = 0.1 can be calculated. Since the authors used self-media for the calorimetric measure-ments, by mixing 0.632 M Ni(ClO4)2 and 2 M NaCl solutions, the β1 value calculated above is only an approximation for the conditions used. Nevertheless, using this con-stant r mH∆ ((A.29), 298 K) = (8.1 ± 2.5) kJ·mol–1 can be derived.

[66KOH/ROD]

The nickel mineral discovered at Lemieux Township, Gaspé Peninsula, Canada was investigated by X-ray, wet chemical, spectrographic, infrared, differential thermal, and petrographic analyses. It was found to be a previously undescribed nickel, magnesium, iron carbonate. It was named after the locality of its discovery.

[66LOF/MCI]

Lofgren and McIver carried out measurements for the cells Mg,MgF2|CaF2|NiF2|Ni and Mg,MgF2|CaF2,YF3,(3% by weight)|NiF2|Ni. The unweighted average of the values from the three sets of measurements judged by the authors to be unaffected by moisture is r mG∆ (A.30)= – (451.90 ± 0.70) kJ·mol–1.

NiF2(cr) + Mg(cr) MgF2(cr) + Ni(cr) (A.30)


327

A third law analysis of the results using auxiliary values from the present re-view, and with heat capacity functions for Mg(cr) and MgF2(cr) consistent with those in the CODATA review [89COX/WAG], leads to f mH ο∆ (NiF2, cr, 298.15 K) = − (657.0 ± 1.4) kJ·mol–1.

The authors also determined potentials for the cells: Mg,MgF2|CaF2,YF3,(3% by weight)|UF4,UF3|Pt and Pt,UF3,UF4|CaF2|NiF2|Ni and Pt,UF3,UF4|CaF2,YF3(3% by weight)|NiF2|Ni.

The combined measurements from these two cells were consistent with, indeed almost identical to, those from the simpler Mg,MgF2|CaF2|NiF2|Ni cell. Results reported for the cell Fe,FeF2|CaF2|NiF2|Ni cannot be used without a proper re-evaluation of the thermodynamic quantities for FeF2(cr), and such an evaluation is outside the scope of the present review.

[66SHI/FUJ]

A solution of K2Ni(CN)4 was irradiated with γ-rays for 15 minutes, then the chemical species in the irradiated sample (labelled with 57Ni by the reaction 58Ni(γ,n)57Ni)) was separated by paper electrophoresis. The authors detected four main peaks. Beside the Ni2+ and 2

4Ni(CN) − ions, the presence of Ni(CN)2 and Ni(CN)3(H2O)– species was also suggested, though the assignment of the peaks is rather speculative.

[66STO/ARC]

Results were reported for low-temperature calorimetric heat capacity measurements of samples of hydrated nickel sulphate with compositions NiSO4·6.010H2O, NiSO4·6.867H2O and NiSO4·7.301H2O. These were combined with results from earlier heat-capacity measurements [41STO/GIA], [64STO/HAD] to determine smoothed sets of entropy and heat capacity values from 1 to 300 K for α-NiSO4·6H2O and NiSO4·7H2O. The authors corrected the measurements for NiSO4·7.301H2O for the heat of fusion of ice (5.879 kJ·mol–1 at 269.75 K) on the basis of solubility data at the tem-perature of the NiSO4·7H2O-ice eutectic [39ROH]. The heat of dehydration,

NiSO4·7H2O NiSO4·6H2O + H2O(NiSO4(aq, sat))

was reported as (7.98 ± 0.21) kJ·mol–1.

The authors also measured the enthalpies of solution of samples of NiSO4·6.000H2O and NiSO4·6.940H2O in water such that the final solutions had essen-tially the same compositions (0.1 M in NiSO4). Thus, in that medium, the enthalpy of the dehydration reaction was (7.68 ± 0.10) kJ·mol–1. The authors assumed that this en-thalpy difference was unaffected by the final medium, and that it was equal to the dif-ference at infinite dilution in water. This value is accepted in the present review, and used in the calculation of the standard enthalpy of solution of the heptahydrate.


328

The authors also carried out vapour pressure measurements to determine the Gibbs energy of hydration of α-NiSO4·6H2O at 293.15, 298.15 and 303.15 K. Their results were within the experimental uncertainties of values determined by others [23SCH], [35BON/BUR]. From the enthalpy of solution measurements and the vapour pressure measurements the difference between mS ο (NiSO4·6H2O, α) and

mS ο (NiSO4·7H2O) was calculated, and compared to the difference in the mS ο values cal-culated from the low-temperature heat capacity measurements.

The differences were of the order of 0.7 J·K–1·mol–1. This is small considering the many experimental quantities used in the calculation. Though they are dependent on fewer assumptions, and are therefore the basis of the selected values in the present re-view, the third-law entropy values for the two hydrated solids may incorporate compen-sating errors. Therefore, an uncertainty of 1.0 J·K–1·mol–1 is estimated in the present review for the values of mS ο for the two solids at 298.15 K. For NiSO4·6H2O, the uncer-tainties in the heat capacity values near 298.15 K are estimated in this review to be ± 1.0 J·K–1·mol–1, and the same uncertainty could be assigned to values for

,mpCο (NiSO4·7H2O) for temperatures below, but close to the NiSO4·7H2O-ice eutectic. However, because of the experimental problems encountered near room temperature, an uncertainty of ± 4.0 J·K–1·mol–1 is assigned to ,mpCο (NiSO4·7H2O, 298.15 K).

[66VOL/KOH]

Vollmer et al. determined the enthalpy of fusion fus mH ο∆ = (17.15 ± 0.42) kJ·mol–1 and the heat capacity of nickel in the liquid region up to 1820 K. Based on the earlier work of one of the authors, R. Kohlhaas, the heat capacity of nickel in the solid state in the temperature range 300 to 1725 K is also reported in this paper. The reported values were used to determine the thermal heat capacity function recommended by this review. The heat capacity of nickel in the liquid region, ,mpCο (Ni, l) = 39.0 J·K–1·mol–1, was found to be independent of the temperature.

[67ANT/WAR]

A mixture of metallic nickel and nickel oxide was treated with a CO2/CO gas mixture at temperatures ranging from 853 to 1289 K. The partial pressure of CO in the equilibrated gas mixture was measured by means of a 14C tracer method after the carbon dioxide had been removed by a liquid nitrogen trap. The authors found for the temperature depend-ence of the partial pressure ratio:

210 CO COlog ( / )p p = – 2381 (K/T) – 0.025. From this

relationship, the Gibbs energy of formation of NiO was derived:

f mG∆ (T) = (– 236.81 + 0.087446 (T/K)) kJ·mol–1, 853 < T/K < 1289.

[67BLA]

Blaszkiewicz determined the solubility of Ni(OH)2 by a 63Ni tracer method and found [Ni(II)] = 1.49 × 10–5 mol·dm–3. Although this value is an order of magnitude lower than Almkvist’s, the mass and charge balance with Ni2+ and NiOH+ still leads to a much too


329

high value of ,0sK or else Ni(OH)2(aq) is unrealistically stable. Consequently this result was also rejected. At room temperature directly measured aqueous solubilities of nickel hydroxide are an insufficient basis for the determination of the respective solubility constant [18ALM], [67BLA]. As the system is not adequately buffered, freshly precipi-tated Ni(OH)2 forms colloidal particles, and the concentration of dissolved Ni(II) at the solubility minimum is at or below the instrumental detection limit [97MAT/RAI].

[67GES/NEU]

The solubility of NiSO4 hydrate is reported for temperatures from 160 to 240°C, and in ammonium nitrate solutions to 0.768 m to 300°C. Activity coefficients over the concen-tration and temperature range were calculated.

[67NAN/TOR]

Differential calorimetry has been used to measure the enthalpies of formation of the MSCN+ complexes (M = Mn(II), Co(II), Ni(II), Cu(II), Zn(II), Cd(II) and Pb(II)) in aqueous solutions at 25°C. No background salt was used to attempt to maintain constant ionic strength. The concentration of nickel(II) perchlorate varied between 5.04 and 8.9 mM, while [SCN–]tot ranged from 2.44 to 4.31 mM. Thus only the mono-thiocyanato complex was formed in the calorimeter:

Ni2+ + SCN– NiSCN+. (A.31)

The concentrations of NiSCN+ formed were calculated by means of the ther-modynamic stability constant reported in [62WIL2] and the Davies equation. For the purposes of the present review, the experimental data given in [67NAN/TOR] have been re-evaluated using the recommended values for 10 1log οb (A.31) and ∆ε((A.31),

4ClO− ). The recalculated reaction enthalpy (– (9.25 ± 0.40) kJ·mol–1) is somewhat higher than the original value (– (10.46 ± 0.40) kJ·mol–1). Although, these values refer to NaClO4 medium, the ionic strength in [67NAN/TOR] was always less than 0.063 M. Therefore, the recalculated reaction enthalpy is accepted with an increased uncertainty (– (9.25 ± 2.00) kJ·mol–1).

[67RUD/DEV]

Rudzitis et al. used fluorine bomb calorimetry to determine the enthalpy of formation of anhydrous nickel fluoride as – 657.7 kJ·mol–1 at 298.15 K. During the reaction, only approximately 50% of the nickel underwent reaction, and the fraction of unreacted ma-terial was determined by mechanical and magnetic separation. Despite these difficulties, the authors estimated that the overall 2σ uncertainty in their value of f mH ο∆ (NiF2, cr, 298.15 K) was 1.7 kJ·mol–1 (0.4 kcal·mol–1), slightly greater than the statistical uncer-tainties of the eight heat of combustion measurements. Because of the need to separate the nickel fluoride from the residual nickel at the end of the experiments, a bias in the final enthalpy of formation value cannot be ruled out. The paper provides an assessment of earlier fluoride enthalpy of formation determinations.


330

[67SIG/BEC]

The determination of the complexation constant between Ni2+ and 24HPO − was almost

incidental to the main set of experiments, which were intended to compare the effects of addition of bipyridyl and organophosphate ligands on nickel complexation. Neverthe-less, this is one of the studies most often cited when an association constant is required for nickel-phosphate complexation (and the values were cited and used by the same group in later publications, e.g., [81BAN/KAD]). Potentiometric titrations were carried out at 25°C using solutions 0.01 or 0.02 M in nickel and 2 × 10–4 M in 2

4HPO − , with 0.02 M NaOH(aq) as titrant, in a 0.1 M perchlorate medium (NaClO4(aq)). To deter-mine the complexation constant, the titration curves for solutions containing the metal and ligand were compared with those with the metal or ligand omitted.

The primary data were not reported except in a figure showing the smoothed data on a coarse scale, and, therefore, the data analysis cannot be properly evaluated. Also, the stoichiometry of the complex was not clearly established. The value

10log K (A.32) = 2.08 is given for the equilibrium constant for

Ni2+ + 24HPO − NiHPO4(aq) (A.32)

but the estimated uncertainty is not reported. In the same paper, the association constant value for the reaction:

H+ + 24HPO − 2 4H PO− (A.33)

is reported as (6.70 ± 0.02) in 0.1 M NaClO4(aq), whereas the value calculated using the SIT (cf. Appendix B) is (6.781 ± 0.015). Correction to zero ionic strength (Appendix B) gives 10log K ο (A.32) = (2.97 ± 0.30), where the uncertainty has been estimated in the present review. The small ∆ε correction has been based on the value for ε(Na+, 2

4HPO − ) from Table B-5 (rather than using the equation in Table B-6), and ε(Ni2+, 4ClO− ) = 0.37 kg·mol–1.

[68AND/KHA2]

This is qualitative work on the complex formation of nickel(II) in concentrated aqueous HBr (0.1 – 18 M) and HI (0.1 – 11 M) solutions, done by comparing the electronic ab-sorption spectra of nickel(II) at varying HX concentrations with those of crystalline Ni(II) complexes. The authors suggested the formation of 2NiBr x

x− (x = 1 – 6) and

2NiI xx

− (x = 1 – 4) complexes. Considering the molar absorption coefficient of 24NiBr −

complex in nitromethane (εmax(698 nm) = 173 [59GIL/NYH]), only the formation of approximately 5 –10 % 2

4NiBr − complex can be estimated in 18 M HBr solution from the published electronic spectra, while the higher bromo complexes can not be justified. On the other hand, the spectra presented suggest, indeed, the formation of a 2

4NiI − com-plex in considerable amounts in 9 – 11 M HI solutions. The absorption coefficients re-ported for 11 M HI solution (εmax(~ 800 nm) = 400, εmax(~510 nm) = 2600) suggest the formation of a tetrahedral complex.


331

[68ARN]

Calorimetric studies by acid titration of hydrolysed nickel(II) solutions have been per-formed at 25°C in 3 M (Na)Cl ionic medium in order to obtain the enthalpy of the hy-drolysis reactions. The experimental data have been interpreted using the hydrolysis mechanisms proposed by Burkov and Lilich [65BUR/LIL2] and Biederman and Ohtaki (cited as a personal communication to the author [68ARN]). The latter values ( 10 4,4

*log b = − (28.57 ± 0.03), 10 1,1*log b = – (10.5 ± 0.1) and 10 1,2

*log b = − (10.3 ± 0.2)) are slightly different from those reported in [71OHT/BIE]. The concen-tration of nickel(II) ranged from 0.055 to 0.750 M (I = 3.055 to 3.750 M). A value of (180 ± 2) kJ·mol–1 was obtained for the heat of formation of the main hydrolysis product

44 4Ni (OH) + , irrespective of which assumptions were made about the minor species. On

the other hand, the rather uncertain enthalpy values for the formation of the possible minor species (NiOH+, Ni2OH3+) strongly depended on the model (formation constants) used. The best fit to the data was obtained using the hydrolysis model proposed by Biederman and Ohtaki: r mH∆ = (64.4 ± 6.0) and (16.0 ± 8.4) kJ·mol–1 for the formation of NiOH+ and Ni2OH3+, respectively.

The formation constants used in the calculations were obtained without consid-ering chloride complexation. Therefore, for the purpose of this review we re-evaluated the reported experimental data, using the selected formation constants for 3 M chloride medium and assuming a value of (8.2 ± 2.5) kJ·mol–1 for the formation of NiCl+ under the conditions used (the latter value was calculated from the data measured at I = 2 M [66KEN/LIS]). The recalculation resulted in the following enthalpy of reaction values, valid for 298 K and I = 3M, r mH∆ = (54.3 ± 2.0) kJ·mol–1, (45.9 ± 1.6) kJ·mol–1 and (192.9 ± 1.0) kJ·mol–1 for the formation of NiOH+, Ni2OH3+ and 4

4 4Ni (OH) + , respec-tively.

[68CAM/ROE]

This paper describes a study of the stability of olivine and pyroxene in the Ni − Mg − Si – O system at temperatures of 1300 – 1500°C at controlled oxygen fugaci-ties using gas mixtures of known mixing ratio. The Gibbs energy of formation of nickel oxide and the Gibbs energy change for the reaction:


determined in this temperature range coincide very well with results of all emf-studies discussed by this review. Moreover, a close approach to ideal behaviour of the magne-sium – nickel olivine solid solutions was reported.

[68CHA/FLE]

A “closed system” solid electrolyte electrochemical cell has been designed to investi-gate the thermodynamic properties of metal oxides. This technique gave a rapid cell response and highly stable potentials through the elimination of mixed potentials. The


332

standard Gibbs energy of formation of NiO was found to be the following function of the temperature: f mGο∆ (T) = (– 233.651 + 0.085 (T/K)) kJ·mol–1. The accuracy calcu-lated from the maximum deviation from the computed least-squares line is equal to ± 0.209 kJ·mol–1. This corresponds to the maximum deviation in cell potential (emf) of ± 1.0 mV.

[68KOL/MAR]

The rate of formation and dissociation of 24Ni(CN) − has been studied at 25°C and I =

0.1 M NaClO4. The rate of formation between pH 5.5 to 7.5 showed a second-order dependence on both [CN–] and [HCN], at higher pH the reaction order in [CN–] was greater than 2, but the total cyanide dependence remained fourth order. The fact that the formation rate depended on [HCN]2 suggested that the reaction product might be Ni(HCN)2(CN)2(aq). Equilibrium measurements, using a spectrophotometric method between pH 4 – 5, resulted in systematic change in the calculated 10 4log b values. This was taken as further proof for the presence of protonated species, although the UV-VIS spectrum of 2

4Ni(CN) − was found to be unaffected by the protonation. The evaluation of combined kinetic and spectrophotometric data resulted in 10 4log b = (30.5 ± 0.3) and three protonation constants of the tetracyanonickelate(II) species: 10 H,0log K = (5.4 ± 0.2), 10 H,1log K = (4.5 ± 0.2), 10 H,2log K = (2.6 ± 0.2), for the reactions:

24Ni(HCN) (CN)x

x x−− + H+ 2 1

1 4 1Ni(HCN) (CN)xx x

− ++ − −

(where x = 0, 1, 2). On the other hand, spectrophotometric [59FRE/SCH] and potenti-ometric [63CHR/IZA], [71IZA/JOH], [74PER] studies performed in the same pH-range did not indicate the existence of protonated species in notable concentration, although their kinetic role in the acid dissociation of 2

4Ni(CN) − was later corroborated [76PER]. The value for 10 4log b is accepted, but with an increased uncertainty (± 0.6).

[68KOS]

Kostryuka and Zarubina measured the heat capacity of NiCl2 between 1.8 and 16 K to ascertain the singularities of the energy spectrum of layered antiferromagnets. The data have been presented graphically only.

[68LAR/CER]

Larson et al. obtained for the reaction:

NiSO4·7H2O(cr) Ni2+ + 24SO − + 7H2O(l) (A.34)

sol mGο∆ (A.34) = 12.929 kJ·mol–1 (3.09 kcal·mol–1) which agrees very well with the value recalculated in Section V.2.1.3.1. A misprint seems to have occured in this paper, as the standard enthalpy of reaction (A.34) is given as sol mH ο∆ (A.34) = 2890 instead of 2986 cal·mol–1. The latter value follows from the data of [66STO/ARC] and the original value given by [66GOL/RID].


333

[68MAL/TUR]

The instability constants of NiSCN+ were determined at 15, 25 and 35°C by a polaro-graphic method in aqueous solutions at I = (0.65 ± 0.03) M using (H)ClO4 as back-ground electrolyte. The thiocyanate concentration ranged from 0.5 mM to 7 mM, with [Ni2+]tot = 55.7 mM, thus the formation of higher complexes (Ni(SCN)x, x > 1) was ne-glected. The determination is based on the measurement of the decrease in catalytic polarographic current of the Ti(IV) – SCN complexes in the absence and in the presence of nickel(II). Although, the authors applied rather acidic conditions, and thus the thiocy-anate ion was partly protonated, the reported values are accepted with increased uncer-tainties. From the temperature dependence of 10 1log b (see Figure A-17) r mH∆ = − (14.5 ± 1.2) kJ·mol–1 can be derived.

Figure A-17: Temperature dependence of 10 1log b values of the NiSCN+ complex re-ported in [68MAL/TUR].

[68NAV/KLE]

The enthalpies of formation from the oxides of thirty-two 2-3 and 2-4 spinels, among others Ni2GeO4, have been determined by solution calorimetry in molten oxide solvents at 970 K. Regularities in the thermodynamics of spinel formation and the olivine – spinel transition for orthosilicates of Mg, Fe, Mn, Co, and Ni have been discussed.

[68PRA]

This is a description of a general method used by Prasad and co-workers to carry out potentiometric measurements to determine complexation constants, and describes how the measurements were extrapolated to I = 0. For NiSO4 only a single value is reported ( 1

1K − = 9.0 × 10–3 at 35°C). This suggests weaker complexation than was found in other

0.0032 0.0033 0.0034 0.0035

0.8

1.0

1.2

1.4

1.6

log 10

β1

K / T


334

work (e.g., [59NAI/NAN]). For experimental details and results, reference was made to unpublished work by S.C. Sircar, P.K. Jena and B. Prasad. However, no related paper on nickel sulphate complexation by these authors seems to have been reported in Chemical Abstracts, and the work may never have been published. The unavailability of the primary data precludes recalculation of the results using the standard SIT proce-dures, and the reported value is not used in the present assessment. The data reported for the nickel(II) – thiocyanate system in this paper, were later published in detail in [71DAS/DAS].

[68VOG]

This paper presented evidence for the existence, in potassium hydroxide melts (0.85 w KOH) between 498 and 523 K, of reversible homogeneous nickel redox couples in which nickel exhibits valencies greater than two. The data can be interpreted by the fol-lowing two simultaneous redox reactions:

Ni2+ Ni3+ + e–

Ni3+ + 4OH– NiO2(s) + 2H2O + e–.

The nature of the black NiO2 crystals (+4-valent nickel or peroxidic oxygen) was left undecided.

[69FUN/TAN]

Kinetic and mechanistic studies of the reaction between 4-(2-pyridylazo)resorcinol (PAR) and (among others) NiOH+ and NiF+ were performed at an ionic strength of 0.10 M (NaClO4) and at (298.15 ± 0.20) K. Nickel(II) perchlorate solutions were mixed with ligand solution, then borate buffer solution (8 × 10–3 M) was added to adjust the pH (7 – 9.5). The concentration ranges were as follows: [Ni2+] ~ 5 – 50 × 10–7 M, [PAR] ~ 10–5 M, [F–] ~ 0.05 M. The influence of the buffer was examined and it was concluded that it had no effect on the rate constants. The kinetic results also provided data for the equilib-rium constant for the reaction:

Ni2+ + OH– NiOH+ (A.35)

10log K (A.35) = (4.3 ± 0.1)) and for

Ni2+ + F– NiF+, (A.36)

10log K (A.36) = (1.1 ± 0.1)).

Using 10 wlog K = – 13.78, 10log K (A.35) can be can be converted to the corre-sponding first hydrolysis constant of nickel(II):

Ni2+ + H2O NiOH+ + H+,

10 1*log K ο = – (9.5 ± 0.1). In this review, these values are accepted, but with the uncer-

tainties increased to ± 0.2.


335

[69IZA/EAT]

Titration calorimetry was used to determine values of 10 r m f mlog , and K H Sο ο∆ ∆ , re-ported for the formation of NiSO4(aq) for 25°C and I = 0. The data analysis is suspect because of a possible correlation between the value for the calculated equilibrium con-stant and the enthalpy of association. A good reassessment was reported by Powell [73POW], who pointed out that if the data of Izatt et al. were used with the association constant at 298.15 K set equal to the value from Nair and Nancollas [59NAI/NAN], then an enthalpy of ion-pair formation of 5.7 to 5.9 kJ·mol–1 was obtained, not the re-ported 1.7 kJ·mol–1 [69IZA/EAT].

In the present review, several approaches were tried to determine the enthalpy of association within the constraints of the SIT formulation in Appendix B. As in earlier reviews [92GRE/FUG], [2001LEM/FUG], for low ionic strengths the assumption was made that r mH∆ is independent of ionic strength. Powell [73POW] suggested that the effect of this approximation for sulphate complexation was < 1 J·K–1·mol–1. Izatt et al. [69IZA/EAT] and Powell [73POW] incorporated values for an association constant and enthalpy of association for the Me4NClO4(aq) ion pair in their calculations. A full evaluation of these values has not been attempted here. Determination of the values for Me4NClO4(aq) required using an explicit value for the association constant between Na+ and 2

4SO − , whereas the SIT, as described in Appendix B, would treat these interactions using an interaction coefficient rather than a specific association constant.

Initially, the effect of the enthalpy of protonation of sulphate has been ignored. Using I = 0 values for the formation constant of NiSO4(aq), as recalculated from the data of Nair and Nancolas [59NAI/NAN] and of Katayama [73KAT], the molar concen-tration of NiSO4(aq) formed at each step of the titration was calculated iteratively. The average enthalpy of ion pair formation is 5.8 kJ·mol–1 and 4.7 kJ·mol–1, respectively. However, it is apparent that the calculated enthalpy of association is strongly dependent on the ionic strength of the solution. If the values at the measured ionic strengths are extrapolated to I = 0, the limiting value (Figure A-18) is 1.2 kJ·mol–1 or 0.4 kJ·mol–1 depending on whether the association constant data of Nair and Nancolas [59NAI/NAN] or from Katayama [73KAT] are used. This cyclic inconsistency may be a result of the neglect of the heats of the secondary association reactions and heats of dilution. It may also represent a failure of the SIT to adequately represent the system. At present, the enthalpy of association as calculated by Powell is accepted, but with an increased uncertainty r mH ο∆ = (5.8 ± 2.0) kJ·mol–1.


336

Figure A-18: Heats of reaction ( ) determined from the calorimetric data of [69IZA/EAT] by using the 298.15 K values of 10 1log K from [73KAT]. The line represents a least-squares fit to these enthalpies as a function of ionic strength.

[69KOL/KIL]

The reaction rate and mechanism of the acid catalysed decomposition of 44 4Ni (OH) +

have been studied by stopped-flow measurements. Potentiometric titrations of 0.46 M nickel(II) perchlorate solutions at 25°C and 1.5 M ionic strength, maintained by Na-ClO4, were also performed to study the hydrolysis under the conditions used for the kinetic measurements. The authors suggested the presence of the 4

4 4Ni (OH) + species, only, and 10 4,4

*log b = – (27.03 ± 0.06) was obtained. This value compares favourably with the corresponding value reported in [65BUR/LIL]. The reported data indicate a slight tendency for a decrease in the calculated 10 4,4

*log b values with decreasing pH, which may be due to the presence a minor hydroxo complex. Therefore, the reported experimental data were re-evaluated for the purpose of this review, but no further spe-cies can be identified with any confidence.

[69KOS]

The measurements carried out in [68KOS] were discussed from the point of view of the magnetic spectrum of layered antiferromagnets. The molar heat capacities of NiCl2 from 1.8 to 16 K have been listed.

0.00 0.02 0.04 0.06 0.08 0.10 0.12 0.140

2000

4000

6000

8000

10000

∆ r Hm /

J·mol

−1

I / mol·kg−1


337

[69KUL/ERS]

On a godlevskite specimen from the Cu – Ni sulphide deposit Noril'sk, Siberia, Russia, Kulagov et al. carried out chemical and mineralogical investigations as well as a crystal structure analysis. Whereas the mean values of the chemical analyses agreed with (Ni, Fe, Co)S the authors maintained that the upper limits of the metal and the lower limits of the sulphur contents are consistent with Ni7S6.

[69LIB/SAD] Osmotic coefficients have been determined by an isopiestic method for (among others) the aqueous solution of Ni(ClO4)2 in the concentration range m = 0.1 – 3.5 mol·kg–1 at 25°C. The mean activity coefficients were calculated from the Gibbs-Duhem equation. For the purposes of this review, these data, together with those listed in [99MAL/CAR], have been used to derive the ion interaction coefficient ε(Ni2+, 4ClO− ), see also the dis-cussion on [99MAL/CAR].

[69MAK/SPI]

Makovskaya and Spivakovskii added sodium hydroxide to 0.001 – 4.7 M NiCl2 solu-tions at 25°C until a slight precipitate formed. The pH and the chloride activity were measured with glass and silver chloride electrodes, respectively. The concentration of Ni(II) in the filtered solutions was determined by complexometric, photometric, and polarographic methods. Ni(OH)2, Ni(OH)1.5Cl0.5, and Ni(OH)1.75Cl0.25 were detected by thermodynamic analysis, but the latter phase disappeared during aging. None of these phases were characterized by X-ray analyses. The solubility product of the freshly pre-cipitated nickel hydroxide turned out to be approximately an order of magnitude higher than after an undefined aging period. So, only the value of the aged phase ( 10 ,0log sK ο = − 15.54) has been accepted for Table V-6 and Figure V-11.

[69MOR/SAT]

The standard Gibbs energy of the reaction:

2Ni + O2 2NiO

was investigated as a function of temperature in the temperature range from 700 to about 1100°C. The Gibbs energy function reads: f mGο∆ (T) = – 476.976 + 0.175770 (T/K)) kJ·mol–1. The uncertainty of emf measurements, caused by different kinds of solid electrolytes, was around ± 3 mV.

[69RAN/BRE]

The potential of nickel sulphide electrodes was measured against the saturated calomel electrode at 298.15 K. The sulphide electrodes were prepared by covering platinum with nickel sulphide which was obtained by precipitation in aqueous solutions of low acidity (1 M Na2S solution was dropped into 0.5 M NiSO4 solution). The composition of the


338

sulphide was altered by anodic as well as cathodic polarisation. Unfortunately, X-ray diffraction analysis revealed that the nickel sulphides were amorphous. Hence, no reli-able thermodynamic data can be evaluated from this study.

[69SOR/KOS]

For the calculation of mS ο (Ni(OH)2, cr, 298.15 K) according to Equation (A.37) the original ,mpCο data have been plotted versus lnT:

m ,m d lnpS C T= ∫ . (A.37)

Figure V-10 shows the experimental ,mpC values of three Ni(OH)2 samples (iii, ii, i) up to the temperature where the data for each have been measured. At T ≤ 4 K se-ries (iii) and (i) essentially coincide. In the range of the magnetic transition the heat ca-pacities decrease in the order ,mpC (iii) > ,mpC (ii) > ,mpC (i), but at temperatures higher than at the minimum of ,mpC = f (lnT), this order is reversed.

Using the CurveExpert routine the functions ,mpC (ii) = f (lnT) and ,mpC (i) = f (lnT) were fitted to linear splines and integrated according to Equation (A.37) [97HYA]. Finally, the Debye T 3 extrapolation down to T = 0 was applied. As no data above T = 104 K were measured for sample (iii), 0.4 and 1.2 J·K–1·mol–1 were sub-tracted from ,mpC (ii) and ,mpC (i) above 104 K, respectively, and the values thus ob-tained were added to series (iii) [69SOR/KOS]. Again the CurveExpert routine was used to integrate the true as well as the dummy part of this ,mpC (iii) function.

[69SUB/COR]

The rate of extraction of nickel(II)- and zinc(II)-dithizonates into CHCl3 has been stud-ied in presence of several auxiliary ligands (L), e.g., thiocyanate, at 25°C and at ionic strength of 0.25 M using NaClO4. The concentrations used were not reported, but were probably sufficiently low compared to the total ionic strength, since the concentration of the Ni(ClO4)2 stock solution was 1 mM. The variation of the apparent rate constant of the extraction as a function of the concentration of the auxiliary ligand, allowed calcula-tion the stability constants of the MLx complexes. The authors reported 10 1log b = (1.3 ± 0.2) for the stability of the species NiSCN+.

[70ASH]

The heat of solution of nickel chloride hexahydrate into 1.0M HCl(aq) was measured, and (6.10 ± 0.22) kJ·mol–1 was reported (2σ uncertainties). The final solution was 0.02 M in Ni. A double correction to 0.0 M HCl(aq) and to I = 0 [95MAN/KOR] would be required to use the reported value to calculate the enthalpy of solution of the salt at infi-nite dilution in water. The corrections themselves have large uncertainties, and this heat of solution value is not used further in the present review.


339

[70BIN/HEB]

The authors reported drop calorimetry results obtained from samples initially at tem-peratures between 381.9 K and 1461.7 K. The scatter in the enthalpy difference results, based on measurements done at similar temperatures, seems to be of the order of 200 kJ·mol–1. Direct fitting of the drop calorimetry results without incorporating some low temperature heat capacity constraint leads to a poor extrapolation for temperatures be-low 400 K. The authors fitted an equation to their values of ( m m( ) (298.15K)H T H ο− ), “weighting the point at 298.15 K for the best fit of all results”. This seems to have en-tailed the use of some of the heat capacity results of Catalano and Stout [55CAT/STO]. The reported value of ,mpCο (298.15 K) (62.93 J·K–1·mol–1) differs by 1.1 J·K–1·mol–1 from the value of [55CAT/STO], though the mS ο value from that source is used. Also, the expression reported by Binford and Hebert

[ m m( ) (298.15K)H T H ο− ] /J·mol–1 = 4.184·(15.677(T/K) + 1.7344 × 10–3 (T/K)2 + 1.4872 × 105(T/K)–1 –5161.4)

is non-zero (693 J·mol–1) for 298.15 K. In the present review, the drop calorimetry re-sults are used together with the heat capacity results of Catalano and Stout [55CAT/STO] to derive a properly constrained expression for ( m m( ) (298.15K)H T H ο− ) and ,mpCο (T). Above 1300 K, the experimental values of ( m m( ) (298.15K)H T H ο− ) appear to be slightly larger (by 0.5 kJ·mol–1 to 1.3 J·K−1·mol−1) than might be expected by extrapolation of values from lower tempera-tures. Because there are only two measurements in that temperature range, these values have not been treated differently from the other measurements in the fitting process.

[70CAR/LAI]

The solubility product of NiS was determined by linear voltage sweep voltammetry using a mercuric sulphide coated electrode (hanging mercury drop). The peak potential for the exchange reaction between the mercuric sulphide coated mercury electrode and Ni2+ ions of a nickel perchlorate solution,

HgS + 2e– + Ni2+ Hg + NiS (A.38)

yields the ratio between the solubility constants of HgS and NiS. At 298.15 K the au-thors obtained 10 ,0log sK ο = – 17.8 for the solubility product of NiS. Unfortunately, nei-ther any further experimental details nor the second dissociation constant for H2S(aq) are indicated in this contribution. Thus, the data cannot be recalculated and compared with results of different studies.

[70EFI/KUD]

The enthalpies of solution in water of nickel chloride, cobalt chloride and the coordina-tion compounds RbNiCl3, RbCoCl3, Rb2CoCl4 and Rb3CoCl5 were determined in a calo-rimeter with an “isothermal jacket”. The authors reported for the reaction,


340

NiCl2(cr) Ni2+ + 2Cl– (A.39)

an enthalpy of solution of − (83.26 ± 0.42) kJ·mol–1 for 10000 ≦ n(H2O) / n(NiCl2, cr) ≦ 40000); this leads to a value of sol mH ο∆ (A.39), between – 83.62 kJ·mol–1 and – 83.93 kJ·mol–1, which is approximately 1 kJ·mol–1 less negative than the value reported in [90EFI/FUR]. The latter (– (84.89 ± 0.27) kJ·mol–1) has been better documented and thus is accepted in this review.

[70HAL/VAN]

The formation of nickel(II) chloro complexes has been studied at 298.15 K in 4 M Na(ClO4,Cl). The solution containing a known total concentration of nickel (0.1, 0.2 or 0.5 M) and chloride (0 – 2.0 M) was saturated with AgCl labelled with radioactive 110Ag. The measured radioactivity was correlated with the free chloride concentration taking into account the solubility of AgCl and the different chloro complexes of sil-ver(I). The experimental data have been interpreted in terms of the unique formation of the bis-complex NiCl2 (i.e., β1 = (0.000 ± 0.015) M–1; β2 = (0.129 ± 0.002) M–2). This interpretation solely in terms of the formation of the NiCl2 complex is questionable, and probably was adopted because of effects from the marked change of the ionic medium during the measurement. The medium effects, however, are less than in most of the other work on chloro complexation of Ni(II). Therefore, we re-evaluated the published experimental data. Taking into account only the bis-complex, a better fit can indeed be achieved ( 10 2log b = – (0.89 ± 0.02)), compared to the result obtained using the assump-tion of the unique formation of NiCl+ ( 10 1log b = – (0.62 ± 0.04)). Nevertheless, in this review we prefer the latter value with a higher uncertainty ( 10 1log b = − (0.62 ± 0.30)). This choice is also confirmed by several equilibrium studies performed in non-aqueous solvents, which have higher precision since the complexes are more stable than in aque-ous solution. For example, in methanol, ethanol and propanol, K1 is much higher than K2 [95KAH/CRO]. Using the recommended values for ε(Na+, Cl–), ε(Ni2+, 4ClO− ) and the ε(NiCl+, 4ClO− ) derived from the data published in [75LIB/TIA], the extrapolation to I = 0 resulted in 10 1log οb = (0.68 ± 0.40).

The experimental work is well documented in this paper, which offers the pos-sibility of using the SIT to assess the medium effect caused by the replacement of the original background electrolyte. The activity of a cation M of charge zM in a mixture of two electrolytes (NX and NY, with a molality mX and mY) at constant ionic strength can be expressed using the SIT as follows:

log10 γM = – 2Mz D + ε(M, X) mX + ε(M, Y) mY

where D is the Debye-Hückel term. The same equation for an anion Y is:

log10 γY = – 2Yz D + ε(N, Y) (mX + mY).

Substituting the log10 γi values in Equations (B.3) (Appendix B) with these ex-pressions, applying it to the equilibrium:


341

Ni2+ + Cl– NiCl+

and rearranging leads to:

10 1log b + 4D – {ε(Na+, Cl–) + ε(Ni2+, 4ClO− )}Im – ε(Ni2+, 4ClO− ) mY – ε(Ni2+, Cl–) mY = 10 1log οb – ε(NiCl+, 4ClO− )Im + ε(NiCl+, 4ClO− ) mY – ε(NiCl+, Cl–) mY

provided that N = Na+, M = Ni2+, X = 4ClO− , Y = Cl– and Im = mX + mY. Since all quan-tities on the left-hand side of this equation are known, it can be written:

Q = 10 1log οb – ε(NiCl+, 4ClO− )Im + ε(NiCl+, 4ClO− ) mY – ε(NiCl+, Cl–) mY

The three unknown quantities on the right-hand side are strongly correlated with each other; thus, to solve this equation we should accept the ε(NiCl+, 4ClO− ) value derived from the data published in [75LIB/TIA]. Then the equation is transformed into

Q’ = 10 1log οb – ε(NiCl+, Cl–) mY (A.40)

Plotting the Q’ values vs. Cl

m − gives a straight line, with the intercept and slope equal to 10 1log οb and – ε(NiCl+, Cl–), respectively.

Before applying this equation to the data published in [70HAL/VAN], two fur-ther simplifications should be used. With increasing chloride concentration the ionic strength of the solution expressed in molarity is constant, but in the molal scale it changes. Since no density data are available for such an electrolyte mixture, Im was cal-culated using the equation:

Cl ClM × + (4 M ) ×

= 4m

a bI

− −−

where MCl– is the concentration of chloride in molar units, and a and b are the molality of 4 M NaCl and NaClO4 solution, respectively. The density of solutions with different compositions was calculated assuming that 4 M NaCl and 4 M NaClO4 solutions mix ideally. With these assumptions, the following dataset can be calculated.

The plot of Q’ versus Cl

m − is depicted in Figure A-19. From this plot 10 1log οb = (0.20 ± 0.03) and ε(NiCl+, Cl–) = – (0.16 ± 0.02) are calculated. This value of 10 1log οb is 0.5 log units lower than that calculated above from the accepted 10 1log b

value at I = 4 M Na(ClO4, Cl) using the conventional SIT extrapolation. This difference is, in great part, due to the medium effect. But it is also the consequence of the sensitiv-ity of function Q’ to the values of different interaction coefficients used (especially on whether the value of ε(NiCl+, 4ClO− ) is correct), and on the validity of the assumptions made concerning Im and mCl–. In this review we accept the value obtained from the plot in Figure A-19, but with a considerably increased uncertainty:

10 1log οb = (0.20 ± 0.50).


342

Table A-14: Dataset.

[Cl–]tot [Cl–]free [Ni2+]tot β1 10 1log b (molal)Cl

m − Im D Q’

0.25 0.2375 0.50 0.1078 – 1.0600 0.3058 4.9139 0.2610 0.2453 0.30 0.2835 0.50 0.1201 – 1.0129 0.3665 4.9067 0.2609 0.2754 0.38 0.3525 0.50 0.1335 – 0.9672 0.4572 4.8958 0.2609 0.2955 0.40 0.3900 0.20 0.1345 – 0.9639 0.4873 4.8922 0.2608 0.2904 0.45 0.4440 0.10 0.1433 – 0.9363 0.5474 4.8850 0.2608 0.3011 0.50 0.4651 0.50 0.1616 – 0.8841 0.6074 4.8778 0.2608 0.3364 0.65 0.6299 0.20 0.1770 – 0.8447 0.7863 4.8561 0.2606 0.3254 0.70 0.6880 0.10 0.1979 – 0.7962 0.8456 4.8489 0.2606 0.3572 0.75 0.7000 0.50 0.1587 – 0.8919 0.9047 4.8416 0.2605 0.2449 0.90 0.8640 0.20 0.2543 – 0.6872 1.0809 4.8200 0.2604 0.4000 1.00 0.9049 0.50 0.2596 – 0.6783 1.1975 4.8055 0.2603 0.3760 1.15 1.0960 0.20 0.3377 – 0.5640 1.3710 4.7838 0.2602 0.4413 1.20 1.1700 0.10 0.3656 – 0.5295 1.4285 4.7766 0.2601 0.4596 1.25 1.1241 0.50 0.2994 – 0.6162 1.4858 4.7694 0.2601 0.3567 1.40 1.3292 0.20 0.4120 – 0.4777 1.6567 4.7477 0.2599 0.4471 1.45 1.4099 0.10 0.4741 – 0.4167 1.7133 4.7405 0.2599 0.4921 1.50 1.3262 0.50 0.4019 – 0.4885 1.7697 4.7333 0.2598 0.4044 1.65 1.5570 0.20 0.5578 – 0.3461 1.9378 4.7116 0.2597 0.4993 1.75 1.5325 0.50 0.5024 – 0.3915 2.0489 4.6971 0.2596 0.4225 1.90 1.7840 0.20 0.7737 – 0.2040 2.2143 4.6755 0.2595 0.5634 1.95 1.8863 0.10 0.9322 – 0.1230 2.2690 4.6682 0.2594 0.6288 2.00 1.7362 0.50 0.6433 – 0.2842 2.3236 4.6610 0.2594 0.4523 2.15 1.9989 0.20 1.5441 0.0961 2.4861 4.6393 0.2592 0.7866 0.20 0.2000 0.10 0.0007 – 3.2532 0.2450 4.9211 0.2610 – 1.9308 0.95 0.8890 0.10 1.7597 0.1529 1.1393 4.8127 0.2603 1.2236 1.70 1.6106 0.10 5.2113 0.6244 1.9934 4.7044 0.2597 1.4541 0.40 0.3940 0.20 0.0785 – 1.1980 0.4873 4.8922 0.2608 0.0563


343

Figure A-19: Extrapolation to I = 0 of the formation constant of NiCl+ using Equation (A.40) and the experimental data published in [70HAL/VAN]. The outlying points (open diamond) were not taken into account in the calculation.

[70MAC]

Potential sweep experiments were performed on film nickel hydroxide electrodes in an effort to determine the redox mechanism. The over-all reaction for α-nickel hydroxide was found to be:

2 2 2 23 7 32[3Ni(OH) ·2H O] + 6OH + KOH 2[3NiOOH· H O· KOH] + 3H O + 6e .2 2 4

− −

The rate-controlling step was assumed to be proton diffusion.

[70MIR/MAK]

This paper reports association constants for the inner- and outer-sphere complexes (β1 and ω1, respectively) of Co(II), Ni(II), and Cu(II) hexa-aqua ions with chloride, bro-mide, thiocyanate, sulphate, and nitrate ions, based on spectrophotometric and solubility measurements in 3 M Li(ClO4, X) media at 298.15 K. The authors used the UV-region of the electronic spectra to determine the sum of β1 and ω1, while the value of β1 was determined from the VIS-region. This is obviously only a rough approximation. Only the thiocyanate ion is proposed to form an inner-sphere complex (approximately 10% in the case of nickel(II)). The sulphate value reported is an average of values determined by UV spectroscopy and from the effects of nickel ions on the solubility of calcium sul-phate. Apparently the two techniques gave reasonable agreement, as the average value of 10 1log οb was reported as (1.8 ± 0.2). This value is smaller than most, and there are insufficient details in the paper to allow the reviewer to properly assess or recalculate the results. In attempting to assess the values for the present review, the SIT was used to

-2.5

-2

-1.5

-1

-0.5

0

0.5

1

1.5

2

0 0.5 1 1.5 2 2.5

m (Cl− ) / mol·kg−1

Q'


344

calculate 10 1log K ο = 1.07, again indicating that the value obtained by Mironov et al. is anomalously low.

The authors provided insufficient experimental details to assess the reliability of their data, although the sum of the β1 and ω1 values are in agreement with other re-ports.

[70NAG/ING]

The partial pressure of sulphur over Ni – S melts was derived from the equilibrium weight of the melt in gas streams of known H2S/H2 compositions. The mass of the melt was monitored continuously, while the composition of the gas stream was maintained constant by appropriate flow meters. Experimental results for melts saturated with nickel as well as for homogeneous melts, with compositions ranging from 33 to 42 at% S, were obtained at temperatures between 973 and 1373 K. The sulphur partial pressure for a given gas composition was calculated from the well-known equilibrium constants for the reaction:

H2S(g) H2(g) + ½S2(g).

[70ROT]

Experimental levels of the configuration 3d94p, 3d84s4p and 3d95p of (NiI) were com-pared with corresponding calculated values. The electrostatic interactions between the configurations were considered explicitly.

[70TUR/RUV]

This paper describes polarographic measurements done at (25.0 ± 0.1)°C, I = 5.0 M (NaClO4), 1 × 10–3 M Ni(NO3)2, 0.0 – 0.2M Na2SO4 and with enough HClO4 to adjust the pH to 3. A value of K1 = 15.6 is reported for nickel sulphate association. The ionic strength is too high (6.6 m) to properly apply the SIT. If the SIT is used with ε(Ni2+, 4ClO− ) equal to (0.37 ± 0.03) kg·mol–1, 10 1log K ο = 1.70 is calculated.

[71ARI/MOR]

The enthalpies of formation were determined for several nickel sulphides, including Ni3S2, NiS (various compositions of the high-temperature form) and NiS2. The results are summarised in Table A-15. No experimental details are given. The mean value for the standard enthalpy of formation of the NiAs-type NiS is calculated to be f mH ο∆ = − (85.8 ± 2.4) kJ·mol–1 which agrees within the limits of uncertainty with the value rec-ommended in the present compilation ( f mH ο∆ = – (88.1 ± 1.0) kJ·mol–1).


345

Table A-15: Experimental results for the enthalpy of formation for various nickel sul-phides determined by calorimetric measurements.

Composition Phases f mH ο∆ (kJ·mol–1)

NiS0.67 Ni3S2 – (202.1 ± 3.9)

NiS0.98 Ni3S2 + NiS(α) – (85.0 ± 2.6)

NiS1.00 NiS(α) – (84.9 ± 2.6)

NiS1.01 NiS(α) – (87.4 ± 2.5)

NiS1.02 NiS(α) – (84.9 ± 2.5)

NiS1.04 NiS(α) – (86.5 ± 2.4)

NiS1.053 NiS(α) – (85.4 ± 2.4)

NiS1.06 NiS(α) + NiS2 – (85.7 ± 2.4)

[71BHA/SUB]

The hydrolysis of nickel(II) was studied in 0.5 M NaClO4 medium at 30°C. The concen-tration of the metal ion ranged from 0.01 to 0.04 M, while the pH varied between 6.8 and 7.1. The graphical analysis of the data, assuming only one hydroxo species, indi-cated the presence of 2

2 6Ni (OH) − with a 10 6,2*log b = – (42.94 ± 0.04). The authors used

too narrow a range of metal ion concentrations (and too narrow a range of pH for each [Ni2+]tot) to allow a satisfactory evaluation of the composition of the oligonuclear spe-cies. The reliability of the results is questionable; therefore, the results from this publi-cation were not used further in this review.

[71BON]

The author earlier suggested a simple calculation method for interpreting metal-ion complexation behaviour from polarographic studies of irreversible electrode processes [70BON]. The present paper was devoted to a test of the validity of the proposed method. One of the major conditions is that the heterogeneous charge-transfer rate con-stant should not change markedly with addition of the ligand to the aquo complex, i.e., the removal of coordinated water also should remain the rate determining step of the reduction in the presence of ligand. This condition was found to be fulfilled in case of simple 1:1 complexes of nickel(II), due to the slow water exchange in both the aquo and [Ni(H2O)5X]+ complexes. Therefore, the proposed simple method, combined with the De Ford-Hume equation, was used to determine the stability of the monofluoro, monochloro and mononitrato complexes of nickel(II). During the measurements, to achieve a measurable effect of complex formation, the background electrolyte (NaClO4) was entirely replaced by the sodium salt of the corresponding anion. The author men-tioned that, taking into account the limits of the method, the strong medium effect, and the weak complexation (especially in the cases of chloride (β1 = 0.6) and nitrate (β1 = 0.4)), the reported values can be considered only as order of magnitude estimates.


346

[71BUR/ZIN]

The hydrolysis of nickel(II) in 3 M (Na)Cl was studied at 60°C by potentiometric titra-tions. The nickel(II) concentration was [NiCl2]tot = 0.2 to 1.5 M and pH range was from 4.9 to 6.3. The reported constants are 10 4,4

*log b = − (25.33 ± 0.02) and 10 1,2*log b =

− (8.5 ± 0.2). The medium effect and thus the change of ionic strength was very impor-tant in these measurements (I varied from 3.2 to 4.5 M, i.e., pure NiCl2 solution was used for the highest nickel(II) concentration). Therefore, the reported experimental data were re-evaluated for the purposes of this review. For the reaction:

Ni2+ + Cl– NiCl+ (A.41)

at Im = 3.2 m, and T = 333.15 K, 10 1log b (A.41) = – 0.44 can be estimated for the NiCl+ complex in NaCl medium, accepting the values of 10 1log b ((A.41), 298.15 K) = – 0.59 (see discussion on [65BUR/LIL2]) and r mH∆ = (8.2 ± 2.5) kJ·mol–1 [66KEN/LIS]. Con-sidering the whole dataset and the same model proposed by the authors, 10 4,4

*log b = − (25.29 ± 0.05) (3σ) and 10 1,2

*log b = – (8.50 ± 0.09) can be calculated, in good agree-ment with the published values. If only the data sets reported for [Ni2+]tot = 0.2, 0.4, 0.6 and 0.8 M are used, taking into account the formation of NiCl+, 10 4,4

*log b = − (24.04 ± 0.06) (3σ) and 10 1,2

*log b = – (7.92 ± 0.10) can be derived. If the species NiOH+ is added to the equilibrium model, a substantial (30%) decrease of the fitting parameter was achieved, and the values 10 4,4

*log b = – (24.09 ± 0.03) (3σ), 10 1,1*log b =

– (8.48 ± 0.10) and 10 1,2*log b = – (8.58 ± 0.40) were obtained. The latter values, with

increased uncertainties, are accepted for use in this review: 10 4,4*log b =

− (24.09 ± 0.30), 10 1,1*log b = – (8.48 ± 0.30) and 10 1,2

*log b = – (8.58 ± 0.60). Combin-ing these constants with the values from [65BUR/LIL2], r mH∆ = (59 ± 10) kJ·mol–1, (20 ± 20) kJ·mol–1 and (186 ± 6) kJ·mol–1 are calculated for the formation of NiOH+, Ni2OH3+ and 4

4 4Ni (OH) + species, respectively (in 3.2 m aqueous NaCl).

[71CON/LOO]

Connelly et al. measured the specific heat of pure single-crystalline nickel near the Cu-rie temperature using a calorimetric technique which permits direct observation of

,mpCο (T) as a continuous function of T, with a temperature resolution of about 10–2 K. Special attention has been devoted to the determination of the analytical form of the magnetic contribution to ,mpCο (T).

[71DAS/DAS]

Potentiometric measurements have been performed at 25, 35 and 45°C to determine thermodynamic parameters of the mono-thiocyanato complexes of five 3d metal ions, among them nickel(II). The activity coefficients of the different species were calculated by the Davies equation. Since the Davies equation is not compatible with the SIT, and no experimental data are provided, the reported thermodynamic functions were not con-sidered in this review.


347

[71HOH/GEI]

A stopped-flow method was used to study the hydrolysis of nickel(II). The ionic strength of the solutions was not controlled ([Ni2+]tot = 5×10–4 to 2.5×10–3 M, [OH]tot ~ 3×10–4 M); the measurement temperature was not specified in the paper. The results indicated, that using different time scales, the deprotonation of 2

2 6Ni(H O) + ion and the polynuclear hydroxo species can be separated. The deprotonation reaction for the mononuclear species is estimated to occur within 10–7 s, while the polynuclear complex formation needs several seconds (20 min. is reported to reach the equilibrium in [65BUR/LIL] and [69KOL/KIL]). The pH of the solutions, measured between 2 – 1000 ms by spectrophotometry using phenolphthalein as the indicator (the value used for the pK of the indicator was not reported), was extrapolated to t = 0 s. From these data

10 1,1*log b = (11.0 ± 0.2) has been derived. This value has not been considered further in

this review, because of the lack of experimental details.

[71ISO]

Freezing point measurements were carried out in a manner similar to that of Brown and Prue [55BRO/PRU], and similar results were obtained. Analysis of the saturated solu-tions was done by comparison of the conductivity with the conductivity of standard so-lutions. An extended Debye-Hückel treatment was used to analyse the data, and values of K1 =118 to 250 dm3·mol–1 were obtained for values of the Debye-Hückel “a” parame-ter of 0.4 to 1.4 nm. The author reported that the calculated association constant also depends on the choice of the maximum solute concentration for which the data analysis can be considered reliable. In the present review the data were reanalysed using the SIT, and only data for total nickel sulphate molalities ≤ 0.03 were used.

[71IZA/JOH]

The thermodynamic parameters for the formation of several metal-cyanide complexes, among others those of 2

4Ni(CN) − , have been determined using pH-metric and calorimet-ric methods at 10, 25 and 40°C. In case of nickel(II), the thermodynamic data were de-termined by titration of Ni(ClO4)2 solutions with NaCN solutions. The ionic strength of the solutions were I < 0.02 M in all cases. The Debye-Hückel equation, related to the SIT model, was used to correct the formation constants to thermodynamic constants valid at I = 0. Since previous experiments indicated that the dependence of r mH ο∆ in the ionic strength in dilute aqueous solutions is small compared to the experimental error, the measured heats of reaction ( r mH∆ = – 189.1 kJ·mol–1 at 10ºC; r mH∆ = – 183.7 kJ·mol–1 at 40ºC) were taken to be valid at I = 0, but the uncertainties were estimated in this review as ± 2.0 kJ·mol–1. From the values of r mH∆ as a function of temperature, average r ,mpCο∆ values were calculated.


348

[71LAN/CUM]

The distribution of 63Ni(II) between anion exchange resin and MSCN (M = Na+, K+, 4NH+ and Li+) solutions was determined using liquid scintillation techniques. The tem-

perature of the measurements was not given in the paper. The thiocyanate concentration ranged between 0.05 and 5.0 M, while [63Ni(II)]tot was varied from 4.26×10–7 to 2.92×10–6 M. No background electrolyte was used, thus the ionic strength varied over a wide range. The authors reported “semi-thermodynamic” formation constants (see [89BJE]) for the four thiocyanato complexes detected ( 2Ni(SCN) x

x− with x = 1 – 4),

using literature data for the main activity coefficients of KSCN and NaSCN solutions, and estimated values for NH4SCN. Since the temperature of the measurements was not specified, and the “semi-thermodynamic” formation constants are not compatible with values extrapolated to zero ionic strength using SIT, the reported data were not consid-ered further in this review.

[71LEB/SAV]

This paper describes the measurements of the heat of fusion of Ni and other refractory metals by an electrical explosion method i.e., the rapid heating by a current of high den-sity. The metal wires 0.05 – 0.1 mm in diameter and about 1 cm in length were heated by single square current pulses of density 1.5 – 4.0×1010 A·cm–2. Under these condi-tions, one can neglect the self-induction, the skin effect, and the energy losses, so the wire resistance and energy supplied can be calculated if the voltage and the time of the pulse are known. The maximum possible systematic error in energy determination was 8%. The authors assume that in the case of Ni this error is 1.5%. Unfortunately,

fus mH ο∆ (Ni) determined by the electrical explosion method fus mH ο∆ = (18.43 ± 0.27) kJ·mol–1 is significantly different from the values obtained by calorimetric studies of Vollmer et al. [66VOL/KOH] fus mH ο∆ = (16.90 ± 0.25) kJ·mol–1, and Geoffray et al. [63GEO/FER] fus mH ο∆ = (17.47 ± 0.23) kJ·mol–1 upon which our selected value is based. One possible reason for the observed differences could be the low-purity of nickel, 99.5%, used in the study of Lebedev and Savvatimskii.

[71NAV]

Navrotsky has measured the enthalpies of formation of Mg2SiO4, Co2SiO4, Ni2SiO4, Zn2SiO4, MgGeO3, Co2GeO4, CoGeO3, Ni2GeO4 and Zn2GeO4 from the component oxides by solution calorimetry in a molten oxide solvent at (965 ± 2) K. Regularities in the thermodynamics of formation of the mentioned silicates and germanates were dis-cussed. In the case of Ni2GeO4, 9PbO·3CdO·4B2O3 was used as a solvent and the en-thalpy of formation from oxides was determined to be – 39.75 kJ·mol–1. This value is identical with that previously given by Navrotsky and Kleppa [68NAV/KLE]. Using the calorimetrically determined enthalpy of formation from oxides for Ni2SiO4, an attempt was made to estimate the temperature of decomposition of Ni2SiO4 into NiO and SiO2. The obtained value, 2065 K, is significantly higher then the value accepted by this re-view (1820 ± 5) K.


349

[71OHT/BIE]

The hydrolysis of nickel(II) ion was studied by potentiometric titrations in 3 M (Na)Cl medium at 25°C. The concentration of nickel(II) ranged between 0.0145 and 1.0 M, thus the ionic strength varied from 3.0145 to 4 M. The evaluation of emf data, per-formed with both graphical and numerical methods, indicated the dominant formation of 4

4 4Ni (OH) + , together with some minor species NiOH+ and Ni2OH3+. The best fit to data was obtained by the following formation constants: 10 4,4

*log b = – (28.55 ± 0.02), 10 1,1

*log b = – (10.5 ± 0.1) and 10 1,2*log b = – (10.5 ± 0.5). The authors attempted to con-

sider the impact of formation of the NiCl+ complex, but they considerably underesti-mated this effect.

Accepting for the reaction,

Ni2+ + Cl– NiCl+ (A.42)

10 1log b (A.42) = – 0.59 in 3.2 m NaCl medium (see the discussion on [65BUR/LIL2]), ca. 38 – 42 % of nickel(II) is present as the chloro complex. Considering this in the mass action equations, the corrected formation constants are as follows: 10 4,4

*log b = − (27.60 ± 0.02), 10 1,1

*log b = − (10.26 ± 0.10) and 10 1,2*log b = – (10.0 ± 0.5).

Taking into account the medium effect caused by replacement of NaCl by NiCl2, the changing ionic strength and the uncertainties of the latter correction

10 4,4*log b = – (27.6 ± 0.3), 10 1,1

*log b = – (10.26 ± 0.20) and 10 1,2*log b = – (10.0 ± 0.6)

are accepted in this review.

[71PAA/HUM]

This is a spectrophotometric study on nickel(II) chloride complexes formed in aqueous NaClO4/LiCl solution. Complex formation was studied first at I = 6 M. The lithium chloride concentration was varied between 0 and 6 M; the ionic strength was maintained by sodium perchlorate. Under these conditions the results indicated the unique forma-tion of the NiCl+ complex and its conditional stability constant is reported to be β1 = (0.32 ± 0.05). During this study, both the anion and the cation of the original back-ground electrolyte were entirely replaced, which induced additional uncertainty in the value determined for the formation constant. The impact of the 4ClO− → Cl– change is difficult to quantify, but the discussion on [70HAL/VAN] may give an approximation. The impact of the Na+ → Li+ change is easier to assess. As discussed in [97GRE/PLY2] the following expression can be applied for 10 1log b using the SIT at constant ionic strength and total perchlorate concentration:

∆ 10 1log b = 410 1,NaClOlog b – 10 1,(Na,Li)ClO4

log b = [ε(Na+, Cl–) – ε(Li+, Cl–)]Li

m +

At I = 6 M, and in case of a complete Na+ → Li+ exchange, ∆ 10 1log b = – 0.48 can be calculated. It means that the replacement of both the cation and the anion of the background electrolyte each induces (as a rough approximation) an increase of 0.5 in the determined value of 10 1log b compared with the “true” value in NaClO4 (it is very


350

difficult to say if these effects are additive or not). Therefore, the reported constant can not be accepted.

At LiCl concentrations above 6 M, the formation of a further complex (NiCl2(aq)) was detected. Since the ionic strength between [LiCl] = 6–13.68 M could not be maintained constant, the activity of chloride ion and water was taken into ac-count, while the ratios

( ) ( )22 26 5Ni H O NiCl H O

/+ +γ γ and ( ) ( )2 2 25 4NiCl H O NiCl H O (aq)

/+γ γ were assumed to be nearly constant with increasing LiCl concentrations and were included in the re-ported “semi-thermodynamic” formation constants ( '

xοb = [ 2

2 6NiCl (H O) xx x

−− ]×aw /

[ 31 2 7NiCl (H O) x

x x−

− − ]×aCl–, where x = 1, 2). The reported values are '1οb = (4.9 ± 0.8)×10–2,

and '2οb = (7.8 ± 4.1)×10–5. Although the use of such constants generally provide higher

correctness in case of weak complexes (see the discussion on [89BJE]), '1οb was ob-

tained from the rather uncertain value of β1 valid for I = 6 M, thus only '2K ο = '

2οb / '

1οb

can be accepted in this review ( '10 2log K ο = – (2.8 ± 0.5)).

[71PIE/HUG]

The equilibrium constant for the reaction, 24Ni(CN) − + CN– 3

5Ni(CN) − K5

was determined by spectrophotometry (400 – 500 nm) at 20°C, I = 2 M (NaClO4(aq) containing 0.01 M NaOH (pH ~ 12)). The formation of the hexacyano complex was not observed even at a 100-fold excess of CN– over 2

4Ni(CN) − . The reported stability con-stant ( 10 5log K = – 0.77) was corrected to 25°C ( 10 5log K = – 0.80), using the selected reaction enthalpy.

[71TUR/MAL]

The effects of temperature (t = 5 to 35°C) and ionic strength (I = 0.075 to 0.2 M, KNO3 + KSCN) on the stability and kinetics of formation of nickel(II)-thiocyanate complexes were investigated at pH = 6.0 – 6.2, by detecting the catalytic polarographic currents in the nickel(II) – thiocyanate system. The metal ion concentration varied from 0.2 to 9.5 mM, while the thiocyanate concentration ranged between 4 and 40 mM. Under the con-ditions used, the formation of Ni(SCN)+ and Ni(SCN)2(aq) was detected. The reported formation constants are accepted, but uncertainties of ± 0.1 and ± 0.2 are assigned to

10 1log b and 10 2log b , respectively. From the temperature dependence of the formation constants (see Figure A-20) r mH∆ = – (14.3 ± 0.5) kJ·mol–1 and r mH∆ = − (29.5 ± 0.5) kJ·mol–1 can be derived for the formation of Ni(SCN)+ and Ni(SCN)2(aq), respectively. These values are selected in this review with both uncertainty values increased to ± 2.0 kJ·mol–1.


351

Figure A-20: Temperature dependence of 10 1log b (filled squares) and 10 2log b (open squares) values for the Ni(II) – (SCN)– system (I = 0.2 M) as reported in [71TUR/MAL].

[72AUF/CAR]

The heats of solution in water of the hexahydrate (actually Ni(NO3)2·6.055H2O) and the tetrahydrate were measured as a function of final solution concentration. Extrapolated “infinite dilution” values of 31.23 kJ·mol–1 and 8.66 kJ·mol–1 were reported. The au-thors did separate extrapolations for the two sets of measurements (and also for the measurements using the dihydrate [72AUF/CAR2]). The enthalpy of dilution correction to obtain a value for the enthalpy of solution at I = 0 should be identical (within the ex-perimental uncertainties) for the hexahydrate, tetrahydrate and dihydrate [72AUF/CAR2], and the correction should be consistent with whatever correction is applied to the heat of solution results of Chukurov et al. [73CHU/DRA] (for the tetra-hydrate). The experimental scatter is fairly large (as much as 0.4 kJ·mol–1) for final nickel concentrations < 0.015 m. It would appear that the enthalpy of dilution curve is similar to those for potassium sulphate and barium nitrate [59LAN].

For each of the three salts [72AUF/CAR], [72AUF/CAR2], the values for the five lowest concentrations were averaged to estimate a value at the third-lowest concen-tration with a solvent:salt ratio of 8000:1. The difference between each experimental heat of solution measurement and the measurement (for the same hydrate) resulting in a solution with a solvent:salt ratio of 8000:1 (0.00694 m) was calculated. Only values from the measurements with final solution concentrations below 0.056 m were used. Different values from 650 J·mol–1 to 200 J·mol–1 were assumed for the heat of dilution to zero molality of the solution with a solvent:salt ratio of 8000:1 (0.00694 m). Chuku-rov et al. [73CHU/DRA] provided an equation for heat of solution of the tetrahydrate in the concentration range 0.007 m to 0.03 m, and difference values based on that equation

0.0032 0.0033 0.0034 0.0035 0.00361.0

1.5

2.0

2.5

log 10

βx

K /T


352

and the same assumed heat of dilution were also calculated. The derived heats of dilu-tion were plotted as a function of m½ (Figure A-21), as was a line with the limiting slope for 1:2 and 2:1 electrolytes [59LAN].

If the enthalpy of dilution from a solution with a solvent:salt ratio of 8000:1 to infinite dilution is assumed to be of the order of (520 ± 100) J·mol–1, all the experimen-tal results enthalpy of solution are in reasonable agreement with each other. However, the values including those of Chukurov et al. [73CHU/DRA] at the lower concentra-tions) are compatible with the limiting-law slopes found for other 1:2 and 2:1 electro-lytes [59LAN].

Figure A-21: Heats of dilution of hydrated nickel nitrate salts based on the experimental heats of solution and an assumed heat of dilution to infinite dilution of 520 J·mol–1 for a solution with a solvent:salt ratio of 8000:1. Heat of solution values for [73CHU/DRA] were based on the equation of the concentration dependence as reported in that paper (see text and Appendix A entry for [73CHU/DRA]).

Using the estimated heat of dilution and average values of the heats of solution at concentrations at or below 0.01 m, the heats of solution corrected to “infinite dilu-tion” for the hexahydrate, tetrahydrate and dihydrate are (30.76 ± 1.00) kJ·mol–1, (8.23 ± 1.00) kJ·mol–1 and – (27.64 ± 1.00) kJ·mol–1. The uncertainties have been esti-mated in the present review based on the values shown in the authors’ Figure 1 and sys-tematic uncertainties generally associated with similar measurements.

0.00 0.05 0.10 0.15 0.20 0.250

500

1000

1500

2000

theoretical limiting slope [73CHU/DRA] tetrahydrate [72AUF/CAR] tetrahydrate [72AUF/CAR] hexahydrate [72AUF/CAR2] dihydrate

heat

of d

ilutio

n / J

·mol

−1

(m / mol·kg−1)0.5


353

The equilibrium water vapour pressure at which the hexahydrate begins to lose water to form a lower hydrate was determined from 15°C to 42°C. The vapour pressure was found to follow the equation:

log10( 2H Op /atm) = ((7.7226 – 3039.2·T–1 ) ± 0.078) or

2H Op /bar = 1.01325 × 1((7.7226 3039.2 ) 0.078)10 T −− ⋅ ±

= 1((7.7283 3039.2 ) 0.078)10 T −− ⋅ ±

and at 298.15 K, if the reaction is assumed to be:

Ni(NO3)2·6.00H2O(cr) Ni(NO3)2·4.00H2O(cr) + 2H2O(g) (A.43)

r mGο∆ ((A.43), 298.15 K) = (28.14 ± 0.89) kJ·mol–1 and r mH ο∆ ((A.43), 298.15 K) = (116.4 ± 10.9) kJ·mol–1. If the heat of vaporisation is calculated from the enthalpy of solution measurements, and the accepted values for the enthalpies of formation of H2O(l, g), r mH ο∆ ((A.43), 298.15 K) = (110.4 ± 1.3) kJ·mol–1. The enthalpy of reaction value from the vapour pressure measurements agrees with the value from calorimetry within the combined uncertainties, though the uncertainty is much lower for the value from calorimetry. Unfortunately, the method used to measure the water vapour pressure does not allow unambiguous identification of the lower hydrate.

[72AUF/CAR2]

The heat of solution in water of the dihydrate (Ni(NO3)2·2.00H2O) was measured as a function of final solution concentration. An extrapolated “infinite dilution” value of − (27.41 ± 1.00) kJ·mol–1 was reported, and a recalculation has been done in the present review (see the Appendix A for [72AUF/CAR]).

The water vapour pressure over mixtures of the tetrahydrate and dihydrate were also measured from 32°C to 73°C. The vapour pressure was found to follow the equation

log10( 2H Op /atm)= ((7.5527 – 3179.4 ·T–1 ) ± 0.094) or

2H Op /bar = 1.01325 × 1((7.5527 3179.4 ) 0.094)10 T −− ⋅ ±

= 1((7.5584 3179.4 ) 0.094)10 T −− ⋅ ±

and at 298.15 K, if the reaction is assumed to be:

Ni(NO3)2·4.00H2O(cr) Ni(NO3)2·2.00H2O(cr) + 2H2O(g) (A.44)

r mGο∆ ((A.44), 298.15 K) = (35.45 ± 1.09) kJ·mol–1 and r mH ο∆ ((A.44), 298.15 K) = (121.7 ± 9.3) kJ·mol–1. If the heat of vaporisation is calculated from the enthalpy of so-lution measurements, and the accepted values for the enthalpies of formation of H2O(l, g), r mH ο∆ ((A.44), 298.15 K) = (124.0 ± 1.3) kJ·mol–1. The enthalpy of reaction value from the vapour pressure measurements agrees with the value from calorimetry


354

within the combined uncertainties, though the uncertainty is much lower for the value from calorimetry. Unfortunately, the method used to measure the water vapour pressure does not allow unambiguous identification of the lower hydrate, and no value for

r mGο∆ ((A.44), 298.15 K) can be accepted based on this work.

[72BON/HEF]

The fluoride complexes of the divalent ions from Mn(II) to Zn(II) have been studied by potentiometric titration, using a fluoride selective electrode at 298.15 K and I = 1.0 M NaClO4. Under the experimental conditions employed (25 cm3 0.05 M metal(II) per-chlorate solution was titrated by 10 cm3 0.004 M NaF solution), only monofluoride complexes are formed but in rather low concentrations (at the end of titrations [NiF+] ~ 2.2×10–6 M can be calculated). This yielded only a moderate effect (1 – 2 mV) in the measured electrode potential, therefore we estimate higher uncertainty in K1 ((2.2 ± 1.0) for NiF+) in the present review, than suggested by the authors. The authors reported an anomalous stability sequence in the Mn(II)-Zn(II) series which does not fit an Irving-Williams type trend. This behaviour was later confirmed by several authors (see e.g., [83SOL/BON]).

[72FLE]

The high temperature phase α-Ni7S6 was synthesised by heating Ni sponge and S crys-tals in an evacuated silica glass tube at 504°C for 9 days; the Ni sponge was reduced by hydrogen at 900°C before use.

[72FRE/STU]

Complexation concentration quotients for 15°C and a medium of 0.1 M KNO3(aq) were reported for the formation of NiHPO4 (K(A.47) = 100), 2

2 7NiP O − ( 10log K (A.45) = 6.22) and 2 7NiHP O− ( 10log K (A.46)= 3.50).

Ni2+ + 42 7P O − 2

2 7NiP O − (A.45)

Ni2+ + 32 7HP O − 2 7NiHP O− (A.46)


The values for the pyrophosphate species were cited and used in a later paper from this group [78FRE/STU]. The potentiometric measurements were carried out by titration of the free acid with KOH(aq) solution in the presence and absence of Ni(NO3)2 in the solutions. The metal:ligand concentrations were varied from 1:2 to 2:1. Ligand concentrations were 2×10–3 to 5×10–3 M. The authors estimated 5% to 10% uncertain-ties in the experimentally determined stability constant values. The electrodes were calibrated against pH buffers rather than concentration standards, and the hydrogen ion activity coefficient, +H

γ was assigned the value of 0.83; the SIT procedure described in Appendix B would have led to +H

γ = 0.79. If the concentration quotients are adjusted


355

for this, the value for formation of NiHPO4 is K = 105) and for 22 7NiP O − ,

10log K (A.45) = 6.24 (the value for 10log K (A.46) is unchanged).

Correction to zero ionic strength (Appendix B) gives 10log K ο (A.47) = (2.88 ± 0.30), where the uncertainty has been assigned in the present review. The small ∆ε correction has been based on the value for ε(Na+, 2

4HPO − ) from Table B-5 (rather than using the equation in Table B-6), and ε(Ni2+, 3NO− ) = 0.18 kg·mol–1.

For Reaction (A.45) in 0.1 M KNO3, 10log K (A.45) at 15°C = (7.94 + 0.1∆ε). For Reaction (A.46) in 0.1 M KNO3, 10log K (A.46) at 15°C = (4.79 + 0.1∆ε).

As the ionic strength is low, the effect of the ∆ε term on the final calculated values of 10log K ο is limited. Rather than use estimated ε values for interactions involv-ing the pyrophosphate moieties (some of which are highly charged), it is simply as-sumed that ∆ε is zero. As a consequence, the uncertainty in the 10log K ο values is in-creased to ± 0.25 for the values from measurements in 0.1 M KNO3 solutions. As dis-cussed in Appendix A for [73PER/SEC], the fact that the calculations were done with-out directly considering the effect of association of K+ with pyrophosphate means that these are “apparent” values (useful for comparisons, but not adequate as selected val-ues). The raw data are unavailable for re-analysis.

[72KLE/KOM]

Based on thermal and X-ray measurements the complete Ni–Te phase diagram was con-structed.

[72POL/HER]

Polgar et al. carried out calorimetric measurements of the specific heat of NiCl2·2H2O(cr) for temperatures between 1.2 and 25.5 K. Peaks were observed at 6.31 and 7.26 K, and at lower temperatures the heat capacity values are not a simple function of T 3. The reasons for these magnetic transitions are not clear. The authors estimated that the contribution to the molar entropy between 0 and 1.2 K is 0.03 R. In the present review, the uncertainty in the extrapolation below 1.5 K is estimated to be of the order of ± 0.2 J·K–1·mol–1.

[72PRE/RUG]

Predel and Ruge reported a value of – 8600 cal·g-atom–1 for the heat of formation of NiAs(cr) based on an unreported value for the heat of solution of NiAs in liquid tin over an unspecified temperature range. Based on another paper from the same group [72PRE/RUG2], it would appear that the reaction was carried out between 600 and 900 K. The value is not claimed to be a value corrected to 298.15 K (though Mah and Pankratz [76MAH/PAN], Barin et al. [77BAR/KNA], and Skeaff et al. [85SKE/MAI] appear to have accepted it as the value for that temperature). If the value is assumed to be an average value obtained at 800 K, and the extrapolated heat capacity function of


356

Muldagalieva et al. [95MUL/CHU] is applied (with values for mH (800 K) – mH ο (298 K) for Ni(cr) from the present review and for As(cr) from Herrick and Feber

[68HER/FEB]), the value:

f mH ο∆ (NiAs, cr, 298.15 K) = – 71.965 – 30.21 + 15.42 + 13.15 kJ·mol–1,

i.e., f mH ο∆ (NiAs, cr, 298.15 K) = – 73.60 kJ·mol–1 can be calculated. The authors claimed that their results were accurate within 5%. In the present review, this is consid-ered to be a 1σ uncertainty, and an overall uncertainty of ± 8.00 kJ·mol–1 is accepted.

[72RET/HUM]

The “semi-thermodynamic” equilibrium constant ( '2K ο ) for the reaction:

2 5NiBr(H O)+ + Br– NiBr2(H2O)4(aq) + H2O(l)

was estimated from spectrophotometric measurements in 8.5 – 11.7 M LiBr solutions. In the evaluation of '

2K ο , several simplifying assumptions were made: (i) the activity coefficient of bromide ion can be approximated by the mean activity coefficients of LiBr, (ii) the ratio of the activity coefficients of the two Ni bromide complexes is con-stant with increasing LiBr concentration, (iii) in 7 – 8 M LiBr solution the only nickel(II)-containing species is the 2 5NiBr(H O)+ complex.

[73CHU/DRA]

The heat of solution of Ni(NO3)2·4H2O in water was found to follow the equation:

sol mH ο∆ = (8.494 + 7.448 m0.5) kJ·mol–1

for final solution concentrations of nickel from 0.007 m to 0.03 m. Although the value as extrapolated to I = 0 (8.49 kJ·mol–1) is very similar to the one obtained by Auffredic et al. [72AUF/CAR] (8.66 kJ·mol–1), the concentration dependence (and hence, it is assumed, the actual values for the measurements at finite concentrations) is less simi-lar—with differences of 0.2 kJ·mol–1 to 0.4 kJ·mol–1 in the concentration range of the measurements (see the Appendix A discussion for [72AUF/CAR] for further details). Using the estimated heat of dilution from 0.007 m, and the value of the heat of solution as calculated using the authors’ equation for that concentration, the heat of solution for the tetrahydrate, corrected to “infinite dilution”, is (8.59 ± 1.00) kJ·mol–1. The uncer-tainty is estimated in the present review.

The authors also reported the heats of dissolution of “anhydrous” Ni(NO3)2 (− (117.2 ± 1.67) kJ·mol–1), water (– (5.31 ± 0.08) kJ·mol–1) and Ni(NO3)2·4H2O (− (63.18 ± 0.84) kJ·mol–1) in dimethyl sulphoxide. This would indicate an enthalpy of (75.23 ± 1.90) kJ·mol–1 for the dehydration reaction:

Ni(NO3)2·4H2O Ni(NO3)2 + 4H2O.

Except for salts containing singly charged cations, synthesis of anhydrous ni-trate salts is very difficult, as the final steps of dehydration tend to cause loss of nitrogen


357

oxides or nitric acid. Only the summary of the paper was available to the reviewers, not the full deposited document, and, in the summary, there were no details on the prepara-tion of the anhydrous salt. The experimental approach used to determine the enthalpy of formation is reasonable, and may lead to a useful value if the deposited paper contains proper analysis results for the solid and solutions. No value for anhydrous Ni(NO3)2 is selected from these data.

[73COR]

Cordfunke measured the solubility of Ni(IO3)2·2H2O between 29°C and 61°C. It is not clear from the text whether a pure sample of the dihydrate was used for the measure-ments or whether the samples were contaminated with the tetrahydrate. Also, the equili-bration times are not reported.

The results were shown graphically, and it was necessary to retrieve values by digitisation of the points plotted in the author’s Figure 2. These values (percent weight as Ni(IO3)2) were converted to molalities, and the apparent solubility products were corrected to I = 0. These were plotted against (T/K)–1 to obtain values for 10 ,0log sK ο at 298.15 K and an average value of r mH ο∆ (A.48) over the temperature range of the meas-urements.

Ni(IO3)2·2H2O Ni2+ + 2 3IO− + 3H2O(l) (A.48)

If only the values to 50°C are included, 10 ,0log sK ο ((A.48), Ni(IO3)2·2H2O, 298.15 K) = – (5.18 ± 0.04) and r mH ο∆ (A.48) = (21.7 ± 4.7) kJ·mol–1, whereas, if the points for temperatures above 50°C also are included, 10 ,0log sK ο ((A.48), Ni(IO3)2·2H2O, 298.15 K) = − (5.17 ± 0.04) and r mH ο∆ (A.48) = (20.5 ± 2.4) kJ·mol–1.

[73FED/SHM]

The solubility of Ni(IO3)2 in nitrate containing LiClO4 solutions (I = 0.5, 1.0, 2.0, 3.0, 4.0 M) has been measured to determine the stability and composition of the nitrato complexes formed with nickel(II). The authors reported the formation of 2

3Ni(NO ) xx

− complexes (x = 1 – 3). Since, during the measurements, the original background electro-lyte (LiClO4) was entirely replaced by LiNO3, the formation of bis- and tris-complexes can not be verified. Therefore, the experimental data published in [73FED/SHM] were re-evaluated taking into account solely the formation of 3NiNO+ species. To minimise the medium effect, only the experimental data [LiNO3]/[LiClO4] ≤ 1 (see Figure A-22) were taken into account. The non-linear regression of these data resulted in the follow-ing values: 10 1log b = − (0.20 ± 0.40) (I = 1.0 M), 10 1log b = − (0.21 ± 0.40) (I = 2.0 M),

10 1log b = − (0.30 ± 0.40) (I = 3.0 M), 10 1log b = − (0.01 ± 0.40) (I = 4.0 M). The uncer-tainties were assigned by the reviewer, taking into account the still substantial medium effect (also see Section V.6.1.2).


358

Figure A-22: Experimental [73FED/SHM] and calculated solubility of nickel(II) in the aqueous phase with increasing nitrate concentrations.

The values for the solubility of Ni(IO3)2·3H2O in aqueous 0.5 M to 4.0 M LiClO4 solutions (i.e., those solutions with no added nitrate) were used to calculate the value of the solubility product of Ni(IO3)2·3H2O (Figure A-23) from a standard SIT plot. This gives the value 10 ,0log sK ο = – (5.09 ± 0.16).

Figure A-23: Experimental [73FED/SHM] and calculated values of the solubility of Ni(IO3)2·3H2O in aqueous LiClO4 solutions.

0.015

0.02

0.025

0.03

0.035

0 0.5 1 1.5 2 2.5 3 3.5 4

[NO3-] / M

s(N

i2+) /

M

I = 1 M

I = 2 M

I = 3 M

I = 4 M

0.00 1.00 2.00 3.00 4.00 5.00-6.60

-6.20

-5.80

-5.40

-5.00

log 10

Ks -

6D

I / mol·kg−1


359

[73FLE]

The data concerning NiO are essentially the same as in [68CHA/FLE].

[73GUR/CAL]

The heat of solution of nickel chloride hexahydrate into 1.0M HCl(aq) was measured, and (6.996 ± 0.133) kJ·mol–1 was reported (2σ uncertainties). The final solution was 0.02 to 0.03 M in Ni. A double correction to 0.0 M HCl(aq) and to I =0 [95MAN/KOR] would be required to use the reported value to calculate the enthalpy of solution of the salt at infinite dilution in water. The corrections themselves have large uncertainties, and this heat of solution value is not used further in the present review.

[73HUT/HIG]

A kinetic method, based on the metal ion catalysed reaction of:

[Co(EDTA)Cl]2– [Co(EDTA)]– + Cl–,

has been used to determine the stability constants of several [MA]– complexes, among others those of the chloro-, bromo-, nitrato- and thiocyanato complexes of nickel(II) (pH = 4.5 – 5.5, T = 298 K, I = 1 M Na(ClO4, X)). The concentration of nickel(II) was varied up to 0.051 M. The highest concentrations of the complex-forming anions (X–) for nickel(II) in a given series were [Br–] = 0.49 M, [Cl–] = 0.278, 0.466 and 0.844 M, [ 3NO− ] = 0.422 M and [SCN–] = 0.15, 0.318 M. The formation constant determined at [Cl–] = 0.844 M was not considered in this review, in order to minimise the need to as-sess the medium effect. The authors acknowledged several disadvantages of the chosen reaction. The most important is that the monocationic complex [NiX]+, not just Ni2+, may catalyse the reaction. This effect was taken into account by a simple correction factor. Therefore, we assigned higher uncertainties to the reported values. The forma-tion constants of chloro-, nitrato- and thiocyanato- complexes also were determined at 318.15 K. The stability of the nitrato complex does not change at higher temperatures, but those of the chloro and thiocyanato complexes decrease moderately; this results in negative r mH∆ values. This observation is the opposite to what was observed in [60LIS/ROS] and [66KEN/LIS] for the NiCl+ complexes, but is in reasonable agree-ment with other reports [67NAN/TOR], [74KUL3] in the case of NiSCN+. Based on the entropies of association, the formation of mostly outer-sphere complexes (ion pairs) was suggested in the case of nickel(II), except for thiocyanate. From the temperature de-pendence of the formation constant for the NiSCN+ species r mH∆ = – (8.3 ± 1.0) kJ·mol−1 can be derived. This value is accepted with an increased uncertainty ( r mH∆ = − (8.3 ± 2.0) kJ·mol–1).

[73KAT]

This is a report on a careful conductometric study. Sets of data were obtained for 5°C intervals for temperatures between 0 and 45°C. The sets of equivalent conductivities are reported for specific concentrations from 0.0002 to 0.003 M, and have the appearance


360

of being smoothed values from unreported primary data. In the data analysis, the ion-size parameter in the Debye-Hückel equation was optimised for each temperature. It is not clear that this is consistent with the use of the Fuoss and Accascina conductance equation [59FUO/ACC]. The result is a slightly greater temperature dependence of the association constant than if a constant value were used for the ion-size parameter. Nev-ertheless, the temperature dependence is calculated to be less than found in the potenti-ometric study of Nair and Nancollas [59NAI/NAN].

For the purposes of the present review, the data have been re-analysed using a version of the Lee–Wheaton equation modified so that the activity-coefficient equation matches the SIT equation usually used in the TDB project (Appendix B). At comparable temperatures, the calculated nickel sulphate association constants are slightly larger than those recalculated from the results of other conductivity studies [1895FRA], [76SHI/TSU], [79FIS/FOX]. In the present review, we estimate the uncertainty in K1 at 25°C to be ± 20. The same uncertainty is estimated for 20 and 30°C, but for tempera-tures above 30°C and below 20°C, an uncertainty of ± 30 is estimated.

The average enthalpy of reaction for 0 to 45°C was recalculated, and found to be 5.3 kJ·mol–1. This value is accepted for 25°C, and an uncertainty of ± 2.0 kJ·mol–1 is assigned, in part by considering separately the average enthalpies of reaction calculated from the association constants for the four highest temperatures (7.5 kJ·mol–1) and for the four lowest temperatures (3.3 kJ·mol–1).

[73KAW/OTS]

The hydrolytic reactions of nickel(II) ion were studied at 25°C in water and dioxane-water mixtures containing 3 M (Li)ClO4 as an ionic medium. The total nickel(II) ion concentration, B, varied between 0.025 – 0.8 M (I = 3.025 – 3.8 M) in water and the average number of hydrogen ions set free per nickel atom, Z, was measured as a func-tion of pH. In the aqueous solutions a range of 5.7 ≤ pH ≤ 7.3 was studied. Approximate values of the solubility constants of Ni(OH)2(s) ( 10 ,0

*log sK = (13.5 ± 0.5)) and Ni(OH)ClO4(s) ( 10 ,0

*log sK = (6.9 ± 0.5)) were estimated from maximum values of Z. The former constant, converted into the molality scale with density data taken from Lobo and Quaresma [89LOB/QUA] (Part B), and extrapolated to zero ionic strength, has been given in Table V-6 and Figure V-11. At lower Z values, the function Z = Z(pH) can be interpreted by the formation of 4

4 4Ni (OH) + as the predominant hydrolysis prod-uct. None of the minor species NiOH+ and Ni2OH3+ could be established from the data in this study. In aqueous solution 10 4,4

*log b = – (27.32 ± 0.08) has been obtained, which is in good agreement with the corresponding value reported in [65BUR/LIL]. This value has been accepted for use in the present review, but with the uncertainty in-creased to ± 0.16.


361

[73NAV]

The enthalpy of the olivine – spinel transition of Ni2SiO4 was obtained by measuring the heat of solution of both polymorphs in a molten oxide solvent, 2PbO·B2O3, at 713°C. The following values were found: enthalpy of the transition Ni2SiO4(oliv) → Ni2SiO4(spin), trs mH∆ (986 K) = (5.9 ± 2.9) kJ·mol–1, the heat content increments,

mH (986 K) – mH ο (298 K), olivine, (107.7 ± 1.8) kJ·mol–1, spinel, (106.2 ± 0.8) kJ·mol−1 and the entropy of olivine – spinel transition trs mS ο∆ = 12.6 to 14.6 J·K–1·mol–1.

[73NAV2]

Navrotsky reported an experimental approach that was used to obtain values of Gibbs energy differences between olivine, spinel, and phenacite forms of silicates and germa-nates from the thermodynamics of terminal solid solutions in ternary systems. This is applied to the systems NiO − MgO – GeO2 and CoO – MgO – GeO2 at 1200°C in air and to the system NiO − MgO – GeO2 at 800°C and 0.57 kbar water pressure. The Gibbs energy of transformation from the olivine to spinel structure for Ni2GeO4 at 1200°C was estimated to be – 34.31 kJ·mol–1.

[73NOV/COS]

The authors point out that Ni(OH)2 phases of different solubilities were obtained de-pending on the speed of precipitation. The solubility products given at 25°C for I = 0.55

M NaCl and I = 1.0 M NaClO4 seem to have been calculated from Equation (A.49):

10 ,0log sK = log10([Ni2+]·[H+]–2) + 2 10 Wlog K . (A.49)

The authors used Ni2+ and presumably H+ concentrations in the first term on the right hand side of Equation (A.49) but the ionic activity product of H2O in the sec-ond term. Thus, adequately corrected numerical values are given in Table V-6 and Fig-ure V-11.

[73PER/SEC]

The complexation of Ni2+ with partially protonated pyrophosphate ion was studied as a function of ionic strength (KNO3, 0.02 M to 0.20 M), temperature (5 to 35°C) for 0.1 M KNO3 in solutions with pH values from 4.5 to 7.0. Values are reported for the formation constants of 2

2 7NiP O − and 2 7NiHP O− .

Ni2+ + 42 7P O − 2

2 7NiP O − (A.50)

Ni2+ + 32 7HP O − 2 7NiHP O− (A.51)

Values for the first and second protonation constants for 42 7P O − also were de-

termined for the same ionic media and temperatures. The ionic strengths are sufficiently low that it is unlikely that meaningful ∆ε values can be established for any of these reac-tions based on the data in this paper alone. Furthermore, the pH electrode appears to have been calibrated against standard buffer solutions, with no allowance being made


362

for different junction potentials in the solutions of different ionic strength. The values of the reported 25°C constants after correction for the extended Debye-Hückel terms (nD, cf. Appendix B), but without consideration of complex formation between K+ and pyro-phosphate ions, are plotted in Figure A-24.

The corrected complexation constants of pyrophosphate with nickel show a marked ionic strength dependence that would correspond to ∆ε values of (2.5 ± 0.7) kg·mol–1 and (1.7 ± 0.7) kg·mol–1. The former, at least, is inconsistent with any reason-able selection of ε values (though the same is not true for the ∆ε values for the first pro-tonation constant for pyrophosphate, which also involves highly charged ions). Based on these calculations, only the results from the higher ionic strength solutions (I = 0.1 M and 0.2 M) are used in the present review.

For Reaction (A.50) in 0.1 M KNO3:

10log K (A.50) = 5.94 and 10log K ο (A.50) = (7.69 + 0.1∆ε) at 25°C, 10log K (A.50) = 5.81 and 10log K ο (A.50) = (7.53 + 0.1∆ε) at 15°C, in 0.2 M KNO3, 10log K (A.50) = 5.60 and 10log K ο (A.50) = (7.78 + 0.2∆ε) at 25°C.

For Reaction (A.51) in 0.1 M KNO3:

10log K (A.51) = 3.71 and 10log K ο (A.51) = (5.02 + 0.1∆ε) at 25°C, 10log K (A.51) = 3.55 and 10log K ο (A.51) = (4.84 + 0.1∆ε) at 15°C, in 0.2 M KNO3, 10log K (A.51) = 3.39 and 10log K ο (A.51) = (5.02 + 0.2∆ε) at 25°C.

As the ionic strength is low, the effect of the ∆ε term on the final calculated values of 10log K ο is limited. Rather than use estimated ε values for interactions involv-ing the pyrophosphate moieties (some of which are highly charged), it is simply as-sumed that ∆ε is zero. As a consequence, the uncertainty in the 10log K ο values is in-creased to ± 0.25 for the values from measurements in 0.1 M KNO3 solutions, and to ± 0.3 for the values from measurements in 0.2 M KNO3 solutions.

There is a further complication. Within the context of a standard (or extended) Debye-Hückel treatment of activity coefficients, the interaction of the highly-charged pyrophosphate ion with alkali metal ions can be considered most easily by assuming weak complex formation [49MON], [57LAM/WAT]. The same is true if the Specific Ion Interaction approach described in Appendix B is used. De Stefano et al. [94STE/FOT] measured the protonation constants for pyrophosphate, as a function of temperature between 5 and 45°C, in aqueous solutions of (CH3)4NCl, NaCl and KCl (0.00 to 0.75 M). The values for a 0.1 M KCl(aq) medium agreed well (within 0.05 in

10log K ) with the values reported for a 0.1 M KNO3(aq) medium in the studies of Perlmutter-Hayman and Secco [73PER/SEC]. From their apparent protonation con-stants, De Stefano et al. [94STE/FOT] calculated values of the formation constants of

32 7KP O − and 2

2 7KHP O − (and of the formation constants for the corresponding sodium-ion species).


363

Selection of values for the association constants for pyrophosphate ion with al-kali metal ions is beyond the scope of the present review. If the formation constants values for 3

2 7KP O − and 22 7KHP O − from [94STE/FOT] are used with the values of the

pyrophosphate protonation constants from the same source, and if the nickel pyrophos-phate complexation constants from Hammes and Morrell [64HAM/MOR] are used, the calculated concentrations of Ni2+ and 2

2 7NiP O − are within 15% of those calculated by Perlmutter-Hayman and Secco (by neglecting the association of K+). This provides strong support for the value of K(A.50) of Hammes and Morrell [64HAM/MOR]. The species 2 7NiHP O− is calculated to be less important (generally < 20% of the total nickel in solution), and its calculated concentration is approximately 50% higher when association is neglected. This would suggest that the formation constant for 2 7NiHP O− reported by Perlmutter-Hayman and Secco is greater than would be the case if association of pyrophosphate with K+ were considered explicitly. The raw data are unavailable for re-analysis.

From the temperature dependencies of the equilibrium constants, the enthalpies of reaction in 0.1 M KNO3(aq) can be calculated as, r mH∆ (A.50) = (47.93 ± 10.16) kJ·mol–1, and r mH∆ (A.51) = (30.56 ± 7.42) kJ·mol–1.

Figure A-24: Variation of formation constants from Perlmutter-Hayman and Secco [73PER/SEC] with ionic strength. For the first and second protonation constants of py-rophosphate, n = – 8 and – 6; for Reactions (A.50) and (A.51), n = – 16 and – 12.

0.00 0.05 0.10 0.15 0.20

5

6

7

8

9

10

11

12

protonation constant for P2O74-

protonation constant for HP2O73-

formation constant for HNiP2O7-

formation constant for NiP2O72-

log 10

K - n

D

I / mol·dm−3


364

[73POW]

This paper presents a reassessment of the data of Izatt et al. [69IZA/EAT]. The author used the K1 value for the nickel sulphate association constant that was determined by Nair and Nancollas [59NAI/NAN], and reinterpreted the results from the later titration calorimetry experiments. The variation of the enthalpy of ion-pair formation with ionic strength was also reported. The recalculated value agrees well with those from some conductance studies, but is still low with respect to the value from [59NAI/NAN].

[73RAU/GUE]

A mixture of metallic nickel and nickel oxide (particle size around 1 µm) was treated with mixtures of water vapour and hydrogen gas at temperatures ranging from 711 to 860 K. The ratio of equilibrium partial pressures

2 2H O H/p p was determined by a ten-

simetric method using a palladium-membrane apparatus. The authors applied a third law analysis to their data and arrived at f mH ο∆ = – (239.5 ± 1.0) kJ·mol–1 for the stan-dard enthalpy of formation of NiO at 298.15 K.

[73VEN/GEI]

The activity of sulphur in dilute solutions of sulphur in pure nickel melts was investi-gated at three different temperatures, viz. 1773, 1823 and 1848 K, for concentrations up to 0.7 wt% S. A gas mixture of H2(g) and H2S(g) was equilibrated with liquid nickel at a constant gas stream through the furnace. The compositions of both the gas mixture and the equilibrated liquid phase were determined by chemical analysis. When S2(g) is the reference state for sulphur, the activities of sulphur in the nickel melt can be calculated from the well-known gas phase equilibrium:

H2S(g) H2(g) + ½S2(g).

Since this contribution is confined to highly diluted nickel melts only, the ex-perimental results are not re-evaluated in the present assessment.

[73WAA/CAL]

Liebenbergite, a natural nickel mineral is described mineralogically in this work.

[74ARU]

The heat of association, in aqueous solution, of fluoride ion with metal ions of the Mn(II)-Zn(II) series has been measured by a direct calorimetric method at I = 0.5 M and 298.15 K. The measurements were performed in a self medium (to 83 cm3 of a 0.1665 M Ni(NO3)2 solution were added three amounts of 5 cm3 0.5 M NaF solution). The mo-lar enthalpies of the reaction were calculated using the stability constants reported in [72BON/HEF]. For the purpose of the present review, the reported experimental data for the reaction,

Ni2+ + F– NiF+ (A.52)


365

have been recalculated using the recommended values for 10 1log οb ((A.52), 298.15 K) and ∆ε (A.52). At I = 0.5 m 10 1log b ((A.52), 298.15 K) = (0.76 ± 0.08) can be calcu-lated. Taking into account the different ionic media used by Aruga, a greater uncertainty in 10 1log b ( ± 0.25) was assigned. The recalculated formation enthalpy of the NiF+ spe-cies is (4.4 ± 2.0) kJ·mol–1.

[74BIX/LAR]

Chloride selective electrode measurements were used to determine the stability con-stants of nickel(II) chloride complexes at 25°C in 1.00 M NaClO4 and 0.10 M HClO4. The total chloride and metal ion concentrations ranged from 0.0003 – 0.01 M and 0.02 – 0.04 M, respectively. The authors noted that the response of the chloride selective elec-trode was not linear below a chloride concentration of 0.2 mM. The effect of complex formation on the measured potential was rather modest (0.4 – 1.1 mV), therefore we assigned a higher uncertainty to the formation constants than was reported by the au-thors.

[74BLO/RAZ]

The enthalpy of formation of M(H2O)6X(2–y)+ complexes (where M = Ni(II), Co(II), or Cu(II) and X–y = Cl–, Br– or 2

4SO − ) was determined, using the stability constants re-ported in [70MIR/MAK] and experimental results on heat of mixing of 0.05 M M(ClO4)2 solution (I = 3.0 M maintained with LiClO4) with the solution of the corre-sponding lithium salt (I = 3 M). The authors did not reporte sufficient experimental de-tails to allow the reliability of their data to be assessed.

[74DIC/HOF]

The kinetics of formation and dissociation of Ni(SCN)2(aq) have been studied in water and in several organic solvents, using the pressure-jump and shock wave relaxation technique at 20°C. The concentration of Ni(SCN)2(aq) ranged between 0.001 and 0.1 M. In water, only the formation of the monothiocyanato complex was observed. No background electrolyte was used, and the activity coefficients were calculated by an extended Debye-Hückel expression. Although, this activity model is not compatible with the SIT, the ionic strengths were low. Therefore, the reported result was corrected to 25°C and the resulting value was accepted with an increased uncertainty ( 10 1log οb = (1.79 ± 0.10)).

[74GRA/WIL]

The ternary and quaternary complexes of nickel(II) and palladium(II) with asparaginate (asn), chloride and hydroxide ions were studied using glass and silver-silver chloride ion selective electrodes (I = 3 M NaClO4, T = 298.15 K). The formation of

(2 )Ni(asn) (OH) x yx y

− − complexes was not reported, only the ternary (quaternary) com-plexes with chloride ions, (1 )Ni(asn) Cl(OH) x y

x y− − , in addition to the binary NiCl+ com-

plex, were found to exist. Taking into account the low affinity of chloride for nickel(II)


366

this is unlikely. Therefore, the reported +10log (NiCl )K = (0.687 ± 0.040), which is

much higher than other reported values, was not taken into account in this review.

[74GRI/FER]

The structure of a millerite crystal from Marbridge Mine, Malartic, Quebec, with the empirical formula Ni0.981Fe0.016Co0.004S has been refined. The hexagonal axes were found to be a0 = 9.607(1) Å and c0 = 3.143(1) Å. Within the lattice each Ni is coordi-nated by five S atoms and two Ni atoms. The observed Ni-S bond lengths are compara-ble to the expected value for a covalent bond. Molecular orbital theory was invoked to show that the millerite structure with five-fold coordination around Ni is more stable than the nickeline structure (α-NiS) with six-fold coordination about each Ni. Thereby it is rationalised that the low temperature phase β-NiS occurs in nature and the high tem-perature phase α-NiS does not.

[74HEF]

This is a review paper on fluoride complexes of some 50 metal ions. The author found that the stability constants of these complexes correlate very well with simple electro-static parameters such as ion size and charge. Only a few metal ions, all found in the lower right hand side of the Periodic Table, seem to be exceptions from this general rule. The enthalpy of formation is positive even for the strongest complexes. Only Be2+ and Ag+ appear to have negative r mH∆ values. Consequently, the formation of the great majority of fluoro complexes is entropy controlled.

[74KAB]

The protonation of tetracyanonickelate(II) ion was investigated by a pH-metric method at temperatures from 25 – 50°C and at ionic strengths of 0.01, 0.02, 0.05, 0.1, 0.2 and 0.5 M, adjusted using NaCl. The author reported two successive protonations of the

24Ni(CN) − ion ( 10 1log K and 10 2log K , respectively). Titrations of aqueous Na2Ni(CN)4

solutions by 0.1 M HCl were performed while nitrogen gas was bubbled through the solution; therefore HCN volatilisation could have been occurring. Several authors have reported slow equilibration in acidic solutions of Na2Ni(CN)4, sometimes lasting several days [63CHR/IZA], [68KOL/MAR], [74PER]. During the titration performed in [74KAB], equilibrium was probably not attained. The dependence of the determined protonation constants on the temperature and ionic strength shows notable deviation from what would be expected (see e.g., Figure A-25 and Figure A-26). Therefore, the reported data were not considered further in this review.


367

Figure A-25: Temperature dependence of 10 2log K for the reaction 24Ni(CN) − + H+

3Ni(HCN)(CN)− at different ionic strengths using NaCl.

Figure A-26: Extrapolation to I = 0 of experimental equilibrium constants for the reac-tion 3Ni(HCN)(CN)− + H+ 2 2Ni(HCN) (CN) (aq) ( 10 2log K ) at different tempera-tures.

4

4.2

4.4

4.6

4.8

5

0.00305 0.00315 0.00325 0.00335

K / T

log 1

0 K2

I = 0.01 M

I = 0.02 M

I = 0.05 M

I = 0.1 M

I = 0.2 M

I = 0.5 M

4.5

4.6

4.7

4.8

4.9

5

0 0.1 0.2 0.3 0.4 0.5 0.6

I / mol.kg-1

log 1

0 K2 +

2D T = 25 °C

T = 30 °C

T = 35 °C

T = 40 °C

T = 45 °C

T = 50 °C


368

[74KUL3]

The stability constants and the enthalpy changes for the formation of nickel(II) (and copper(II)) thiocyanate complexes were determined by calorimetric titrations at 25°C in 1 M (NaClO4) aqueous solution. The concentrations of nickel(II) and thiocyanate ranged from ~ 0.01 – 0.1 M and 0 – 0.8 M, respectively. The formation of NiSCN+, Ni(SCN)2(aq) and 3Ni(SCN)− was assumed according to [53FRO2]. Both graphical and computer evaluations of the experimental data were performed. The formation constants obtained by the computer programme "Kalori" are as follows: 10 1log b = (1.12 ± 0.04) (3σ), 10 2log b = (1.57 ± 0.14), 10 3log b = (1.28 ± 1.00) (two years later the author re-ported [76KAR/KUL] somewhat different constants: 10 1log b = (1.12 ± 0.02) (3σ),

10 2log b = (1.57 ± 0.04), 10 3log b = (1.25 ± 0.20)). The reported "best" values are slightly different from both above mentioned sets, since Fronaeus' data [53FRO2] were also taken into account: 10 1log b = (1.14 ± 0.02) (3σ), 10 2log b = (1.58 ± 0.05),

10 3log b = (1.6 ± 0.2). The latter constants have been used to calculate the correspond-ing stepwise enthalpy values: r mH∆ = – (12.02 ± 0.15), – (8.9 ± 1.0) and − (8.2 ± 4.0) kJ·mol–1 for the three successively formed complexes, respectively. The enthalpy of reaction values corresponding to the three cumulative formation constants are r mH∆ = − (12.0 ± 0.5), – (20.9 ± 1.2) and – (29.1 ± 5.2) kJ·mol–1, respectively, where the uncer-tainties have been estimated.

[74MA]

In this work the olivine – spinel transformation in Ni2SiO4 at high pressures was inves-tigated over a p – t range of 20 – 40 kbar and 650 – 1200°C, using a piston-cylinder apparatus. The transformation curve is essentially a straight line and can be expressed by an equation: p(bar) = 23300 + 11.8·t (°C). The melting behaviour of Ni2SiO4 olivine was investigated in a p – t range of 5 – 13 kbar and 1600 – 1700°C. The results show that Ni2SiO4, olivine melts incongruently at high pressures to NiO plus melt. The re-ported dissociation of Ni2SiO4, olivine, to NiO plus cristobalite prior to melting [63PHI/HUT] was not observed. The melting curve has a p – t slope of approximately 105 bar·K–1. It can be extrapolated to atmospheric pressure to give a melting tempera-ture of approximately 1575°C. This value is 75°C lower than that given by [63PHI/HUT].

[74MAK/MAS]

Values for the association of nickel with sulphate for temperatures from 25 to 95°C and ionic strengths to 2.6 M were measured using a variety of pH titration, conductivity and transport number measurement techniques. Although the reported association constants are in rough agreement with those from other studies, there are insufficient experimental details (or results) to use this work in deriving a recommended value for K1.


369

[74NET/BAT]

Complex formation of Ni(II), Pr(III) and Ho(III) with thiocyanate was studied in aque-ous solutions (I = 3 M (LiClO4)) by IR spectroscopy. The concentration of thiocyanate was 0.36 M, the concentration of metal ions ranged from 0.36 to 0.86 M. The formation constants of the mono-complexes (MSCN+) were calculated by three different methods, which resulted in considerably different values. The temperature of the measurements is not given in the paper. Therefore, the reported data were not considered further in this review.

[74PER]

The formation of complexes between nickel(II) and cyanide ion was studied by spectro-photometric and potentiometric measurements with a glass electrode at 25°C, at an ionic strength of 3 M using aqueous sodium perchlorate as the ionic medium. The potentiometric data are best described by the formation of two complexes NiCN+ and

24Ni(CN) − ( 10 1log b = (7.03 ± 0.2) and 10 4log b = (31.06 ± 0.03)). On the other hand,

the spectrophotometric measurements gave no positive evidence for the formation of NiCN+; therefore the 10 1log b value given above should be regarded as a maximum value. If only formation of 2

4Ni(CN) − was postulated, 10 4log b = (31.12 ± 0.08) was obtained, with a somewhat higher uncertainty in the fitted parameter. This value is ac-cepted in the present review, but the uncertainty has been estimated here as ± 0.15. The potentiometric investigation was performed at 4 < pH < 5.5, but protonated tetracyano complexes, as reported in [68KOL/MAR], were not detected.

[74RAJ/PRE]

The structure of a single crystal of millerite from Quebec, Canada, with hexagonal axes a0 = 9.6190(5) Å, c0 = 3.1499(3) Å and composition Ni1.03Fe0.016Co0.004S has been re-fined. The Ni and S atoms are in five-fold coordination (tetragonal pyramidal) with each other. The average Ni-S distance, 2.310 Å, is consistent with a divalent Ni in five-fold coordination. Rajmani and Prewitt suggested that the metal-metal bonding and the for-mation of the trinuclear cluster stabilise the millerite structure.

[74RUT/HAU]

This paper describes vapour pressure measurements over Ni in the range 1277 – 1658 K as obtained using the Langmuir technique. The results yield the enthalpy of sublimation (447.7 ± 24.3) kJ·mol–1 and (431.8 ± 16.7) kJ·mol–1 according to 2nd and 3rd law analy-sis, respectively. In addition Rutner and Haury [74RUT/HAU] determined the “best values” of the parameters in the vapour – pressure equations of solid and liquid nickel, and the heat of vaporisation and sublimation of nickel by a statistical treatment of data available in the literature and their own results. The results of Rutner and Haury [74RUT/HAU] were not used by [98CHA] for the determination of the enthalpy of for-mation for Ni(g) because of an inconsistency with the melting data.


370

[74SLO/JON]

In their Figure 2 Slough and Jones plotted z·r–1 versus r mH ο∆ of reactions,

MO(cr) + 2B2O3(cr) MB4O7(s) (A.53) and

M2O(cr) + 2B2O3(cr) M2B4O7(s) (A.54)

where r are presumably the Pauling radii [40PAU] and r mH ο∆ were determined experi-mentally by reactions of the crystalline oxides, see Figure A-27. The value for NiB4O7, given by [74SLO/JON], can only be estimated when the Goldschmidt radius of Ni2+ (r = 0.078 nm) [26GOL] and not the Pauling radius (r = 0.070 nm) is employed. The effec-tive ionic radii of Shannon (for coordination number 6 [76SHA]) lead to an inferior correlation and a completely different r mH ο∆ for the respective nickel reaction, even on the plot based on the Pauling radii. This example demonstrates that methods of thermo-dynamic data estimation also must be reviewed critically.

Figure A-27: r mH ο∆ (298 K) of reaction of B2O3(cr) with MO or M2O.

[74WOR/COW]

Heat capacities from 10 K to 300 K were determined by adiabatic calorimetry, and by differential scanning calorimetry (DSC) from 300 K to 550 K. In the paper, only rounded values are supplied for the adiabatic calorimetry results for NiI2(cr) except for values near the (structural phase [81KUI/SAN]) transition at approximately 59 K. The authors also supply an equation (A.55)

0 -100 -200 -300 -400

5

10

15

20

25

30 2 B2O3 (cr) + MO (cr) MB4O7 (s)2 B2O3 (cr) + M2O (cr) M2B4O7 (s)

K+

Na+

Li+

Ca2+

z · r

−1 /

nm−1

∆rH°m/ kJ·mol−1

Ni2+ [26GOL] Ni2+ [76SHA]

Pauling ionic radii [40PAU] Effective ionic radii [76SHA]

∆rH°m / kJ·mol–1


371

550K,m 200K[ ]pCο (NiI2, cr) / J·K–1·mol–1 = 65.5 + 0.0515 T/K – 0.0000397 (T/K)2 (A.55)

that provides a fit to the combined heat capacity results from adiabatic calorimetry and differential scanning calorimetry between 200 K and 550 K, but do not separately pro-vide measured values above 300 K. This equation is not in a form normally used for the temperature dependence of heat capacities of solids near room temperature [93KUB/ALC].

The adiabatic calorimetry heat capacity data were available as supplementary material from the British Library Lending Division (SUP 21075), and were used in the preparation of the current review. The authors integrated the low-temperature heat ca-pacity measurements and reported 138.7 J·K–1·mol–1 for mS ο (NiI2, cr). No contribution was added for the magnetic phase transition at 75 K to 76 K reported by Billerey et al. [77BIL/TER] and Kuindersma et al. [81KUI/SAN], and this may have led to an under-estimation of between 0.3 J·K–1·mol–1 and 0.5 J·K–1·mol–1 in the value of mS ο above 80 K. Indeed, the two most relevant values from the authors “Run I” (for 77.69 K and 79.36 K) show no evidence whatsoever of an anomaly. However, the authors’ “Run I” values also do not provide a well-defined peak for the anomaly near 59 K. The “Run II” experiments covered only the temperature range of the 59 K anomaly, and appear to have been done with much more care and with mush smaller temperature increments. The values of entropies of reaction as derived from adiabatic calorimetry and from measured dissociation pressures and heats of solution of NiI2, Ni(NH3)2I2 and Ni(NH3)6I2 were found to be in good agreement. This indicates that it is likely that as-suming a smooth extrapolation of ,mpC between 10 K and 0 K does not introduce any substantial error in the calculated mS ο value. The adiabatic calorimetry results between 200 K and 300 K showed a random scatter of approximately ± 0.4 J·K–1·mol–1, unusu-ally large for this type of measurement.

The experimental scatter for the differential scanning calorimetry values was reported to be within 4%, but it is not reported how well the authors’ equation corre-sponded to the measured values, or how well measurements obtained from the two techniques agree near 300 K. Unfortunately, the DSC data were not part of the supple-mentary material.

To obtain an equation for ,mpC (T) in a more standard form, values of the heat capacity were calculated at 25 K intervals between 300 K and 550 K using the authors’ equation (A.55). An equation of the form (A + B(T/K) + C(T/K)–2) was fitted to these values and to the experimental low-temperature adiabatic calorimetry values from 200 K to 300 K. Values from the adiabatic calorimetry and from the equation based on the DSC measurements were weighted in a ratio of 0.4:2.0. Because the number of DSC measurements used in the original fitting by Worswick et al. is unknown, these weight-ings are arbitrary by definition. The obviously erroneous datum at 229.92 K was not used. The resulting equation (A.56),

550K,m 200K[ ]pCο (NiI2, cr) / J·K–1·mol–1 = 73.5775 + 0.016425T – 1.057753×105 (T/K)–2 (A.56)


372

fit the low temperature data as well as the authors’ equation (and well within the ex-perimental scatter). As would be expected, agreement between values from the two equations was poorer for 300 K to 550 K, but the differences were always less than 0.5 J·K–1·mol–1 —again much less than the experimental uncertainties.

The values from equations (A.55) and (A.56) for 298.15 K agree well with the reported smoothed value from the adiabatic calorimetry results (78.4 J·K–1·mol–1) and with each other (within 0.05 J·K–1·mol–1). Because the actual DSC results are unavail-able, and there is considerable scatter in the adiabatic calorimetry results above 200 K, an uncertainty of ± 1.0 J·K–1·mol–1 is assigned to the heat capacity value for nickel io-dide at 298.15 K.

[74YAG/MAR]

Yagi et al. have investigated structures of the spinel polymorphs of Fe2SiO4 and Ni2SiO4 in detail using single crystals synthesised at high pressures and temperatures. Ni2SiO4 spinel with a = 8.044(1) Å had a normal spinel structure.

[75ARN/MAL]

The sulphur-rich part of the system Ni – S above 1253 K was re-investigated by metal-lographic methods. Vaesite (NiS2) was prepared from NiS and sulphur by solid state reaction. Mixtures of NiS2 and S were heated in evacuated silica glass tubes. After quenching, polished samples were examined microscopically. In contrast to Kullerud and Yund [62KUL/YUN], who reported a congruent melting point of vaesite at 1280 K and a monotectic reaction at 1264 K, the heating experiments indicated that NiS2 melts syntectically, forming two immiscible liquids at (1295 ± 3) K, i.e., NiS2 L1 + L2. The liquid L2 was estimated to contain 0.5 wt% Ni, while the composition of the liquid L1 amounted to (52.0 ± 1.8) wt% Ni.

[75ARU]

Enthalpies of association of nitrate and chlorate with Mn(II), Co(II), Ni(II), Cu(II) and Zn(II) have been determined by direct calorimetry in an aqueous medium at 298 K and an ionic strength I = 1 M. In a typical experiment, 2.5 and 5.0 cm3 of 1 M NaNO3 solu-tion were added to 93 cm3 of 0.34 M Ni(ClO4)2 solution, and this was repeated twice under identical conditions. The association between Na+ and nitrate ion was considered during the calculations (Kass = 0.25 M–1). The molar enthalpies of the reaction:

Ni2+ + 3NO− 3NiNO+ (A.57)

were calculated using the stability constants reported in [73HUT/HIG]. For the purpose of the present review, the reported experimental data have been recalculated using the tentative values for 10 1log οb ((A.57), 298.15 K) and ∆ε((A.57), (NiClO4)2) = − (0.11 ± 0.15) kg·mol–1. At I = 1.05 m 10 1log b ((A.57), 298.15 K) = – (0.21 ± 1.00) can be calculated. The recalculated formation enthalpy of the 3NiNO+ species is (6 ± 4)


373

kJ·mol–1. The relatively high uncertainty is assigned due to the large error in 10 1log οb ((A.57), 298.15 K).

[75BAR/MAS]

Barvinok et al. measured the enthalpy of solution of NiCl2(cr) in 2.0 M HCl at 298.15 K. A value for the molar enthalpy of solution, sol mH∆ = – (74.43 ± 0.17) kJ·mol−1, was reported. As other experimental details are lacking, this single value was disregarded in this review.

[75CLA/KEP]

The kinetics of hydrolysis of nickel(II) were studied by mixing nickel(II) sulphate or nickel(II) perchlorate solutions with NaOH solutions in a stopped-flow apparatus at 25 and 41.7°C. The ionic strength was maintained at 0.8 M by additions of sodium sulphate or perchlorate. The hydrolysis occurred in two stages, as already mentioned in [71HOH/GEI]: (i) the first stage (probably the formation of NiOH+) was complete within the mixing time (2 ms), (ii) the second stage resulted in the formation of

44 4Ni (OH) + within a few seconds (after a longer time interval, slight precipitation of

nickel hydroxide was also noted). In perchlorate medium some runs were made to fol-low the changes of pH, after addition of base, using Neutral Red as an indicator (the value used for the pK of the indicator was not reported) at [Ni2+]tot = 0.203 M. The pH was measured just after mixing and at equilibrium. From the pH values measured at “t = 0 s”, 10 1,1

*log b = – (8.93 ± 0.10) was obtained for the reaction :

Ni2+ + H2O NiOH+ + H+. (A.58)

Since this value is considerably lower than other literature values, the authors concluded that some oligomerisation occurred even during mixing. From the pH values at equilibrium, the authors calculated the formation constants of 4

4 4Ni (OH) + assuming it as the only hydrolysis species that was formed. Taking into account the relatively low pH values at equilibrium (pH = 6.50 – 6.75), the formation of NiOH+ should also be considered. Since only 6 experimental data points were reported, the value of 10 1,1

*log b was fixed at – 9.87 (calculated from the selected 10 1,1

*log b ((A.58), 298.15 K) for Im = 0.82), and only the 10 4,4

*log b value was refined during our re-assessment. With this constraint, 10 4,4

*log b = − (26.98 ± 0.20) was obtained, a value close to the reported value. Taking into account the limited number of experimental data, 10 4,4

*log b = − (27.0 ± 0.4) is accepted for use in this review.

[75KAK/GIE]

The heat of solution of nickel chloride hexahydrate into 4.36 M HCl(aq) was measured ((21.32 ± 0.13) kJ·mol–1). Few experimental details are provided, and a double correc-tion (to 0.0 M HCl(aq) and to I = 0) with large uncertainties would be required to use the reported value to calculate the enthalpy of solution of the salt at infinite dilution in water. This heat of solution value is not used further in the present review.


374

[75KOS]

The molar heat capacities of NiCl2 from 2 to 30 K have been investigated. Smoothed values in this region were listed. The results were compared with the predictions of the spin-wave theory.

[75LEV/GOL]

The authors investigated the equilibrium potentials of the following galvanic cell with a solid electrolyte: Pt | (Ni + SiO2) Ni2SiO4 | O2– | Ni, NiO | Pt. Based on the Gibbs energy of the cell reaction determined in the temperature range of 1228 to 1355 K, the standard Gibbs energy of formation and the standard molar enthalpy of formation of nickel or-thosilicate at 298.15 K were calculated. We used the raw data of this work to represent the Gibbs energy of the reaction:


as a function of temperature.

[75LIB/TIA]

Emf and spectrophotometric methods have been used to determine the stability con-stants of MCl+ complexes in M(ClO4)2 medium (M = Mn2+, Co2+, Ni2+ and Zn2+) at dif-ferent concentrations.

The following cell was used for the potentiometic determination of formation constants: Ag | AgCl,HCl(m1),Mg(ClO4)2(m) || M(ClO4)2(m),HCl(m1), AgCl | Ag, where m1 was fixed at approximately 0.01 m, while the molalities of the metal perchlorate (m) were 1.0, 1.5, 2, 2.5 and 3 m. The emf of the cells varied between 1.0 and 26.4 mV. The calculation was based on the following plausible assumptions: (i) the activity coefficient of the chloride is identical in the equimolar solution of Mg(ClO4)2 and M(ClO4)2, (ii) the liquid-junction potential of the above cell is negligible, (iii) Mg(II) does not form complexes or ion pairs with chloride. In the spectrophotometric study, the strong ab-sorption of the CuCl+ complex at 272 nm was used to assess the formation constants of the MCl+ complexes. At this wavelength the molar absorptivities of both the M(H2O)2+ and the MCl+ complexes were assumed to be negligible. The free chloride concentration and, thus, the formation constants of the MCl+ complexes were calculated using the formation constant and the ελ=272nm value of the CuCl+ complex determined earlier [73LIB]. The two methods resulted in nearly coincident values.

The main advantage of the above experimental procedure is that in Ni(ClO4)2 media the medium effect can be eliminated. In the case of nickel, the results are inter-preted in terms of the existence of [Ni(H2O)5]Cl+ and [NiCl(H2O)5]+ complexes. The overall stability constant of the NiCl+ species (β1 = 0.37 at I = 0) is due mainly to the stability of the outer-sphere ion pair. The β1 value for the inner-sphere complex was reported to be around 0.01 M–1.


375

[75MEY/WAR]

The equilibrium partial pressure of sulphur over Ni – S melts was determined by bub-bling H2/H2S gas mixtures through the melts at temperatures between 1373 and 1873 K. The compositions of the melt varied from nickel saturation to 27 wt% sulphur. The composition of the gas phase was monitored with the help of an optical interferometer, and the composition of the melt was determined by chemical analysis. The sulphur par-tial pressure for a given gas composition was calculated from the well-known equilib-rium constants for the reaction:

H2S(g) H2(g) + ½S2(g).

In addition, the liquidus line of the binary phase diagram Ni – S was obtained in the region from 0 to 20 wt% sulphur.

[75MOS/FIT]

The data concerning NiO are essentially the same as in [75MOS/FIT2].

[75MOS/FIT2]

The least-squares line representing the temperature dependence of f mGο∆ (NiO) was determined to be f mG∆ (T) = (– 244.019 + 0.092 (T/K)) kJ·mol–1. However, no detailed information concerning uncertainties of the measured values was available.

[75RAU2]

The partial pressure of sulphur in equilibrium with the non-stoichiometric high-temperature modification of Ni1–xS was measured as a function of composition in the temperature range from 781 to 1250 K. The sulphur pressure was determined directly by a tensimetric method. In addition, the solid sample was reduced with hydrogen gas. The amount of H2S(g) formed during the reduction as well as the ratio between H2S(g) and H2(g) in the gas phase were analysed by tensimetric measurements. From these data both the composition of the solid phase and the temperature dependence of the sulphur activity in Ni1–xS, with S2(g) being the reference state for sulphur, could be calculated from the well-known equilibrium:

H2S(g) H2(g) + ½S2(g).

Moreover, phase boundaries of the homogeneity range of Ni1–xS are likewise obtained from this equilibrium study.

[75TAM]

This is a reanalysis of some of the data of Katayama [73KAT] using different activity coefficient assumptions.


376

[75WEE/KOE]

Weenk et al. determined the enthalpies of solution of a series of alkalitrichloronicke-lates ANiCl3 (A = Cs, Rb, K, Tl, NH4) and their constituting chlorides (NiCl2 and ACl) in a medium of 0.1 M HCl at 298.15 K with n(H2O) / n(NiCl2) = 1200. A solution calo-rimeter of the isoperibol type was used. The molar enthalpy of solution of NiCl2 was found to be sol mH ο∆ = – (76.82 ± 0.42) kJ·mol–1 for the reaction:

NiCl2 Ni2+ + 2 Cl–.

This value refers to an ionic strength I ≈ 0.24 mol·kg–1, which is too high for the usual extrapolation to I = 0 to be applicable. Consequently this single value was disregarded in this review.

[76BAE/MES]

This is a well-known standard work concerning metal ion hydrolysis in water. Selected experimental values for 298.15 K were converted to zero ionic strength using the equa-tions from Pitzer and Brewer [61LEW/RAN]. For the water soluble nickel(II) hydroxo complexes, 2Ni OH x y

x y− the following 10 ,

*log y xοb values are reported: 10 1,1

*log οb = − (9.86 ± 0.03), 10 2,1

*log οb = – (19 ± 1) 10 3,1*log οb = – (30.0 ± 0.5), 10 4,1

*log οb < – 44, 10 1,2

*log οb = – (10.7 ± 0.5) and 10 4,4*log οb = – (27.74 ± 0.02). No thermodynamic data

have been taken directly from this reference.

[76BER2]

A high quality data set was obtained by accurate measurements of the Gibbs energy of formation of NiO by using a solid state emf method with a gaseous hydrogen/water buffer atmosphere. The result of the weighed linear regression for the temperature range from 825 to 1675 K was given by the equation:

f mGο∆ (NiO) = – (235.071 ± 0.141) + (0.08620 ± 0.00010)·(T/K) kJ·mol–1.

[76BIA/CAR]

A method using variation of solvent composition to study the formation equilibria of complexes and its application to the nickel(II)-chloride system is reported. In pure ace-tone the formation of the complex 2

4NiCl − was detected, and this species readily de-composed with increasing water content. The extrapolation of experimental results, de-termined between 0 and 0.05 mole fraction of water, to pure water, resulted in a value of

210 4log (NiCl )−b ~ – 14 as a rough approximation.

[76GEE/SHE]

The emf measured for solid electrolyte cells of the type:

Pt | M’, M’Cl2 | electrolyte | M, MCl2 | Pt

provided Gibbs energies of formation for CrCl2, MnCl2, CoCl2 and NiCl2 at


377

470 ≤ T / K ≤ 910. Depending on the temperature range either PbCl2 or BaCl2 were used as the solid electrolyte. The Fe, FeCl2 electrode served as reference electrode. From results two-parameter Gibbs energy of formation equations have been derived, which obviously contain the uncertainties of the thermodynamic parameters ascribed to the substances of the reference half-cell. The Gibbs energy equation related to NiCl2 has been plotted in Figure V-18.

[76KAR/KUL]

This is a description of the "Kalori" computer programme (see the discussion of [74KUL3]). As one of the applications, the nickel(II) – thiocyanate system was men-tioned. Using the identical dataset, the authors reported stability constants and enthalpy values that differ slightly from those reported for the same complexes in the original paper [74KUL3].

[76KUL/BLO]

See comments under [81KUL/BLO].

[76MUR/KUR]

A solvent extraction method was used to determine, the stability of NiCl+ and NiSCN+ complexes, among others. The aqueous phase was 1 M NaClO4 solution, while the or-ganic phase was CCl4 containing 0.05 M thenoyltrifluoroacetone and 0.03 M trioctyl-phosphine. The concentration of nickel(II) was ~ 1.0 × 10–4 M, while that of the com-plex forming anion varied up to 0.4 M for thiocyanate and ca 0.85 M in case of chlo-ride. As the amount of metal extracted was very low in the case of chloride, even at high chloride concentrations, only the formation constant reported for the thiocyanate com-plex is considered in this review.

[76NAV/KAS]

The enthalpy of formation from the oxides of Mg2SnO4 and Co2SnO4 were found by oxide solution calorimetry. Using thermochemical data for the formation of olivine, for olivine – spinel transitions and for the transformation of quartz to stishovite, pressures for the disproportionation of silicate spinels were calculated to be in the range 150 – 200 kbar.

[76PAN/PAT]

Pant and Pathak measured the solubilities of Ni3(PO4)2·7H2O and NiHPO4 in water from 30 to 60°C, and reported the metal concentration in solution (0.012 M to 0.027 M), the metal to phosphate ratio in solution (near 2.0 in all cases) and the “pH” (between 7.1 and 7.7, without allowance for the change in Kw with temperature). In the absence of a good set of values for the complexation constants for Ni2+ with phosphate ligands as a function of temperature, and good analyses of the final solids, no useful thermodynamic quantities can be extracted from the information in this paper.


378

[76PER]

The kinetics of formation of 24Ni(CN) − have been studied spectrophotometrically at

367.5 nm in the pH-range 4.7 – 7.0, at I = 3 M NaClO4(aq) and t = 25°C. The author found that, at the equilibrium, only Ni2+ and 2

4Ni(CN) − exist in substantial concentra-tions, although the complexes NiCN+, Ni(CN)2(aq) and 3Ni(CN)− are kinetically signifi-cant. In contrast to the acidic dissociation of 2

4Ni(CN) − [76PER/EKS], its association from nickel(II) can be described, within the experimental errors, without the protonated complexes. This is probably due to the higher pH range used in this study. From the kinetic data presented in this paper, the author determined the formation constant of

24Ni(CN) − ( 10 4log b = (31.08 ± 0.05), at I = 3 M NaClO4 and t = 25°C, using HCNpK =

9.484 [71PER]). This value of 10 4log b is retained in the present review, but with the uncertainty increased to ± 0.15.

[76PER/EKS]

The acidic dissociation of 24Ni(CN) − was studied spectrophotometrically in the pH

range 0 < pH < 3, at I = 3 M NaClO4 and t = 25°C, in order to elucidate the presence of protonated 2

4H Ni(CN)xx

− species. Persson [74PER] reported 24Ni(CN) − as the only

complex existing in the nickel(II) – cyanide system. On the other hand, Kolski and Margerum, in [68KOL/MAR], reported extensive formation of protonated complexes. Persson and Ekström [76PER/EKS] acknowledged the formation of protonated species and their role in the acidic decomposition, but their data indicated that the protonated species have steady state concentrations during the decomposition. Therefore, according to the authors, the pK values of the protonated complexes should be much lower than suggested in [68KOL/MAR].

[76RAU]

The activity of sulphur in the high temperature modification of Ni3S2 was determined as a function of composition at temperatures between 815 and 1044 K. The solid sample was reduced with hydrogen gas in a stepwise manner. The amount of H2S(g) formed during the reduction as well as the ratio between H2S(g) and H2(g) in the gas phase were obtained from tensimetric measurements. From these data both the composition of the solid phase and the temperature dependence of the sulphur activity in Ni3S2, with S2(g) being the reference state for sulphur, could be calculated from the well-known equilib-rium H2S(g) H2(g) + ½S2(g). Neither the initial compound nor the solid products formed owing to reduction with hydrogen gas were checked by X-ray diffraction analy-sis.

[76SHI/TSU]

This was a conductance study done at pressures from 0.1 to 160 MPa and at 15, 25 and 40°C. The sets of equivalent conductivities are reported for a very limited range of con-centrations from 0.0001 to 0.0003 M, and have the appearance of being smoothed val-


379

ues from unreported primary data. The data analysis used the simple Onsager equation to extrapolate the NiSO4(aq) dissociation constant to I = 0, and values for r mH∆ and ∆V are reported. The reported association constants are greater than those found in many other studies. Re-analysis of the “1 atm.” results (assumed to be identical at 0.1 MPa) using a version of the Lee-Wheaton equation that is consistent with the SIT equation used in the present review (Appendix B) leads to results that are consistent with the other studies. In the present review, we estimate the uncertainty in K1 at 25°C (0.1 MPa) to be ± 30. The high-pressure results are consistent with values found in a later study of Fisher and Fox [79FIS/FOX].

The average enthalpy of reaction, calculated from the equally weighted values for the temperature range 15 to 40°C, is 6 kJ·mol–1, in excellent agreement with the value obtained from the data of Katayama [73KAT]. An uncertainty of ± 2.5 kJ·mol–1 is estimated in the present review.

[76SMI/MAR]

This is a well-known compilation of stability constants including a number of metal ions and a variety of inorganic ligands. For the water soluble nickel(II) hydroxo com-plexes the following 10 ,log y xb for the reaction:

xNi2+ + yOH– 2Ni (OH) x yx y

−

are reported: 10 1,1log οb = 4.1, 10 2,1log οb = 8, 10 3,1log οb = 11, 10 1,2log b = 4.2 (I = 3 M) and 10 4,4

*log οb = 28.3 (I = 0), (29.37 ± 0.03) (I = 3 M). No thermodynamic data have been taken directly from this reference.

[76TAY/DIE]

This paper reports a pH study of complexation between 24HPO − and Ni2+ at 25°C in

0.1 M NaClO4:


Nickel(II) concentrations were between 7 × 10–3 and 3.3 × 10–2 M, and the ini-tial ligand concentration was approximately 2 × 10–2 M. The association constant be-tween Ni2+ and 2

4HPO − was reported as (130 ± 6) dm3·mol–1, and the association con-stant between Ni2+ and 2 4H PO− to form 2 4NiH PO+ as (3.5 ± 1) dm3·mol–1. There is in-sufficient information to recalculate the results. For the protonation constants for

24HPO − , the reported 10log K values are within ± 0.2 of values from other studies, and

in the present review a similar uncertainty is assigned to the reported NiHPO4(aq) for-mation constant ( 10log K = (2.11 ± 0.20)). Correction to zero ionic strength (Appen-dix B) gives 10log K ο (A.59) = (3.00 ± 0.20), where the uncertainty has been estimated in the present review. The small ∆ε correction has been based on the value for ε(Na+, 2

4HPO − ) from Table B-5 (rather than using the equation in Table B-6), and ε(Ni2+, 4ClO− ) = 0.37 kg·mol–1. The evidence for 2 4H NiPO+ is not considered to be suf-ficient to assign a formation constant for this species.


380

[77ASH/HAN]

UV spectroscopic measurements (I = 5 mol·dm–3, NaClO4/Na2SO4) were compared for Na2SO4 concentrations of 0.1 mol·dm–3 and 1.0 mol·dm–3. Measurements for 47, 60 and 70°C were reported. The authors extrapolated their measurements to 25°C, assuming the effects of temperature were small ( r mH∆ was assumed to be 0 kJ·mol–1), and reported

10 1log K (298.15 K, I = 5 M NaClO4) = 0.77. The assumption concerning r mH∆ is rea-sonable in a general sense, but probably introduces an uncertainty of the order of ± 0.2 in the value of 10 1log K (298.15 K). The ionic strength (6.6 m) is too high to properly apply the SIT. If the SIT is used with ε(Ni2+, 4ClO− ) equal to (0.37 ± 0.03), 10 1log K ο = 1.3 is calculated.

[77BAR/KNA]

This volume contains the Planck function, the enthalpy and the Gibbs energy tabulated at 100 K intervals for pure substances. It continued the work started by [73BAR/KNA].

[77BIL/TER]

In this paper, the heat capacity and magnetisation of NiI2 were reported from measure-ments done between 20 K and 150 K. The heat capacity measurements showed peaks at 59.5 K and 76 K (later ascribed by Kuindersma et al. [81KUI/SAN] to a structural phase transition and a magnetic phase transition, respectively), and a shoulder and a peak were found at the corresponding temperatures in the magnetic measurements. The measurements are reported only in graphical form. In the present review, the values of

,mpC (T) from the authors’ Figure 2(a) were digitized and used to calculate values of ,mpC /T. Equations were fit to the values over short ranges of temperature and to a set of

values including points for temperatures above and below the anomalies. If the contri-butions from the two transitions are neglected, the contribution to the entropy between 50 K and 85 K is calculated to be 27.2 J·K–1·mol–1. The additional contributions to the entropy from the transitions near 59.5 K and 76 K are estimated to be 0.40 J·K−1·mol–1 and 0.37 J·K–1·mol–1, respectively.

[77CHE/BRO]

The 14N nuclear quadripole resonance spectra were measured for several Pd(II) and Ni(II) thiocyanate complexes. The thiocyanate is N-bonded in all nickel(II) compounds.

[77FLE]

Fleet investigated synthetic α-Ni3S2 and demonstrated unambiguously that there are four similar Ni–Ni ‘bonds’ per Ni atom. The short Ni–Ni distances in heazlewoodite are only slightly greater than those for metallic Ni thus indicating metallic bonding.


381

[77KAS/SAK]

The chloride-formation reaction of nickel oxide and other metal oxides with HCl gas were studied at temperatures from 25 to 500°C and at 1.0 atm. The experiments were carried out by using gas analysis, a “thermobalance” and X-ray analysis to determine compositions for the reaction:

NiO (cr) + 2 HCl (g) NiCl2 (cr) + H2O (g).

A third-law analysis of the reported equilibrium constants at 298.15 and 773.15 K resulted in f mH ο∆ (NiCl2, cr, 298.15 K) = – (316.2 ± 6.4) kJ·mol–1. This value has a comparatively large uncertainty, does not overlap with the values obtained by the definitive studies of [53BUS/GIA] and [84LAV/TIM], and is therefore rejected.

[77KAT]

This paper (in Japanese) uses the same nickel sulphate conductiometric data reported previously by the same author [73KAT], and compares the results with those for other electrolytes.

[77NIC/ROB]

The otwayite was investigated with respect to its chemical composition, its optical and physical properties and powder diffraction patterns. No synthesis of otwayite has been attempted so far.

[77OGA]

The heat capacities of FeS2, CoS2, NiS2 and binary solid solutions of these sulphides were measured over the temperature range from 15 to 350 K in order to investigate the magnetic contribution to the specific heat at low temperatures. The solid samples were not characterised by X-ray diffraction analysis.

The temperature dependence of the heat capacity data for NiS2 is depicted in Figure A-28. Between 0 and 20.5 K, a Debye – T 3 function is fitted to the data, and at higher temperatures a spline fit of polynomials of the general form (a + bT + cT2 + dT –2 + eT −½) is applied, leading to the solid line in Figure A-28. The integration

298.15K

,m0( / )dpC T T∫ results in mS ο = (79.6 ± 2.4) J·K–1·mol–1. The integration of the ,mpC

function of FeS2 yields mS ο = 54.4 J·K–1·mol–1 which deviates from the value listed in the JANAF tables [98CHA], mS ο = 52.9 J·K–1·mol–1, by 3%. Thus, the same relative error is assumed for the standard entropy of NiS2, leading to the uncertainty of ± 2.4 J·K−1·mol−1.


382

Figure A-28: Heat capacity of NiS2 plotted versus temperature. ○ [77OGA], solid line: polynomial fit to the experimental data.

[77OSW/ASP]

Oswald and Asper reviewed the preparation and properties of bivalent metal hydrox-ides.

[78BEC/LIU]

In this study, the effect of temperature (0 to 180°C) on formation of inner-sphere and outer-sphere nickel sulphate complexes was investigated by 17O nmr spectroscopy. The solutions used were 0.097 m NiSO4 in 1.07 m and 2.34 m Li2SO4(aq). Results indicated that inner-sphere complexation increased with increasing temperature, and that at least two different inner-sphere complexes were formed, involving displacement of one and two inner-sphere water molecules, respectively. It was not clearly established whether the second complex was formed by bidentate complexation of a single inner-sphere sulphate ion, or by unidentate complexation of two sulphate ions. Near 25°C, the esti-mated value of the formation constant for the first inner-sphere complex was approxi-mately an order of magnitude less than the value of the formation constant as measured by methods [59NAI/NAN], [73KAT] that did not distinguish between inner- and outer-sphere complexation. The complex formed by displacement of two inner-sphere water molecules became predominant at temperatures above 160°C.

0 50 100 150 200 250 300 350

0

10

20

30

40

50

60

70

80

C p,m /

J·K−1

·mol

−1

T / K

C° p,m

/ J·K

–1·m

ol–1


383

[78BUR/KAM]

A potentiometric study of the hydrolysis of nickel(II), similar to the studies of [65BUR/LIL], [65BUR/LIL2], [66BUR/IVA] and [71BUR/ZIN], at 25°C, and in 3 M (Na)Br medium is described. The metal ion concentration ranged from 0.2 to 1.4 M (I = 3.2 – 4.4 M). The best fit to the experimental data was obtained assuming the presence of 4

4 4Ni (OH) + and Ni2OH3+ species, with formation constants 10 4,4*log b =

− (28.18 ± 0.05) and 10 1,2*log b = – (9.5 ± 0.1). During these measurements, a large part

of the Na+ ion is replaced by Ni2+; consequently the constancy of activity coefficients at different nickel(II) concentrations is questionable. Furthermore, the formation of bromo complexes (both binary and ternary) was not considered.

Therefore, the reported experimental data were re-evaluated for the purposes of this review. At I = 3.0 M, equilibrium constant for the reaction:

Ni2+ + Br– NiBr+ (A.60)

10 1log b (A.60) = – 0.49 can be calculated for the NiBr+ complex in NaBr medium, us-ing the recommended value 10 1log οb (A.60) and ε(NiBr+, Br–) = ½{ε(Ni2+, Br–) + ε(Na+, Br−)} = (0.16 ± 0.03) kg·mol–1. Considering the whole dataset and the same model as proposed, 10 4,4

*log b = − (28.22 ± 0.04) (3σ) and 10 1,2*log b = – (9.34 ± 0.08)

can be calculated, in good agreement with the published values. Using only the data sets reported for [Ni2+]tot = 0.2, 0.4 and 0.8 M (to reduce the influence of medium effects), and taking into account the formation of NiBr+, 10 4,4

*log b = – (27.19 ± 0.06) (3σ) and 10 1,2

*log b = − (8.82 ± 0.10) can be derived. If the formation of NiOH+ is also consid-ered, the fit to the data does not improve. Thus, the values for the other two constants are selected in this review, with increased uncertainties: 10 4,4

*log b = − (27.19 ± 0.30) and 10 1,2

*log b = – (8.82 ± 0.40).

[78ENO/TSU]

Enoki and Tsujikawa measured the heat capacity of hydrothermally aged, bulky Ni(OH)2 crystals from 4.2 to 35 K. Whereas the magnetic entropy was found to be higher by 0.37 J·K–1·mol–1 than that of Sorai et al. [69SOR/KOS], ,mpC values above the transition temperature were lower. However, as the results were presented graphi-cally in the rather narrow temperature range mentioned above, they were taken into ac-count only by increasing the error limits.

[78FOR]

The equilibria in the nickel(II)-imidazole-chloride ternary system have been investi-gated by means of pH-metric and spectrophotometric titrations in 3 M Na(ClO4, Cl) solution. The latter method was used to determine the stability constant of the binary NiCl+ complex in the nickel(II)-chloride system. The concentration range studied was [Ni2+] = 0.096 to 0.300 M and [Cl–] = 1.25 to 3.00 M. The 10 1log b value obtained (− (0.47 ± 0.10)) is in good agreement with that determined from potentiometric data in


384

an indirect way from the ternary system (– (0.53 ± 0.10)). In this review, the mean value of these constants is accepted with an increased uncertainty ( 10 1log b = – (0.50 ± 0.40)).

[78KAR/HUB]

The stability constant of a complex of pyrophosphate with Ni2+ (presumably 22 7NiP O −

though this is not specified by the authors) was measured by an amperometric titration procedure. Small amounts of the nickel nitrate salt were titrated with the pyrophosphate in an ammonium nitrate buffer solution (0.1 M), adjusted to a pH value of 8 with aque-ous ammonia. A PbO2 indicating electrode was used with a platinum foil counter elec-trode and a saturated calomel reference electrode. The procedure was found to be very sensitive to pyrophosphate even in the presence of phosphate.

At 30°C, the measured value of 10log K was 5.35, and the authors used esti-mated activity coefficients to obtain 10log K = 6.35 at I = 0. The original data are not reported and cannot be re-analysed. For reaction:

Ni2+ + 42 7P O − 2

2 7NiP O −

in 0.1 M KNO3, 10log K at 30°C = (7.11 + 0.1∆ε)

As the ionic strength is low, the effect of the ∆ε term on the final calculated value of 10log K ο is limited. Rather than use estimated ε values for interactions involv-ing the pyrophosphate moieties (some of which are highly charged), it is simply as-sumed that ∆ε is zero. As a consequence, the uncertainty in the 10log K ο value is in-creased to ± 0.30 for the values from measurements in 0.1 M NH4NO3 solutions.

[78LIB/KOW]

The experimental procedure is very similar to that used in [75LIB/TIA]. In the 1978 work, a potentiometric method was used to determine the stability constants of MBr+ complexes in M(ClO4)2 medium (M = Mn2+, Co2+, Ni2+ and Zn2+) at different concentra-tions.

The following cell was used for the potentiometric determination of formation constants: Ag|AgCl,HBr(m1),Mg(ClO4)2(m) || M(ClO4)2(m),HBr(m1),AgCl|Ag, where m1 was fixed at around 0.01 m, while the molalities of the metal perchlorate (m) were 1.5, 2, 2.5 and 3 m. The emf of the cells varied between 1.6 and 12 mV. Among the metal ions studied, nickel(II) forms the least stable bromo complex. The variation of the stability constants within the series is dictated by the formation of inner-sphere com-plexes to different degrees.

[78LIN/HU]

The activity of sulphur in the high-temperature modification of Ni3S2 was determined as a function of composition of the solid solution at various temperatures ranging from 823 to 1023 K. The sulphide was equilibrated with gas mixtures of H2(g) and H2S(g) and the


385

composition of the gas phase was obtained from the flow rates of the gas streams. When S2(g) is the standard state for sulphur the activity of sulphur can be calculated from the well-known thermodynamic data for the gas phase equilibrium H2S(g) H2(g) + 1/2 S2(g). The composition dependence of the sulphur activity at constant temperature revealed that the high-temperature modification of Ni3S2 is composed of two non-stoichiometric solid solutions, viz. β1-Ni3S2 and β2-Ni4S3. The homogeneity range as well as the phase relationships of these phases were described by means of a sub-regular model. Based on the thermodynamic properties of these mixture phases (β1-Ni3S2 and β2-Ni4S3) a new phase diagram for the binary system Ni – S was calculated, opposing the nickel-rich portion of the phase diagram proposed by Kullerud and Yund [62KUL/YUN] as well as Rau [76RAU].

[78SCH/MIL]

The phase relations in the binary system Ni – S were investigated in the composition range from xNi = 0.33 to xNi = 0.83 at temperatures between 509 and 1287 K by employ-ing isothermal sublimation in Knudsen effusion cells. The authors derived Gibbs ener-gies of formation for Ni3S2, Ni7S6, NiS, Ni3S4, and NiS2 as a function of temperature as well as enthalpies of formation for T = 298.15 K. No original experimental data are re-ported. The results given by the authors (see Table A-16) are not consistent with any other thermodynamic study on nickel sulphides, published so far. Thus, the data given in this paper are discarded.

Table A-16: Gibbs energy functions for various nickel sulphides.

Reaction r mGο∆ (kJ·mol–1)

Temperature range (K)

r mH ο∆ (298.15 K) (kJ·mol–1)

3Ni + S2 Ni3S2 152.297 – 0.065689 T 925 – 1070 – 65.27

73 Ni + S2 ⅓Ni7S6 167.36 – 0.09037 T 640 – 840 – 90.79

2Ni +S2 2NiS 184.93 – 0.11004 T 670 – 1200 – 72.38

32 Ni + S2 ½Ni3S4 212.97 – 0.08703 T 600 – 625 – 79.08

Ni + S2 NiS2 180.33 – 0.12761 T 509 – 730 – 87.86

[78STU/FER]

For anhydrous NiBr2, the authors report low temperature ,mpC measurements between 7.9 K and 303 K, drop calorimetry measurements to 1150 K, and heat of solution meas-urements at 298 K. Samples of NiBr2(cr), NiSO4(cr) and H2SO4·6H2O(l) were dissolved in 4.36 m HCl(aq). Prior to use, the NiBr2 was heated for several weeks at 200°C, and the X–ray diffraction pattern was claimed to be in good agreement with the ASTM pat-tern. Nevertheless, it is not certain which crystalline form of the solid was used [34KET]. The values were combined with the results of earlier measurements to deter-mine r mH∆ (A.61) = (12.320 ± 0.319) kJ·mol–1.


386

2[HCl·12.731H2O](l) +2NaBr(cr) +NiSO4(cr) 19.462H2O(l) +2NaCl(cr) +H2SO4·6H2O(l) +NiBr2(cr) (A.61)

The auxiliary data used by the authors were updated using the values in Table A-17.

Table A-17: Auxiliary data used in the recalculation of NaBr heat of formation values with the data from [78STU/FER].

Reaction # Equation r mH ο∆ Reference

(A.62) 0.5H2(g) + 0.5Cl2(g) + 12.731H2O(l) HCl·12.731H2O – (162.344 ± 0.167) a

(A.63) Na(cr) + 0.5Br2(l) NaBr(cr) – (361.160 ± 0.200) b

(A.64) Ni(cr) + S(rh) + 2O2(g) NiSO4(cr) – (873.280 ± 1.570) c

(A.65) Na(cr) + 0.5Cl2(g) NaCl(cr) – (411.260 ± 0.120) d

(A.66) H2(g) + 2O2(g) + 6H2O + S(rh) H2SO4·6H2O(l) – (874.295 ± 0.418) a

(a) [82WAG/EVA] adjusted to be consistent with values in Table IV.1 (b) [82GLU/GUR] (selected auxiliary datum) (c) This review. Section V.5.1.2.2.1.3. (d) Table IV.1 ([82GLU/GUR], selected in [2001LEM/FUG])

Thus, for Ni(cr) +Br2(l) NiBr2(cr) (A.67)

r mH ο∆ (A.67) = r mH ο∆ (A.61) + 2 r mH ο∆ (A.62) + 2 r mH ο∆ (A.63) + r mH ο∆ (A.64) – 2 r mH ο∆ (A.65) – r mH ο∆ (A.66),

and f mH ο∆ (NiBr2, cr) = – (211.214 ± 1.752) kJ·mol–1.

The method used is a reasonable way to determine f mH ο∆ (NiBr2, cr), but any consequent determination of the value of the enthalpy of formation of Ni2+ with this value for f mH ο∆ (NiBr2, cr) is linked to the value from the sulphate cycle. The heat of solution of Ni(c) in the sulphuric acid solution is included in part of the cycle to deter-mine the required f mH ο∆ (NiSO4, cr) auxiliary value.

The authors’ integration of their low-temperature heat capacity measurements (adiabatic calorimetry) led to mS ο (NiBr2, cr, 298.15 K) = (122.42 ± 0.25) J·K–1·mol–1. The reported value of ,mpCο (NiBr2, cr, 298.15 K) was (75.40 ± 0.02) J·K–1·mol–1. The heat capacity results showed no sharp transition above 7.8 K, though there was a small excess heat capacity between 40 and 55 K, near the Néel temperature. There was no indication of the transition near 20 K [81ADA/BIL], [82WHI/STA], and the reported heat capacity values for temperatures between 7.8 K and 15 K are markedly higher than the corresponding values measured by White and Staveley [82WHI/STA]. Also, as shown in Figure V-19, for the purposes of extrapolation to 0 K, an equation in T 3 does not provide a satisfactory representation of the experimental ,mpC values between 7.89 K and 18 K. It appears that the values reported for temperatures below 20 K are


387

unreliable. In the present review, the contribution to the entropy from integration of ,mpC /T between 22 K and 45.4 K was calculated to be only 0.4 J·K–1·mol–1 greater than

the contribution as calculated from the results of White and Staveley; and the contribu-tion to the entropy between 46.5 K and 250.0 K was calculated to be 94.17 J·K−1·mol–1, in good agreement with 94.31 J·K–1·mol–1 from the later work [82WHI/STA].

The authors’ fitted heat capacity equation reproduces all the high-temperature experimental m m( ) (298.15)H T H ο− values to better than 0.5% and is fixed to the 298.15 K value from the low-temperature measurements. The values from the equation deviate slightly from the experimental values below 298.15 K.

In the same report, the authors reported low-temperature calorimetry heat ca-pacity measurements (9 to 70 K) for anhydrous NiSO4(cr). Limited sets of drop-calorimetry measurements were carried out for the same solid (403 to 1001.5 K), and these were used to generate equations for thermodynamic functions for NiSO4 (cr) from 298 to 1200 K.

[79FIS/FOX]

This paper presents a thorough discussion of the pressure dependence of ion association parameters for several 2:2 electrolytes, including NiSO4(aq). New conductance data are reported for 15 and 25°C to 2000 atm (analysed using the Fernandez-Prini modification of the Fuoss-Hsai equation). The trends in the association constants for Ni2+ with sul-phate are in agreement with earlier work [76SHI/TSU]. Reanalysis of the 1 atm results (assumed to be identical at 0.1 MPa) using a version of the Lee-Wheaton equation that is consistent with the SIT equation used in the present review (Appendix B) leads to results that are in good agreement with the other studies. The calculated association constants have a rather large (10%) statistical uncertainty, but it does not appear that the data are unusually badly scattered. In the present review, we estimate the uncertainty in K1 at 25°C (0.1 MPa) to be ± 60. The high-pressure results are consistent with values found in the study of Shimizu et al. [76SHI/TSU].

[79GOL/NUT]

Goldberg et al. reviewed the available literature data, and provided recommended val-ues for the activity coefficients and osmotic coefficients for solutions of Ni(NO3)2 in water at 298.15 K. The authors commented on the “unusually large amount of scatter” between (and within) the osmotic coefficient measurements. The tables extend only to solutions 4.623 m in Ni(NO3)2 and, based on the study of Sieverts and Schreiner [34SIE/SCH], the saturation solubility is 5.47 m at 298.15 K. Therefore, in the present review, no thermodynamic quantities for Ni(NO3)2 solids are based on this careful as-sessment.


388

[79KEM/KAT]

Emf measurements on an electrochemical cell involving the Ni – NiO system were car-ried out in the temperature range from 1191 to 1823 K. No values of f mGο∆ (NiO) have been reported hitherto at temperatures above 1726 K (the melting point of nickel). For this review, however, the results in the lower temperature range are of primary interest. The results were expressed by the authors as f mGο∆ (NiO) = ((− (232.463 ± 0.251) + (0.083596 ± 0.000167) (T/K)) kJ·mol–1. The standard deviation of the f mGο∆ values amounted to ± 0.209 kJ·mol–1.

[79RIC/VRE]

The standard molar entropy of solid NiB4O7 was estimated by the linear dependence of entropies on the radius of the compound’s cation constituent. The extrapolation to r(cation) → 0 led to the borate contribution of the standard molar entropy depending on the cationic charge.

[79WEI/HER]

Proton and deuteron relaxation rates in aqueous solutions of NiX2 (X = Cl, Br) in the concentration range 0.2 – 4 m are reported. The deuteron relaxation rates were found to depend substantially linearly upon the NiX2 concentration. Comparison of the concen-tration dependences of the deuteron relaxation rate in these solutions with those in dia-magnetic MgX2 solutions yielded information about the average number (nX) of halide ions bound in the first coordination sphere of Ni2+. For NiCl2 solutions nCl = (0.12 ± 0.02)mNiCl2, for NiBr2 nBr = (0.085 ± 0.02)mNiBr2 was found. Some additional measurements with solutions containing 0.02 m NiX2 and 1 – 4 m KX were also per-formed, and nCl = (0.045 ± 0.04)mKCl, and nBr = (0.06 ± 0.04)mKBr were found. From the latter data the equilibrium constant for the reaction:

Ni2+ + Cl– NiCl+ (A.68)

10 1log b ((A.68), 298.15 K) = – (1.26 ± 1.00) and for:

Ni2+ + Br– NiBr+, (A.69)

10 1log b ((A.69), 298.15 K) = – (1.11 ± 0.60) can be derived in 4.0 m KX solutions. Using the recommended ion interaction coefficients to calculate ∆ε (A.68) and ∆ε (A.69) for KX background electrolytes, 10 1log οb ((A.68), 298.15 K) = (0.18 ± 1.00) and

10 1log οb ((A.69), 298.15 K) = (0.8 ± 0.6) can be derived. For the sake of comparison, we also calculated the “semi-thermodynamic” formation constants ( '

10 1log οb ) as de-fined in [88BJE], using γ± (4 m KCl) = 0.577 and γ± (4 m KBr) = 0.608 [49ROB/STO]:

'10 1log οb ((A.68), 298.15 K) = – (1.03 ± 1.00) and '

10 1log οb ((A.69), 298.15 K) = − (0.89 ± 0.60).


389

[80BAR/RAN]

Reversible potentials of partially charged α- and β-Ni(OH)2 electrodes have been meas-ured up to an oxidation state of ≈ 2.5 over a range of KOH concentrations from 0.01 – 10.0 M. Couples derived from the parent α- and β-Ni(OH)2 systems can be distin-guished by the relative change in KOH level on oxidation and reduction as well as the difference in the formal potentials with respect to Hg | HgO | KOH.

[80CHI/SAB]

The solubility of Ni(OH)2 in pure water was measured from 298.15 to 313.15 K. Only the solubility products calculated by the assumption 2[Ni(II)] = [OH–] have been given (at 25°C 10 ,0log sK ο = – 11.82). Thus, the conclusion has been confirmed that by direct reaction of Ni(OH)2(cr) with water no reliable solubility data can be obtained [18ALM], [67BLA]. Thus, these results were not included in either Table V-6 or Figure V-11.

[80DUB/CLA]

The phase transition between millerite, NiS(β), and the NiAs-type high-temperature modification, NiS(α), was studied by calorimetric and emf measurements. NiS(α) was prepared by solid state reaction from the elements. Millerite was obtained by annealing NiS(α) at 627 K for 3 – 4 days. Both solid compounds were checked by X-ray diffraction analysis. The calorimetric measurements of enthalpy increments during heating or cooling the sample yielded the temperature and the enthalpy of transition (β → α), respectively, Ttrs = (667 ± 2) K, trs mH ο∆ = (5.86 ± 0.36) kJ·mol–1. In addition, the measurement of the emf of the electrochemical cell C (graphite) | Ni | NiCl2 | NiS | C (graphite) as a function of temperature allowed the determination of an independent value for the entropy of transition, trs mS ο∆ = (9.62 ± 0.84) J·K–1·mol–1.

[80LIB/SAD]

Isopiestic results (relative to KCl(aq)) are reported for NiSO4 at 25°C for nickel sul-phate molalities between 0.1271 and 1.8526 mol·kg–1, and changes in the UV and visi-ble spectra as a function of the concentration of nickel sulphate are discussed. The au-thors only present semi-quantitative estimates for the degree of association, and con-clude that except for concentrations above 1 M, complexation between Ni2+ and 2

4SO − is primarily outer-sphere complexation.

[80MIL/BUG]

The hydrolysis of nickel(II) in (K)Cl medium was studied by emf titrations at 25°C. In the first series of titrations, the chloride concentration was held constant at 2.0 M, and [Ni2+]t was varied from 0.05 – 0.50 M (I = 2.05 – 2.5 M). In the second series, the [Ni2+]t was held constant at 0.10 M while the chloride concentration was varied from 0.5 to 3.0 M (I = 0.6 – 3.1 M). Taking into account the smaller medium effect, the results of


390

the second series are more reliable. The experimental data were interpreted by assuming that 4

4 4Ni (OH) + was the main hydrolysis product formed. At I = 0.5 and 1.5 M the value of 10 1,1

*log b was also estimated, although the authors noted that the formation of NiOH+ is not completely certain. The formation of the NiCl+ complex was not consid-ered in the calculation; therefore a correction was applied for the purposes of this re-view, using the recommended 10 1log οb (NiCl+) and ∆ε(298.15 K, KCl) = − (0.085 ± 0.05) kg·mol–1 (see the discussion of [71OHT/BIE]). The corrected forma-tion constants are as follows:

Table A-18: Corrected formation constants.

I (M) 0.5 1.0 1.5 2.0 3.0

10 1,1*log b reported – (9.87 ± 0.40) – – (10.06 ± 0.40) – –

10 4,4*log b reported – (28.04 ± 0.10) – (28.14 ± 0.06) – (28.28 ± 0.06) – (28.24 ± 0.04) – (28.58 ± 0.06)

10 1log b (NiCl+) – 0.58 – 0.65 – 0.67 – 0.67 – 0.63

10 1,1*log b corrected – (9.82 ± 0.50) – – (9.95 ± 0.50) – –

10 4,4*log b corrected – (27.83 ± 0.30) – (27.80 ± 0.30) – (27.80 ± 0.30) – (27.63 ± 0.30) – (27.66 ± 0.30)

The SIT-plot of the recalculated data, together with the accepted values for [65BUR/LIL2] and [71OHT/BIE], as well as the selected 10 4,4

*log οb for the reaction:

4Ni2+ + 4H2O 44 4Ni (OH) + + 4H+, (A.70)

is presented in Figure A-29. As can be seen, the analysis of the data from [80MIL/BUG] results in the determination of a considerably different 10 4,4

*log οb value than the value selected in this review. The different medium cation can not be the reason of this devia-tion, since both the aquo ion and hydroxo complex of nickel(II) are positively charged. Therefore, the above data were not considered further in this review.


391

Figure A-29: SIT-plot of recalculated data from [80MIL/BUG] (open diamond), [65BUR/LIL2] and [71OHT/BIE] (open squares), and the selected 10 4,4

*log οb ((A.70), 298.15 K) (filled square).

[80OPP/TOS]

Oppermann and Toschew studied the thermal decomposition and the sublimation of crystallised nickel(II) iodide using a membrane manometer and transport ampoules. The temperature functions of the total pressure (membrane manometer), as well as of the sublimation pressure (transport ampoules), in a range from 800 to 1000 K, were ther-modynamically analysed, and the following quantities were obtained: f mH ο∆ (NiI2, cr, 298.15 K) = – (83.7 ± 8.4) kJ·mol–1, f mH ο∆ (NiI2, g, 298.15 K) = (130.5 ± 20.9) kJ·mol−1, mS ο (NiI2, cr, 298.15 K) = (146 ± 8) J·K–1·mol–1·, mS ο (NiI2, g, 298.15 K) = (335 ± 8) J·K–1·mol–1·, and mS ο (Ni2I4, g, 298.15 K) = (536 ± 13) J·K–1·mol–1·. In this review the uncertainties given by the authors were accepted.

The recalculation of these quantities suffered from the total pressure function having been presented only graphically. For the sake of comparison with independent measurements, an attempt was made to recalculate f mH ο∆ (NiI2, cr, 298.15 K) and

mS ο (NiI2, cr, 298.15 K), regardless of the uncertainty implicit in reading the relevant coordinates from a Figure. Based on the auxiliary data used throughout this review the second-law and approximate third-law methods employed by Oppermann and Toschew resulted in f mH ο∆ (NiI2, cr, 298.15 K) = – (91.45 ± 4.18) kJ·mol–1 and f mH ο∆ (NiI2, cr, 298.15 K) = – (91.00 ± 8.37) kJ·mol–1, respectively. Within the uncertainty limits, both of these values overlap with those reported by the original authors. The recalculated enthalpies of formation also agree with the value selected in this review (− (96.42 ± 0.84) kJ·mol–1), but their uncertainty is 5 to 10 times greater.

0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5-29.0

-28.5

-28.0

-27.5

-27.0

log 10

* β 4,4 −

4D

− 4

log 10

aw

I / mol·kg−1


392

As a reliable value for mS ο (NiI2, cr, 298.15 K) is a prerequisite for the third-law analysis, this quantity only can be determined by the second-law method. The standard entropy of nickel(II) iodide is recalculated to be mS ο (NiI2, cr, 298.15 K) = (138.0 ± 8.4) J·K–1·mol–1, which, within the large uncertainty limits, agrees very well with the value selected in this review ( mS ο (NiI2, cr, 298.15 K) = (139.1 ± 1.0) J·K–1·mol–1).

[80REI]

In this work neutral anhydrous transition metal carbonates MCO3 (M = Mn, Fe, Co, Ni, Zn) were synthesised by a hydrothermal autoclave method. The solubility of the well crystallised pure carbonates (size of rhombohedral crystals 10 – 100 µm) in HClO4/NaClO4 solutions of constant ionic strength (1.0 mol·kg–1) and at various partial pressures of CO2 was studied by means of the pH variation method. At T = 323 K equi-librium was usually attained within a few days. Only NiCO3 was so inert that it had to be equilibrated at 363 K ≥ T ≥ 348 K. When the Gibbs energy of formation was calcu-lated from the solubility constant obtained it was noted that it is remarkably more nega-tive than the then-accepted value [77BAR/KNA].

[80SHA/CHA]

The thermodynamic properties of all phases in the binary system Ni – S were optimised for the first time, allowing the calculation of the phase diagram for this system. The liquid phase was modelled by applying the three suffix Margules equation, the high-temperature phases β1-Ni3S2 and β2-Ni4S3 were described using sub-regular models, and the non-stoichiometric hexagonal modification of nickel monosulphide was modelled by application of a statistical model [69LIG/LIB] adopted previously for pyrrhotite, Fe1–

xS, [77LIN/IPS], [79SHA/CHA]. The phase Ni7S6, which shows a narrow homogeneity range [94STO/FJE], was assumed to be a stoichiometric compound, neglecting its trans-formation into Ni9S8 at 675 K [94STO/FJE]. The phases Ni3S2, NiS (low-temperature form), Ni3S4, NiS2 were likewise assumed to be stoichiometric compounds. The calcu-lated phase diagram enables excellent predictions of available experimental phase dia-gram data [62KUL/YUN], [75ARN/MAL].

However, the phase relations of Ni7S6 and Ni9S8, respectively, were not con-sidered properly. Furthermore, it should be mentioned that the thermodynamic data for NiS2, the low- and high-temperature forms of NiS are not in perfect coincidence with decomposition pressure measurements [54ROS], [64LEE/ROS] as well as heat capacity measurements published recently [95GRO/STO].

[80TRE/LEB2]

The solubility of NiO, characterised by X-ray diffraction analysis, was measured in a flow apparatus from 423 to 573 K. The solid phase was treated with acidic solutions, containing HCl with molalities ranging from 10–5 to 5 × 10–4 mol·kg–1, or basic solu-tions, containing NaOH with molalities ranging from 10–6 to 4 × 10–2 mol·kg–1. An ex-


393

tended Debye-Hückel expression was used by the authors in order to calculate the activ-ity coefficients. The pH of the leach solutions was measured with a glass electrode, and varied between 3 – 12. The nickel content of solutions equilibrated with NiO was de-termined continuously by means of an atomic-absorption spectrometer. As dilute solu-tions were investigated, only the formation of mononuclear hydroxo complexes (NiOH+, Ni(OH)2 and 3Ni(OH)− ) was considered. The values for the hydrolysis con-stants at 25°C were taken from [64PER] and [49GAY/GAR]. From the temperature dependence of the solubility data the authors derived the first, second and third hydroly-sis constants, see Table A-19.

Table A-19: Stability constants for the complexes NiOH+, Ni(OH)2 and 3Ni(OH)−

t /°C 10 1,1*log b (A.71) 10 2,1

*log b (A.72) 10 3,1*log b (A.73)

298.15 – (9.86 ± 0.03)(a) – (12.01 ± 1.64)(c)

< 19(b) – (20.26 ± 1.77)(c)

– (29.03 ± 0.50)(b) – (29.78 ± 1.97)(c)

373 – (9.37 ± 0.85)(c) – (16.86 ± 0.95)(c) – (26.03 ± 1.12)(c) 423 – (8.05 ± 0.48) – (15.18 ± 0.58) – (24.19 ± 0.76) 473 – (6.95 ± 0.22) – (13.81 ± 0.35) – (22.68 ± 0.55) 523 – (6.03 ± 0.20) – (12.65 ± 0.32) – (21.42 ± 0.50) 573 – (5.23 ± 0.36) – (11.67 ± 0.43) – (20.35 ± 0.57)

a: Taken from [64PER]. b: Taken from [49GAY/GAR]. c: Extrapolated from 150 – 300°C.

From these data the authors obtained:

Ni2+ + H2O Ni(OH)+ + H+ (A.71) r mH ο∆ (A.71) = (87 ± 15) kJ mol–1,

Ni2+ + 2 H2O Ni(OH)2(aq) + 2 H+ (A.72) r mH ο∆ (A.72) = (109 ± 15) kJ mol–1

Ni2+ + 3 H2O 3Ni(OH)− + 3 H+ (A.73) r mH ο∆ (A.73) = (119 ± 15) kJ mol–1.

The experimental data from this paper have been re-evaluated in the present assessment. New values for the third hydrolysis constant, valid from 423 to 573 K, were determined by non-linear curve fitting. The equilibrium values of the time-independent solubility for each run are plotted versus the molality of HCl or NaOH of the starting solutions in Figure A-30 and Figure A-31. The solid lines in Figure A-30 correspond to the prediction of the solubility of NiO in acidic solutions according to the thermody-namic model of the present compilation including the data for NiOH+. In the range [HCl]initial > 10–4.9 mol·kg–1 the predominant species is Ni2+. The calculated lines are biased to somewhat higher nickel concentrations which may be caused (i) by insuffi-


394

cient equilibration of the solid phase (NiO) with the aqueous media or (ii) by uncertain-ties of the solubility calculations for hydrothermal conditions owing to the lack of heat capacity functions for the aqueous species. However, the deviation of the predicted lines from the experimental data is less than 0.15 logarithmic units.

Figure A-30: Solubility of NiO in acidic solutions as a function of initial molality of HCl. Symbols refer to experimental data [80TRE/LEB2], the solid lines correspond to the thermodynamic model of the present assessment.

The results in basic solutions indicate the occurrence of the hydrolysis species 3Ni(OH)− owing to the increase of the solubility at higher NaOH concentrations, see

Figure A-31. Due to the scatter of solubility data, the stability constants for Ni(OH)2(aq) were not determined. For [NaOH]initial < 10–4 mol·kg–1, the predominant species is NiOH+. The stability constants of the complex 3Ni(OH)− are listed in Table A-20 for the temperature range 423 – 573 K, neglecting the formation of Ni(OH)2(aq).

-5.0 -4.8 -4.6 -4.4 -4.2 -4.0 -3.8 -3.6 -3.4 -3.2-8

-7

-6

-5

-4

-3

573 K 523 K 473 K 423 K

log 10

([Ni(I

I)]to

t / m

ol·k

g−1)

log10

([HCl]initial

/ mol·kg−1)


395

Figure A-31: Solubility of NiO in basic solutions as a function of initial molality of NaOH. Symbols refer to experimental data [80TRE/LEB2], the lines have been ob-tained from non-linear least squares fits to the experimental data in order to optimise the stability constants for 3Ni(OH)− ;

, : 423 K; , : 473 K; , : 523 K; , : 573 K.

Table A-20: Stability constants 10 3,1*log οb (A.73) for 3Ni(OH)− , I = 0

T/K 10 3,1*log οb

this assessment10 3,1

*log οb [80TRE/LEB2]

423 – 24.61 – 24.2

473 – 22.97 – 22.7

523 – 21.83 – 21.4

573 – 20.64 – 20.3

From the temperature dependence of the optimised data in Table A-20 the lin-ear extrapolation of the stability constant to 298.15 K ( 10 3,1

*log b plotted versus 1/T) results in:

10 3,1*log οb = – (30.9 ± 1.3).

-6.5 -6.0 -5.5 -5.0 -4.5 -4.0 -3.5 -3.0 -2.5 -2.0 -1.5 -1.0-10

-9

-8

-7

-6

-5

-4

log10

([NaOH]initial

/ mol·kg−1)

log 10

([Ni(I

I)]to

t / m

ol·k

g−1)


396

The corresponding enthalpy value for 298.15 K is found to be:

r mH ο∆ ((A.73), 298.15K) = (121.2 ± 6.5) kJ·mol–1.

[81BAR/RAN]

Reversible potentials of partially charged α- and β-Ni(OH)2 electrodes have been meas-ured. Experimental emf-oxidation state measurements for both the β | β- and the α | γ-phase couples agree well with theoretically derived expressions. From considerations of the homogeneous emf-composition regions it is deduced that the oxidised species are dissociated in both the β | β- and the α | γ-phase systems.

[81KUL/BLO]

The authors have reported the formation and thermodynamic parameters (K, r mH∆ , r mS∆ ) of fluoride complexes of the Mn(II) – Zn(II) series in three papers

[76KUL/BLO], [81KUL/BLO], [83AVR/BLO]. In the first paper [76KUL/BLO] they reported the stability constants and thermodynamic functions of complex formation determined at 298.15 K and at I = 3 M NaClO4, by means of a fluoride-selective elec-trode and a differential double calorimeter. The experiments were performed in the presence of an excess of metal ions (0.1 M Me(ClO4)2) to assure the formation of only monofluoro complexes.

In the second paper [81KUL/BLO] the authors reported the ionic strength dependence of thermodynamic quantities using the same experimental techniques. The initial concentration of the metal ions was varied with varying ionic strength maintained with NaClO4 ([M] = 0.02 M for I = 0.1 M, [M] = 0.05 M for I = 0.25 M and [M] = 0.1 M for I = 0.5, 1.0, 2.0 and 3.0 M). The data revealed that the stability of MF+ complexes varies in the following order: Mn > Co ~ Ni < Cu > Zn, which is not consistent with the Irving-Williams series. The authors also reported thermodynamic properties extrapo-lated to I = 0, but a reassessment of these data using the SIT was done in this review. The authors classified NiF+ as an outer-sphere complex, since the addition of fluoride ions to a solution of nickel perchlorate is not accompanied by any changes in the elec-tronic spectrum, and the thermodynamic quantities are very similar to those of the com-plex [Co(NH3)5NO2]F+, which is clearly an outer-sphere complex.

In the third paper [83AVR/BLO], the stability constant values for NiF+ have been reported for 0 – 60 vol. % H2O/EtOH mixture (I = 0.25, 0.5 and 1.0 M NaClO4), using a fluoride-selective electrode. In contrast with the earlier conclusions, the authors suggest inner-sphere association for NiF+. This assumption was based on the determina-tion of the so called complete equilibrium constant, in which the water is also taken into consideration in the equilibrium process:

22 6Ni(H O) + + F– 2 5NiF(H O)+ + H2O(l).


397

The reported thermodynamic quantities seem to have been re-determined in each study, since the three papers report somewhat different values for the parameters referring to identical conditions. To avoid the over-representation of the data from this laboratory during our SIT analysis, we used an average value for each ionic strength studied.

[81MAR/ECO]

Marcopoulos and Economou found almost pure nickel hydroxide in the Vermion region of northern Greece, described it comprehensively, named it theophrastite [81MAR/ECO], and studied its genesis [83ECO/MAR]. The new mineral and its name were approved of by the Commission on New Minerals and Mineral Names, the Inter-national Mineralogy Association, IMA.

[81NIC/BER]

Crystallography and composition of nullaginite, found as nodular grains and cross-fibre vein lets in a nickel-rich assemblage from the Otway deposit, Nullagine district, West-ern Australia were investigated. The close relationship between nullaginite, Ni2(OH)2CO3, and rosasite, (Cu, Zn)2(OH)2CO3, as well as glaukosphaerite, (Cu, Ni)2(OH)2CO3, was pointed out.

[81SCH]

The partial pressure of oxygen for the equilibrium between β2-Ni4S3 and NiO at 2SOp =

1 atm was measured as a function of temperature between 968.6 and 1053.7 K. The emf of the electrochemical cell Pt, β2-Ni4S3, NiO, SO2 (1 atm) | ZrO2 | O2 (0.0122 atm), Pt can be easily transformed into the equilibrium partial pressure of oxygen. The solid phases were checked by X-ray diffraction analysis. During equilibration sulphur was incorporated in excess in the non-stoichiometric high-temperature phase β2-Ni4S3 due to reaction with SO2 of the gas phase. The sulphur content of quenched samples was de-termined by standard analytical procedures. A composition of Ni2.515S2 was calculated from these chemical analyses. However, it is not expected that the sulphur content of β2-Ni4S3 remains constant when the temperature is varied. As the variation of the com-position of β2-Ni4S3 in equilibrium with NiO and SO2 was not investigated as a function of temperature, the results of this work are not suitable for derivation of the Gibbs en-ergy function of this phase.

[81YOK/YAM]

The data of Katayama [73KAT] were re-interpreted and compared to a calculated theo-retical value. No new data are presented.

[82CHI/SAB]

Chickerur et al. [82CHI/SAB] reported solubility product values for Ni3(PO4)2·8H2O of 4.07 × 10–25 and 1.95 × 10–21 at 30 and 35°C, respectively. The paper provides few de-


398

tails, though there is a suggestion that the measurements were done in solutions with “pH” values < 5.75. At higher temperatures, the authors indicated that Ni3(PO4)2·8H2O converted to “NiHPO4·H2O”. No allowance was made for complexation, no raw data were presented, and no values can be recalculated from the sparse information in this paper.

[82FER/GOK]

Enthalpy increments from 298.15 to 1197 K were measured on Ni3S2 by means of drop calorimetry, leading to heat capacity values in this temperature region. The samples were prepared by solid state reaction from the elements and characterised properly (X-ray diffraction). The transition temperature for the phase transformation between the high- and low-temperature form was found to be 834 K. The heat capacity at 298.15 K amounts to 117.7 J·K–1·mol–1, which is in fair agreement with that obtained by Stølen et al. [91STO/GRO], ,mpCο (298.15 K) = 118.2 J·K–1·mol–1.

[82GAM/REI]

The solubility constants measured at 363 K ≥ T ≥ 348 K were extrapolated to T = 298.15 K implying a linear dependence of 10 , ,0

*log p sK for the reaction:

NiCO3(s) + 2H+ Ni2+ + CO2(g) + H2O(l)

on 1/T [80REI]. Data obtained at I = 1.0 and 0.2 mol·kg–1 were extrapolated to I = 0 with the Davies approximation [62DAV]. On this basis f mGο∆ (NiCO3, cr, 298.15 K) was calculated.

[82KLO]

The solubility of NiCO3(cr) synthesised by the hydrothermal autoclave method was determined at T = 363.15 K and I = 0.2 mol·kg–1 NaClO4.

[82LIV/BIS]

Livingstone and Bish reported X-ray powder diffraction, infrared, thermal, chemical, optical, and physical data for a Mg-bearing theophrastite from Hagdale Quarry, Unst, held in the Scottish Mineral Collection. These authors endeavoured to establish a min-eral name for the compound (Ni, Mg)(OH)2 by a submission to the IMA in late 1979. Two months later IMA received another submission for pure Ni(OH)2 [81MAR/ECO]. In the vote the Unst material and name were narrowly defeated in favour of the Greek Ni(OH)2.

The values of c determined by Livingstone and Bish are intermediate between those for pure Ni(OH)2 and the isomorphous Mg(OH)2, brucite, indicating that a solid-solution series exists between these end-members.


399

[82WAT]

Watanabe carried out measurements of the molar heat capacities of high pressure com-pounds relevant to the earth’s mantle by differential scanning calorimetry between 350 and 700 K. The investigated compounds were, among others, α-Ni2SiO4 and γ-Ni2SiO4. In general, the heat capacity for the high pressure phase was found to be smaller than that for the isochemical low-pressure phase i.e., ,mpC decreases in the phase sequence olivine (α), modified spinel (β), and spinel (γ). The calculated standard entropies at 298 K for high-pressure phases are presented and entropy changes for various transforma-tions of silicate compounds were discussed. The heat capacities of Ni2SiO4, olivine, reported by Watanabe are 4.6% larger than the heat capacities obtained by [84ROB/HEM] with the more accurate low-temperature calorimeter. Therefore, the results of Watanabe are not used by this review.

[82WHI/STA]

The authors report low-temperature heat capacity measurements (adiabatic calorimetry) from 8.03 K to 299.83 K. A sample of NiBr2 that had been resublimed at 600°C was used. When plotted against temperature, the results show a shoulder at 46 K, and a spike at 19.5 K (Figure A-32). The values between 20 K and 250 K are in reasonable agree-ment with the earlier work of Stuve et al. [78STU/FER] (cf. discussion in Appendix A). At lower temperatures, the values are lower, and the value of ,mpC (NiBr2, cr, 8.0 K) is only 30% of the value reported by Stuve et al. [78STU/FER]. As shown in Figure V-19, for the purposes of extrapolation to 0 K, an equation in T 3 provides a satisfactory repre-sentation of the experimental ,mpC values between 8.03 K and 18 K. In the present re-view, polynomials were fitted to ,mpC /T, and integration resulted in mS (NiBr2, cr, 250 K) = 108.23 kJ·mol–1. The authors reported difficulties with their measurements between 250 K and 300 K, and the values of ,mpC increase more rapidly than expected. Integration of ,mpC /T between 250 K and 300 K gives a contribution to mS (NiBr2,cr) that is 0.61 J·K–1·mol–1 greater than the contribution obtained by integration of values from Stuve et al. over the same temperature range. The results from White and Staveley for the temperature range 250 K to 300 K are not accepted in the present review.

[83AVR/BLO]

See comments under [81KUL/BLO].

[83ECO/MAR]

Marcopoulos and Economou found almost pure nickel hydroxide in the Vermion region of northern Greece, described it comprehensively, named it theophrastite [81MAR/ECO], and studied its genesis [83ECO/MAR]. The new mineral and its name were approved of by the Commission on New Minerals and Mineral Names, the Inter-national Mineralogy Association, IMA.


400

Figure A-32: Values of ,mpC (NiBr2, cr) as a function of temperature as reported by White and Staveley [82WHI/STA].

[83GOW/DOD]

A 35Cl NMR study of NiCl2 solutions has shown that at – 5 and + 1°C the system is in the slow-exchange region. This allowed the concentration of the bound chloride ion to be calculated, and thus the stability constant for the inner-sphere 2 5Ni(H O) Cl+ complex at different concentrations of the salt. The average number of inner-sphere-bound chlo-rides varied between

Cln − = 0.07 and 0.27, depending on the concentration of NiCl2. The

determined stability constants do not show notable temperature or ionic strength de-pendence, and 10 1log b = (0.034 ± 0.020) was determined from the whole dataset. The applied ionic strength is out of the validity range of the SIT. Nevertheless, if it is used,

10 1log οb for the reaction:

Ni2+ + Cl– NiCl+

can be calculated to be – (0.95 ± 0.50). The method applied has substantial inherent inaccuracy, but since outer-sphere interaction is unlikely to produce slow ligand-exchange, the results verify the existence of an inner-sphere complex.

5 10 15 20 25 0

2

4

6

8

10

0 50 100 150 200 250 3000

20

40

60

80

100

Cp

/ J·

K-1·m

ol-1

T / KT / K

C° p,m

/ J·K

–1·m

ol–1


401

[83HOL/MES]

The authors carried out isopiestic measurements on aqueous solutions of nickel sulphate at 383.14 K and 413.22 K for nickel sulphate molalities between 1.52 and 5.66. The results were interpreted using a modified Pitzer equation without invoking formation of NiSO4(aq) as discrete species. The results at these high ionic strengths were not used in the present assessment.

[83SOL/BON]

The thermodynamics (K, r mH∆ , r mS∆ ) for the formation of monofluoride complexes of the Mn(II) – Zn(II) series have been studied in methanol at I = 0.05 M (CH3)4NClO4 and 298.15 K with use of fluoride selective electrode potentiometry. Additionally, the K values have been determined in aqueous solution (I = 0.05 M (CH3)4NClO4 and 298.15 K). In a typical measurement, a solution containing 2 – 40×10–5 M fluoride was titrated with the solution of the given metal ion, to reach a final metal ion concentration of 2 – 80×10–4 M. This work was devoted to gaining further insight into the anomalous stability order of the monohalide complexes of 3d transition metal ions first reported in [72BON/HEF]. The results indicated that the monofluoride complexes are a clear and distinct exception to the Irving-Williams sequence; the stability of MF+ complexes was found to be Mn ~ Fe > Co > Ni < Cu > Zn in both water and methanol. This stability order does not follow even the trend predicted by ionic radii. Therefore the authors sug-gested that the monofluoride complexes have some characteristics of ion pairs rather than those of inner-sphere complexes. The thermodynamic data indicated that the com-plex (ion pair) formation reactions are all distinctly entropy controlled.

[84BRA/DEL]

Sodium nickelate, NaNiO2, was prepared by direct reaction of sodium oxide and nickel oxide at 600 – 800°C under a stream of oxygen. This compound was hydrolysed and reduced by so-called soft chemical reactions in aqueous media at room temperature.

[84COM/PRA]

The standard molar Gibbs energy of formation of NiO was determined at temperatures as low as 760 K up to 1275 K with an excellent accuracy of ± 0.039 kJ·mol–1 (which is ± 0.2 mV for the standard deviation of the emf). The obtained results were represented by the equation: f mGο∆ (T) = – 232.450 + 0.083435 (T/K) kJ·mol–1.

[84LAV/TIM]

High-purity nickel (wimpurities < 6 ×10–5) was subjected to chlorination in a calorimetric bomb with a microfurnace for sample heating. The microfurnace developed a tempera-ture of about 1070 K over approximately 10 min at a power input of 75 W. Under these conditions nickel chlorination proceeded practically to completion.


402

[84ROB/HEM]

Robie et al. have carried out excellent measurements of the heat capacity of Ni2SiO4 olivine between 5 and 387 K by cryogenic adiabatic-shield calorimetry and between 360 and 1000 K by differential scanning calorimetry. At 298.15 K the molar heat capac-ity and entropy of Ni2SiO4 olivine were determined. This molar heat capacity and en-tropy were accepted by this review. The thermal heat capacity function proposed by Robie et al. for the temperature range between 300 and 1300 K was, however, not adopted by this review, because it is based on measurements which were made only at temperatures up to 1000 K. Combining the heat capacity measurements with results of molten salt calorimetry, thermal decomposition of Ni2SiO4 olivine into its constituent oxides, and equilibrium studies, both by CO reduction and solid state electrochemical cell measurements for the reaction:

2Ni + SiO2 + O2 Ni2SiO4,

Robie et al. calculated for Ni2SiO4 olivine f mGο∆ (298.15 K) = – (1289.0 ± 3.1) kJ·mol–1 and f mH ο∆ (298.15 K) = – (1396.5 ± 3.0) kJ·mol–1.

[84ROG/BOR]

The authors carried out emf-measurements using a galvanic cell with a Ni2+ exchanged β-alumina solid electrolyte to determine the Gibbs energy, enthalpy and entropy of for-mation of nickel orthosilicate from the constituent oxides. The Gibbs energy of the cell reaction as a function of temperature, given by Róg and Borchard for the temperature range 970 to 1370 K, coincides very well with results of other emf-studies of Ni2SiO4, as shown in this review. The estimated temperature of decomposition of Ni2SiO4 to NiO and SiO2, 1790 K, is significantly lower than the value recommended by this review (1820 ± 5) K.

[84VAS/VAS]

Vasil’ev et al. determined r mH∆ of Reaction (A.74) calorimetrically at 298.15 K. Me-tallic nickel was dissolved in 2.0 to 6.0 M HCl or HClO4 solutions containing 1.0 to 1.5% H2O2.

Ni(cr) + 2H+ + H2O2(aq) Ni2+ + 2H2O(l) (A.74)

When the authors plotted the measured enthalpies minus the Debye-Hückel term versus the ionic strength (which appears to be dominated by the acid concentra-tion) the data fell on straight lines with slopes depending on the H2O2 content, but the intercepts coincided within the error limits. Thus, the results were extrapolated semi-empirically to I = 0, resulting in r mH ο∆ (A.74) = – (432.59 ± 0.43) kJ·mol–1. As the SIT model cannot predict the H2O2 dependence of the slopes this evaluation has been ac-cepted as is. With the standard enthalpies of f mH ο∆ (H2O, l) = – (285.830 ± 0.040) kJ·mol–1 and f mH ο∆ (H2O2, aq) = – (191.170 ± 0.100) kJ·mol–1 taken from CODATA


403

and the NBS Tables, respectively, f mH ο∆ (Ni2+) = – (52.10 ± 0.45) kJ·mol–1 has been obtained [82WAG/EVA], [89COX/WAG].

[85DRA/MAD]

Apparent molar heat capacities were determined for nickel sulphate solutions at 50, 70 and 90°C. The lowest molalities used were 0.3 mol·kg–1, and the extrapolation to I = 0 is problematic, especially if a model that invokes association between nickel and sulphate ions is to be used. The values from this study are not used in the present assessment.

[85MEH/TAR]

The partial pressure of oxygen for the solid phase equilibrium β2-Ni4S3 / NiO at 2SO

p = 1 atm was determined in the temperature range 973 – 1173 K by performing emf meas-urements on the electrochemical cell Pt, O2 (air) | YSZ | NiO, β2-Ni4S3, SO2 (1 atm), Pt or Au. Yttria-stabilised zirconia served as the solid electrolyte. The solid phases were prepared by solid state reactions from the elements. Phase identification was carried out by means of X-ray diffraction analysis and the composition of the solid nickel sulphide was checked by chemical analysis. The authors found a composition of Ni2.666S2. The variation of this composition with temperature was reported to be not appreciable, but was not checked carefully by the authors. Hence, the results of this work are apparently not suitable for the derivation of the Gibbs energy function for this phase.

[85SKE/MAI]

Skeaff et al. [85SKE/MAI] used an electrochemical technique to determine the Gibbs energies of reaction of NiAs(cr) with oxygen to form NiO (bunsenite) and As4O6(g) from 711 K to 833 K. Similar measurements were carried out with Ni11As8(cr) (805 K to 991 K) and Ni5As2(cr) (937 K to 1139 K). Experiments with NiO-NiAs2 and NiAs-NiAs2 mixtures were unsuccessful in that there was no indication of any change in potential with changes in partial pressure of As4O6(g).

Based on the approximation that r ,mpC∆ = 0 for

xNiO(cr) + 4y As4O6(g) NixAsy(cr) +

4(2 3 )x y+ O2(g)

the authors’ estimated heat capacity equations (in J·K–1·mol–1) were:

,mpCο (NiAs, cr) = 63.7532 + 6.4616 ×10–3 (T/K) – 8.6016 ×105 (T/K)–2

,mpCο (Ni11As8, cr) = 605.483 + 70.7846 ×10–3 (T/K) – 66.3005 ×105 (T/K)–2

,mpCο (Ni5As2, cr) = 222.965 + 32.0067 ×10–3 (T/K) – 14.6911 ×105 (T/K)–2

and the following values were reported: f mH ο∆ (NiAs, cr, 298.15 K) = – 73.29 kJ·mol–1, mS ο (NiAs, cr, 298.15 K) = 45.4 J·K–1·mol–1, f mH ο∆ (Ni11As8, cr, 298.15 K) = – 777.04

kJ·mol–1, mS ο (Ni11As8, cr, 298.15 K) = 468.9 J·K-1·mol-1, f mH ο∆ (Ni5As2, cr, 298.15 K) = – 251.10 kJ·mol–1, and mS ο (Ni5As2, cr, 298.15 K) = 190.6 J·K–1·mol–1.


404

The authors used values for the heat capacity and entropy of As4O6(g) from the analysis of Behrens and Rosenblatt [72BEH/ROS], and tied the value for f mH ο∆ (As4O6) to the vapour pressure measurements of [72BEH/ROS] and Murray et al. [74MUR/POT]. Behrens and Rosenblatt, in turn, relied on values calculated from the thorough sets of low-temperature heat capacity measurements of Chang and Bestul [71CHA/BES] for the As4O6 solids. As was discussed by Grenthe et al. [92GRE/FUG], auxiliary data for arsenic compounds have not been critically reassessed in the TDB Project, and at 298.15 K several of the self-consistent values from the U.S. National Bureau of Standards compilation [82WAG/EVA] are accepted as an interim measure. The values f mH ο∆ (As4O6, g, 298.15 K) = – (1196.2 ± 16.0) kJ·mol–1 and mS ο (As4O6, g, 298.15 K) = (408.6 ± 6.0) J·K–1·mol–1 were obtained by adjustment of the corresponding values reported by Behrens and Rosenblatt [72BEH/ROS], and have been used in the present review.

The reported reaction enthalpies and entropies for Reactions (A.75), (A.76), (A.77) are listed in Table A-21

NiO(cr) + 14

As4O6(g) NiAs(cr) + 54

O2(g) (A.75)

11 NiO(cr) + 2 As4O6(g) Ni11As8(cr) + 232

O2(g) (A.76)

5 NiO(cr) + 12

As4O6(g) Ni5As2(cr) + 4 O2(g) (A.77)

Table A-21: Enthalpies and entropies of reaction from Skeaff et al. [85SKE/MAI].

arsenide r mH ο∆ (kJ·mol–1)

r mS ο∆ (J·K–1·mol–1)

Tr * (K)

NiAs(cr) 465.458 159.057 772

Ni11As8(cr) 4254.395 1564.75 800

Ni5As2(cr) 1544.845 604.452 1038

* approximate mid-temperature for the measurements

Equations for ,mpCο (NiAs) and ,mpCο (Ni11As8) from work of Muldagalieva et al. [93MUL/ISA], [95MUL/CHU] are used with auxiliary data for As4O6(g) (see Chapter IV), O2(g) [98CHA], and values for NiO(cr) from the present review to obtain

f mH ο∆ (NiAs, cr, 298.15 K) = – (69.73 ± 5.00) kJ·mol–1, mS ο (NiAs, cr, 298.15 K) = (50.76 ± 6.00) J·K–1·mol–1, f mH ο∆ (Ni11As8, cr, 298.15 K) = – (743.0 ± 35.0) kJ·mol–1 and mS ο (Ni11As8, cr, 298.15 K) = (518 ± 60) J·K–1·mol–1. In the absence of the original data except as points on a graph, but with the reported uncertainties in the values for

r mG∆ from the equations, the uncertainties here are estimated by the reviewer. For both solids, the heat capacity equations are used slightly beyond the range suggested by the original authors, but the ,mpC (T) curves appear to be monotonic, and there are no better


405

data. For Ni5As2(cr), no heat capacity equation was available, and based on the ap-proximation r ,mpC∆ = 0 used by the original authors, and auxiliary data from the present review, f mH ο∆ (Ni5As2, cr, 298.15 K) = – (244.6 ± 20.0) kJ·mol–1, and mS ο (Ni5As2, cr, 298.15 K) = (196.2 ± 30.0) J·K–1·mol–1 is obtained. The uncertainties have been in-creased to account for the lack of proper heat capacity data.

[86BHA/MUK]

Bhattacharya et al. report the electrochemical generation of the first discrete NiIIIO6 species, 3

2 3Ni(bpyO ) + . The cation is low spin and rhombic. This paper is interesting in the present context, because the redox potential of 3

2 3Ni(bpyO ) + | 22 3Ni(bpyO ) + can be

measured and correlated to the redox potential of Ni3+ | Ni2+.

[86CEM/KLE]

The standard enthalpies of formation of various phases in the system Ni – S were de-termined by means by drop calorimetry. Synthetic NiS was prepared by solid state reac-tion from the elements. X-ray diffraction and microprobe analysis revealed that the high-temperature modification of NiS was obtained. The minerals vaesite (NiS2) and heazlewoodite (Ni3S2) were synthesised by reaction of NiS with sulphur and metallic nickel, respectively. All samples were proved to be well-crystallised material. The en-thalpy of formation for the stoichiometric composition of the hexagonal NiAs-type nickel monosulphide, NiS(α), at 298.15 K was determined by dropping evacuated silica ampoules containing a stoichiometric mixture of powdered metallic nickel and sulphur from room temperature (298.15 K) into the hot calorimeter (1021 K) where the forma-tion of NiS occurred. In a second experiment the silica ampoules containing NiS formed during the first experiment were dropped from room temperature into the hot calorime-ter. Obviously, the difference of the enthalpy effects corresponding to the dropping ex-periments directly yields the enthalpy of formation of the high-temperature form of NiS at 298.15 K, f mH ο∆ = – (88.1 ± 1.0) kJ·mol–1.The enthalpy of formation of the low-temperature modification, NiS(β), cannot be obtained directly because in the high tem-perature calorimeter NiS(α) always is formed, and this solid does not transform into NiS(β) upon quenching.

The enthalpy of formation of heazlewoodite was obtained by dropping evacu-ated silica capsules containing a mixture of powdered metallic nickel and NiS from room temperature into the calorimeter at 873 K. Again, in a second experiment the sil-ica capsules containing the product of the first reaction (N3S2) were dropped into the calorimeter. Taking into account the enthalpy of formation of NiS determined earlier, the difference of the enthalpy changes associated with these experiments leads to the enthalpy of formation of heazlewoodite at 298.15 K, f mH ο∆ = – (217.24 ± 1.60) kJ·mol–1.

The enthalpy of formation of vaesite (NiS2) at 298.15 K was derived from the reaction:


406

NiS2 + Ni 2NiS(α)

at 873 K in the calorimeter. In a second experiment the heat content of the silica am-poule and NiS formed during the reaction was measured. The difference between the enthalpy changes associated with these experiments results in the enthalpy of formation of vaesite at 298.15 K, f mH ο∆ = – (124.9 ± 1.0) kJ·mol–1.

The enthalpy of formation for Ni7S6 was likewise determined by reaction of NiS with Ni at 833 K in the calorimeter. At these conditions the non-stoichiometric high-temperature modification of Ni7S6 is formed. However, when this product is quenched, metastable modifications are obtained depending on the cooling rate [96SEI/FJE]. Therefore, the enthalpy of formation reported by the authors instead should be assigned to formation of metastable phases of Ni7S6 or to a mixture of Ni3S2 and Ni9S8.

All experimental results of this interesting paper are listed in Table A-22.

Table A-22: Enthalpies of formation for solid nickel sulphides at 298.15 K.

Solid phase f mH ο∆ (kJ·mol–1)

NiS(α), high-temperature form – (88.1 ± 1.0)

Ni3S2 – (217.2 ± 1.6)

NiS2 – (124.9 ± 1.0)

Ni7S6, maybe ⅓Ni3S2 + ⅔Ni9S8 – (582.8 ± 5.7)

[86HOL/NEI]

The standard Gibbs energy of formation of NiO has been experimentally determined over the temperature range from 900 to 1400 K using a galvanic cell with the solid elec-trolyte made of 15% calcia-stabilised zirconia. The measured value of f mGο∆ (NiO) at 1300 K was – 123.555 kJ·mol–1 with a precision of ± 0.057 kJ·mol–1 and an estimated accuracy not worse than 0.200 kJ·mol–1. This precision is equivalent to an error of only 0.2 – 0.3 mV in the cell potential (emf). In comparison, most previous studies have re-ported a precision of 1 – 2 mV. Using the third law analysis, the authors obtained for the enthalpy of formation of NiO, f mH ο∆ (NiO, 298.15 K) = – 240.110 kJ·mol–1.

[86HSI/CHA]

Equilibria between two solid phases in the ternary system Ni – S – O at 2SO

p = 1 atm were studied by using the electrochemical cell Pt, O2(air) | YSZ or CSZ | mixture of phases, SO2 (1 atm), Au, where yttria- or calcia-stabilised zirconia served as the solid electrolyte. The following phase equilibria were studied: NiS(α) / β2-Ni4S3; β2-Ni4S3 / NiO; NiS(α) / NiO; and NiO / NiSO4. While the solid sulphides and NiO were prepared by solid state reaction from the elements, NiSO4 was synthesised by dehydrating NiSO4·6H2O. All solid phases were identified by X-ray diffraction analysis. The emf of


407

the electrochemical cell, recorded as a function of temperature, corresponds directly to the partial pressure of oxygen at

2SOp = 1 atm for the phase equilibria mentioned above.

Figure A-33 shows the temperature dependence of the oxygen partial pressure for NiS(α) / β2-Ni4S3 as well as β2-Ni4S3 / NiO, where the experimental results of [81SCH], [85MEH/TAR] are likewise included. It can be deduced from Figure A-33 that the thermodynamic models of both [80SHA/CHA] and the present data assessment are in close agreement with the experimental results for the phase equilibrium NiS(α) / β2-Ni4S3. The experimental data from various literature sources [81SCH], [85MEH/TAR], [86HSI/CHA] for the equilibrium β2-Ni4S3 / NiO coincide remarkably well with each other despite the fact that Mehrotra et al. [85MEH/TAR] and Schaefer [81SCH] gave constant but different values for the composition of the non-stoichiometric phase β2-Ni4S3. However, it is to be expected that the composition of β2-Ni4S3, which is fixed by the phase equilibrium with NiO at

2SOp = 1 atm, should change

with temperature. Moreover, it is worth mentioning that the temperature dependence of the sulphur content in β2-Ni4S3 in equilibrium with NiO has not yet been investigated experimentally.

In addition, the variation of the oxygen partial pressure with temperature for the equilibrium NiS(α) / NiO at

2SOp = 1 atm is depicted in Figure A-34. The experi-

mental data agree reasonably with equilibrium calculations using the present thermody-namic model.

Figure A-33: Plot of logarithm of the oxygen partial pressure versus temperature for the phase equilibria α-NiS / β2-Ni4S3 and β2-Ni4S3 / NiO, respectively, at

2SOp = 1 atm. ●, ■

[86HSI/CHA]; [81SCH]; dotted line: [85MEH/TAR]. Dashed line: thermodynamic model according to [80SHA/CHA]; solid line: thermodynamic model of the present assessment.

920 940 960 980 1000 1020 1040 1060 1080 1100

-14.0

-13.5

-13.0

-12.5

-12.0

-11.5

-11.0

NiS(α)

β2-Ni4S3

NiO

log 10

[p(O

2) / b

ar]

T / K


408

Figure A-34: Logarithm of the oxygen partial pressure plotted as a function of tempera-ture for the equilibrium α-NiS / NiO at p(SO2) = 1 atm. ● [86HSI/CHA]. Solid line: thermodynamic model of the present assessment.

[86JAC/KAL]

The emf-study of Ni2SiO4 done by Jacob et al. was not used for the evaluation of ther-modynamic data by this review, because the results were presented only in graphical form, making accurate recovery of experimental values difficult.

[86VAS/DMI]

Vasil’ev et al. reported another value of f mH ο∆ (Ni2+) = – (51.83 ± 0.88) kJ·mol–1 deter-mined calorimetrically by i) dissolving Ni (cr) in solutions containing 0.3 M Br2, 0.05 M HClO4, 0.95 to 3.95 M NaBr, ii) accounting for the enthalpy of mixing NiBr2 as well as iii) Br2 to make up the final solutions, and iv) extrapolating to I = 0. The re-calculation with the CODATA value for f mH ο∆ (Br–) = – (121.41 ± 0.15) kJ·mol–1 resulted in a minute correction only, but the error limits appear too optimistic in view of four differ-ent contributions to the final result, and thus were set to 2σ = ± 2.0 kJ·mol–1. Thus,

f mH ο∆ (Ni2+) = – (51.85 ± 2.00) kJ·mol–1 has been assigned to this experimental infor-mation.

870 880 890 900 910 920 930 940

-15.0

-14.8

-14.6

-14.4

-14.2

-14.0

-13.8

NiS(α)

NiO

T / K

log 10

[p(O

2) / b

ar]


409

[87BEC]

The formation constants of the proton and metal ion complexes of cyanide ion have been critically surveyed. The recommended value for HCNpKο is correct, but the tentative values for higher ionic strengths seem to be erroneous [92BAN/BLI]. Several literature values for the formation constants of 2

4Ni(CN) − and 24Ni(CN) − as well as reaction en-

thalpies for the formation of 24Ni(CN) − were reported, but no recommended values

were given. In addition, the formation constants of several NiX(CN)x complexes (X = multidentate organic ligands, x = 1 or 2) were listed.

[87BJE]

These “semi-thermodynamic” formation constants are not equivalent to the value de-termined by extrapolation to zero ionic strength using the SIT, except for the case when the ratio:

( ) ( )2 1 26 7[NiX H O ] [NiX H O ]/n nn n−− −

γ γ = 1

in the whole range of the ionic strength studied.

[87CEM/KLE]

The standard enthalpy of formation of synthetic pentlandite, Fe4.5Ni4.5S8, and natural violarite, (Fe0.2941Ni0.7059)3S4, were measured by solution drop calorimetry. Synthetic pentlandite was prepared from the elements and checked by X-ray diffraction analysis. First a stoichiometric mixture of FeS, NiS, Ni, and Fe was dropped from 298.15 K into a melt of Ni0.6S0.4 at 1100 K. In a second experiment synthetic pentlandite was dropped into the Ni0.6S0.4 melt. The enthalpy of formation of pentlandite from its elementary and binary compounds is obtained from the difference of the enthalpy effects associated with these two experiments, r mH ο∆ = – (78.41 ± 14.45) kJ·mol–1 for:

4 NiS + 4 FeS + 0.5 Ni + 0.5 Fe Ni4.5Fe4.5S8.

Taking f mH ο∆ = – (101.67 ± 0.20) kJ·mol–1 for the standard enthalpy of forma-tion of FeS from the JANAF tables [98CHA] and f mH ο∆ = − (88.1 ± 1.0) kJ·mol–1 for the high-temperature modification of NiS [86CEM/KLE], the standard enthalpy of for-mation of pentlandite at 298.15 K is calculated to be f mH ο∆ = − (837 ± 15) kJ·mol–1. By using a similar procedure, the authors arrive at f mH ο∆ = − (378 ± 12) kJ·mol–1 in the case of violarite with the composition (Fe0.2941Ni0.7059)3S4.

[87EMA/FAR]

Emara et al. gave an example of how they made up their solutions. For a solution in-tended to contain 2.4×10–4 M 3HCO− they started with 6.18×10–4 M NaHCO3, but in reality this leads to 6.05×10–4 M 3HCO− . Thus the results of this paper are entirely unre-liable.


410

[87FLE]

On a specimen which also originated from the Cu-Ni sulphide deposit Noril'sk, Siberia, Russia, Fleet carried out a definitive crystal structure analysis of godlevskite. Discrep-ancies between the unit-cell parameters of [69KUL/ERS] and [87FLE] are attributable to incorrectly indexed powder data.

Discrepancies between the chemical composition used by [69KUL/ERS] and [87FLE] are attributable to advances in electron microprobe technology and calibration procedures applied in the latter study. Although there are similarities in the unit-cells, powder patterns, and S-atom arrays of godlevskite and pentlandite, godlevskite does not have a pentlandite derivative crystal structure.

[87GAR/PAR]

The CODATA Task Group on Chemical Thermodynamic Tables presented recom-mended values for thermodynamic properties of selected compounds of calcium and their mixtures in this prototype set of tables.

[87NEI]

This is an excellent study of NiO, CoO, Ni2SiO4 and Co2SiO4 using an electrochemical technique. Cells were used with calcia-stabilised zirconia as an oxygen specific electro-lyte. The fact that not only the overall cell reaction:

2 Ni + SiO2 + O2 Ni2SiO4

but also the equilibria of Ni + NiO and Fe + FeO at the reference electrode were inves-tigated in detail, make the obtained results very reliable. Moreover, some information is also given for the two equilibria involving SiO2 both quartz and cristobalite. The Gibbs energy of the above cell reaction measured between 972 and 1326 K was compared by O’Neill with results of previous work and fairly good agreement has been found be-tween all emf results. However, the discrepancy between the emf and calorimetric data of [84ROB/HEM] remains unresolved. In order to approach this problem, O’Neill made additional experiments to confirm the decomposition temperature of Ni2SiO4 to the ox-ides and found that Ni2SiO4 breaks down to NiO and SiO2-cristobalite at (1820 ± 5) K. Taking the Gibbs energy of the cell reaction and the proper decomposition temperature of Ni2SiO4, a consistent value of the standard molar enthalpy of the reaction under con-sideration has been ascertained.

[88BJE]

The formation constant of the NiCl+ complex has been determined by a spectrophotometric study at 298.15 K in concentrated HCl (4.4 – 10.67 M) and LiCl (2.94 – 12.46 M) solutions, and the formation constant of the NiCl2 complex was tentatively estimated. The method used by the author results in a “semi-thermodynamic” constant, in which the activity of chloride ion is considered, but the ratios 2+ +Ni NiCl

/γ γ and + 2NiClNiCl/γ γ were assumed to be nearly constant with


411

+ 2NiClNiCl/γ γ were assumed to be nearly constant with increasing MCl concentration,

and were included in the reported formation constants (see discussion on [89BJE]). The method used by the author is more appropriate to the study of weak complex formation than the constant ionic medium principle if very high and varying concentrations of a complex-forming anion is used. (see discussion on [89BJE]). Nevertheless, the forma-tion constants determined in this way are not compatible with the 10log οb values ex-tracted from the SIT analysis

[88BJE2]

The “semi-thermodynamic” stability constant (see [89BJE]) for the reaction:

3Ni(SCN)− + SCN– 24Ni(SCN) −

was determined by a spectrophotometric method at 298.15 K in concentrated aqueous NaSCN (2.14 – 10.5 M) solutions, using the main activity coefficients of NaSCN solu-tions reported in [56MIL/SHE]. The reported 10 4log K value is − (0.92 ± 0.18). Al-though, the method used by the author is probably more appropriate to the study of weak complex formation than the constant ionic medium principle, if high and varying complex-forming anion concentrations are needed (see the discussion of [89BJE]), the constants determined in this way are not compatible with the 10log οb values extracted from the SIT analysis. Therefore, the reported constant was not considered further in this review.

[88EFI/EVD]

Efimov et al. devised an ingenious thermodynamic cycle to measure the enthalpy of Reaction (A.78) at 298.15 K by solution calorimetry:

Br2(l) + Ni(cr) + 2 KCl(cr) NiCl2(cr) + 2 KBr(cr). (A.78)

Five calorimetric experiments were designed in such a way that the aqueous species of the solutions used cancelled when the individual reactions were combined according to Equation (A.78) and r mH ο∆ (A.78) = – (218.38 ± 0.36) kJ·mol–1. Then,

f mH ο∆ (NiCl2, cr) = r mH ο∆ (A.78) – 2 ( f mH ο∆ (KBr, cr) – f mH ο∆ (KCl, cr) )

= – (303.71 ± 1.10)

can be recalculated using the selected standard enthalpies of formation for crystalline KCl and KBr (cf. Table IV-1 and Chapter VI). The assigned uncertainty limits corre-spond to twice the standard deviation given by the authors. The value is consistent with, but less precise than the value determined by Lavut et al., [84LAV/TIM] and the value derived in this review from Busey and Giauque's equilibrium studies [53BUS/GIA].

[88LIC2]

Based on the Gibbs energy of formation for I, II, and III-NiS, derived from the experi-mental investigation of the solubility of nickel sulphide performed by Thiel and Gessner


412

[14THI/GES], and a new value for the second dissociation constant of H2S, the solubil-ity products for I, II, and III-NiS were recalculated. The author arrived at the following values for the pertinent solubility products:

10 ,0log sK ο (I-NiS) = – 24.3, 10 ,0log sK ο (II-NiS) = – 29.5, 10 ,0log sK ο (III-NiS) = – 31.3.

[88NIS/ITO]

Nishimura et al. measured the solubility for Ni3(AsO4)2·8H2O and, based on measure-ments at 25°C over a range of pH from 4 to 8, reported 10log K = 9.9 for

Ni3(AsO4)2·8H2O + 4H+ 3Ni2+ +2 2 4H AsO− + 8H2O.

Langmuir et al. [99LAN/MAH], using this equilibrium constant, but different auxiliary data, calculated 10 ,0log sK = – 27.02 for

Ni3(AsO4)2·8H2O 3 Ni2+ + 2 34AsO − + 8 H2O (A.79)

and 10 ,0log sK (A.79) = – 28.38 when corrected for associative complex formation. In the present review, the correction to 10 ,0log sK for complexation is accepted. Then, if the auxiliary data in Table IV-2 are used, 10 ,0log sK (A.79) = – 28.12 is calculated. The authors did not propose uncertainty limits but, based on the figures in the original paper, an uncertainty of ± 0.50 in the equilibrium constant seems reasonable and is accepted here.

The authors also reported a value of 10log K = 1.8 for the solubility product for NiHAsO4·2H2O from measurements done from pH 2 to 3.

NiHAsO4·2H2O + H+ Ni2+ + 2 4H AsO− + 2 H2O.

This value is accepted in the present review, again with an uncertainty of ± 0.5.

[89BJE]

The limitations inherent in the constant ionic medium principle have been discussed with respect to the determination of weak stability constants for metal-anion complexes. The author suggests that the use of the constant medium principle to study weak com-plexes, for which the background electrolyte should be almost entirely replaced by the complex-forming electrolyte to enhance the complex formation, introduces an inherent and electrolyte-dependent uncertainty, due to the change in activity coefficients. Ac-cording to Harned's rule, the logarithms of the activity coefficients in a mixture of two electrolytes vary linearly with their mole fraction. This variation, which depends mainly on the hydration of the ions involved, results in an uncertainty that can be very large relative to the stability constants of weak complexes. For this reason, the author pro-posed a different method to study weak complex formation. In a series of papers (see e.g., [88BJE], [90BJE]) he published spectrophotometric studies of such systems, using relatively small concentrations of the metal ion in solution with up to the highest possi-


413

ble concentrations of highly soluble (pseudo)halogenides such as HCl, LiCl, LiBr or NaSCN. His data are based on the following three approximations:

(i) the logarithm of the main activity coefficient of the complex-forming electrolyte (NX) increases linearly with the concentration above 3 M ( NX

10log γ ± = A + B[X]),

(ii) the ratio of the activity coefficients of MXn and MXn–1 does not change with in-creasing salt concentration (MX) and thus the following “semi-thermodynamic” mass action expression can be used '

nοb = [MXn]/[MXn–1] × aX (where aX = [X]

× NXγ ± )

(iii) the absorption spectrum of a given complex does not change with respect to changes in medium.

The latest isopiestic studies confirm the validity of assumption (i), while as-sumption (iii) is certainly more valid for d-d transitions than for charge-transfer bands.

In this respect, it is instructive to follow the scientific debate between J. Bjer-rum and R.W. Ramette concerning the correctness of stability constants of copper(II)-chloride complexes [83RAM/FAN], [86RAM], [86BJE/SKI], [87BJE]. Bjerrum stated that maintaining a high constant ionic strength, but changing the background electrolyte, results in unacceptable uncertainty for K3 and K4. On the other hand, Ramette claimed that spectrophotometric data determined by Bjerrum et al. are also subject to a medium effect.

These “semi-thermodynamic” formation constants (at I = 0) are not equivalent to the values obtained by extrapolation to zero ionic strength using the SIT, except for the case when the ratio:

2 6

1 2 7

[NiX (H O) ]

[NiX (H O) ] = 1n n

n n

γ

γ−

− −

in the whole range of the ionic strength studied.

[89BRA]

Standard electrode potentials of the elements and their temperature coefficients in water at 298.15 K were listed for nearly 1700 half-reactions at pH = 0.000 and pH =13.996. General and specific methods of estimation of thermodynamic quantities were de-scribed, but the original papers, from which the individual entries were derived, have not been cited.

[89EVD/EFI]

The enthalpy of formation of NiBr2(cr) was determined from the heats of solution of Ni(cr) and NiBr2(cr) in KBr-HBr(aq)-Br2(aq) mixtures (1.00KBr, 0.40Br2, 0.12HBr, 50.87H2O - probably molar ratios). The sample preparation appears to have been done carefully, but the resultant solid probably had the “Wechselstruktur” [34KET].


414

With solution 1 defined as (1.00KBr, 0.40Br2, 0.12HBr,50.87H2O), the reac-tions were:

Br2(l) + soln 1 Br2(soln 2) (A.80)

Ni(cr) + Br2(soln 2) NiBr2 (soln 3) (A.81)

NiBr2(soln 1) NiBr2 (soln 3) (A.82)

Table 1 in the translation contains some obvious typographical errors. For reac-tion (A.81), the last four reported heats are a factor of 10 too low. For reaction (A.82), the reported heats are too low by a factor of 103. The reported uncertainties seem to be average uncertainties, close to 1σ uncertainties. The reported values have been re-averaged, and the values are – (8.239 ± 0.216) kJ·mol–1, – (289.442 ± 0.712) kJ·mol–1, and – (85.285 ± 0.311) kJ·mol–1 for the authors’ reactions 1, 2 and 3 respectively. The correction of (0.462 ± 0.001) J·K–1·mol–1 applied by the authors to allow for the pres-ence of some bromine in the gaseous state in the measurements of the heat of reaction (A.80) is accepted. From these values, f mH ο∆ (NiBr2, cr) = − (211.93 ± 0.90) kJ·mol–1, 2σ uncertainties.

[89IUL/POR]

Chloride complexation with nickel(II) and copper(II) ions has been studied at 298.15 K by solubility determinations of M(IO3)2 in aqueous solutions of 6 M Na(ClO4,Cl). The chloride concentration varied from 0 to 6 M. Medium changes, caused by the replace-ment of 4ClO− by Cl– ions, were accounted for by the SIT. The solubility product of Ni(IO3)2 is reported to be (5.1 ± 0.2) × 10–5, under the conditions used. The formation of two complexes, NiCl+ and NiCl2, was detected. The authors also reported equilibrium constants for infinite dilution by means of the SIT by using estimated interaction coeffi-cients, ε(NiCl+, 4ClO− ) = 0.202 and ε(NiCl2(aq), NaClO4) = – 0.10. In the present re-view a considerably different value for ε(NiCl+, 4ClO− ) has been derived ((0.47 ± 0.06) kg·mol−1); thus we calculated a markedly higher value of 10 1log οb from the data [89IUL/POR] than the authors themselves. The reported value of K2 is higher than K1, which is rather unlikely considering the low affinity of nickel(II) for chloride. There-fore, only the value of K1 is considered in this review.

[89ZIE/JON]

Ziemniak, Jones and Combs reported solubility measurements of NiO(cr) in aqueous sodium phosphate solutions (pH values between 10.2 and 12.2) for temperatures from 290 K to 560 K. The experiments appear to have been done carefully, and it might be expected that the measured solubilities should be suitable for derivation of chemical thermodynamic parameters ( r mGο∆ , r mS ο∆ and r ,mpCο∆ ) for the dissolution reactions. The authors analysed the data in terms of formation of nickel(II) hydroxo and mixed phosphato-hydroxo complexes. Only mononuclear species of nickel were considered due to the low nickel(II) concentrations.


415

The authors concluded that a layer of the more stable hydrous nickel oxide (Ni(OH)2) existed on the NiO surface below 468 K, and only at higher temperatures did the NiO phase control the solubility behaviour of nickel(II). The data analyses provided by the authors contained a number of inconsistencies [95LEM]. In the present evalua-tion, because the solid phase was poorly defined in the study of Ziemniak et al., any value used must be assigned a rather large uncertainty. The values 10 2,1

*log οb = − (19.77 ± 1.00) and 10 3,1

*log οb = – (30.86 ± 1.00) can be derived based on these data and the analyses in [95LEM]. However, these values are not accepted in the present review.

Ziemniak et al. proposed two aqueous nickel phosphate species, 2

3 2 4Ni(OH) H PO − and 22 4Ni(OH) HPO − . These differ only by a (dissociated) water mole-

cule in the inner sphere. The concentration of 23 2 4Ni(OH) H PO − was reported to be

small compared to the concentration of 22 4Ni(OH) HPO − for all solutions studied at

298 K, and the reverse was reported for solutions at 560 K. If the Gibbs energies of re-action proposed by the authors are used, the concentrations of these nickel-phosphate species are approximately equal at 450 K. Even if both species exist, quantitative spec-troscopic data for the relative concentrations of 2

3 2 4Ni(OH) H PO − and 22 4Ni(OH) HPO −

in a solution would be required to determine formation constants for the two species. Lacking such information, the solubility data cannot be used to calculate independent concentrations for the two species. From the point of view of the available variables (total nickel concentration, phosphate concentration, sodium to phosphate ratio and pH) the species 2

3 2 4Ni(OH) H PO − and 22 4Ni(OH) HPO − are mathematically redundant.

A re-analysis of the data was done by Lemire [95LEM]. He concluded that the total equilibrium solution concentrations of nickel species over NiO(cr) (or β-Ni(OH)2 at low temperature) are higher in solutions containing substantial quantities of phos-phate than they are in dilute aqueous solutions, and that this can be accounted for by assuming formation of complexes of the form 2

2 4Ni(OH) HPO − , as proposed by Ziem-niak et al.. However, in their experiments, the ionic strength increased (from 0.002 m to 0.6 m) in proportion to the added phosphate concentration and, from the available data, it is not possible to separate the effects of the ionic medium and complexation. Ionic strength effects seem to be less important than experimental uncertainties and model errors. Less hydrolysed phosphato complexes (e.g., NiHPO4(aq)) do not appear to be important under the conditions of the experiments. The re-analysis suggested that for models that assume the stable solid at 298.15 K is β-Ni(OH)2, and that the solution spe-cies include only hydrolysis species and one nickel phosphate complex,

22 4Ni(OH) HPO − , formed according to

β-Ni(OH)2(cr) + 24HPO − 2

2 4Ni(OH) HPO − , (A.83)

the Gibbs energy of complex formation is approximately (27 ± 4) kJ·mol–1, and 10log K (A.83) = – (5 ± 1).


416

[90BJE]

This is very similar experimental work to that of [88BJE]. In this paper the formation of bromo complexes of nickel(II) was studied in 3.5 – 11.1 M LiBr solution. The “semi-thermodynamic” formation constant of NiBr+ was determined, and that of the complex NiBr2(aq) was estimated (see Table V-13).

[90BJE2]

The author demonstrated that from an aqueous solution that is dilute in nickel(II) ions, but concentrated in NaSCN, the tetrahedral 2

4Ni(SCN) − complex can be extracted with hexanone. This shows that the octahedral [Ni(SCN)4(H2O)]2– complex, which forms in presence of thiocyanate ion in very high excess, must be in equilibrium with small con-centrations of the tetrahedral species.

[90EFI/EVD]

Efimov and Evdokimova determined the enthalpy of formation of NiI2(cr) at 298.15 K by two related thermodynamic cycles involving dissolution of a series of compounds in a mixture with the composition 1.00 KBr·0.40 Br2·0.12HBr·50.87H2O. In one cycle the enthalpy of Reaction (A.84) was calculated in a manner similar to the enthalpy from the cycle devised in [88EFI/EVD]:

Br2(l) + Ni(cr) + 2 KI(cr) NiI2(cr) + 2 KBr(cr). (A.84)

The authors made a correction of (0.462 ± 0.001) kJ·mol–1 to allow for the amount of Br2 in the calorimeter sample bulb that was Br2(g) rather than Br2(l).

The other cycle was more direct:

I2(cr) + Ni(cr) NiI2(cr). (A.85)

Five calorimetric experiments were necessary to give r mH ο∆ (A.84) = − (224.76 ± 0.38) kJ·mol–1, but only three values were required to determine

r mH ο∆ (A.85) = – (96.42 ± 0.84) kJ·mol–1 (2σ uncertainties). The value given as the en-thalpy of solution of KBr(cr) in the solvent mixture was (18.123 ± 0.039) kJ·mol–1, whereas, in the earlier paper [88EFI/EVD] the value is reported as (18.145 ± 0.027) kJ·mol–1. The value in [90EFI/EVD] is presumed here to be a deliberate (but minor) correction, and r mH ο∆ (A.84) = – (224.76 ± 0.44) kJ·mol–1 is calculated,

f mH ο∆ (NiI2, cr) = r mH ο∆ (A.85) = – (96.42 ± 0.84) kJ·mol–1 (A.86)

f mH ο∆ (NiI2, cr) = r mH ο∆ (A.84) – 2 ( f mH ο∆ (KBr, cr) – f mH ο∆ (KI, cr)) = – (96.40 ± 1.00) kJ·mol–1. (A.87)

The two values are virtually identical. The calculations from the two cycles both use exactly the same solution enthalpies for Ni(cr) and NiI2(cr), and the bulk of the


417

uncertainty in either cycle can be attributed to the measurements of those quantities. Differences in the f mH ο∆ (NiI2, cr) values result only from the values for the enthalpies of formation of KBr(cr) and KI(cr) and the measured heats of solution of the elemental halides and the potassium salts. Therefore, in this review, only the more precise of the two values is accepted for further use:

f mH ο∆ (NiI2, cr, 298.15 K) = – (96.42 ± 0.84) kJ·mol–1 (A.88)

This is virtually identical to the value given by the authors [90EFI/EVD] (if it is assumed that their error limits correspond to one standard deviation).

[90EFI/FUR]

Efimov and Furkalyuk determined the enthalpy of Reactions:

NiCl2(cr) Ni2+ + 2Cl–, (A.89)

NiBr2(cr) Ni2+ + 2Br–, (A.90)

NiI2(cr) Ni2+ + 2I– (A.91)

calorimetrically by dissolving crystallised NiCl2, NiBr2 and NiI2 in 0.001 M HClO4 at 298.15 K. For the experiments with NiCl2 they used samples prepared as described in [88EFI/EVD]. According to the analytical composition given there (w(Ni) = (0.4529 ± 0.14) , w(Cl) = (0.5465 ± 0.10)), these samples contained less moisture than those used in [58MUL] and [95MAN/KOR]. When Efimov and Furkalyuk’s experimen-tal data were extrapolated to I = 0, like those listed in [58MUL] and [95MAN/KOR] shown in Table A-23 and Table A-25, their results had to be corrected slightly. They used perchloric acid but neglected its contribution to the ionic strength. The following values were obtained: sol mH ο∆ (A.89) = – (84.89 ± 0.27) kJ·mol–1, which is about 1.5 kJ·mol–1 more negative than those found in [58MUL] and [95MAN/KOR],

sol mH ο∆ (A.90) = – (86.18 ± 0.38) kJ·mol–1 and sol mH ο∆ (A.91) = – (73.38 ± 0.12) kJ·mol−1.

Table A-23: Enthalpy of solution of NiCl2 in H2O at 298.15 K [90EFI/FUR].

m (NiCl2) (kg·mol–1)


ΦL (kJ·mol–1)


0.005495 – 84.169 0.6864 – 84.8554 0.005213 – 84.366 0.6718 – 85.0378 0.004584 – 84.233 0.6372 – 84.8702 0.004422 – 84.057 0.6278 – 84.6848 0.004410 – 84.354 0.6271 – 84.9811


418

[90GRZ]

The stability of the monohalide complexes of the Mn2+–Zn2+ series was surveyed. The order of constants both in water and organic solvents is distinctly different from the stability order suggested by the Irving-Williams series. This behaviour seems to be characteristic of complexes for which the stabilities are entropy controlled. The order of stability constants has been correlated with the water-exchange rate and electrostriction partial molar volumes of the hydrated cations.

[90JUR/DOM]

Using an adiabatic calorimeter, Juraitis et al. [90JUR/DOM] measured the heat capacity of the solid between 80 K and 281 K, with emphasis on the effects of the structural tran-sition near 220 K. The heat capacity values were recovered from the authors’ graph, and polynomials were fitted to the ,m /pC T values to determine the entropy increments over the temperature ranges 100 K to 200 K, 200 K to 220 K, 220 K to 240 K, and 240 K to 281 K. The results between 100 K and 200 K show considerable scatter, and it is not clear from these results (only reported graphically) whether the values from 250 K to 280 K can be extrapolated smoothly to 298.15 K The experimental values from 250 K to 273 K increase monotonically from 174 kJ·mol–1 to 209 J·K–1·mol–1, but the reported value for 281 K is 207 J·K–1·mol–1. The highest-temperature calorimetry measurement at 281 K, closest to room temperature, is likely to have the greatest uncertainty. From the simple quadratic equation fitted to the unweighted ,m /pC T values for temperatures from 240 K to 281 K, ,mpCο (NiCl2·2H2O, cr, 298.15 K) = 222.3 J·K–1·mol–1. In the pre-sent review, it is assumed that the ,m ( )pC T function has no anomalous behaviour above 240 K, and that differences from a monotonic curve are the result of experimental scat-ter. Based on that premise, ,mpCο (NiCl2·2H2O, 298.15 K) = (230 ± 15) J·K–1·mol–1 is selected in the present review, where the uncertainty is an estimate. These uncertainties in the ,mpCο values above 240 K introduce an uncertainty of ± 3 J·K–1·mol–1 in the con-tribution to the calculated value of mS ο (298.15 K).

[90LEE/NAS]

The assessed Ni–Te phase diagram of the Handbook “Binary Alloy Phase Diagrams” is essentially the same phase diagram as that published by Klepp and Komarek [72KLE/KOM].

[90OSA/ROS]

The phase equilibria between β1-Ni3S2, β2-Ni4S3, and NiS(α) were studied by measuring the equilibrium partial pressure of oxygen as a function of temperature.

The following galvanic cells were employed:

Pt, Au, NiS(α), β2-Ni4S3, NiO | CSZ | Fe, FeO, Pt

Pt, Au, β2-Ni4S3, β1-Ni3S2, NiO | CSZ | Fe, FeO, Pt


419

Pt, Au, SO2 (0.010 atm), NiS(α), β2-Ni4S3 | CSZ | O2 (0.21 atm), Pt

where calcia stabilised zirconia (CSZ) served as the solid electrolyte. Figure A-35 shows the temperature dependence of the partial pressure of oxygen derived from these measurements. The calculated lines, using the present thermodynamic model, agree remarkably well with the experimental data which provides further evidence for the reliability of the thermodynamic data recommended in this assessment. Furthermore, it should be mentioned that the solid sulphides were prepared by solid state reaction from the elements and NiO was obtained by thermal decomposition of NiCO3. All solid phases were identified by X-ray powder diffraction analysis.

Figure A-35: Plot of logarithm of the oxygen partial pressure versus temperature for the phase equilibria NiS(α) / β2-Ni4S3 and β2-Ni4S3 / NiO, respectively, at

2SOp = 1 atm. ▲

equilibrium between O2(g), NiO, NiS(α) and β2-Ni4S3; ● equilibrium between O2(g), NiO, β2-Ni4S3 and β1-Ni3S2. Dashed line: thermodynamic model according to [80SHA/CHA]; solid lines: thermodynamic model of the present assessment.

[91KNA/KUB]

This is the second updated and computerised edition of [73BAR/KNA], [77BAR/KNA].

800 850 900 950 1000 1050-20

-19

-18

-17

-16

-15

-14

-13

-12

-11

lo

g 10 [p

(O2) /

bar

]

T / K


420

[91SPI/TRI]

The hydrolysis of aqueous nickel(II) chloride solutions ([NiCl2]tot = 0.03 to 1.5 M) has been studied by potentiometric titrations at 25°C. The ionic strength was not controlled; instead, the activity coefficients, based on literature values, were taken into account. The formation of a single hydroxo complex was assumed. The evaluation of experimen-tal data indicated the presence of the 4

4 4Ni (OH) + tetrameric species. A value of 10 4,4

*log οb = − (27.0 ± 0.1) was reported for its thermodynamic hydrolysis constant. The formation of chloro complexes was not taken into account, although these may only have made a small contribution to the total solution concentration of nickel at [NiCl2]tot ≤ 0.1 M. Therefore, the reported value is accepted for use in this review, but with an increased uncertainty: 10 4,4

*log οb = – (27.0 ± 0.3).

[91STO/GRO]

The heat capacity function of Ni3S2 was measured by adiabatic-shield calorimetry in the temperature range from 5 to 350 K. High temperature measurements were carried out up to a temperature of 1000 K. The solid phase heazlewoodite, Ni3S2, was synthesised directly from the elements and characterised by X-ray diffraction analysis. The integra-tion of the low temperature results yields the standard entropy of Ni3S2 at 298.15 K,

mS ο = 133.2 J·K–1·mol–1. The high-temperature data are depicted in Figure A-36, where the results of Ferrante and Gokcen [82FER/GOK] are likewise included. The heat ca-pacity function:

835K,m 298.15K[ ]pCο (Ni3S2, cr) = (165.79352 – 0.10316 (T/K) + 0.00012(T/K)2

– 2444624(T/K)–2) J·K–1·mol–1 (A.92)

valid from 298.15 to 835 K, was fitted to both sets of data mentioned above.

The steep increase of the ,mpC values of [91STO/GRO] at approximately 800 K can be attributed to the onset of the phase transition into the high-temperature modification of Ni3S2 which was observed to occur at 834 K.


421

Figure A-36: Heat capacity data for Ni3S2 as a function of temperature. ○ [91STO/GRO], ▲ [82FER/GOK]. Solid line: optimised heat capacity function (A.92).

[91SWI/PAW]

This polarographic study reports 10log K = 3.26, 25°C, 0.2 M NaClO4 for an ill-defined Ni2+ complex with “phosphate” at pH 6.0. The concentration of nickel ion in solution was 5 × 10–5 M and the ligand concentration was varied between 5 × 10–5 and 2.5 × 10–3 M. Figure 3 of the paper shows the effect of changes in “ 3

4PO − ” on the polarograms (in 0.1 M NaClO4(aq)), but the paper does not state the chemical form of the added phos-phate nor provide information as to the stoichiometry of the complex. The information provided is not sufficient to recalculate a meaningful complexation constant, and the results of the study are not used in the present review.

[91TAR/FAZ]

Tareen et al. determined T, p data for the decomposition equilibrium of NiCO3 to NiO,

NiCO3(s) NiO(s) + CO2(g), (A.93)

in the pressure range 300 – 2000 bar and at temperatures up to 700°C. The results of this investigation relevant for the calculation of f mH ο∆ (NiCO3, 298.15 K) are listed in Table A-24. The re-evaluation was based on the following set of equations:

200 300 400 500 600 700 800 900100

120

140

160

180

200

T / K

C° p,m

/ J·K

–1·m

ol–1


422

f mH ο∆ (NiCO3, T) = f mH ο∆ (NiO, T) + f mH ο∆ (CO2, T) – T× r mS ο∆ ((A.93), T)

+ (Vm(NiO, 298.15 K) – Vm(NiCO3, 298.15 K))×(p – 1)

+ RT ln f(CO2) + RT ln y(CO2) (A.94)

RT ln f(CO2) = RT ln p + B(CO2)p (A.95)

B(CO2) (cm3·mol–1) = – 127 – 288×298.15 1

T −

– 118×2298.15 1

T −

(A.96)

f mH ο∆ (NiCO3, 298.15 K) = f mH ο∆ (NiCO3, T) – f ,m298.15

T

pC∆∫ (NiCO3) dT (A.97)

At first Equations (A.94), (A.95) and (A.96) have been used to calculate f mH ο∆ (NiCO3, T), where Equation (A.96) was taken from [96KEH]. Then f mH ο∆ (NiCO3, T) was extrapolated to 298.15 K with Equation (A.97), see Table A-24.

As already noted by Tareen et al. the f mH ο∆ values obtained were slightly more nega-tive at higher pressures than at lower ones [91TAR/FAZ]. This may either be due to experimental uncertainties or to incorrect ,mpC and ln f(CO2) functions or both. When the temperature and pressure dependence of the solid phase volumes, as in Equation (A.94), is ignored, this certainly plays a minor role. By averaging the individual results the error limit of the standard enthalpy has been estimated to be 2σ = ± 6.0 kJ·mol–1, see Section V.7.1.1.1.1.3.1 and Table A-24 respectively.

Table A-24: f mH ο∆ (NiCO3, 298.15 K) from decomposition data [91TAR/FAZ].

T (K)

p (bar)

y(CO2) f mH ο∆ (NiCO3, T)(kJ·mol–1)

f mH ο∆ (NiCO3, 298.15 K) (kJ·mol–1)

593.15 1020 0.43 – 675.51 – 704.73 623.15 1400 0.50 – 673.92 – 706.42 643.15 1550 0.46 – 673.84 – 708.54 668.15 1900 0.49 – 672.60 – 710.08

isobars T low

673.15 2070 0.45 – 672.81 – 710.85

608.15 1020 0.50 – 674.75 – 705.60 648.15 1400 0.43 – 674.33 – 709.59 653.15 1550 0.46 – 673.66 – 709.47 688.15 1900 0.42 – 673.01 – 712.74

isobars T high

708.15 2070 0.45 – 672.01 – 714.01

– (709.2 ± 6.0)


423

[92BAL/DIC]

In this report1 the crystal structure, the phase transitions, the heat capacity function, the entropy and the standard enthalpy of formation of the solid phases in the Ni–Te system are comprehensively discussed and critically assessed. Thermodynamic functions in the range of 298.15 to 1100 K were listed in intervals of 100 K for the following Ni–Te phases with selected stoichiometries Ni0.333Te0.667(cr), Ni0.4Te0.6(cr), Ni0.476Te0.524(cr), Ni0.54Te0.46(cr), Ni0.563Te0.437(cr), Ni0.6Te0.4(cr), NiTe2(g).

[92RAR2]

Isopiestic vapour-pressure experiments were performed for aqueous solutions of high-purity NiCl2 at 25°C. Values of the water activity, aW, the osmotic coefficient, φ, and the mean activity coefficient, γ±, were recommended for this system. The solubility of NiSO4·6H2O(cr) was determined definitively. Older solubility mesurements [37PEA/ECK], [79OIK], [79OJK/MAK], [79PET/SHE], [83KIM/IMA], [85FIL/CHA], [86FIL/CHA], [87RAR3] were reviewed and the reason for the discrepancies were dis-cussed. Standard Gibbs energy of reaction,

NiCl2·6H2O(cr) Ni2+ + 2Cl– + 6H2O(l), (A.98)

sol mGο∆ (A.98) = – (17.38 ± 0.04) kJ·mol–1 was based on the statistically weighted aver-age of Rard’s as well as Pearce and Eckstrom’s [37PEA/ECK] solubility of NiSO4·6H2O(cr).

[92VID/KOR]

The standard enthalpy of formation of heazlewoodite, Ni3S2, was measured directly by reaction of elemental nickel with sulphur in a calorimeter. The reaction was initiated by the energy of a light pulse. The heat of reaction at 298.15 K was determined to be

f mH ο∆ = – (226.0 ± 2.2) kJ·mol–1. This value deviates by about 10 kJ·mol–1 from results obtained by equilibrium experiments [54ROS] as well as drop calorimetry [86CEM/KLE]. The products formed during the calorimetric process were shown to be Ni3S2 and metallic nickel by X-ray diffraction analysis. Thus, a certain amount of sul-phur evaporated during the reaction. It may be difficult to properly correct for the en-thalpy associated with this process, and thus the results may be erroneous.

[93BAL/DIV]

Balej and Divisek re-evaluated qualitatively equilibrium or rest potentials of solid nickel oxide hydroxides with 2 < z(Ni) ≤ 3.6 measured at ambient temperatures. The experimental data were converted to standard conditions i.e., 25°C, aKOH =

2H Oa = 1.0, E versus SHE. From the functions E = f (z(Ni)), thus obtained, the authors concluded that the so-called “β-NiOOH | β-Ni(OH)2” systems with z(Ni) < 3 are thermodynamically 1 In this report, Fig. 16 is erroneously titled ‘Phase diagram of the Ni–Te system’, it should read ‘⋅⋅⋅ the

Fe–Te system’ instead.


424

unstable or metastable. In the whole range investigated (z(Ni) > 3.5) the so-called “γ−NiOOH | α−Ni(OH)2” systems exhibit very flat S-shaped functions E = f (z(Ni)) with the lowest E values at the given z(Ni) and were regarded as thermodynamically stable solid solutions of Ni(II) + Ni(IV) components only.

[93BAL/DIV2]

Balej and Divisek based a quantitative re-evaluation of critically selected data on the equilibrium or rest potentials of the so-called “γ-NiOOH | α-Ni(OH)2” systems on the assumption that these systems consist of homogeneous solid solutions with the end-members Ni(OH)2 and NiO2·xH2O. Whereas the Ni(II) component is considered to be completely undissociated, the Ni(IV) component fully deprotonates to give H+ and NiO2·(x–1)H2O·OH–. While these authors convincingly describe the redox behaviour of Ni(II, IV) oxide hydroxides, the stoichiometric and structural characterisation of the hypothetical end-members is still lacking. So the thermodynamic quantities derived from this model were not accepted for the present review.

[93KUB/ALC]

In Chapter 3 of this textbook the estimation of thermodynamic data is described for compounds for which no experimental data are available. For the heat capacity of pre-dominantly ionic, solid compounds at 298 K cationic and anionic group contributions can be added. Moreover a set of equations is given to estimate the constants in the equa-tion ( 700K

,m 298K[ ]pCο (NiCO3, cr) = (88.701 + 38.91×10–3 (T/K) – 12.34×105 (T/K)–2) J·K−1·mol−1). The variables are the sum of the cationic and anionic contributions, the melting temperature of the compound, and the number of atoms in the molecule.

[93MAK/VOE]

Based on thermodynamic quantities, derived from the experimental investigation of the solubility of nickel sulphide performed by Thiel and Gessner [14THI/GES], solubility constants as well as Pourbaix diagrams (redox potential plotted vs. pH) have been calcu-lated.

[93MUL/ISA]

Muldagalieva et al. [93MUL/ISA] reported the equation: 700K

,m 298.15K[ ]pCο (Ni11As8, cr) = 380.12 + 367.27×10–3 (T/K) + 0.018×105 (T/K)–2 J·K–1·mol–1

for 298.15 K to 700 K based on dynamic calorimetric measurements ( ,mpCο (Ni11As8, cr, 298.15 K) = 490 J·K–1·mol–1). The experimental data were not provided, nor the good-ness of fit of the equation, though the authors claimed that the experimental uncertain-ties were “no greater than 8%”, and this uncertainty is accepted in the present review to be a 2σ value. In the temperature range of the equation, it would seem that the T –2 term in the equation makes a negligible contribution to the value, suggesting that the fitting was not done properly. The authors also reported an estimated standard entropy at


425

298.15 K of 580.0 kJ·mol–1 for the same compound (that estimate, by Kumok, was made in a reference unavailable to the present reviewers). The heat capacity values from this study near 700 K are in rough agreement with those estimated by Skeaff et al. [85SKE/MAI], though the variation with temperature is markedly different. For 298.15 K, the values from the two equations for ,mpCο (Ni11As8) differ by 62 J·K–1·mol–1, which is slightly greater than the experimental uncertainties.

[93NEI/POW]

Thermodynamic data of high accuracy and reliability for the Ni – NiO system at high temperatures were obtained from electrochemical cells with yttria- or calcia-stabilised zirconia electrolytes. The experimental results are in excellent agreement with available calorimetric data as well as emf measurements of previous publications. The recom-mended Gibbs energy equation for the reaction:

2 Ni + O2 2 NiO is f mG∆ (NiO) = (– 478.967 + 0.248514 (T/K) – 0.0097961 (T/K) ln(T/K)) kJ·mol–1, valid from 700 to 1700 K. The uncertainty in f mG∆ was reported to be ± 0.200 kJ·mol−1.

[93NIC/ROB]

Widgiemoolthalite was chemically and crystallographically investigated. The diffuse X-ray diffraction pattern suggests a high degree of structural disorder, and the pseu-doorthorhombic cell indicates the possibility of twinning. Thus a definitive crystallo-graphic characterisation was prevented, but the close similarity of its X-ray diffraction pattern to that of hydromagnesite indicates that it is the nickel analogue of the latter.

[93OBE/KAS]

Obendrauf et al. have measured the thermophysical properties of nickel by a fast pulse heating technique using a newly developed highly sensitive pyrometer. The values ob-tained for fus mH ο∆ (Ni) are slightly greater than, and the values for ,mpCο (Ni, l) are sig-nificantly greater than the calorimetrically determined values on which the quantities in this review are based.

[93PRZ/WIS]

The value 15.644 kJ·mol–1 was reported by Przepiera et al. [93PRZ/WIS] as the heat of solution of nickel sulphate heptahydrate in water. Although Archer [99ARC] assumed that this was an infinite dilution value, that information is not stated explicitly in the original paper, and it might also be concluded that the value represents the heat of solu-tion to produce a solution with a 5×10–4 mass ratio of salt to water. If the value is as-sumed to be for the enthalpy of solution to this stated dilution (approximately 0.002 m), the application of a heat of dilution correction [66GOL/RID] to I = 0 gives a value (12.5 kJ·mol–1) that is in good agreement with the value based on results of Goldberg et al.


426

[66GOL/RID] and Stout et al. [66STO/ARC] (also see [98PLY/ZHA]). As this ambigu-ity cannot be resolved, the value from this paper [93PRZ/WIS] is not used in determin-ing thermodynamic quantities selected in the present review.

[93SHE]

The paper provides the results of a conductance study. The results were analysed using the equation of d’Aprano [71APR]. No values are reported for the nickel sulphate sys-tem other than the limiting “equivalent conductance at infinite dilution” and the associa-tion constant. These values are not inconsistent with values from other conductance studies, but reanalysis of the data to ensure consistency with the SIT model is not possi-ble.

[93UVA/TIM]

The compositions of solid nickel chloride hydrates were determined at different water vapour pressures. The water vapour pressures were set using an isopiestic arrangement with aqueous H2SO4 solutions. At 25 and 35°C, the transition from the dihydrate to the tetrahydrate was found to occur at 0.998 kPa and 1.70 kPa, respectively. At 25 and 35°C, the transition from the tetrahydrate to the hexahydrate was found to occur at 1.19 kPa and 2.61 kPa, respectively. The uncertainty in the reported transition water vapour pressures is estimated in the present review as ± 0.05 kPa.

[94STE/FOT]

De Stefano et al. [94STE/FOT] measured the protonation constants for pyrophosphate, as a function of temperature between 5 and 45°C, in aqueous solutions of (CH3)4NCl, NaCl and KCl (0.00 to 0.75 M). These constants agree well (within 0.11 in 10log K ) with the values reported in the studies of Perlmutter-Hayman and Secco [73PER/SEC] (KNO3(aq)) and Hammes and Morrell [64HAM/MOR] ((CH3)4NCl(aq)) in their studies of nickel pyrophosphate complexation, and of Edwards et al. [73EDW/FAR] ((CH3)4NCl(aq)). From the apparent protonation constants, De Stefano et al. [94STE/FOT] calculated values of the formation constants of 3

2 7KP O − and 22 7KHP O −

(and of the formation constants for the corresponding sodium-ion species). The pyro-phosphate protonation constant values at I = 0, differ by less than 0.2 in 10log K from those selected by Grenthe et al. [92GRE/FUG].

[94STO/FJE]

The phase relations of the system Ni – S in the compositional range from 45 to 48 at% S were studied by application of high temperature X-ray diffraction analysis, differential thermal analysis (DTA), and adiabatic calorimetry. The solid compounds were prepared by solid state synthesis from the elements. The structure of the orthorhombic high tem-perature modification of Ni7S6 was resolved by high temperature X-ray diffraction analysis. This disordered high temperature phase transforms into metastable phases on quenching [96SEI/FJE]. In contrast to the results of Kullerud and Yund [62KUL/YUN]


427

the composition of the orthorhombic low temperature modification was found to be Ni9S8 instead of Ni7S6. The stable non-stoichiometric high temperature phase is formed eutectoidally from Ni3S2 and Ni9S8 at (675 ± 3) K and the decomposition into Ni3+xS2 and Ni1–xS occurs at (850 ± 2) K. Based on high temperature X-ray diffraction and DTA, the decomposition of Ni9S8 into Ni7S6 and Ni1–xS was observed at (709 ± 5) K, contradicting the findings of [62KUL/YUN] who obtained 673 K for the peritectoidal decomposition of the low temperature phase of Ni7S6.

The heat capacity of a sample with a nominal composition of Ni7S6 was measured by means of adiabatic calorimetry from 5 to 970 K. Below 675 K and above 850 K the specimen consisted of a mixture of two phases, viz. Ni9S8 + Ni3S2 and Ni3+xS2 + Ni1–xS, respectively. From the heat capacity function for Ni3S2, determined earlier [91STO/GRO], the heat capacity of Ni9S8 can be calculated for the temperature range from 298.15 to 640 K:

640K,m 298.15K[ ]pCο (Ni9S8, cr) = (1079.2452 – 2.05994·(T/K) + 0.00198·(T/K)2

– 20839639·(T/K)–2 ) J·K–1·mol–1. (A.99)

The exothermic and endothermic enthalpy effects occurring at 415 and 515 K, respectively, are attributed to metastable phases of Ni7S6 which are formed during the preparation of the samples. These enthalpy effects disappeared after annealing the sam-ples at 810 K for two days.

Figure A-37: Heat capacity of “Ni7S6” consisting of a mixture of Ni3S2 and Ni9S8 as a function of temperature. The solid line is derived from the heat capacity functions for Ni3S2 and Ni9S8 recommended in this compilation. The various symbols refer to sam-ples subjected to different heat treatment.

250 300 350 400 450 500 550 600 650 700

280

300

320

340

360

380

400

420

2/3 Ni9S8 + 1/3 Ni3S2

T / K

C° p,m

/ J·K

–1·m

ol–1


428

[94ZHA/MIL]

The stability constant of the nickel sulphide complex NiHS+ in sea water (I = 0.7 mol·kg–1 NaCl) is determined at 298.15 K by applying cathodic stripping square wave voltammetry. Typically, a nickel(II) standard solution was added to sea water, leading to molalities of total nickel(II) ranging between 500 and 1500 nmol·kg–1. These solu-tions were titrated with Na2S solutions, where successive standard additions of sulphide were in the range from 250 to 500 nmol·kg–1. The measurement of the stripping current of free sulphide, deposited on the hanging mercury drop electrode, yielded the molality of HS–. A proper evaluation of the experimental titration curves resulted in 10 1log b = (5.3 ± 0.1) for the formation constant of NiHS+ in sea water at 298.15 K.

[95GRO/STO]

The heat capacity of synthetic NiS was measured by adiabatic-shield calorimetry from 260 to 1000 K. The solid phase was prepared from the elements by solid state reaction and checked by X-ray diffraction analysis. The rhombohedral low temperature modifi-cation (millerite) was found to transform to the hexagonal non-stoichiometric NiAs polytype at 660 K where the heat of transition amounts to trs mH ο∆ = (6591 ± 50) J·mol–1. The high temperature form can easily be quenched to room temperature. Thus, the heat capacity function for the stoichiometric composition of this phase can be derived from 298.15 to 900 K:

660K,m 298.15K[ ]pCο (NiS, α) = (46.676 + 0.0199807 (T/K)

– 255499.4 (T/K)–2) J·K–1·mol–1 (A.100) 900K

,m 660K[ ]pCο (NiS, α) = (50.07833 + 0.0149 (T/K) – 336916.92 (T/K)–2) J·K–1·mol–1. (A.101)

In the case of millerite the heat capacity function is obtained for temperatures up to 660 K:

660K,m 298K[ ]pCο (NiS, β) = (54.008 – 0.01437 (T/K) + 3.2282 × 10–5 (T/K)2

– 490363 (T/K)–2) J·K–1·mol–1. (A.102)

Figure A-38 shows the temperature dependence of the heat capacity of NiS, where the dashed line corresponds to the metastable NiAs-type phase below 660 K. The steep increase of ,mpC at temperatures above 900 K can be attributed to the formation of small traces of Ni3+xS2.


429

Figure A-38: Heat capacity of NiS as a function of temperature. ○, solid lines: stable NiS phases. , dashed line: metastable NiAs-type modification below 660 K.

[95MAN/KOR]

Manin and Korolev reported having measured the enthalpy of solution of NiCl2(cr) into water and into aqueous solutions of other electrolytes. The measurements were per-formed with an isoperibol calorimeter with a claimed accuracy of ± 0.6%. Their results for the enthalpy of solution into H2O are listed in Table A-25. The mean value of

sol mH ο∆ for the reaction:

NiCl2(cr) Ni2+ + 2 Cl– (A.103)

agrees nicely with the one derived from Muldrow’s data. According to the Cl content (54.0%) determined by argentometric titration, the sample of NiCl2 used was not com-pletely anhydrous (n(H2O) / n(NiCl2) ≈ 0.09). This may lead to systematic errors similar to those discussed in the Appendix A entry for [58MUL].

200 300 400 500 600 700 800 900 100040

50

60

70

80

T / K

C° p,m

/ J·K

–1·m

ol–1


430

Table A-25: Enthalpy of solution of NiCl2 in H2O at 298.15 K [95MAN/KOR].

m (NiCl2) (kg·mol–1)

sol mH∆ (A.103) (kJ·mol–1)

ΦL (kJ·mol–1)


0.004230 – 82.67 0.5961 – 83.2661 0.004570 – 82.46 0.6169 – 83.0769 0.006320 – 82.92 0.7111 – 83.6311 0.009250 – 83.08 0.8371 – 83.9171 0.009260 – 82.83 0.8375 – 83.6675 0.009630 – 82.87 0.8515 – 83.7215 0.012410 – 82.10 0.9465 – 83.0465 0.014380 – 82.63 1.0054 – 83.6354 0.027340 – 82.10 1.2957 – 83.3957 0.030460 – 82.14 1.3499 – 83.4899

The variation of the enthalpy of solution of the anhydrous solid should parallel the enthalpy of solution of the hydrated solid as a function of salt molality. The experi-mental values for dissolution in water drift to slightly more negative values at lower NiCl2(aq) concentrations, but the measurements themselves have fairly large uncertain-ties of approximately ± 0.5 kJ·mol–1. The solution concentrations are not sufficiently low to provide a good extrapolation to I = 0, but values based on a theoretical extrapola-tion [97GRE/PLY2] are provided in Table A-25. The authors’ Figure 1 compares the values for the heats of solution in water to those taken from two compilations. However, there is no evidence as to how the values from the compilations were based on experi-mental heats of dilution. (There is also a typographical error in the key to the points in the authors’ Figure 1. The points designated “3” appear to be the data from [95MAN/KOR]).

The heats of dissolution of anhydrous nickel chloride in aqueous solutions con-taining several different concentrations of HCl(aq) also were determined. For measure-ments with final solutions having nickel chloride molalities of ca. 0.007 m, the enthalpy of solution of anhydrous nickel chloride for each mol% of HCl is about 3 kJ·mol–1 more positive than in water.

[95MUL/CHU]

This paper is similar to the earlier paper on Ni11As8(cr) [93MUL/ISA], but contains more details. The heat capacity equation, from dynamic calorimetric measurements, is reported as:

675K,m 298K[ ]pCο (NiAs, cr) = (36.54 + 41.87×10–3(T/K) + 1.58×105 (T/K)–2) J·K–1·mol–1.


431

[96BRU/PIA]

The authors reported an extensive set of torsion balance and Knudsen cell measure-ments for the determination of the enthalpy of sublimation of NiF2 between 950 and 1250 K. Third law and second law values are in much better agreement if the values of

m m( ( ) (298.15 K)) /G T H Tο− for the sublimation reaction for CoF2(cr) are substituted for those in the literature for NiF2(cr). The work appears to have been carefully carried out, and if the interpretation were correct, this would indicate that (a) mS ο (Ni, cr, 298.15 K) is 8 – 11 J·K–1·mol–1 less than based on [55CAT/STO], or (b) ,mpC (T) from [70BIN/HEB] is incorrect (although the values used for CoF2 are from work by the same group [67BIN/STR]) or (c) the value of f mH ο∆ (NiF2, cr, 298.15 K) based on Rudzitis et al. [67RUD/DEV] is too negative by about 8 kJ·mol–1.

[96KEH]

The table presented gives second virial coefficients of about 110 inorganic and organic gases as a function of temperature. Selected data from the literature have been fitted by least squares to the equation:

B (cm3·mol–1) = [ ] 10

1

( ) ( / ) 1n

i

i

a i T T −

=

−∑

where T0 = 298.15 K. For the present purpose, use has been made of the a(i) values of CO2.

[96LUT/RIC]

The stability constants for the aqueous species NiHS+, Ni2HS3+ and Ni3HS5+ at 298.15 K in seawater as well as sodium chloride solutions (artificial seawater) have been deter-mined by application of square wave voltammetry. Solutions of sodium sulphide (1 – 10 µmol·kg–1) in diluted seawater of constant ionic strength ranging from 0.07 to 0.7 mol·kg–1 NaCl were titrated with nickel (II) solutions, where care was taken to remove traces of oxygen from all solutions by purging with high purity argon. The molality of free sulphide was determined from both the peak potential and the peak current at con-stant pH for the electrochemical reaction:

HS– + Hg HgS + H+ + 2e–

occurring at the hanging mercury drop electrode. The titration curves were evaluated by means of a standard method [51DEF/HUM], [77HEA/HEF], leading to stability con-stants for NiHS+, Ni2HS3+ and Ni3HS5+, respectively. Plots of the peak potential versus pH indicate the existence of bisulfide (HS–) complexes at pH values above 7.

The ionic strength dependence of the nickel bisulphide stability constants is depicted in Figure A-39. By applying the SIT concept [97GRE/PLY2] the respective stability constants are extrapolated to zero ionic strength. As well as the stability con-stants, the optimisation procedure also yields the SIT interaction parameters ∆ε in chlo-ride media, T = 298.15 K:


432

Ni2+ + HS– NiHS+ (A.104) 10 1log οb = (5.18 ± 0.20); ∆ε(A.104) = – (0.97 ± 0.39) kg·mol–1.

2Ni2+ + HS– Ni2HS3+ (A.105) 10 2log οb = (9.92 ± 0.10); ∆ε(A.105) = – (0.05 ± 0.22) kg·mol–1.

3Ni2+ + HS– Ni3HS5+ (A.106) 10 3log οb = (14.008 ± 0.099); ∆ε(A.106) = (0.59 ± 0.22) kg·mol–1.

Figure A-39: Logarithm of stability constants of nickel bisulphide complexes in sea-water (NaCl solutions) plus the Debye-Hückel term for ionic strength correction plotted as a function of ionic strength, T = 298.15 K. The solid lines correspond to least squares fits of the SIT model to the experimental data, ● [96LUT/RIC], ▲ [94ZHA/MIL].

(a) 10 1log οb = (5.18 ± 0.20); ∆ε(A.104) = – (0.97 ± 0.39) kg·mol–1.

(b) 10 2log οb = (9.92 ± 0.10); ∆ε(A.105) = – (0.05 ± 0.22) kg·mol–1.

(c) 10 3log οb = (14.008 ± 0.099); ∆ε(A.106) = (0.59 ± 0.22) kg·mol–.

0.0 0.2 0.4 0.6 0.8 1.0

4.6

4.8

5.0

5.2

5.4

5.6

5.8

6.0

6.2

6.4

(a)

log

β 1I + D

H

I / mol·kg-1

0.0 0.2 0.4 0.6 0.89.80

9.85

9.90

9.95

10.00

10.05

10.10

10.15

10.20

(b)

log

β 2I + D

H

I / mol·kg-1

0.0 0.2 0.4 0.6 0.8 1.013.0

13.2

13.4

13.6

13.8

14.0

14.2

14.4

(c)

log

β 3I + D

H

I / mol·kg-1


433

[96POU/DRE]

Nickel hydroxide purchased from Johnson Matthey Chemical Co. was equilibrated with 0.002 to 0.200 mol·kg–1 NaNO3 over a pH range of approximately 5.5 – 6.5 at 22°C. The experimental data were evaluated using the Davies equation in MINTEQA2 at these ionic strengths [93ALL/BRO]. The solubilities obtained at I = 0.2 mol·kg–1 were apparently lower than those at I = 0.002 – 0.02 mol·kg–1. The authors preferred the higher values, because their model might become incorrect at I > 0.02 mol·kg–1. When the activity data were recalculated assuming that Equation (A.107) is valid a mean value of 10 ,0log sK ο = – (15.96 ± 0.10) was obtained at 295.15 K. 10 ,0

*log sK ο = 12.24 recorded in Table V-6 and Figure V-11 was based on this evaluation.

2+10 10 ,0 10Ni OHlog log 2logsa K a −

ο= − (A.107)

[96SAH/SAH]

This paper reports 10 1log K = (2.20 ± 0.02) for the association constant of Ni2+ with 24HPO − at 25°C in 0.1 M NaNO3. The study appears to have been done carefully,

though several of the minor difficulties found for an earlier paper [67SIG/BEC] from the same group still occur (see Appendix A for [67SIG/BEC]). Correction to zero ionic strength (Appendix B) gives 10 1log K ο = (3.07 ± 0.10), where the uncertainty has been assigned in the present review. The small ∆ε correction has been based on the value for ε(Na+, 2

4HPO − ) from Table B-5 (rather than using the equation in Table B-6), and ε(Ni2+, 3NO− ) = 0.18 kg·mol–1.

[96SEI/FJE]

The relative stability and structural descriptions of four metastable phases with compo-sition close to Ni7S6 were obtained from X-ray diffraction, transmission electron mi-croscopy, calorimetry, and thermal analysis. The structures of these phases are closely related to that of the disordered high-temperature phase Ni7S6 which is formed eutec-toidally from Ni3S2 and Ni9S8 at 675 K [94STO/FJE]. When Ni7S6 is cooled to room temperature, this phase decomposes into the stable compounds Ni3S2 and Ni9S8 or a transformation of Ni7S6 into metastable modifications of similar composition occurs.

[97BAH/PAR]

This is a recent compilation of thermodynamic data of metal thiocyanate complexes. The paper listed the majority of available thermodynamic data for the nickel(II) – thio-cyanate system, but recommended values are given only for 10 1log b at I = 1.0 ( 10 1log b = (1.14 ± 0.09)) and 1.5 M ( 10 1log b = (1.14 ± 0.02)).

[97BAL/DIV]

Revised E–pH and E–m(KOH) diagrams for nickel at 25°C are calculated and graphi-cally represented using the thermodynamic data derived for hypothetical NiO2·xH2O


434

[93BAL/DIV], [93BAL/DIV2]. This E–pH diagram is meant to supersede previous ver-sions [63DEL/ZOU], [81SIL].

[97MAT/RAI]

Nickel hydroxide was precipitated by adding NaOH to a 0.2 M NiCl2 stock solution. The final pH of the Ni(OH)2 suspension was approximately seven. The precipitate was washed with water until it was free of chloride. The X-ray diffraction data of the freeze dried precipitate matched with JCPDS data file 14-117 for the Ni(OH)2 mineral, theophrastite. Approximately 300 mg of Ni(OH)2 were suspended in 30 cm3 of 0.01 M NaClO4. The pH was adjusted between 7 and 12 by HClO4 and NaOH, respectively. Thus, only a minute Ni(OH)2 portion, doubtlessly having consisted of the finest particles, was actually dissolved. When equilibrium was approached from oversaturation again finely dispersed Ni(OH)2 will have been precipitated. As solubilities of metal oxides and hydroxides depend on the particle size [65SCH/ALT], the numerical value of the solubility constant will probably be overestimated by this method, see Figure V-11.

[97PAN/CAM]

This is probably the most comprehensive set of heat capacity results available for any nickel salt in aqueous solution. Apparent molar heat capacities of aqueous Ni(ClO4)2 were measured calorimetrically from 25 to 85°C over a molality range of 0.02 to 0.80 mol·kg–1. Standard molar heat capacities of Ni2+ for the same temperature range were obtained by using the additivity rule and data for HClO4(aq), given in literature. The results for ,mpCο (Ni2+) can be fitted with a conventional heat capacity model valid from 298.15 to 358.15 K using ,mpCο (HClO4, aq, T) from Lemire et al. [96LEM/CAM]:

358.15K,m 298.15K[ ]pCο (Ni2+) = (842.427 – 1.8166 (T/K) – 3.06653 × 107/(T/K)2) J·K–1·mol–1.

The values change only slowly with temperature in this range, with an apparent shallow maximum value near 50°C. The authors extrapolated their results to 300°C us-ing the Helgeson-Kirkham-Flowers equation [74HEL/KIR2], [81HEL/KIR], [97PAN/CAM]. The mean value for the standard partial molar heat capacity of Ni2+ at 25°C in the authors’ Table III is – 44.0 kJ·mol–1 whereas the result in the authors’ Table IV is ,mpCο (Ni2+, 298.15 K) = – 49.6 J·K–1·mol–1. The difference reflects the sensitivity of the ,mpCο value to the use of two different, but reasonable, auxiliary values for the partial molar heat capacity for HClO4(aq). As noted in Section V.2.1.4, results for other nickel salt solutions lead to similar uncertainties in the value of ,mpCο (Ni2+, 298.15 K).

[97RIC]

This monograph describes the synthesis, structure and reactivity of aqua ions. Only a short paragraph deals with Ni3+, because this ion occurs only as a ‘transient’ species.


435

[98GAM/KON]

A plot of 2

2+10 CO

log {[Ni ] }p× versus pH showed that the corresponding solubility con-stant is almost independent of temperature at least in the range of 348.15 to 363.15 K. Apparently other neutral transition metal carbonates behave similarly.

[98JAI/ELM]

The redox reaction of nickel hydroxide and nickel oxide hydroxide, the electrochemi-cally active compounds at the positive electrode of a nickel battery, was investigated. The thermodynamics of non-ideal solid solutions were applied to the reversible poten-tial as a function of the state-of-discharge. In a temperature range 5 to 55°C two pa-rameter activity coefficient models perform significantly better than one parameter models.

[98PLY/ZHA]

Plyasunova et al. critically evaluated the standard thermodynamic quantities of Ni2+, its hydrolysis reactions and hydroxo-complex formation on the basis of published experi-mental studies and the specific interaction theory (SIT) for activity coefficient model-ling [97GRE/PLY2]. Recommended thermodynamic functions and interaction coeffi-cients relevant for the present review and its compounds were presented, see Table A-26.

Table A-26: Recommended formation constants ( ,*

n mοb ) and reaction enthalpies for the

hydroxo complexes formed in the acidic region, at 298.15 K, mNi2+ + nH2O(l) 2Ni (OH) m n

m n− + nH+

Species 10 ,*log n m

οb r mH ο∆ (kJ·mol–1)

NiOH+ – (9.50 ± 0.36) (50 ± 21)

Ni2OH3+ – (9.8 ± 1.2) (43.3 ± 3.0) 4

4 4Ni (OH) + – (27.9 ± 1.0) (180 ± 2)

Table A-27: Recommended equilibrium constants ( , ,

*s n mK ο ) and reaction enthalpies for

the hydroxo complexes formed in the alkaline region, at 298.15 K, mNi(OH)2(cr) + nOH– 2Ni (OH) n

m m n−

+

Species 10 , ,*log s n mK ο r mH ο∆ (kJ·mol–1)

Ni(OH)2(aq) – (7.52 ± 0.80) –

3Ni(OH)− – (3.7 ± 1.8) (7.1 ± 3.0) 24Ni(OH) − (6.43 ± 0.23) – (14 ± 8)


436

Table A-28: Recommended ion interaction coefficients ε(j, k). k j Cl– 4ClO− Na+

NiOH+ ≥ 0.31 (0.30 ± 0.18) –

Ni2OH3+ (0.41 ± 0.56) (0.59 ± 0.16) – 4

4 4Ni (OH) + (0.68 ± 0.38) (0.90 ± 0.32) –

Ni(OH)2(aq) – – – (0.07 ± 0.03)

3Ni(OH)− – – (0.18 ± 0.05) 24Ni(OH) − – – (0.27 ± 0.06)

Table A-29: Standard thermodynamic functions of the aqueous species.

Species f mGο∆ (kJ·mol–1) f mH ο∆ (kJ·mol–1) mS ο (J·K–1·mol–1)

Ni2+ – (45.5 ± 3.4) – (54.1 ± 2.5) – (130 ± 3)

NiOH+ – (228.4 ± 5.7) – (290 ± 24) – (74 ± 81)

Ni2OH3+ – (272 ± 14) – (359 ± 20) – (260 ± 86)

Ni4(OH) 44

+ – (971 ± 19) – (1190 ± 27) – (203 ± 90)

Ni(OH)2 (aq) – (417 ± 11) – (540 ± 13) – (48 ± 32)

Ni(OH) 3− – (587 ± 15) – (791 ± 18) – (85 ± 47)

Ni(OH) 24

− – (734.0 ± 8.6) – –

The analysis was useful in considering some of the references, but the values from [98PLY/ZHA] are not used in the present review.

[98ROG/KOZ]

Exchange of sodium ions in β’’-alumina, performed in molten salts, made possible a preparation of a series of divalent β’’-aluminas. In this way the solid electrolytes con-ducting zinc, cobalt, nickel, calcium, manganese and copper ions were prepared. Their application to the solid-state galvanic cells have been presented. The Gibbs energy of formation of nickel orthosilicate, from oxides, determined by [98ROG/KOZ] in the range 973 – 1323 K was used by this review to fit best experimental values in the tem-perature range 970 – 1770 K.

[99ARC]

Archer detailed the sources of previous thermodynamic property values for nickel and some of its compounds being of importance to environmental processes. This review is particularly relevant to the origin of values contained in the NBS Thermodynamic Ta-bles [82WAG/EVA].


437

[99KON/KON]

Königsberger et al. presented a low-temperature thermodynamic model for the Na2CO3–MgCO3–CaCO3–H2O system. The model was based on calorimetrically deter-mined f mH ο∆ (298 K) values, mS ο (298 K) values and ,mpCο (T) functions taken from the literature as well as on οµ (298 K) values of solids derived from solubility measure-ments.

[99MAL/CAR]

The activity coefficients of (among others) aqueous Ni(ClO4)2 solutions (4 2Ni(ClO )m =

7.3×10–5 – 0.89 mol·kg–1) have been determined from the emf of liquid-membrane cells. For the purposes of this review, these data, together with the activity and osmotic coef-ficients listed in [69LIB/SAD], have been used to derive the ion interaction coefficient ε(Ni2+, 4ClO− ). The selected value is ε(Ni2+, 4ClO− ) = (0.37 ± 0.03) kg·mol–1 (see Section V.4.3). Ciavatta [80CIA] proposed to use ionic strength dependent ion interaction coef-ficients if the uncertainty is ± 0.03 or greater (Appendix B, Section B.3). The above uncertainty is at the limit. Accepting Ciavatta’s proposal ε(Ni2+, 4ClO− ) = ε1 + ε2 log10 Im, where ε1 = (0.287 ± 0.02) kg·mol–1 , ε2 = (0.121 ± 0.03) kg·mol–1 can be calculated. However, these values provide a better fit to the experimental data only for m > 2.5 mol·kg–1 (I > 7.5 mol·kg–1), and lead to negative ion interaction coefficients for m < 0.0015 mol·kg–1. Therefore, the use of an ionic strength dependent ε(Ni2+, 4ClO− ) was rejected in this review.

[99MCB]

McBreen comprehensively reviewed nickel hydroxide battery electrodes, the solid state chemistry of nickel hydroxides, and the electrochemical reactions of the Ni(OH)2 | NiOOH couple. Any critical discussion of the thermodynamic data of nickel oxide hydroxides with higher oxidation states has to refer to this splendidly written ac-count of nickel solid state electrochemistry.

[2000ROG/KOZ]

This paper of Róg et al. is a description of emf-measurements of Ni2SiO4 and Co2SiO4 with a solid electrolyte. A calcium fluoride-based composite containing SiO2 was ap-plied as a solid electrolyte. The results obtained in the temperature range 823 to 1273 K coincide perfectly with results of other emf-studies of Ni2SiO4 carried out by Róg et al. [84ROG/BOR], [98ROG/KOZ] and also by other authors quoted in this review.

[2000TSI/MOL]

The main object of this conductance study was to obtain information on changes in as-sociation at a function of solvent composition. However, the value for the nickel sul-phate association constant in water (calculated from the data using the Lee-Wheaton equation) at 20°C is reported. For the purposes of the present review, the data were re-


438

analysed with the distance parameter fixed to the value consistent with the SIT proce-dure described in Appendix B. The recalculated value of K1 is (118.6 ± 7.2) compared to (139 ± 1) reported in the original paper. This value is markedly lower than the results from other studies [73KAT], [76SHI/TSU], [79FIS/FOX].

[2001GAM/PRE]

In order to synthesise hellyerite, Rossetti-François’ method was applied and an aqueous solution of nickel chloride was reacted with ammonium carbonate. Fragmented crystals (d ≤ 2µm) and amorphous particles were precipitated [52ROS]. Therefore a new method was developed. The initially precipitated amorphous particles were dissolved and larger crystals (d = 0.05 mm) obtained when Ni2+ was kept in excess and the solution was slowly acidified by CO2. The X-ray data of hellyerite prepared by this new method agreed satisfactorily with the data given in the JCPDS-ICDD card 12- 276.

The aqueous solubility of hellyerite, then believed to have the stoichiometric formula NiCO3·6H2O, was studied at varying temperatures and constant ionic strength (I = 1.0 mol·kg–1). The crystals used were X-rayed before and after the dissolution ex-periment, but each time only the characteristic peaks of hellyerite were found. Data of

2

2+10 CO

log {[Ni ] }p× plotted versus pH fell on straight lines with slopes of approxi-mately – 2.0 for all temperatures.

From the experimental data obtained a preliminary set of the thermodynamic quantities f mGο∆ , f mH ο∆ and mS ο for hellyerite was derived using the ChemSage opti-miser routine [95KON/ERI].

[2001HEN/RED]

This is a study of the non-convergent ordering of Mg and Ni in synthetic olivines by means of the neutron powder diffraction and X-ray absorption spectroscopy (EXAFS). The implications of the intracristalline ordering to olivine crystal chemistry and their application to the consideration of thermodynamic relations, geothermometry, geo-speedometry, and minor or trace element distribution control is discussed in detail.

[2001MOL/TSI]

The same data for nickel sulphate are reported as in [2000TSI/MOL]. However, values for the specific conductance are also reported.

[2002GAM/WAL]

The re-evaluation of f mGο∆ and f mH ο∆ of β-Ni(OH)2 discussed in Section V.3.2.2.3. has been based on the results of this paper.

Theophrastite, β-Ni(OH)2, was synthesised by an improved method based on the hydrolysis of Na2[Ni(OH)4]. Pure single crystals were obtained with sizes up to 0.25 mm [2002WAL/GAT]. The solubility of β-Ni(OH)2 was measured at different


439

temperatures and ionic strengths, and the data observed were thermodynamically ana-lysed to obtain the respective standard molar quantities of formation f mGο∆ and f mH ο∆ . For parameter optimisation non-linear least squares routines were used.

The tendency of precipitated Ni(OH)2 to form larger, well-shaped crystallites is extraordinarily small (see Figure 13 of [77OSW/ASP]). Consequently it seemed doubt-ful that solubility equilibria according to the following equations

β-Ni(OH)2(cr) Ni2+ + 2 OH–

β-Ni(OH)2(cr) + 2 H+ Ni2+ + 2 H2O(l)

can be approached both from over- and undersaturation. Thus the method of pH-variation (i.e., the initial acidity was varied from 0.005 to 0.100 mol·kg–1 HClO4) was employed [63SCH], but met in the case of β-Ni(OH)2 with two difficulties:

• The equilibrium pH in the neutral range is poorly buffered by solute species.

• At 25°C Ni(OH)2 is notoriously inert, thus the solid-solute buffer system doesn’t work either.

However, when the experiments were carried out in a temperature range be-tween 35 and 80°C constant pH values indicating equilibrium were attained within peri-ods from 3 weeks to as little as 48 hours.

In a first series of solubility experiments on theophrastite the temperatures were varied, and the ionic strength was kept constant at 1.0 mol·kg–1 NaClO4. Figure A-40 shows results typical for the pH-variation method. Data of log10[Ni2+]tot plotted ver-sus pH fall on straight lines with the theoretical slope of – 2.0.

At 50°C a second series of solubility measurements was carried out at different ionic strengths varying from 0.5 to 3.0 mol·kg–1 NaClO4. Solubility constants were ex-trapolated to infinite dilution using the SIT [97GRE/PLY2]. Figure A-41 shows that the experimental data of β-Ni(OH)2 coincide reasonably well with the optimised curve ob-tained when the SIT equation is employed.


440

Figure A-40: Solubility of theophrastite as determined by the pH variation method; t(°C): 80 ▲, 70 □, 60 ▼, 50 ○, 35 ■; solid straight lines; calculated with mean values of

10 ,0*log sK and theoretical slope = – 2.0; I = 1.0 mol·kg–1 (Na)ClO4.

Figure A-41: Ionic strength dependence of theophrastite solubility. Inert electrolyte: NaClO4. I (mol·kg–1): 0.5 ♦, 1.0 ●, 2.0 ■, 3.0 ▲, + mean value and error bar (∆ 10

*log sK = ± 0.2). Solid curve: calculated according to the SIT model [97GRE/PLY2] with the interaction parameter ε(Ni2+, 4ClO− ) = 0.37 kg·mol–1 selected in this review.

5.0 5.2 5.4 5.6 5.8 6.0 6.2 6.4 6.6 6.8-2.8

-2.6

-2.4

-2.2

-2.0

-1.8

-1.6

-1.4

-1.2

-1.0

log 10

([N

i2+] to

t/mol

·kg−1

)

pH

0.0 0.2 0.4 0.6 0.8 1.0 1.2 1.4 1.6 1.8

9.8

10.0

10.2

10.4

10.6

10.8

T/K = 323.15

log 10

Ks

(I / mol·kg−1)0.5


441

[2002MON/HEL]

A kinetic study of cyanide exchange in [M(CN)4]2– square-planar complexes (M = Ni, Pd and Pt) was performed as a function of pH using 13C NMR. The kinetic data con-firmed the presence of the pentacyano complex at a higher excess of cyanide, and the protonated species below pH 7. Surprisingly, the protonation does not affect the chemi-cal shift of the 13C NMR complexed-cyanide signal (Kolski and Margerum detected identical UV-VIS spectra for 2

4Ni(CN) − and its protonated derivatives [68KOL/MAR]). The observed kinetic behaviour was described using 10 4log b = 31.0 for the formation constant of 2

4Ni(CN) − [59FRE/SCH], 10 5log K = – 0.77 for the reaction: 24Ni(CN) − + CN– 3

5Ni(CN) −

[71PIE/HUG], and 10 1,Hlog K = 5.4 and 10 2,Hlog K = 4.5 for the two successive protona-tions of 2

4Ni(CN) − [68KOL/MAR]. Under the conditions used (25°C, Im = 0.6 m), 10 4log b = (29.75 ± 0.20) seems to be more reliable, but the impact of the erroneous β4

on the conclusions (e.g., the presence of protonated species) is rather difficult to judge.

[2002WAL/GAT]

A new method to synthesize β-Ni(OH)2 has been described. It is based on the hydrolysis of Na2[Ni(OH)4] and leads to pure single crystalline nickel hydroxide with crystal sizes up to 0.25 mm. The influence of temperature and concentration on the crystal size was studied.

[2002WAL/PRE]

Solubility measurements of hellyerite were carried out at different ionic strengths of NaClO4 (I = 0.5, 1.0, 2.0 and 3.0 mol·kg–1). These and all other solubility data on gaspé-ite [80REI], [82KLO] and hellyerite [11AGE/VAL], [30MUL/LUB], [2001GAM/PRE] available were thermodynamically analysed to obtain the respective standard thermody-namic quantities f mGο∆ and f mH ο∆ . For hellyerite, mS ο was also determined. Note that in this paper all calculations on hellyerite were based on the stoichiometric formula NiCO3·6H2O.

[2003BAE/BRA]

The stability constants of nickel carbonato complexes were investigated by measuring the solid/liquid distribution ratio of Ni in the absence and presence of varying carbonate concentrations. For this purpose a 63Ni tracer method was employed. A considerable part of the uncertainty ( 10 1log K ο , for the reaction Ni2+ + 2

3CO − NiCO3(aq), is calcu-lated to be (4.2 ± 0.4)) inherent in this method arises from the poorly known formation constants of neutral and anionic hydroxo complexes of nickel.


442

[2003HUM/CUR]

Hummel and Curti comprehensively reviewed complex formation between nickel and carbonato and hydrogen carbonato ligands. Not only was the basis for experimental measurements discussed, but also methods used for estimation of the formation con-stants [76ZHO/BEZ], [77MAT/SPO], [79MAT/SPO], [81TUR/WHI], [84FOU/CRI], [87EMA/FAR]. Therefore a short summary of their arguments suffices.

As suggested by Garrels and Christ linear relationships between – 10 1log K ο or – 10 1log οb and the electronegativity of the metals were used by Zhorov et al. to deter-mine the respective constants for Ni(II) [65GAR/CHR], [76ZHO/BEZ]. The former case was calibrated with CuCO3(aq), MgCO3(aq) and CaCO3(aq), the latter with

3MgHCO+ and 3CaHCO+ only, although Garrels and Christ warned against an overes-timation of the then available data on alkaline earth hydrogen carbonato complexes. Mattigod and Sposito correlated the complex formation constants with the sum of the radii of the metal ion and the ligand forming the ion-pair. Their estimation formula for the nickel carbonate system was calibrated to a single experimental data point (CuCO3(aq)) [77MAT/SPO]. Two years later, Mattigod and Sposito correlated these stability constants with a multiple parameter involving ionic charge, electronegativity, integral weighting factors, and a contribution calculated from atomic shielding constants [79MAT/SPO]. The numerical values obtained by the two estimation methods differ considerably, but their empirical bases remained equally doubtful. Turner et al. used the following correlation between the stability constants of carbonato and oxalato com-plexes:

310 MCOlog b = – 0.011 + 1.0422 410 MC Olog b (A.108)

Hummel and Curti succeeded in revealing the rationale of Equation (A.108) and pointed out that it leads to 10 1log K ο (NiCO3, aq) = (5.4 ± 1.5) rather than 10 1log K ο = 5.36 [81TUR/WHI]. Fouillac and Criaud [84FOU/CRI] applied different versions of electrostatic models, but finally resorted to the numbers of Zhorov et al. [76ZHO/BEZ] which they mistook for experimental data and therefore corrected to I = 0. This vicious circle was probably one of the reasons which led to the subtitle “A thermodynamic el-egy” for Hummel and Curti’s review [2003HUM/CUR].

B Appendix B

Ionic strength corrections1

Thermodynamic data always refer to a selected standard state. The definition given by IUPAC [82LAF] is adopted in this review as outlined in Section II.3.1. According to this definition, the standard state for a solute B in a solution is a hypothetical solution, at the standard state pressure, in which mB = m = 1 mol·kg−1, and in which the activity coefficient γB is unity. However, for many reactions, measurements cannot be made accurately (or at all) in dilute solutions from which the necessary extrapolation to the standard state would be simple. This is invariably the case for reactions involving ions of high charge. Precise thermodynamic information for these systems can only be ob-tained in the presence of an inert electrolyte of sufficiently high concentration that en-sures activity factors are reasonably constant throughout the measurements. This ap-pendix describes and illustrates the method used in this review for the extrapolation of experimental equilibrium data to zero ionic strength.

The activity factors of all the species participating in reactions in high ionic strength media must be estimated in order to reduce the thermodynamic data obtained from the experiment to the state I = 0. Two alternative methods can be used to describe the ionic medium dependence of equilibrium constants:

• One method takes into account the individual characteristics of the ionic media by using a medium dependent expression for the activity coefficients of the species involved in the equilibrium reactions. The medium dependence is de-scribed by virial or ion interaction coefficients as used in the Pitzer equations [73PIT] and in the specific ion interaction theory.

1 This Appendix essentially contains the text of the TDB-2 Guideline written by Grenthe and Wanner

[2000GRE/WAN], earlier versions of which have been printed in the previous NEA TDB reviews [92GRE/FUG], [95SIL/BID], [99RAR/RAN], [2001LEM/FUG] and [2003GUI/FAN]. The equations pre-sented here are an essential part of the review procedure and are required to use the selected thermody-namic values. The contents of Tables B.4 and B.5 have been revised.

443

B Ionic strength corrections

444

• The other method uses an extended Debye-Hückel expression in which the ac-tivity coefficients of reactants and products depend only on the ionic charge and the ionic strength, but it accounts for the medium specific properties by in-troducing ion pairing between the medium ions and the species involved in the equilibrium reactions. Earlier, this approach has been used extensively in ma-rine chemistry, cf. Refs. [79JOH/PYT], [79MIL], [79PYT], [79WHI2].

The activity factor estimates are thus based on the use of Debye-Hückel type equations. The “extended” Debye-Hückel equations are either in the form of specific ion interaction methods or the Davies equation [62DAV]. However, the Davies equa-tion should in general not be used at ionic strengths larger than 0.1 mol · kg−1. The method preferred in the NEA Thermochemical Data Base review is a medium-dependent expression for the activity coefficients, which is the specific ion interaction theory in the form of the Brønsted-Guggenheim-Scatchard approach. Other forms of specific ion interaction methods (the Pitzer and Brewer “B-method” [61LEW/RAN] and the Pitzer virial coefficient method [79PIT]) are described in the NEA Guidelines for the extrapolation to zero ionic strength [2000GRE/WAN].

The specific ion interaction methods are reliable for intercomparison of ex-perimental data in a given concentration range. In many cases this includes data at rather low ionic strengths, I = 0.01 to 0.1 M, cf. Figure B-1, while in other cases, nota-bly for cations of high charge ( ≥ + 4 and ≤ − 4), the lowest available ionic strength is often 0.2 M or higher, see for example Figures V.12 and V.13 in [92GRE/FUG]. It is reasonable to assume that the extrapolated equilibrium constants at I = 0 are more pre-cise in the former than in the latter cases. The extrapolation error is composed of two parts, one due to experimental errors, and the other due to model errors. The model errors seem to be rather small for many systems, less than 0.1 units in 10log K ο . For reactions involving ions of high charge, which may be extensively hydrolysed, one cannot perform experiments at low ionic strengths. Hence, it is impossible to estimate the extrapolation error. This is true for all methods used to estimate activity corrections. Systematic model errors of this type are not included in the uncertainties assigned to the selected data in this review.

It should be emphasised that the specific ion interaction model is approximate. Modifying it, for example by introducing the equations suggested by Ciavatta ([90CIA], Eqs. (8−10), cf. Section B.1.4), would result in slightly different ion interac-tion coefficients and equilibrium constants. Both methods provide an internally consis-tent set of values. However, their absolute values may differ somewhat. Grenthe et al. [92GRE/FUG] estimate that these differences in general are less than 0.2 units in

10log K ο , i.e., approximately 1 kJ·mol−1 in derived f mGο∆ values.

B.1 The specific ion interaction equations 445

B.1 The specific ion interaction equations

B.1.1 Background In the following discussion of the SIT model strict quantity calculus was sacrified in order to improve readability (cf. [91STO]). The Debye-Hückel term, which is the domi-nant term in the expression for the activity coefficients in dilute solution, accounts for electrostatic, non-specific long-range interactions. At higher concentrations, short range, non-electrostatic interactions have to be taken into account. This is usually done by adding ionic strength dependent terms to the Debye-Hückel expression. This method was first outlined by Brønsted [22BRO], [22BRO2] and elaborated by Scatchard [36SCA] and Guggenheim [66GUG]. Biedermann [75BIE] highlighted its practical value, especially for the estimation of ionic medium effects on equilibrium constants. The two basic assumptions in the specific ion interaction theory are described below.

• Assumption 1: The activity coefficient jg of an ion j of charge zj in the solu-

tion of ionic strength Im may be described by Eq. (B.1): 2

10log = ( , , )j j m kk

z D j k I m− + ε∑g (B.1)

D is the Debye-Hückel term:

= 1 +

m

j m

A ID

B a I (B.2)

where Im is the molal ionic strength:

21 = 2m i i

iI m z∑

A and B are constants which are temperature and pressure dependent, and aj is an ion

size parameter (“distance of closest approach”) for the hydrated ion j. The Debye-Hückel limiting slope, A, has a value of (0.509 ± 0.001) 1 1

2 2kg mol−⋅ at 25°C and 1 bar, (cf. Section B.1.2). The term Baj in the denominator of the Debye-Hückel term has been assigned a value of Baj = 1.5 1 1

2 2kg mol−⋅ at 25°C and 1 bar, as proposed by Scatchard [76SCA] and accepted by Ciavatta [80CIA]. This value has been found to minimise, for several species, the ionic strength dependence of ( , , )ε mj k I between Im = 0.5 m and Im = 3.5 m. It should be mentioned that some authors have proposed different values for Baj ranging from Baj = 1.0 [35GUG] to Baj = 1.6 [62VAS]. However, the parameter Baj is empirical and as such is correlated to the value of ( , , )ε mj k I . Hence, this variety of values for Baj does not represent an uncertainty range, but rather indi-cates that several different sets of Baj and ( , , )ε mj k I may describe equally well the experimental mean activity coefficients of a given electrolyte. The ion interaction coef-ficients at 25°C listed in Table B-4, Table B-5 and Table B-6 have thus to be used with Baj = 1.5 1 1

2 2kg mol−⋅ .


446

The summation in Eq. (B.1) extends over all ions k present in solution. Their molality is denoted by mk, and the specific ion interaction parameters, ( , , )ε mj k I , in general depend only slightly on the ionic strength. The concentrations of the ions of the ionic medium are often very much larger than those of the reacting species. Hence, the ionic medium ions will make the main contribution to the value of log10γj for the react-ing ions. This fact often makes it possible to simplify the summation ( , , )ε∑ m k

kj k I m ,

so that only ion interaction coefficients between the participating ionic species and the ionic medium ions are included, as shown in Eqs. (B.4) to (B.8).

• Assumption 2: The ion interaction coefficients, ( , , )ε mj k I are zero for ions of the same charge sign and for uncharged species. The rationale behind this is that ε, which describes specific short-range interactions, must be small for ions of the same charge since they are usually far from one another due to electrostatic repulsion. This holds to a lesser extent also for uncharged species.

Eq. (B.1) will allow fairly accurate estimates of the activity coefficients in mixtures of electrolytes if the ion interaction coefficients are known. Ion interaction coefficients for simple ions can be obtained from tabulated data of mean activity coef-ficients of strong electrolytes or from the corresponding osmotic coefficients. Ion inter-action coefficients for complexes can either be estimated from the charge and size of the ion or determined experimentally from the variation of the equilibrium constant with the ionic strength.

Ion interaction coefficients are not strictly constant but may vary slightly with the ionic strength. The extent of this variation depends on the charge type and is small for 1:1, 1:2 and 2:1 electrolytes for molalities less than 3.5 m. The concentration de-pendence of the ion interaction coefficients can thus often be neglected. This point was emphasised by Guggenheim [66GUG], who has presented a considerable amount of experimental material supporting this approach. The concentration dependence is larger for electrolytes of higher charge. In order to reproduce accurately their activity coeffi-cient data, concentration dependent ion interaction coefficients have to be used, cf. Lewis et al. [61LEW/RAN], Baes and Mesmer [76BAE/MES], or Ciavatta [80CIA]. By using a more elaborate virial expansion, Pitzer and co-workers [73PIT], [73PIT/MAY], [74PIT/KIM], [74PIT/MAY], [75PIT], [76PIT/SIL], [78PIT/PET], [79PIT] have managed to describe measured activity coefficients of a large number of electrolytes with high precision over a large concentration range. Pitzer’s model gener-ally contains three parameters as compared to one in the specific ion interaction theory. The use of the theory requires knowledge of all these parameters. The derivation of Pitzer coefficients for many complexes, such as those of the actinides would require a very large amount of additional experimental work, since few data of this type are cur-rently available.


The way in which the activity coefficient corrections are performed in this re-view according to the specific ion interaction theory is illustrated below for a general case of a complex formation reaction. Charges are omitted for brevity.

+2M + L + H O(l) M L (OH) + Hm q nm q n n

The formation constant of , ,*M L (OH) , m q n q n mb , determined in an ionic me-

dium (1:1 salt NX) of the ionic strength Im , is related to the corresponding value at zero ionic strength, , ,

*q n mοb by Eq.(B.3).

2

+

10 , , 10 , , 10 M 10 L 10 H O

10 , , 10 H

* *log = log + log + log + log

log log q n m q n m

q n m

m q n a

n

ο

− −

b b g gg g

(B.3)

The subscript (q,n,m) denotes the complex ion, M L (OH)m q n . If the concentra-tions of N and X are much greater than the concentrations of M, L, M L (OH)m q n and H+, only the molalities mN and mX have to be taken into account for the calculation of the term, ( , , )ε∑ m k

kj k I m in Eq. (B.1). For example, for the activity coefficient of the

metal cation M, γM, Eq. (B.4) is obtained at 25°C and 1 bar. 2M

10 M X

0.509log = + (M,X, )

1 + 1.5m

mm

z II m

I−

εg (B.4)

Under these conditions, Im ≈ mX = mN Substituting the log10γ j values in Eq. (B.3) with the corresponding forms of Eq. (B.4) and rearranging leads to:

2

210 , , 10 H O 10 , ,

* *log z log = log q n m q n m mD n a Iο− ∆ − − ∆ εb b (B.5)

where, at 25°C and 1 bar: 2 2 2 2

M L M L = ( ) + z m z q z n n mz q z∆ − − − − (B.6)

0.509 =

1 + 1.5m

m

ID

I (B.7)

= ( , , , N or X) (H, X) (N, L) (M, X)∆ε ε + ε − ε − εq n m n q m (B.8)

Here M L( )m z q z n− − , zM and zL are the charges of the complex, M L (OH)m q n , the metal ion M and the ligand L, respectively.

Equilibria involving H2O(l) as a reactant or product require a correction for the activity of water,

2H Oa . The activity of water in an electrolyte mixture can be calcu-lated as:

2

2

10 H O

H O

log = 1ln(10)

kk

ma

M

−

⋅

∑φ (B.9)

where φ is the osmotic coefficient of the mixture and the summation extends over all solute species k with molality mk present in the solution. In the presence of an ionic


448

medium NX as the dominant species, Eq. (B.9) can be simplified by neglecting the contributions of all minor species, i.e., the reacting ions. Hence, for a 1:1 electrolyte of ionic strength Im ≈ mNX, Eq. (B.9) becomes:

2

2

NX10 H O

H O

2log =

1ln(10)

ma

M

φ−

⋅ (B.10)

Alternatively, water activities can be taken from Table B-1. These have been calculated for the most common ionic media at various concentrations applying Pitzer’s ion interaction model and the interaction parameters given in [91PIT]. Data in italics have been calculated for concentrations beyond the validity of the parameter set applied. These data are therefore extrapolations and should be used with care.

Table B-1: Water activities 2H Oa at 298.15 K for the most common ionic media at vari-

ous concentrations applying Pitzer’s ion interaction approach and the interaction pa-rameters given in [91PIT]. Data in italics have been calculated for concentrations be-yond the validity of the parameter set applied. These data are therefore extrapolations and should be used with care.

Water activities 2H Oa at 298.15 K

c (M) HClO4 NaClO4 LiClO4 NH4ClO4 Ba(ClO4)2 HCl NaCl LiCl

0.10 0.9966 0.9966 0.9966 0.9967 0.9953 0.9966 0.9966 0.9966

0.25 0.9914 0.9917 0.9912 0.9920 0.9879 0.9914 0.9917 0.9915

0.50 0.9821 0.9833 0.9817 0.9844 0.9740 0.9823 0.9833 0.9826

0.75 0.9720 0.9747 0.9713 0.9769 0.9576 0.9726 0.9748 0.9731

1.00 0.9609 0.9660 0.9602 0.9694 0.9387 0.9620 0.9661 0.9631

1.50 0.9357 0.9476 0.9341 0.9542 0.8929 0.9386 0.9479 0.9412

2.00 0.9056 0.9279 0.9037 0.8383 0.9115 0.9284 0.9167

3.00 0.8285 0.8840 0.8280 0.7226 0.8459 0.8850 0.8589

4.00 0.7260 0.8331 0.7309 0.7643 0.8352 0.7991

5.00 0.5982 0.7744 0.6677 0.7782 0.7079

6.00 0.4513 0.7075 0.5592 0.6169



Table B-1: (continued)

c (M) KCl NH4Cl MgCl2 CaCl2 NaBr HNO3 NaNO3 LiNO3

0.10 0.9966 0.9966 0.9953 0.9954 0.9966 0.9966 0.9967 0.9966

0.25 0.9918 0.9918 0.9880 0.9882 0.9916 0.9915 0.9919 0.9915

0.50 0.9836 0.9836 0.9744 0.9753 0.9830 0.9827 0.9841 0.9827

0.75 0.9754 0.9753 0.9585 0.9605 0.9742 0.9736 0.9764 0.9733

1.00 0.9671 0.9669 0.9399 0.9436 0.9650 0.9641 0.9688 0.9635

1.50 0.9500 0.9494 0.8939 0.9024 0.9455 0.9439 0.9536 0.9422

2.00 0.9320 0.9311 0.8358 0.8507 0.9241 0.9221 0.9385 0.9188

3.00 0.8933 0.8918 0.6866 0.7168 0.8753 0.8737 0.9079 0.8657

4.00 0.8503 0.8491 0.5083 0.5511 0.8174 0.8196 0.8766 0.8052

5.00 0.8037 0.3738 0.7499 0.7612 0.8446 0.7390

6.00 0.6728 0.7006 0.8120 0.6696

c (M) NH4NO3 Na2SO4 (NH4)2SO4 Na2CO3 K2CO3 NaSCN

0.10 0.9967 0.9957 0.9958 0.9956 0.9955 0.9966

0.25 0.9920 0.9900 0.9902 0.9896 0.9892 0.9915

0.50 0.9843 0.9813 0.9814 0.9805 0.9789 0.9828

0.75 0.9768 0.9732 0.9728 0.9720 0.9683 0.9736

1.00 0.9694 0.9653 0.9640 0.9637 0.9570 0.9641

1.50 0.9548 0.9491 0.9455 0.9467 0.9316 0.9438

2.00 0.9403 0.9247 0.9283 0.9014 0.9215

3.00 0.9115 0.8735 0.8235 0.8708

4.00 0.8829 0.8050 0.7195 0.8115

5.00 0.8545 0.5887 0.7436

6.00 0.8266 0.6685


450

Values of osmotic coefficients for single electrolytes have been compiled by various authors, e.g., Robinson and Stokes [59ROB/STO]. The activity of water can also be calculated from the known activity coefficients of the dissolved species. In the presence of an ionic medium,

+N X ,

−ν ν of a concentration much larger than those of the reacting ions, the osmotic coefficient can be calculated according to Eq. (B.11) (cf. Eqs. (23−39), (23−40) and (A4−2) in [61LEW/RAN]).

3

NX

ln(10) 11 = 1 2ln(1 )( ) 1

ln(10) (N,X)

j m j mm j j m

A z zB a I B a I

I B a B a I

m

+ −

+ −

+ −

− + − + −

+ ν ν

− ε ν + ν

f (B.11)

where and + −ν ν are the number of cations and anions in the salt formula ( = + + − −ν νz z ) and in this case:

NX1 = ( )2mI z z m+ − + −ν + ν

The activity of water is obtained by inserting Eq. (B.11) into Eq. (B.10). It should be mentioned that in mixed electrolytes with several components at high con-centrations, it might be necessary to use Pitzer’s equation to calculate the activity of water. On the other hand,

2H Oa is nearly constant in most experimental studies of equi-libria in dilute aqueous solutions, where an ionic medium is used in large excess with respect to the reactants. The medium electrolyte thus determines the osmotic coefficient of the solvent.

In natural waters the situation is similar; the ionic strength of most surface wa-ters is so low that the activity of H2O(l) can be set equal to unity. A correction may be necessary in the case of seawater, where a sufficiently good approximation for the os-motic coefficient may be obtained by considering NaCl as the dominant electrolyte.

In more complex solutions of high ionic strengths with more than one electro-lyte at significant concentrations, e.g., (Na+, Mg2+, Ca2+) (Cl−, 2

4SO − ), Pitzer’s equation (cf. [2000GRE/WAN]) may be used to estimate the osmotic coefficient; the necessary interaction coefficients are known for most systems of geochemical interest.

Note that in all ion interaction approaches, the equation for the mean activity coefficients can be split up to give equations for conventional single ion activity coeffi-cients in mixtures, e.g., Eq. (B.1). The latter are strictly valid only when used in combi-nations which yield electroneutrality. Thus, while estimating medium effects on stan-dard potentials, a combination of redox equilibria with, +

21H H (g)2

−+ e , is nec-essary (cf. Example B.3).


B.1.2 Ionic strength corrections at temperatures other than 298.15 K

Values of the Debye-Hückel parameters A and B in Eqs. (B.2) and (B.11) are listed in Table B-2 for a few temperatures at a pressure of 1 bar below 100°C and at the steam saturated pressure for t ≥ 100°C. The values in Table B-2 may be calculated from the static dielectric constant and the density of water as a function of temperature and pres-sure, and are also found for example in Refs. [74HEL/KIR], [79BRA/PIT], [81HEL/KIR], [84ANA/ATK], [90ARC/WAN].

The term, Baj, in the denominator of the Debye-Hückel term, D, cf. Eq. (B.2), has been assigned in this review a value of 1.5 1 1

2 2kg mol−⋅ at 25°C and 1 bar, cf. Sec-tion B.1.1. At temperatures and pressures other than the reference and standard state, the following possibilities exist:

• The value of Baj is calculated at each temperature assuming that ion sizes are independent of temperature and using the values of B listed in Table B-2.

• The value Baj is kept constant at 1.5 1 12 2kg mol−⋅ . Due the variation of B with

temperature, cf. Table B-2, this implies a temperature dependence for ion size parameters. Assuming for the ion size is in reality constant, then it is seen that this simplification introduces an error in D, which increases with temperature and ionic strength (this error is less than ± 0.01 at t ≤ 100°C and I < 6 m, and less than ± 0.006 at t ≤ 50°C and I ≤ 4 m).

• The value of Baj is calculated at each temperature assuming a given temperature variation for aj and using the values of B listed in Table B-2. For example, in the aqueous ionic model of Helgeson and co-workers ([88TAN/HEL], [88SHO/HEL], [89SHO/HEL], [89SHO/HEL2]) ionic sizes follow the relation:

( ) = (298.15 K, 1 bar) + ( , )j j ja T a z g T p [90OEL/HEL], where g(T, p) is a temperature and pressure function which is tabulated in [88TAN/HEL], [92SHO/OEL], and is approximately zero at temperatures below 175°C.

The values of ( , , )ε mj k I , obtained with the methods described in Section B.1.3 at temperatures other than 25°C, will depend on the value adopted for Baj

.. As long as a consistent approach is followed, values of ( , , )ε mj k I absorb the choice of Baj, and for moderate temperature intervals (between 0 and 200°C) the choice Baj =

1.5 1 12 2kg mol−⋅ .is the simplest one and is recommended by this review.

The variation of ( , , )ε mj k I with temperature is discussed by Lewis et al. [61LEW/RAN], Millero [79MIL], Helgeson et al. [81HEL/KIR], [90OEL/HEL], Gif-faut et al. [93GIF/VIT2] and Grenthe and Plyasunov [97GRE/PLY]. The absolute val-ues for the reported ion interaction parameters differ in these studies due to the fact that the Debye-Hückel term used by these authors is not exactly the same. Nevertheless, common to all these studies is the fact that values of ( / )∂ ε ∂ pT are usually ≤ 0.005 kg·mol−1·K−1 for temperatures below 200°C. Therefore, if values of ( , , )ε mj k I obtained


452

at 25°C are used in the temperature range 0 to 50°C to perform ionic strength correc-tions, the error in 10log /j mIg will be ≤ 0.13. It is clear that in order to reduce the un-certainties in solubility calculations at t ≠ 25°C, studies on the variation of ( , , )ε mj k I values with temperature should be undertaken.

Table B-2: Debye-Hückel constants as a function of temperature at a pressure of 1 bar below 100°C and at the steam saturated pressure for t ≥ 100°C. The uncertainty in the A parameter is estimated by this review to be ± 0.001 at 25°C, and ± 0.006 at 300°C, while for the B parameter the estimated uncertainty ranges from ± 0.0003 at 25°C to ± 0.001 at 300°C.

t(°C) p(bar) A ( 1 12 2kg mol−⋅ ) B × 10−10 ( 1 1

2 2 1kg mol mol− −⋅ ⋅ )

0 1.00 0.491 0.3246

5 1.00 0.494 0.3254

10 1.00 0.498 0.3261

15 1.00 0.501 0.3268

20 1.00 0.505 0.3277

25 1.00 0.509 0.3284

30 1.00 0.513 0.3292

35 1.00 0.518 0.3300

40 1.00 0.525 0.3312

50 1.00 0.534 0.3326

75 1.00 0.564 0.3371

100 1.013 0.600 0.3422

125 2.32 0.642 0.3476

150 4.76 0.690 0.3533

175 8.92 0.746 0.3593

200 15.5 0.810 0.365

250 29.7 0.980 0.379

300 85.8 1.252 0.396


B.1.3 Estimation of ion interaction coefficients

B.1.3.1 Estimation from mean activity coefficient data Example B.1:

The ion interaction coefficient +(H ,Cl )−ε can be obtained from published values of ±, HCl HCl versus mg :

+

+

10 , HCl 10 10H Cl+ +

Cl H+

10 , HCl HCl

2 log = log log

= + (H ,Cl ) (Cl ,H )

log = + (H ,Cl )

D m D m

D m

−

−

±

− −

−±

+

− ε − + ε

− ε

g g g

g

By plotting 10 ,HCl(log )D± +g versus mHCl a straight line with the slope +(H ,Cl )−ε is obtained. The degree of linearity should in itself indicate the range of

validity of the specific ion interaction approach. Osmotic coefficient data can be treated in an analogous way.

B.1.3.2 Estimations based on experimental values of equilibrium constants at different ionic strengths

Example B.2:

Equilibrium constants are given in Table B-3 for the reaction: 2+ +2 2UO + Cl UO Cl− (B.12)

The following formula is deduced from Eq. (B.5) for the extrapolation to I = 0:

10 1, 10 1,log + 4 = logm m mD Iο − ∆εb b (B.13)

The linear regression is done as described in Appendix C. The following re-sults are obtained:

10 1,log = (0.170 0.021)mο ±b

∆ε(B.12) = 1(0.248 0.022) kg mol−− ± ⋅ .

The experimental data are depicted in Figure B-1, where the dashed area represents the uncertainty range that is obtained by using the results in 10 1log οb and ∆ε and correcting back to I ≠ 0.


454

Table B-3: The preparation of the experimental equilibrium constants for the extrapola-tion to I = 0 with the specific ion interaction method at 25°C and 1 bar, according to reaction (B.12). The linear regression of this set of data is shown in Figure B-1.

Ic Im 10 1log (exp)b (a) 10 1,log mb (b)

10 1,log 4+m Db

0.100 0.101 − (0.17 ± 0.10) − 0.173 (0.265 ± 0.100)

0.200 0.202 − (0.25 ± 0.10) − 0.254 (0.292 ± 0.100)

0.254 0.258 − (0.35 ± 0.04) − 0.356 (0.230 ± 0.040)

0.303 0.308 − (0.39 ± 0.04) − 0.398 (0.219 ± 0.040)

0.403 0.412 − (0.41 ± 0.04) − 0.419 (0.246 ± 0.040)

0.500 0.513 − (0.32 ± 0.10) − 0.331 (0.372 ± 0.100)

0.552 0.569 − (0.42 ± 0.04) − 0.433 (0.288 ± 0.040)

0.651 0.674 − (0.34 ± 0.04) − 0.355 (0.394 ± 0.040)

0.852 0.886 − (0.42 ± 0.04) − 0.437 (0.358 ± 0.040)

1.000 1.052 − (0.31 ± 0.10) − 0.332 (0.491 ± 0.100)

1.000 1.052 − (0.277 ± 0.260) − 0.298 (0.524 ± 0.260)

1.500 1.617 − (0.24 ± 0.10) − 0.273 (0.618 ± 0.100)

2.000 2.212 − (0.15 ± 0.10) − 0.194 (0.743 ± 0.100)

2.000 2.212 − (0.12 ± 0.10) − 0.164 (0.773 ± 0.100)

2.500 2.815 − (0.06 ± 0.10) − 0.112 (0.860 ± 0.100)

3.000 3.503 (0.04 ± 0.10) − 0.027 (0.973 ± 0.100)

(a) Equilibrium constants for reaction (B.12) with assigned uncertainties, corrected to 25°C where necessary.

(b) Equilibrium constants corrected from molarity to molality units, as described in Section II.2


Figure B-1: Plot of 10 1,log 4m D+b versus Im for reaction (B.12), at 25°C and 1 bar.

The straight line shows the result of the weighted linear regression, and the dotted lines represent the uncertainty range obtained by propagating the resulting uncertainties at I = 0 back to I = 4 m.

Example B.3:

When using the specific ion interaction theory, the relationship between the redox po-tential of the couple, 2+ 4+

2PuO /Pu , in a medium of ionic strength, Im, and the corre-sponding quantity at I = 0 should be calculated in the following way. The reaction in the galvanic cell:

Pt H2(g, r) H+(r) 2+ 4+ +2 2PuO , Pu , H , H O(l) Pt

is:

2+ + + 4+2 2 2PuO + H (g, r) + 4H 2 H (r) Pu 2H O(l)− + (B.14)

where "r" is used to indicate that H2(g) and H+ are at the chemical conditions in the ref-erence electrode compartment, i.e., standard conditions when the reference electrode is the SHE. However, activities of H+, H2O(l) and the ratio of activity of 2+

2PuO to 4+Pu

depend on the conditions of the experimental measurements (i.e., non-standard condi-tions, usually high ionic strength to improve the accuracy of the measurement).

0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.00.0

0.2

0.4

0.6

0.8

1.0

1.2

1.4

log10

β1

0 = (0.17 ± 0.02)∆ε = - (0.25 ± 0.02)

log 10

β1+

4D

Ionic strength, molal

10 1log b ο = (0.17 ± 0.02)

∆ε = – (0.25 ± 0.02) kg·mol–1

I (mol · kg–1)

log 1

0 β1,

m +

4 D


456

For reaction (B.14):

4+ +2

2+ + 22

2H OPu H (r)

10 10 4H (r)PuO H

log = loga a a

Ka a f

ο ⋅ ⋅ ⋅ ⋅

.

Since by definition of the SHE, +2H (r) H (r)1 and 1= γ =f ,

4+ 2+ + 2210 10 10 10 10 10 H OPu PuO H

log = log + log log 4log + 2 logK K aο − −g g g ,

and 4+4

4+10 4Pu ClO

log = 16 + (Pu ,ClO )D m −−− εg

2+2 4

2+10 2 4PuO ClO

log = 4 + (PuO ,ClO )D m −−− εg

+4

+10 4H ClO

log = + (H ,ClO )D m −−− εg

Hence,

24

4+ 2+10 10 4 2 4

+4 10 H OClO

log = log 8 + ( (Pu , ClO ) (PuO , ClO )

4 (H , ClO ) ) 2log

K K D

m a−

ο − −

−

− ε − ε

− ε + (B.15)

The relationship between the equilibrium constant and the redox potential is:

ln = Rn FK E

Tο′ (B.16)

ln = .R

ο οn FK ET

(B.17)

Eο′ is the redox potential in a medium of ionic strength I, οE is the corre-sponding standard potential at I = 0, and n is the number of transferred electrons in the reaction considered. Combining Eqs. (B.15), (B.16) and (B.17) and rearranging them leads to Eq.(B.18):

2 410 H O ClO

R ln(10) R ln(10)(8 2log ) T TE D a E mn F n F

−ο ο′

− − = − ∆ε

(B.18)

For n = 2 in the present example and T = 298.15 K, Eq.(B.18) becomes:

2 410 H O ClO[mV] 236.6 59.16log = [mV] 29.58 E D a E m −

ο ο′ − + − ∆ε

where 4+ 2+ +

4 2 4 4 = (Pu , ClO ) (PuO , ClO ) 4 (H , ClO )− − −∆ε ε − ε − ε .

The value of 2H Oa can be taken from experimental data or calculated from

equations (B.10) and (B.11).

In general, formal potentials are reported with reference to the standard hydro-gen electrode, cf. Section II.1.6.5, as exemplified in Tables V.2 and V.3 of the uranium


NEA review [92GRE/FUG]. In that case, the H+ appearing in the reduction reaction is already at standard conditions. For example, experimental data are available on the for-mal potentials for reactions:

2+ + 4+2 2PuO 4H + 2 Pu 2H O(l)−+ +e (B.19)

and

2+ +2 2PuO PuO−+ e . (B.20)

While reaction (B.19) corresponds to (B.14), reaction (B.20) is equivalent to:

2+ + +2 2 2

1PuO H (g) PuO H2

+ + (B.21)

where the designator "(r)" has been omitted, since in these equations only the H+ in the reference compartment is relevant.

The cations in reaction (B.14) represent aqueous species in the ionic media used during the experiments. In reaction (B.21) H+ represents the cation in the standard hydrogen electrode, and therefore it is already in standard conditions, and its activity coefficient must not be included in any extrapolation to I = 0 of experimental values for Reaction (B.20). Reactions (B.20) and (B.21) are equivalent, as are reactions (B.14) and (B.19), as can be seen if any of these equations are combined with reaction (II.26). Hence Eq. (B.18) can be obtained more simply by using Eq. II.33 for reaction (B.19).

B.1.4 On the magnitude of ion interaction coefficients Ciavatta [80CIA] made a compilation of ion interaction coefficients for a large number of electrolytes. Similar data for complexations of various kinds were reported by Spahiu [83SPA] and Ferri et al. [83FER/GRE]. These and some other data for 25°C and 1 bar have been collected and are listed in Section B.3.

It is obvious from the data in these tables that the charge of an ion is of great importance for determining the magnitude of the ion interaction coefficient. Ions of the same charge type have similar ion interaction coefficients with a given counter-ion. Based on the tabulated data, Grenthe et al. [92GRE/FUG] proposed that it is possible to estimate, with an error of at most ± 0.1 kg · mol−1 in ε, ion interaction coefficients for cases where there are insufficient experimental data for an extrapolation to I = 0. The error that is made by this approximation is estimated to ± 0.1 kg · mol−1 in ∆ε in most cases, based on comparison with ∆ε values of various reactions of the same charge type.

Ciavatta [90CIA] has proposed an alternative method to estimate values of ε for a first or second complex, ML or ML2, in an ionic media NX, according to the fol-lowing relationships:

(ML,N or X)ε ≈ ( )(M,X) (L,N) / 2ε + ε (B.22)


458

2(ML ,N or X) ε ≈ ( )(M,X) 2 (L,N) / 3ε + ε (B.23)

Ciavatta obtained [90CIA] an average deviation of ± 0.05 kg · mol−1 between ε estimates according to Eqs. (B.22) and (B.23), and the ε values at 25°C obtained from ionic strength dependency of equilibrium constants.

B.2 Ion interaction coefficients versus equilibrium con-stants for ion pairs

It can be shown that the virial type of activity coefficient equations and the ionic pairing model are equivalent provided that the ionic pairing is weak. In these cases the distinc-tion between complex formation and activity coefficient variations is difficult or even arbitrary unless independent experimental evidence for complex formation is available, e.g., from spectroscopic data, as is the case for the weak uranium(VI) chloride com-plexes. It should be noted that the ion interaction coefficients evaluated and tabulated by Ciavatta [80CIA] were obtained from experimental mean activity coefficient data with-out taking into account complex formation. However, it is known that many of the metal ions listed by Ciavatta form weak complexes with chloride and nitrate ion. This fact is reflected by ion interaction coefficients that are smaller than those for the non-complexing perchlorate ion, cf. Table B-4. This review takes chloride and nitrate com-plex formation into account when these ions are part of the ionic medium and uses the value of the ion interaction coefficient, +

4(M ,ClO ),n −ε as a substitute for +(M ,Cl )n −ε and +

3(M , NO )n −ε . In this way, the medium dependence of the activity co-efficients is described with a combination of a specific ion interaction model and an ion pairing model. It is evident that the use of NEA recommended data with ionic strength correction models that differ from those used in the evaluation procedure can lead to inconsistencies in the results of the speciation calculations.

It should be mentioned that complex formation may also occur between nega-tively charged complexes and the cation of the ionic medium. An example is the stabili-sation of the complex ion, 5

2 3 3UO (CO ) − , at high ionic strength, see for example Section V.7.1.2.1.d (p. 322) in the uranium review [92GRE/FUG].

B.3 Tables of ion interaction coefficients Table B-4, Table B-5, Table B-6 and Table B-7 contain the selected specific ion interac-tion coefficients used in this review, according to the specific ion interaction theory described. Table B-4 contains cation interaction coefficients with 4 3Cl , ClO and NO− − − , Table B-5 anion interaction coefficients with Li+, Na+ (or +

4NH ) and K+, and Table B-7 neutral species—electroneutral combination of ions. The coefficients have the units of kg·mol−1 and are valid for 298.15 K and 1 bar. The species are ordered by charge and appear, within each charge class, in standard order of arrangement, cf. Section II.1.8.

B.3 Tables of ion interaction coefficients 459

In some cases, the ionic interaction can be better described by assuming ion interaction coefficients as functions of the ionic strength rather than as constants. Ciavatta [80CIA] proposed the use of Eq. (B.24) for cases where the uncertainties in Table B-4 and Table B-5 are ± 0.03 kg·mol−1 or greater.

1 2 10 = + logε ε ε mI (B.24)

For these cases, and when the uncertainty can be improved with respect to the use of a constant value of ε , the values 1ε and 2ε given in Table B-6 should be used.

It should be noted that ion interaction coefficients tabulated in Table B-4, Table B-5 and Table B-6 may also involve ion pairing effects, as described in Section B.3. In direct comparisons of ion interaction coefficients, or when estimates are made by analogy, this aspect must be taken into account.


460

Table B-4: Ion interaction coefficients ( , )ε j k (kg·mol−1) for cations j with k = Cl−, 4ClO− and 3NO− , taken from Ciavatta [80CIA], [88CIA] unless indicated otherwise. The

uncertainties represent the 95% confidence level. The ion interaction coefficients marked with † can be described more accurately with an ionic strength dependent func-tion, listed in Table B-6. The coefficients +(M ,Cl )−ε n and +

3(M , NO )−ε n reported by Ciavatta [80CIA] were evaluated without taking chloride and nitrate complexation into account, as discussed in Section B. 2.

j k → ↓ Cl− 4ClO− 3NO−

+H (0.12 ± 0.01) (0.14 ± 0.02) (0.07 ± 0.01) +4NH − (0.01 ± 0.01) − (0.08 ± 0.04)† − (0.06 ± 0.03)†

+2H gly − (0.06 ± 0.02) +Tl − (0.21 ± 0.06)†

+3ZnHCO 0.2(a)

+CdCl (0.25 ± 0.02) +CdI (0.27 ± 0.02)

+CdSCN (0.31 ± 0.02) +HgCl (0.19 ± 0.02)

+Cu (0.11 ± 0.01) +Ag (0.00 ± 0.01) − (0.12 ± 0.05)†

+NiOH − (0.01 ± 0.07)(K) (0.14 ± 0.07)(L) +NiF (0.34 ± 0.08)(M)

+NiCl (0.47 ± 0.06)(N) +3NiNO (0.44 ± 0.14)(O)

+NiBr (0.59 ± 0.10)(P) +NiHS − (0.85 ± 0.39)(Q)

+NiSCN (0.31 ± 0.04)(R)

3YCO+ (0.17 ± 0.04)(b)

2Am(OH)+ − (0.27 ± 0.20)(q) (0.17 ± 0.04)(c)

2AmF+ (0.17 ± 0.04)(c)

4AmSO+ (0.22 ± 0.08)(d)

3AmCO+ (0.01 ± 0.05)(r) (0.17 ± 0.04)(c)

2PuO+ (0.24 ± 0.05)(e)

2PuO F+ (0.29 ± 0.11)(f)

2PuO Cl+ (0.50 ± 0.09)(g)



Table B-4 (continued)


2NpO+ (0.09 ± 0.05) (0.25 ± 0.05)(h)

2NpO OH+ − (0.06 ± 0.40)(i)

2 3 5(NpO ) (OH)+ (0.45 ± 0.20)

2NpO F+ (0.29 ± 0.12)(j)

2NpO Cl+ (0.50 ± 0.14)(k)

2 3NpO IO+ (0.33 ± 0.04)(l)

3Np(SCN)+ (0.17 ± 0.04)(m)

2UO+ (0.26 ± 0.03)(n)

2UO OH+ − (0.06 ± 0.40)(n) (0.51 ± 1.4)(n)

2 3 5(UO ) (OH)+ (0.81 ± 0.17)(n) (0.45 ± 0.15)(n) (0.41 ± 0.22)(n)

3UF+ (0.1 ± 0.1)(o) (0.1 ± 0.1)(o)

2UO F+ (0.04 ± 0.07)(p) (0.28 ± 0.04)

2UO Cl+ (0.33 ± 0.04)(n)

2 3UO ClO+ (0.33 ± 0.04)(o)

2UO Br+ (0.24 ± 0.04)(o)

2 3UO BrO+ (0.33 ± 0.04)(o)

2 3UO IO+ (0.33 ± 0.04)(o)

2 3UO N+ (0.3 ± 0.1)(o)

2 3UO NO+ (0.33 ± 0.04)(o)

2UO SCN+ (0.22 ± 0.04)(o) 2+Pb (0.15 ± 0.02) − (0.20 ± 0.12)†

2+AlOH 0.09(s) 0.31(s) 2

2 3 2Al CO (OH) + 0.26(s) 2+Zn (0.33 ± 0.03) (0.16 ± 0.02)

23ZnCO + (0.35 ± 0.05)(a)

2+Cd (0.09 ± 0.02) 2+Hg (0.34 ± 0.03) − (0.1 ± 0.1)†

2+2

Hg (0.09 ± 0.02) − (0.2 ± 0.1)†

2+Cu (0.08 ± 0.01) (0.32 ± 0.02) (0.11 ± 0.01)

2+Ni (0.17 ± 0.02) (0.370 ± 0.032)(S) (0.182 ± 0.010)(T)

2+Co (0.16 ± 0.02) (0.34 ± 0.03) (0.14 ± 0.01)



462



2FeOH + 0.38(b) 2FeSCN + 0.45(b)

2+Mn (0.13 ± 0.01)

23YHCO + (0.39 ± 0.04)(b) 2AmOH + − (0.04 ± 0.07)(q) (0.39 ± 0.04)(c)

2AmF + (0.39 ± 0.04)(c) 2AmCl + (0.39 ± 0.04)(c)

23AmN + (0.39 ± 0.04)(c)

22AmNO + (0.39 ± 0.04)(c) 23AmNO + (0.39 ± 0.04)(c)

22 4AmH PO + (0.39 ± 0.04)(c)

2AmSCN + (0.39 ± 0.04)(c) 22PuO + (0.46 ± 0.05)(t) 22PuF + (0.36 ± 0.17)(j)

2PuCl + (0.39 ± 0.16)(u) 2PuI + (0.39 ± 0.04)(v)

22NpO + (0.46 ± 0.05)(w)

22 2 2(NpO ) (OH) + (0.57 ± 0.10)

22NpF + (0.38 ± 0.17)(j)

24NpSO + (0.48 ± 0.11)

22Np(SCN) + (0.38 ± 0.20)(j)

22UO + (0.21 ± 0.02)(x) (0.46 ± 0.03) (0.24 ± 0.03)(v)

22 2 2(UO ) (OH) + (0.69 ± 0.07)(n) (0.57 ± 0.07)(n) (0.49 ± 0.09)(n)

22 3 4(UO ) (OH) + (0.50 ± 0.18)(n) (0.89 ± 0.23)(n) (0.72 ± 1.0)(n)

22UF + (0.3 ± 0.1)(o)

24USO + (0.3 ± 0.1)(o)

23 2U(NO ) + (0.49 ± 0.14)(y)

2Mg + (0.19 ± 0.02) (0.33 ± 0.03) (0.17 ± 0.01) 2Ca + (0.14 ± 0.01) (0.27 ± 0.03) (0.02 ± 0.01) 2Ba + (0.07 ± 0.01) (0.15 ± 0.02) − (0.28 ± 0.03)





3Al + (0.33 ± 0.02) 3

2Ni OH + (0.59 ± 0.15)(U) 3Fe + (0.56 ± 0.03) (0.42 ± 0.08) 3Cr + (0.30 ± 0.03) (0.27 ± 0.02) 3La + (0.22 ± 0.02) (0.47 ± 0.03) 3 3La Lu+ +→ 0.47 → 0.52(b)

3Am + (0.23 ± 0.02)(G) (0.49 ± 0.03)(c) 3Pu + (0.49 ± 0.05)(z)

3PuOH + (0.50 ± 0.05)(i) 3PuF + (0.56 ± 0.11)(i)

3PuCl + (0.85 ± 0.09)(#) 3PuBr + (0.58 ± 0.16)(A)

3PuSCN + (0.39 ± 0.04)(B) 3Np + (0.49 ± 0.05)(z)

3NpOH + (0.50 ± 0.05)(i) 3NpF + (0.58 ± 0.07)(C)

3NpCl + (0.81 ± 0.09)(D) 3NpI + (0.77 ± 0.26)(E)

3NpSCN + (0.76 ± 0.12)(j) 3U + (0.49 ± 0.05)(y)

3UOH + (0.48 ± 0.08)(y) 3UF + (0.48 ± 0.08)(o)

3UCl + (0.50 ± 0.10)(k) 3UBr + (0.52 ± 0.10)(o)

3UI + (0.55 ± 0.10)(o) 33UNO + (0.62 ± 0.08)(y)

32Be OH + (0.50 ± 0.05)(F)

33 3Be (OH) + (0.30 ± 0.05)(F) (0.51 ± 0.05)(y) (0.29 ± 0.05)(F)

43 3 4Al CO (OH) + 0.41(s)

44 4Ni (OH) + (1.08 ± 0.08)(V)

42 2Fe (OH) + 0.82(b)



464

Table B-4: (continued)


42 3Y CO + (0.80 ± 0.04)(b)

4Pu + (0.82 ± 0.07)(H) 4Np + (0.84 ± 0.06)(I)

4U + (0.76 ± 0.06)(J)(o) 4Th + (0.25 ± 0.03) (0.11 ± 0.02)

53 4Al (OH) + 0.66(s) 1.30(s)

(a) Taken from Ferri et al. [85FER/GRE].

(b) Taken from Spahiu [83SPA].

(c) Estimated in [95SIL/BID].

(d) Evaluated in [95SIL/BID].

(e) Derived from 22 4 2 4 = (PuO ,ClO ) (PuO ,ClO )+ − + −∆ε ε − ε = (0.22 ± 0.03) kg·mol−1 [95CAP/VIT]. In

[92GRE/FUG], 2 4(PuO ,ClO )+ −ε = (0.17 ± 0.05) kg·mol−1 was tabulated based on [89ROB], [89RIG/ROB] and [90RIG]. Capdevila and Vitorge’s data [92CAP], [94CAP/VIT] and [95CAP/VIT] were unavailable at that time.

(f) Estimated in [2001LEM/FUG] by analogy with ∆ε of the corresponding Np(IV) reaction.

(g) From ∆ε evaluated by Giffaut [94GIF].

(h) As in [92GRE/FUG], derived from ∆ε = 22 4 2 4 (NpO ,ClO ) (NpO ,ClO )+ − + −ε − ε = (0.21 ± 0.03) kg·mol−1

[87RIG/VIT], [89RIG/ROB] and [90RIG].

(i) Estimated in [2001LEM/FUG].

(j) Estimated in [2001LEM/FUG] by analogy with ∆ε of the corresponding U(IV) reaction.

(k) Estimated in [2001LEM/FUG] by analogy with ∆ε of the corresponding P(VI) reaction.

(l) Estimated in [2001LEM/FUG] by assuming 2 3 4 2 3 4 (NpO IO ,ClO ) (UO IO ,ClO )+ − + −≈ε ε .

(m) Estimated in [2001LEM/FUG] by assuming 3 4 2 4 (Np(SCN) ,ClO ) (AmF ,ClO )+ − + −≈ε ε .

(n) Evaluated in the uranium review [92GRE/FUG], using 22(UO , X)+ε = (0.46 ± 0.03) kg·mol−1, where X =

Cl−, 4ClO− and 3NO− .

(o) Estimated in the uranium review [92GRE/FUG].

(p) Taken from Riglet et al. [89RIG/ROB], where the following assumptions were made: 3 3

4 4(Np ,ClO ) (Pu ,ClO )+ − + −ε ≈ ε = 0.49 kg·mol−1 as for other (M3+, 4ClO− ) interactions, and 2 2 22 4 2 4 2 4 (NpO ,ClO ) (PuO ,ClO ) (UO ,ClO )+ − + − + −≈ ≈ε ε ε = 0.46 kg·mol−1.

(q) Evaluated in [2003GUI/FAN] (cf. Section 12.3.1.1) from ∆ε (in NaCl solution) for the reactions 3+ (3 ) +

2An + H O(l) An(OH) + Hnnn n− .

(r) Evaluated in [2003GUI/FAN] (cf. Section 12.6.1.1.1) from ∆ε (in NaCl solution) for the reactions 3+ 2 (3 )

3 3An + CO An(CO ) nnn − − (based on 3(Am ,Cl )+ −ε = (0.23 ± 0.02) kg·mol−1 and 2

3(Na ,CO )+ −ε = − (0.08 ± 0.03) kg·mol−1.

(s) Taken from Hedlund [88HED].

(t) By analogy with 2+2 4(UO ,ClO )−ε as derived from isopiestic measurements in [92GRE/FUG]. The uncer-

tainty is increased because the value is estimated by analogy.


(u) Estimated in [2001LEM/FUG] by analogy with ∆ε of the corresponding Am(III) reaction.

(v) Estimated in [2001LEM/FUG] by assuming 2+4(PuI ,ClO )−ε ≈ 2+

4(AmSCN ,ClO )−ε and +4(I , NH )−ε ≈

+(SCN , Na )−ε .

(w) By analogy with 2+2 4(UO ,ClO )−ε as derived from isopiestic measurements noted in [92GRE/FUG]. The

uncertainty is increased because the value is estimated by analogy.

(x) These coefficients were not used in the NEA-TDB uranium review [92GRE/FUG] because they were evaluated by Ciavatta [80CIA] without taking chloride and nitrate complexation into account. Instead, Grenthe et al. used 2+

2(UO ,X)ε = (0.46 ± 0.03) kg·mol−1, for X = Cl−, 4ClO− and 3NO− .

(y) Evaluated in the uranium review [92GRE/FUG] using 4+(U ,X)ε = (0.76 ± 0.06) kg·mol−1.

(z) Estimated by analogy with 3+4(Ho ,ClO )−ε [83SPA] as in previous books in this series [92GRE/FUG],

[95SIL/BID]. The uncertainty is increased because the value is estimated by analogy.

(#) Derived from the ∆ε evaluated in [2001LEM/FUG].

(A) Estimated in [2001LEM/FUG] by analogy with ∆ε of the corresponding U(IV) reaction, and by assum-ing + +(Br , H ) (Br , Na )− −ε ≈ ε .

(B) Estimated in [2001LEM/FUG] by assuming 2+4(PuSCN ,ClO )−ε ≈ 2+

4(AmSCN ,ClO )−ε .

(C) Evaluated in [2001LEM/FUG].

(D) Derived from the ∆ε selected in [2001LEM/FUG].

(E) Estimated in [2001LEM/FUG] by analogy with ∆ε of the corresponding Np(IV) chloride reaction, and by assuming + +(I , H ) (I , Na )− −ε ≈ ε .

(F) Taken from Bruno [86BRU], where the following assumptions were made: 24(Be ,ClO )+ −ε = 0.30

kg·mol−1 as for other 2 24(M ,ClO ); (Be ,Cl )+ − + −ε ε = 0.17 kg·mol−1 as for other 2(M ,Cl )+ −ε and

23(Be , NO )+ −ε = 0.17 kg·mol−1 as for other 2

3(M , NO )+ −ε .

(G) The ion interaction coefficient 3+(An ,Cl )−ε for An = Am and Cm is assumed to equal to 3(Nd ,Cl )+ −ε which is calculated from trace activity coefficients of Nd3+ ion in 0− 4 m NaCl. These trace activity coef-ficients are based on the ion interaction Pitzer parameters evaluated in [97KON/FAN] from osmotic co-efficients in aqueous NdCl3 − NaCl and NdCl3 − CaCl2.

(H) Derived from 4+ 3+4 4∆ = (Pu ,ClO ) (Pu ,ClO )− −ε ε − ε = (0.33 ± 0.035) kg·mol−1 [95CAP/VIT]. Uncertainty

estimated in [2001LEM/FUG] (see Appendix A). In the first book of this series [92GRE/FUG], 3+

4(Pu ,ClO )−ε = (1.03 ± 0.05) kg·mol–1 was tabulated based on references [89ROB], [89RIG/ROB], [90RIG]. Capdevila and Vitorge’s data [92CAP], [94CAP/VIT] and [95CAP/VIT] were unavailable at that time.

(I) Derived from 4+ 3+4 4∆ = (Np ,ClO ) (Np ,ClO )− −ε ε − ε = (0.35 ± 0.03) kg·mol−1 [89ROB], [89RIG/ROB],

[90RIG].

(J) Using the measured value of 4+ 3+4 4∆ = (U ,ClO ) (U ,ClO )− −ε ε − ε = (0.35 ± 0.06) kg·mol−1 p.89 [90RIG],

where the uncertainty is recalculated in [2001LEM/FUG] from the data given in this thesis, and 3+

4(U ,ClO )−ε = (0.49 ± 0.05) kg·mol−1 (see footnote (y)), a value for 4+4(U ,ClO )−ε can be calculated in

the same way as is done for 4+4(Np ,ClO )−ε and 4+

4(Pu ,ClO )−ε . This value, 4+4(U ,ClO )−ε = (0.84 ± 0.06)

kg·mol−1 is consistent with that tabulated 4+4(U ,ClO )−ε = (0.76 ± 0.06) kg·mol–1, since the uncertainties

overlap. The authors of [2001LEM/FUG] do not believe that a change in the previously selected value for 4+

4(U ,ClO )−ε is justified at present.

(K) Evaluated in Section V.3.1.1 from ∆ε in chloride media for the reaction Ni2+ + H2O NiOH+ + H+.

(L) Evaluated in Section V.3.1.1 from ∆ε in perchlorate media for the reaction Ni2+ + H2O NiOH+ + H+.

(M) Derived from + 2+ +4 4∆ = (NiF ,ClO ) (Ni ,ClO ) (Na ,F )− − −ε ε − ε − ε = – (0.049 ± 0.060) kg·mol−1 (see Section

V.4.2.3).


466

(N) See details in Section V.4.2.4.

(O) See details in Section V.6.1.2, specially sub-section V.6.1.2.1 for an alternative treatment of this system.

(P) See details in Section V.4.2.5, specially sub-section V.4.2.5.1 for an alternative treatment of this system.

(Q) See details in Section V.5.1.1.2.

(R) Derived from 2+4∆ = (NiSCN , Na ) (SCN , Na ) (Ni ,ClO )− + − + −ε ε − ε − ε = – (0.109 ± 0.025) kg·mol−1 (see

Section V.7.1.3.1).

(S) Derived from the ionic strength dependence of the osmotic and mean activity coefficient of Ni(ClO4)2 solution (Section V.4.3).

(T) Derived from the ionic strength dependence of the osmotic and mean activity coefficient of Ni(NO3)2 solution (Section V.6.1.2.1).

(U) By assuming 3 32 4 2 4(Ni OH ,ClO ) (Be OH ,ClO ),+ − + −ε ≈ ε see Section V.3.1.1.

(V) Derived from + 4 2+4 4 4 4 4∆ = 4 (H ,ClO ) (Ni OH ,ClO ) 4 (Ni ,ClO )− + − −ε ⋅ ε − ε − ⋅ ε = (0.16 ± 0.05) kg·mol−1 (see

Section V.3.1.1.1).


Table B-5: Ion interaction coefficients, ε(j,k) kg · mol–1, for anions j with k = Li, Na and K, taken from Ciavatta [80CIA], [88CIA] unless indicated otherwise. The uncertainties represent the 95% confidence level. The ion interaction coefficients marked with † can be described more accurately with an ionic strength dependent function, listed in Table B-6.

j k → ↓

Li+ Na+ K+

OH− − (0.02 ± 0.03)† (0.04 ± 0.01) (0.09 ± 0.01) F− (0.02 ± 0.02)(a) (0.03 ± 0.02)

2HF− − (0.11 ± 0.06)(a) Cl− (0.10 ± 0.01) (0.03 ± 0.01) (0.00 ± 0.01)

3ClO− − (0.01 ± 0.02)

4ClO− (0.15 ± 0.01) (0.01 ± 0.01) Br− (0.13 ± 0.02) (0.05 ± 0.01) (0.01 ± 0.02)

3BrO− − (0.06 ± 0.02) I− (p) (0.16 ± 0.01) (0.08 ± 0.02) (0.02 ± 0.01)

3IO− − (0.06 ± 0.02)(b)

4HSO− − (0.01 ± 0.02)

3N− (0.0 ± 0.1)(b)

2NO− (0.06 ± 0.04)† (0.00 ± 0.02) − (0.04 ± 0.02)

3NO− (0.08 ± 0.01) − (0.04 ± 0.03)† − (0.11 ± 0.04)†

2 4H PO− − (0.08 ± 0.04)† − (0.14 ± 0.04)†

3HCO− (0.00 ± 0.02)(d) − (0.06 ± 0.05)(i) CN− (0.07 ± 0.03)(x) SCN− (0.05 ± 0.01) − (0.01 ± 0.01) HCOO− (0.03 ± 0.01)

3CH COO− (0.05 ± 0.01) (0.08 ± 0.01) (0.09 ± 0.01)

3SiO(OH)− − (0.08 ± 0.03)(a)

2 2 5Si O (OH)− − (0.08 ± 0.04)(b)

4B(OH)− − (0.07 ± 0.05)†

3Ni(SCN)− (0.66 ± 0.13)(t)

4 2Am(SO )− − (0.05 ± 0.05)(c)

3 2Am(CO )− − (0.14 ± 0.06)(r)

2 3PuO CO− − (0.18 ± 0.18)(o)

2 2NpO (OH)− − (0.01 ± 0.07)(q)

2 3NpO CO− − (0.18 ± 0.15)(f)

2 2 3 3(NpO ) CO (OH)− (0.00 ± 0.05)(k)

2 3UO (OH)− − (0.09 ± 0.05)(b)

2 3UO F− − (0.14 ± 0.05)(w)

2 3 3UO (N )− (0.0 ± 0.1)(b)



468


j k → ↓

Li+ Na+ K+

2 2 3 3(UO ) CO (OH)− (0.00 ± 0.05)(b) 23SO − − (0.08 ± 0.05)† 24SO − − (0.03 ± 0.04)† − (0.12 ± 0.06)† − (0.06 ± 0.02)

22 3S O − − (0.08 ± 0.05)†

24HPO − − (0.15 ± 0.06)† − (0.10 ± 0.06)†

23CO − − (0.08 ± 0.03)(d) (0.02 ± 0.01)

22 2SiO (OH) − − (0.10 ± 0.07)(a)

22 3 4Si O (OH) − − (0.15 ± 0.06)(b)

24Ni(CN) − (0.185 ± 0.081)(u)

24CrO − − (0.06 ± 0.04)† − (0.08 ± 0.04)†

22 4 2NpO (HPO ) − − (0.1 ± 0.1)

22 3 2NpO (CO ) − − (0.02 ± 0.14)(k)

22 4UO F − − (0.30 ± 0.06)(w)

22 4 2UO (SO ) − − (0.12 ± 0.06)(b)

22 3 4UO (N ) − − (0.1 ± 0.1)(b)

22 3 2UO (CO ) − − (0.02 ± 0.09)(d)

22 2 2 4 2(UO ) (OH) (SO ) − − (0.14 ± 0.22)

34PO − − (0.25 ± 0.03)† − (0.09 ± 0.02)

33 6 3Si O (OH) − − (0.25 ± 0.03)(b)

33 5 5Si O (OH) − − (0.25 ± 0.03)(b)

34 7 5Si O (OH) − − (0.25 ± 0.03)(b)

35Ni(CN) − (0.25 ± 0.14)(v)

33 3Am(CO ) − − (0.23 ± 0.07)(r)

33 3Np(CO ) − − (0.15 ± 0.07)(n)

32 3 2NpO (CO ) − − (0.33 ± 0.17)(f)

42 7P O − − (0.26 ± 0.05) − (0.15 ± 0.05)

46Fe(CN) − − (0.17 ± 0.03)

42 3 3NpO (CO ) − − (0.40 ± 0.19)(e) − (0.62 ± 0.42)(g)(h)

42 3 2NpO (CO ) OH − − (0.40 ± 0.19)(m)

43 4U(CO ) − − (0.09 ± 0.10)(b)(d)

42 3 3UO (CO ) − − (0.01 ± 0.11)(d)

42 3 4 4 3(UO ) (OH) (SO ) − (0.6 ± 0.6)

52 3 3NpO (CO ) − − (0.53 ± 0.19)(f) − (0.22 ± 0.03)(s)

52 3 3UO (CO ) − − (0.62 ± 0.15)(d)




j k → ↓

Li+ Na+ K+

63 5Np(CO ) − − (0.73 ± 0.68)(j)

62 3 3 6(NpO ) (CO ) − − (0.46 ± 0.73)(e)

63 5U(CO ) − − (0.30 ± 0.15)(d) − (0.70 ± 0.31)(i)

62 3 3 6(UO ) (CO ) − (0.37 ± 0.11)(d)

62 2 2 3 6(UO ) NpO (CO ) − (0.09 ± 0.71)(l)

62 5 8 4 4(UO ) (OH) (SO ) − (1.10 ± 0.5)

72 4 7 4 4(UO ) (OH) (SO ) − (2.80 ± 0.7)

(a) Evaluated in the NEA-TDB uranium review [92GRE/FUG]. (b) Estimated in the NEA-TDB uranium review [92GRE/FUG]. (c) Estimated in the NEA-TDB americium review [95SIL/BID]. (d) These values differ from those reported in the NEA−TDB uranium review. See the discussion in

[95GRE/PUI]. Values for 23CO − and 3HCO− are based on [80CIA].

(e) Calculated in [2001LEM/FUG] (Section 12.1.2.1.2). (f) Calculated in [2001LEM/FUG] (Section 12.1.2.1.3). (g) Calculated in [2001LEM/FUG] (Section 12.1.2.2.1). (h) 4

2 3 3 4(NpO (CO ) , NH )− +ε = – (0.78 ± 0.25) kg · mol−1 is calculated in [2001LEM/FUG] (Section 12.1.2.2.1).

(i) Calculated in [2001LEM/FUG] from Pitzer coefficients [98RAI/FEL]. (j) Calculated in [2001LEM/FUG] (Section 12.1.2.1.4). (k) Estimated by analogy in [2001LEM/FUG] (Section 12.1.2.1.2). (l) Estimated by analogy in [2001LEM/FUG] (Section 12.1.2.2.1). (m) Estimated in [2001LEM/FUG] by analogy with 4

32 3NpO (CO ) − . (n) Estimated by analogy in [2001LEM/FUG] (Section 12.1.2.1.5). (o) Estimated in [2001LEM/FUG] by analogy with 2 3(NpO CO , Na )− +ε . (p) 4(I , NH ) (SCN , Na )− + − +ε ≈ ε = (0.05 ± 0.01) kg · mol−1. (q) Estimated in [2001LEM/FUG] (Section 8.1.3). (r) Evaluated in [2003GUI/FAN], Section 12.6.1.1.1 from ∆εn (in NaCl solution) for the reactions

3+ 2 (3 )3 3An CO An(CO ) n

nn − −+ (based on 3+(Am ,Cl )−ε = (0.23 ± 0.02) kg · mol−1 and 23(Na , CO )+ −ε = – (0.08 ± 0.03) kg · mol−1.

(s) Evaluated in Appendix A, discussion of [98ALM/NOV] from ∆ε for the reactions 2 3KNpO CO (s) 23+ 2 CO − 5

3+

2 3 NpO (CO ) + K− (in K2CO3−KCl solution) and K3NpO2(CO3)2(s) + 23CO −

52 3 3NpO (CO ) − + 3 K+ (in K2CO3 solution) (based on 2

3(K , CO )+ −ε = (0.02 ± 0.01) kg·mol−1). (t) Evaluated in Section V.7.1.3.1 from ∆εn in perchlorate media for the reaction

2+3Ni 3SCN Ni(SCN)− −+ .

(u) Evaluated in Section V.7.1.2.1.1 from ∆εn in perchlorate media for the reaction 2+ 2

4Ni 4 CN Ni(CN)− −+ . (v) Evaluated in Section V.7.1.2.1.1 from ∆εn in perchlorate media for the reaction

2+ 35Ni 5 CN Ni(CN)− −+ .

(w) Evaluated in [2003GUI/FAN], Section 9.4.2.2.1.1. (x) As reported in [92BAN/BLI].


470

Table B-6: Ion interaction coefficients, ε(1,j,k) and ε(2,j,k) kg · mol–1, for cations j with k = Cl−, 4ClO− and 3NO− (first part), and for anions j with k = Li+, Na+ and K+ (second part), according to the relationship ε = ε1 + ε2 log10Im. The data are taken from Ciavatta [80CIA], [88CIA]. The uncertainties represent the 95% confidence level.

Cl− 4ClO− 3NO− j k → ↓ ε1 ε2 ε1 ε2 ε1 ε2

+4NH −(0.088±0.002) (0.095±0.012) −(0.075±0.001) (0.057±0.004)

+Tl −(0.18±0.02) (0.09±0.02) +Ag −(0.1432±0.0002) (0.0971±0.0009)2+Pb −(0.329±0.007) (0.288±0.018) 2+Hg −(0.145±0.001) (0.194±0.002) 2+2Hg −(0.2300±0.0004) (0.194±0.002)

+Li +Na +K j k → ↓ ε1 ε2 ε1 ε2 ε1 ε2

OH− −(0.039±0.002) (0.072±0.006)

2NO− (0.02±0.01) (0.11±0.01)

3NO− −(0.049±0.001) (0.044±0.002) −(0.131±0.002) (0.082±0.006)

2 4H PO− −(0.109±0.001) (0.095±0.003) −(0.1473±0.0008) (0.121±0.004)

4B(OH)− −(0.092±0.002) (0.103±0.005) 23SO − −(0.125±0.008) (0.106±0.009) 24SO − −(0.068±0.003) (0.093±0.007) −(0.184±0.002) (0.139±0.006)

22 3S O − −(0.125±0.008) (0.106±0.009)

24HPO − −(0.19±0.01) (0.11±0.03) −(0.152±0.007) (0.123±0.016)

24CrO − −(0.090±0.005) (0.07±0.01) −(0.123±0.003) (0.106±0.007)

34PO − −(0.29±0.02) (0.10±0.01)

Table B-7: SIT interaction coefficient ε(j,k) kg · mol–1 for neutral species, j, with k, elec-troneutral combination of ions.

j k → ↓

Na+ + 4ClO−

2Ni(SCN) (aq) (0.38 ± 0.06) see discussion in Section V.7.1.3.1

471

C Appendix C

Assigned uncertainties1

This Appendix describes the origin of the uncertainty estimates that are given in the TDB tables of selected data. The original text in [92GRE/FUG] has been retained in [95SIL/BID], [99RAR/RAN], [2001LEM/FUG] and [2003GUI/FAN] except for some minor changes. Because of the importance of the uncertainty estimates, the present re-view offers a more comprehensive description of the procedures used.

C.1 The general problem The focus of this section is on the uncertainty estimates of equilibria in solution, where the key problem is analytical, i.e., the determination of the stoichiometric composition and equilibrium constants of complexes that are in rapid equilibrium with one another. We can formulate analyses of the experimental data in the following way: From N measurements, yi, of the variable y we would like to determine a set of n equilibrium constants kr, r = 1, 2,…, n, assuming that we know the functional relationship:

y = f(k1, k2, …kr...kn; a1, a2,….) (C.1)

where a1, a2 , etc. are quantities that can be varied but whose values (a1i; a2i; etc.) are assumed to be known accurately in each experiment from the data sets (yi, a1i, a2i,…), i = 1, 2, …N. The functional relationship (C.1) is obtained from the chemical model pro-posed and in general several different models have to be tested before the "best" one is selected. Details of the procedures are given in Rossotti and Rossotti [61ROS/ROS].

When selecting the functional relationship (C.1) and determining the set of equilibrium constants that best describes the experiments one often uses a least-squares method. Within this method, the “best” description is the one that will minimise the residual sum of squares, U:

[ ]21 1 2 ( ... ; , ...)i i n i i

i

U w y f k k a a= −∑ (C.2)

where wi is the weight of each experimental measurement yi.

1 This Appendix essentially contains the text of the TDB-3 Guideline, [99WAN/OST], earlier versions of

which have been printed in the previous NEA TDB reviews [92GRE/FUG], [95SIL/BID], [99RAR/RAN], [2001LEM/FUG] and [2003GUI/FAN]. Because of its importance in the selection of data and to guide the users of the values in Chapters III and IV, the text is reproduced here with minor revisions.

C. Assigned uncertainties

472

The minimum of the function (C.2) is obtained by solving a set of normal equations:

0, 1,..... r

U r nk

∂= =

∂ (C.3)

A "true" minimum is only obtained if:

• the functional relationship (C.1) is correct, i.e., if the chemical model is cor-rect.

• all errors are random errors in the variable y, in particular there are no sys-tematic errors.

• the random errors in y follow a Gaussian (normal) distribution.

• the weight wi(yi, a1i, a2i,….) of an experimental determination is an exact measure of its inherent accuracy.

To ascertain that the first condition is fulfilled requires chemical insight, such as information of the coordination geometry, relative affinity between metal ions and various donor atoms, etc. It is particularly important to test if the chemical equilibrium constants of complexes that occur in small amounts are chemically reasonable. Too many experimentalists seem to look upon the least-squares refinement of experimental data more as an exercise in applied mathematics than as a chemical venture. One of the tasks in the review of the literature is to check this point. An erroneous chemical model is one of the more serious types of systematic errors.

The experimentalist usually selects the variable that he/she finds most appro-priate to fulfill the second condition. If the estimated errors in a1i, a2i … are smaller than the error in yi, the second condition is reasonably well fulfilled. The choice of the error-carrying variable is a matter of choice based on experience, but one must be aware that it has implications, especially in the estimated uncertainty.

The presence of systematic errors is, potentially, the most important source of uncertainty. There is no possibility to handle systematic errors using statistics; statistical methods may indicate their presence, no more. Systematic errors in the chemical model have been mentioned. In addition there may be systematic errors in the methods used. By comparing experimental data obtained with different experimental methods one can obtain an indication of the presence and magnitude of such errors. The systematic errors of this type are accounted for both in the review of the literature and when taking the average of data obtained with different experimental methods. This type of systematic error does not seem to affect the selected data very much, as judged by the usually very good agreement between the equilibrium data obtained using spectroscopic, potenti-ometric and solubility methods.

C.2 Uncertainty estimates in the selected thermodynamic data. 473

The electrode calibration, especially the conversion between measured pH and − log10[H+] is an important source of systematic error. The reviewers have when possi-ble corrected this error, as seen in many instances in Appendix A.

The assumption of a normal distribution of the random errors is a choice made in the absence of better alternatives.

Finally, a comment on the weights used in least-squares refinements; this is important because it influences the uncertainty estimate of the equilibrium constants. The weights of individual experimental points can be obtained by repeating the experi-ment several times and then calculating the average and standard deviation of these data. This procedure is rarely used, instead most experimentalists seem to use unit weight when making a least-squares analysis of their data. However, also in this case there is a weighting of the data by the number of experimental determinations in the parameter range where the different complexes are formed. In order to have comparable uncertainty estimates for the different complexes, one should try to have the same num-ber of experimental data points in the concentration ranges where each of these com-plexes is predominant; a procedure very rarely used.

As indicated above, the assignment of uncertainties to equilibrium constants is not a straightforward procedure and it is complicated further when there is lack of pri-mary experimental data. The uncertainty estimates given for the individual equilibrium constants reported by the authors and for some cases re-estimated by this review are given in the tables of this and previous reviews. The procedure used to obtain these es-timates is given in the original publications and in the Appendix A discussions. How-ever, this uncertainty is still a subjective estimate and to a large extent based on "expert judgment".

C.2 Uncertainty estimates in the selected thermodynamic data.

The uncertainty estimate in the selected thermodynamic data is based on the uncertainty of the individual equilibrium constants or other thermodynamic data, calculated as de-scribed in the following sections. A weighted average of the individual log10K values is calculated using the estimated uncertainty of the individual experimental values to as-sign its weight. The uncertainty in this average is then calculated using the formulae given in the following text. This uncertainty depends on the number of experimental data points − for N data points with the same estimated uncertainty, σ, the uncertainty in the average is / Nσ . The average and the associated uncertainty reported in the tables of selected data are reported with many more digits than justified only in order to allow the users to back-track the calculations. The reported uncertainty is much smaller than the estimated experimental uncertainty and the users of the tables should look at the discussion of the selected constants in order to get a better estimate of the uncertainty in an experimental determination using a specific method.


474

One of the objectives of the NEA Thermochemical Data Base (TDB) project is to provide an idea of the uncertainties associated with the data selected in this review. As a rule, the uncertainties define the range within which the corresponding data can be reproduced with a probability of 95% at any place and by any appropriate method. In many cases, the statistical treatment is limited or impossible due to the availability of only one or few data points. A particular problem has to be solved when significant dis-crepancies occur between different source data. This appendix outlines the statistical procedures, which were used for fundamentally different problems, and explains the philosophy used in this review when statistics were inapplicable. These rules are fol-lowed consistently throughout the series of reviews within the TDB Project. Four fun-damentally different cases are considered:

1. One source datum available

2. Two or more independent source data available

3. Several data available at different ionic strengths

4. Data at non-standard conditions: Procedures for data correction and recalculation.

C.3 One source datum The assignment of an uncertainty to a selected value that is based on only one experi-mental source is a highly subjective procedure. In some cases, the number of data points, on which the selected value is based, allows the use of the “root mean square” [82TAY] deviation of the data points, Xi, to describe the standard deviation, sX, associ-ated with the average, X :

2X

1

1 = ( )1 =

−− ∑

N

ii

s X XN

(C.4)

The standard deviation, sX, is thus calculated from the dispersion of the equally weighted data points, Xi, around the average X , and the probability is 95% that an Xi is within X ± 1.96 sX, see Taylor [82TAY] (pp. 244-245). The standard deviation, sX, is a measure of the precision of the experiment and does not include any systematic errors.

Many authors report standard deviations, sX, calculated with Eq. (C.4) (but often not multiplied by 1.96), but these do not represent the quality of the reported val-ues in absolute terms. Therefore, it is thus important not to confuse the standard devia-tion, sX, with the uncertainty, σ. The latter reflects the reliability and reproducibility of an experimental value and also includes all kinds of systematic errors, sj, that may be involved. The uncertainty, σ, can be calculated with Eq. (C.5), assuming that the sys-tematic errors are independent.

2 2= + ( )X X jj

s sσ ∑ (C.5)

C.4 Two or more independent source data

475

The estimation of the systematic errors sj (which, of course, have to relate to X and be expressed in the same units), can only be made by a person who is familiar

with the experimental method. The uncertainty, σ, has to correspond to the 95% confi-dence level preferred in this review. It should be noted that for all the corrections and recalculations made (e.g., temperature or ionic strength corrections) the rules of the propagation of errors have to be followed, as outlined in Section C.6.2.

More often, the determination of sX is impossible because either only one or two data points are available, or the authors did not report the individual values. The uncertainty σ in the resulting value can still be estimated using Eq. (C.5) assuming that

2Xs is much smaller than 2( )∑ j

js , which is usually the case anyway.

C.4 Two or more independent source data Frequently, two or more experimental data sources are available, reporting experimental determinations of the desired thermodynamic data. In general, the quality of these de-terminations varies widely, and the data have to be weighted accordingly for the calculation of the mean. Instead of assigning weight factors, the individual source data, Xi, are provided with an uncertainty, σi, that also includes all systematic errors and represents the 95% confidence level, as described in Section C.3 The weighted mean X and its uncertainty, σX , are then calculated according to Eqs. (C.6) and (C.7).

21

21

1=

=

σ ≡ σ

∑

∑

Ni

i i

N

i i

X

X (C.6)

21

11

=

σ = σ

∑X N

i i

(C.7)

Eqs. (C.6) and (C.7) may only be used if all the Xi belong to the same parent distribu-tion. If there are serious discrepancies among the Xi, one proceeds as described below under Section C.4.1. It can be seen from Eq. (C.7) that σX is directly dependent on the absolute magnitude of the σi values, and not on the dispersion of the data points around the mean. This is reasonable because there are no discrepancies among the Xi, and be-cause the σi values already represent the 95% confidence level. The selected uncer-tainty, σX , will therefore also represent the 95% confidence level.

In cases where all the uncertainties are equal, σi = σ, Eqs. (C.6) and (C.7) re-duce to Eqs. (C.8) and (C.9).

1

1=

= ∑N

ii

X XN

(C.8)


476

σσ =X N

(C.9)

Example C.1:

Five data sources report values for the thermodynamic quantity, X. The reviewer has assigned uncertainties that represent the 95% confidence level as described in Section C.3.

i Xi σι

1 25.3 0.5

2 26.1 0.4

3 26.0 0.5

4 24.85 0.25

5 25.0 0.6

According to Eqs.(C.6) and (C.7), the following result is obtained:

X = (25.3 ± 0.2).

The calculated uncertainty, σX = 0.2, appears relatively small, but is statisti-cally correct, as the values are assumed to follow a Gaussian distribution. As a conse-quence of Eq. (C.7), σX will always come out smaller than the smallest σi. Assuming σ4 = 0.10 instead of 0.25 would yield X = (25.0 ± 0.1) and σ4 = 0.60 would result in X = (25.6 ± 0.2). In fact, the values (Xi ± σi) in this example are at the limit of consis-

tency, i.e., the range (X4 ± σ4) does not overlap with the ranges (X2 ± σ2) and (X3 ± σ3). There might be a better way to solve this problem. Three possible choices seem more reasonable:

i. The uncertainties, σi, are reassigned because they appear too optimistic after fur-ther consideration. Some assessments may have to be reconsidered and the uncer-tainties reassigned. For example, multiplying all the σi by 2 would yield X = (25.3 ± 0.3).

ii. If reconsideration of the previous assessments gives no evidence for reassigning the Xi and σi (95% confidence level) values listed above, the statistical conclusion will be that all the Xi do not belong to the same parent distribution and cannot therefore be treated in the same group (cf. item iii below for a non−statistical ex-planation). The values for i =1, 4 and 5 might be considered as belonging to Group A and the values for i = 2 and 3 to Group B. The weighted average of the values in Group A is XA (i = 1, 4, 5) = (24.95 ± 0.21) and of those in Group B, XB (i = 2, 3) = (26.06 ± 0.31), the second digit after the decimal point being carried over to avoid loss of information. The selected value is now determined as de-scribed below under “Discrepancies” (Section C.4.1, Case I). XA and XB are aver-

C.4 Two or more independent source data

477

aged (straight average, there is no reason for giving XA a larger weight than XB), and σX is chosen in such a way that it covers the complete ranges of expectancy of XA and XB. The selected value is then X = (25.5 ± 0.9).

iii. Another explanation could be that unidentified systematic errors are associated with some values. If this seems likely to be the case, there is no reason for split-ting the values up into two groups. The correct way of proceeding would be to calculate the unweighted average of all the five points and assign an uncertainty that covers the whole range of expectancy of the five values. The resulting value is then X = (25.45 ± 1.05), which is rounded according to the rules in Section C.6.3 to X = (25.4 ± 1.1).

C.4.1 Discrepancies Two data are called discrepant if they differ significantly, i.e., their uncertainty ranges do not overlap. In this context, two cases of discrepancies are considered. Case I: Two significantly different source data are available. Case II: Several, mostly consistent source data are available, one of them being significantly different, i.e., an “outlier”.

Case I. Two discrepant data: This is a particularly difficult case because the number of data points is obviously insufficient to allow the preference of one of the two values. If there is absolutely no way of discarding one of the two values and selecting the other, the only solution is to average the two source data in order to obtain the se-lected value, because the underlying reason for the discrepancy must be unrecognised systematic errors. There is no point in calculating a weighted average, even if the two source data have been given different uncertainties, because there is obviously too little information to give even only limited preference to one of the values. The uncertainty, σX , assigned to the selected mean, X , has to cover the range of expectation of both source data, X1, X2, as shown in Eq.(C.10),

maxσ = − + σiX X X (C.10)

where i =1, 2, and maxσ is the larger of the two uncertainties σi, see Example C.1.ii and Example C.2.

Example C.2:

The following credible source data are given:

X1 = (4.5 ± 0.3)

X2 = (5.9 ± 0.5).

The uncertainties have been assigned by the reviewer. Both experimental methods are satisfactory and there is no justification to discard one of the data. The se-lected value is then:


478

X = (5.2 ± 1.2).

Figure C-1: Illustration for Example C.2

Case II. Outliers: This problem can often be solved by either discarding the outlying data point, or by providing it with a large uncertainty to lower its weight. If, however, the outlying value is considered to be of high quality and there is no reason to discard all the other data, this case is treated in a way similar to Case I. Example C.3 illustrates the procedure.

Example C.3:

The following data points are available. The reviewer has assigned the uncertainties and sees no justification for any change.

i Xi σi

1 4.45 0.35 2 5.9 0.5 3 5.7 0.4 4 6.0 0.6 5 5.2 0.4

There are two data sets that, statistically, belong to different parent distribu-tions, A and B. According to Eqs. (C.6) and (C.7), the following average values are found for the two groups: XA(i =1) = (4.45 ± 0.35) and XB(i = 2, 3, 4, 5) = (5.62 ± 0.23). The selected value will be the straight average of XA and XB, analogous to Example C.1:

X = (5.0 ± 0.9).

4 4.5 5 5.5 6 6.5

X1 X2

X

X

C.5 Several data at different ionic strengths 479

C.5 Several data at different ionic strengths The extrapolation procedure for aqueous equilibria used in this review is the specific ion interaction model outlined in Appendix B. The objective of this review is to provide selected data sets at standard conditions, i.e., among others, at infinite dilution for aque-ous species. Equilibrium constants determined at different ionic strengths can, according to the specific ion interaction equations, be extrapolated to I = 0 with a linear regression model, yielding as the intercept the desired equilibrium constant at I = 0, and as the slope the stoichiometric sum of the ion interaction coefficients, ∆ε . The ion interaction coefficient of the target species can usually be extracted from ∆ε and would be listed in the corresponding table of Appendix B.

The available source data may sometimes be sparse or may not cover a suffi-cient range of ionic strengths to allow a proper linear regression. In this case, the correc-tion to I = 0 should be carried out according to the procedure described in Section C.6.1.

If sufficient data are available at different ionic strengths and in the same inert salt medium, a weighted linear regression will be the appropriate way to obtain both the constant at I = 0, X ο , and ∆ε . The first step is the conversion of the ionic strength from the frequently used molar (mol·dm−3, M) to the molal (mol·kg−1, m) scale, as described in Section II.2. The second step is the assignment of an uncertainty, σi, to each data point Xi at the molality, mk,i, according to the rules described in Section C.3. A large number of commercial and public domain computer programs and routines exist for weighted linear regressions. The subroutine published by Bevington [69BEV] (pp.104 − 105) has been used for the calculations in the examples of this appendix. Eqs. (C.11) through (C.15) present the equations that are used for the calculation of the intercept X ο and the slope − ∆ε :

2, , ,

2 2 2 21 1 1 1

1 N N N Nk i k i k i ii

i i i ii i i i

m m m XXX ο

= = = =

= − ∆ σ σ σ σ

∑ ∑ ∑ ∑ (C.11)

, ,2 2 2 2

1 1 1 1

1 1N N N Nk i i k i i

i i i ii i i i

m X m X= = = =

−∆ε = − ∆ σ σ σ σ

∑ ∑ ∑ ∑ (C.12)

2,

21

1 N k iX

i i

mο

=σ =

∆ σ∑ (C.13)

21

1 1N

i i∆ε

=σ =

∆ σ∑ (C.14)

where 22

, ,2 2 2

1 1 1

1N N Nk i k i

i i ii i i

m m= = =

∆ = − σ σ σ

∑ ∑ ∑ . (C.15)


480

In this way, the uncertainties, σi, are not only used for the weighting of the data in Eqs. (C.11) and (C.12), but also for the calculation of the uncertainties,

X οσ and ∆εσ , in Eqs. (C.13) and (C.14). If the σi represents the 95% confidence level,

X οσ and ∆εσ will also do so. In other words, the uncertainties of the intercept and the slope do not depend on the dispersion of the data points around the straight line, but rather directly on their absolute uncertainties, σi.

Example C.4:

Ten independent determinations of the equilibrium constant, 10*log b , for the reaction:

2+ + +2 2UO HF(aq) UO F H+ + (C.16)

are available in HClO4/NaClO4 media at different ionic strengths. Uncertainties that represent the 95% confidence level have been assigned by the reviewer. A weighted linear regression, ( 10

*log b + 2D) vs. mk, according to the formula, 10*log b (C.16) +

2D = 10*log οb (C.16) − ∆ε mk, will yield the correct values for the intercept,

10*log οb (C.16), and the slope, ∆ε . In this case, mk corresponds to the molality of

4 ClO− . D is the Debye-Hückel term, cf. Appendix B.

i 4ClO , im − 10

*log b + 2D σi

1 0.05 1.88 0.10

2 0.25 1.86 0.10

3 0.51 1.73 0.10

4 1.05 1.84 0.10

5 2.21 1.88 0.10

6 0.52 1.89 0.11

7 1.09 1.93 0.11

8 2.32 1.78 0.11

9 2.21 2.03 0.10

10 4.95 2.00 0.32

The results of the linear regression are:

intercept = (1.837 ± 0.054) = 10*log οb (C.16)

slope = (0.029 ± 0.036) = − ∆ε

Calculation of the ion interaction coefficient +2 4(UO F ,ClO )−ε = ∆ε +

2+2 4(UO ,ClO )−ε − +

4(H ,ClO )−ε : from 2+2 4(UO ,ClO )−ε = (0.46 ± 0.03) kg·mol−1,

+4(H ,ClO )−ε = (0.14 ± 0.02) kg·mol−1 (see Appendix B) and the slope of the linear re-

gression, ∆ε = − (0.03 ± 0.04) kg·mol−1, it follows that +2 4(UO F ,ClO )−ε =


(0.29 ± 0.05) kg·mol−1. Note that the uncertainty (± 0.05) kg·mol−1 is obtained based on the rules of error propagation as described in Section C.6.2:

2 2 2(0.04) (0.03) (0.02)σ = + +

The resulting selected values are thus:

10*log οb (C.16) = (1.84 ± 0.05)

+2 4(UO F ,ClO )−ε = (0.29 ± 0.05) kg·mol−1.

C.5.1 Discrepancies or insufficient number of data points Discrepancies are principally treated as described in Section C.4. Again, two cases can be defined. Case I: Only two data points are available. Case II: An “outlier” cannot be discarded. If only one data point is available, the procedure for correction to zero ionic strength outlined in Section C.4 should be followed.

Case I. Too few molalities: If only two source data are available, there will be no straightforward way to decide whether or not these two data points belong to the same parent distribution unless either the slope of the straight line is known or the two data refer to the same ionic strength. Drawing a straight line right through the two data points is an inappropriate procedure because all the errors associated with the two source data would accumulate and may lead to highly erroneous values of 10log K ο and ∆ε . In this case, an ion interaction coefficient for the key species in the reaction in question may be selected by analogy (charge is the most important parameter), and a straight line with the slope ∆ε as calculated may then be drawn through each data point. If there is no reason to discard one of the two data points based on the quality of the underlying experiment, the selected value will be the unweighted average of the two standard state data points obtained by this procedure, and its uncertainty must cover the entire range of expectancy of the two values, analogous to Case I in Section C.4. It should be mentioned that the ranges of expectancy of the corrected values at I = 0 are given by their uncertainties, which are based on the uncertainties of the source data at I ≠ 0 and the uncertainty in the slope of the straight line. The latter uncertainty is not an estimate, but is calculated from the uncertainties in the ion interaction coefficients in-volved, according to the rules of error propagation outlined in Section C.6.2. The ion interaction coefficients estimated by analogy are listed in the table of selected ion inter-action coefficients (Appendix B), but they are flagged as estimates.

Case II. Outliers and inconsistent data sets: This case includes situations where it is difficult to decide whether or not a large number of points belong to the same parent distribution. There is no general rule on how to solve this problem, and decisions are left to the judgment of the reviewer. For example, if eight data points follow a straight line reasonably well and two lie way out, it may be justified to discard the “out-liers”. If, however, the eight points are scattered considerably and two points are just a


482

bit further out, one can probably not consider them as “outliers”. It depends on the par-ticular case and on the judgment of the reviewer whether it is reasonable to increase the uncertainties of the data to reach consistency, or whether the slope, ∆ε , of the straight line should be estimated by analogy.

Example C.5:

Six reliable determinations of the equilibrium constant, 10log b , of the reaction:

2+2 2UO + SCN UO SCN− + (C.17)

are available in different electrolyte media:

Ic = 0.1 M (KNO3) 10log b (C.17) = (1.19 ± 0.03)

Ic = 0.33 M (KNO3) 10log b (C.17) = (0.90 ± 0.10)

Ic = 1.0 M (NaClO4) 10log b (C.17) = (0.75 ± 0.03)

Ic = 1.0 M (NaClO4) 10log b (C.17) = (0.76 ± 0.03)

Ic = 1.0 M (NaClO4) 10log b (C.17) = (0.93 ± 0.03)

Ic = 2.5 M (NaNO3) 10log b (C.17) = (0.72 ± 0.03)

The uncertainties are assumed to represent the 95% confidence level. From the values at Ic = 1 M, it can be seen that there is a lack of consistency in the data, and that a linear regression similar to that shown in Example C.4 would be inappropriate. Instead, the use of ∆ε values from reactions of the same charge type is encouraged. Analogies with ∆ε are more reliable than analogies with single ε values due to canceling effects. For the same reason, the dependency of ∆ε on the type of electrolyte is often smaller than for single ε values.

A reaction of the same charge type as Reaction (C.17), and for which ∆ε is well known, is:

2+2 2UO + Cl UO Cl− + . (C.18)

The value of ∆ε (C.18) = − (0.25 ± 0.02) kg · mol–1 was obtained from a linear regression using 16 experimental values between Ic = 0.1 M and Ic = 3 M Na(Cl,ClO4) [92GRE/FUG]. It is thus assumed that:

∆ε (C.17) = ∆ε (C.18) = − (0.25 ± 0.02) kg · mol–1.

The correction of 10log b (C.17) to Ic = 0 is done using the specific ion interac-tion equation, cf. TDB-2, which uses molal units:

10log b + 4D = 10log οb − ∆ε Im. (C.19)

D is the Debye-Hückel term in molal units and Im the ionic strength converted to molal units by using the conversion factors listed in Table II-5. The following list


gives the details of this calculation. The resulting uncertainties in 10log b are obtained based on the rules of error propagation as described in Section C.6.2.

Table C-1: Details of the calculation of equilibrium constant corrected to I = 0, using (C.19).

Im electrolyte 10log b 4D ∆ε Im. 10log οb

0.101 KNO3 (1.19 ± 0.03) 0.438 − 0.025 (1.68 ± 0.03)(a)

0.335 KNO3 (0.90 ± 0.10) 0.617 − 0.084 (1.65 ± 0.10)(a)

1.050 NaClO4 (0.75 ± 0.03) 0.822 − 0.263 (1.31 ± 0.04)

1.050 NaClO4 (0.76 ± 0.03) 0.822 − 0.263 (1.32 ± 0.04)

1.050 NaClO4 (0.93 ± 0.03) 0.822 − 0.263 (1.49 ± 0.04)

2.714 NaNO3 (0.72 ± 0.03) 0.968 − 0.679 (1.82 ± 0.13)(a)

(a) These values were corrected for the formation of the nitrate complex, +2 3UO NO , by using

+10 2 3log (UO NO )K = (0.30 ± 0.15) [92GRE/FUG].

As was expected, the resulting values, 10log οb , are inconsistent and have therefore to be treated as described in Case I of Section C.4. That is, the selected value will be the unweighted average of 10log οb , and its uncertainty will cover the entire range of expectancy of the six values. A weighted average would only be justified if the six values of 10log οb were consistent. The result is:

10log οb = (1.56 ± 0.39).

C.6 Procedures for data handling

C.6.1 Correction to zero ionic strength The correction of experimental data to zero ionic strength is necessary in all cases where a linear regression is impossible or appears inappropriate. The method used throughout the review is the specific ion interaction equations described in detail in Appendix B. Two variables are needed for this correction, and both have to be provided with an un-certainty at the 95% confidence level: the experimental source value, 10log K or

10log b , and the stoichiometric sum of the ion interaction coefficients, ∆ε . The ion interaction coefficients (see Tables B.3 and B.4 of Appendix B) required to calculate ∆ε may not all be known. Missing values therefore need to be estimated. It is recalled that the electric charge has the most significant influence on the magnitude of the ion interaction coefficients, and that it is in general more reliable to estimate ∆ε from known reactions of the same charge type, rather than to estimate single ∆ε values. The uncertainty of the corrected value at I = 0 is calculated by taking into account the


484

propagation of errors, as described below. It should be noted that the ionic strength is frequently given in moles per dm3 of solution (molar, M) and has to be converted to moles per kg H2O (molal, m), as the model requires. Conversion factors for the most common inert salts are given in Table II.5.

Example C.6:

For the equilibrium constant of the reaction:

M3+ + 2 H2O(l) + +2M(OH) + 2 H , (C.20)

only one credible determination in 3 M NaClO4 solution is known to be, 10*log b (C.20)

= − 6.31, to which an uncertainty of ± 0.12 has been assigned. The ion interaction coef-ficients are as follows:

3+4( , )M ClO−ε = (0.56 ± 0.03) kg·mol−1,

+2 4(M(OH) ,ClO )−ε = (0.26 ± 0.11) kg·mol−1,

+4(H ,ClO )−ε = (0.14 ± 0.02) kg·mol−1.

The values of ∆ε and ∆εσ can be obtained readily (cf. Eq. (C.22)): + + 3+2 4 4 4( , ) ( , )(M(OH) ,ClO ) 2 H ClO M ClO− − −∆ε = ε + ε − ε = – 0.02 kg·mol–1,

2 2 2 1 = (0.11) (2 0.02) (0.03) 0.12 kg mol−∆εσ + × + = ⋅ .

The two variables are thus:

10*log b (C.20) = − (6.31 ± 0.12),

∆ε = − (0.02 ± 0.12) kg·mol−1.

According to the specific ion interaction model the following equation is used to correct for ionic strength for the reaction considered here:

10*log b (C.20) + 6D = 10

*log οb (C.20) − 4ClO−∆ε m

D is the Debye-Hückel term:

0.509(1 1.5 )

m

m

ID

I=

+.

The ionic strength, Im, and the molality, 4ClO

m − (4ClOmI m −≈ ), have to be ex-

pressed in molal units, 3 M NaClO4 corresponding to 3.5 m NaClO4 (see Section II.2), giving D = 0.25. This results in:

10*log οb (C.20) = − 4.88.

The uncertainty in 10*log οb is calculated from the uncertainties in 10

*log b and ∆ε (cf. Eq. (C.22)):

C.6 Procedures for data handling 485

* * -10 10 4

2 2 2 2log log ClO

= + ( ) (0.12) (3.5 0.12) = 0.44.mο ∆εσ σ σ = + ×b b

The selected, rounded value is:

10*log οb (C.20) = − (4.9 ± 0.4).

C.6.2 Propagation of errors Whenever data are converted or recalculated, or other algebraic manipulations are per-formed that involve uncertainties, the propagation of these uncertainties has to be taken into account in a correct way. A clear outline of the propagation of errors is given by Bevington [69BEV]. A simplified form of the general formula for error propagation is given by Eq.(C.21), supposing that X is a function of Y1, Y2,…,YN .

2

2

1i

N

X Yi i

XY=

∂σ = σ ∂

∑ (C.21)

Eq. (C.21) can be used only if the variables, Y1, Y2,…,YN , are independent or if their uncertainties are small, i.e., the covariances can be disregarded. One of these two assumptions can almost always be made in chemical thermodynamics, and Eq. (C.21) can thus almost universally be used in this review. Eqs. (C.22) through (C.26) present explicit formulas for a number of frequently encountered algebraic expressions, where c, c1, c2 are constants.

X = c1Y1 ± c2Y2 : 1 2

2 2 21 2( ) ( )X Y Yc cσ = σ + σ (C.22)

X = ± cY1Y2 and X = 1

2

cYY

± : 1 2

2 22

1 2

Y YX

X Y Y σ σ σ

= + (C.23)

X = 21

cc Y ± : 2X Yc

X Yσ σ

= (C.24)

X = 21

c Yc e± : 2X

YcX

σ= σ (C.25)

X = 1 2ln( )c c Y± : 1Y

X cYσ

σ = (C.26)

Example C.7:

A few simple calculations illustrate how these formulas are used. The values have not been rounded.

Eq. (C.22): r mG∆ = 2·[ − (277.4 ± 4.9)] kJ·mol−1 − [ − (467.3 ± 6.2)] kJ·mol−1

= − (87.5 ± 11.6) kJ·mol−1.

Eq. (C.23): (0.038 0.002) (8.09 0.96)(0.0047 0.0005)

K ±= = ±

±


486

Eq. (C.24): K = 4·(3.75 ± 0.12)3 = (210.9 ± 20.3)

Eq. (C.25): r m

R ;GTK e

ο−∆ο = r mGο∆ = − (2.7 ± 0.3) kJ·mol−1

R = 8.3145 J·K−1·mol−1

T = 298.15 K

K ο = (2.97 ± 0.36).

Note that powers of 10 have to be reduced to powers of e, i.e., the variable has to be multiplied by ln(10), e.g.,

10 10log (ln(10) log )10log (2.45 0.10); 10 (282 65).K KK K e ⋅= ± = = = ±

Eq. (C.26) : r m R ln ;G T Kο ο∆ = − K ο = (8.2 ± 1.2) × 106

R = 8.3145 J·K−1·mol−1

T = 298.15 K

r mGο∆ = − (39.46 ± 0.36) kJ·mol−1

ln K ο = (15.92 ± 0.15)

10log ln / ln(10) (6.91 0.06).K Kο ο= = ±

Again, it can be seen that the uncertainty in 10log K ο cannot be the same as in ln K ο . The constant conversion factor of ln(10) = 2.303 is also to be applied to the un-certainty.

C.6.3 Rounding The standard rules to be used for rounding are:

1. When the digit following the last digit to be retained is less than 5, the last digit retained is kept unchanged.

2. When the digit following the last digit to be retained is greater than 5, the last digit retained is increased by 1.

3. When the digit following the last digit to be retained is 5 and

a) there are no digits (or only zeroes) beyond the 5, an odd digit in the last place to be retained is increased by 1 while an even digit is kept unchanged;

b) other non-zero digits follow, the last digit to be retained is increased by 1, whether odd or even.

This procedure avoids introducing a systematic error from always dropping or not dropping a 5 after the last digit retained.

C.6 Procedures for data handling 487

When adding or subtracting, the result is rounded to the number of decimal places (not significant digits) in the term with the least number of places. In multiplica-tion and division, the results are rounded to the number of significant digits in the term with the least number of significant digits.

In general, all operations are carried out in full, and only the final results are rounded, in order to avoid the loss of information from repeated rounding. For this rea-son, several additional digits are carried in all calculations until the final selected data set is developed, and only then are data rounded.

C.6.4 Significant digits The uncertainty of a value basically defines the number of significant digits a value should be given.

Example: (3.478 ± 0.008)

(3.48 ± 0.01)

(2.8 ± 0.4)

(10 ± 1)

(105 ± 20).

In the case of auxiliary data or values that are used for later calculations, it is often inconvenient to round to the last significant digit. In the value (4.85 ± 0.26), for example, the “5” is close to being significant and should be carried along a recalculation path in order to avoid loss of information. In particular cases, where the rounding to significant digits could lead to slight internal inconsistencies, digits with no significant meaning in absolute terms are nevertheless retained. The uncertainty of a selected value always contains the same number of digits after the decimal point as the value itself.

489

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[2002MON/HEL] Monlien, F. J., Helm, L., Abou-Hamdan, A., Merbach, A. E.,

Mechanistic diversity covering 15 orders of magnitude in rates: cyanide exchange on [M(CN)4]2- (M = Ni, Pd and Pt), Inorg. Chem., 41, (2002), 1717-1727. Cited on pages: 231, 441.

[2002PAT/WOO] Patterson, B. A., Woolley, E. M., Thermodynamics of ionization of

water at temperatures 278.15 K < T/K < 393.15 K and at the pressure p = 0.35 MPa: apparent molar volumes and apparent molar heat capacities of aqueous of potassium and sodium nitrates and nitric acid, J. Chem. Thermodyn., 34, (2002), 535-556. Cited on pages: 87, 88.

BIBLIOGRAPHY 583

[2002SCH] Schlesinger, M. E., The thermodynamic properties of phosphorus and solid binary phosphides, Chem. Rev., 102, (2002), 4267-4301. Cited on page: 204.

[2002WAL/GAT] Wallner, H., Gatterer, K., Growth of pure Ni(OH)2 single crystals

from solution-control of the crystal size, Z. Anorg. Allg. Chem., 628, (2002), 2818-2820. Cited on pages: 112, 438, 441.

[2002WAL/PRE] Wallner, H., Preis, W., Gamsjäger, H., Solid-solute phase equilibria

in aqueous solutions XV. Thermodynamic analysis of the solubility of nickel carbonates, Thermochim. Acta, 382, (2002), 289-296. Cited on pages: 218, 220, 221, 222, 263, 310, 441.

[2003BAE/BRA] Baeyens, B., Bradbury, M. H., Hummel, W., Determination of

aqueous nickel-carbonate and nickel-oxalate complexation constants, J. Solution Chem., 32, (2003), 319-339. Cited on pages: 221, 224, 225, 441.

[2003GAM/KON] Gamsjäger, H., Königsberger, E., In Hefter, G. T., 4.2 Solubility of

sparingly soluble solids in liquids, The experimental determination of solubilities, Tomkins, R. P. T., Ed., pp.315-358, John Wiley & Sons, Ltd, (2003). Cited on page: 111.

[2003GUI/FAN] Guillaumont, R., Fanghänel, T., Fuger, J., Grenthe, I., Neck, V.,

Palmer, D. A., Rand, M. H., Update on the chemical thermodynamics of Uranium, Neptunium, Plutonium, Americium and Technetium, Nuclear Energy Agency Data Bank, Organisation for Economic Co-operation and Development, Ed., vol. 5, Chemical Thermodynamics, North Holland Elsevier Science Publishers B. V., Amsterdam, The Netherlands, (2003). Cited on pages: 5, 6, 7, 53, 54, 58, 64, 65, 66, 443, 464, 469, 471.

[2003HUM/CUR] Hummel, W., Curti, E., Nickel aqueous speciation and solubility at

ambient conditions: A thermodynamic elegy, Monatsh. Chem., 134, (2003), 941-976. Cited on pages: 223, 224, 225, 442.

[2004BRO/MER] Brown, B. R., Merkley, E. D., McRae, B. R., Origlia-Luster, M. L.,

Woolley, E. M., Apparent molar volumes and apparent molar heat capacities of aqueous nickel(II) nitrate, copper(II) nitrate, and zinc(II) nitrate at temperatures (from 378.15 to 393.15) K at the pressure 0.35 MPa, J. Chem. Thermodyn., 36, (2004), 437-446. Cited on page: 88.

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[2005OLI/NOL] Olin, Å., Noläng, B., Öhman, L-O., Osadchii, E. G., Rosén, E.,

Chemical thermodynamics of Selenium, Nuclear Energy Agency Data Bank, Organisation for Economic Co-operation and Development, Ed., vol. 7, Chemical Thermodynamics, North Holland Elsevier Science Publishers B. V., Amsterdam, The Netherlands, (2005). Cited on pages: 6, 53, 61, 71, 195.

585

List of cited authors This chapter contains an alphabetical list of the authors of the references cited in this book. The reference codes given with each name corresponds to the publications of which the person is the author or a co-author. Note that inconsistencies may occur due to variations in spelling between different publications. No attempt has been made to correct for such inconsistencies in this volume.

Author Reference Abbattista, F. [63BUR/ABB] Abou-Hamdan, A. [2002MON/HEL] Accascina, F. [59FUO/ACC] Achenza, F. [59ACH] Acs, V. [41KIS/ACS] Adam, A. [80ADA/BIL], [81ADA/BIL] Adami, L. H. [65ADA/KIN] Aditya, S. [61MOH/DAS] Ageno, F. [11AGE/VAL] Ahrland, S. [56AHR/LAR], [56AHR/ROS] Ajdinjan, N. H. [53GRI/SLU] Akimoto, S.-I. [65AKI/FUJ], [74YAG/MAR] Aksel'rud, N. V. [50AKS/FIA] Al Mahamid, I. [98ALM/NOV] Alcock, C. B. [64ALC/BEL], [64ALC/SUD], [65STE/ALC],

[93KUB/ALC] Allard, B. [97ALL/BAN] Allison, J. D. [93ALL/BRO] Almqvist, G. [18ALM] Althaus, H. [65SCH/ALT] Alyamovskaya, K. V. [61CHU/ALY2] Ananthaswamy, J. [84ANA/ATK] Anderegg, G. [99RAR/RAN] Anderson, C. T. [35KEL/AND] Anderson, G. M. [91AND/CAS] Andreev, S. N. [68AND/KHA], [68AND/KHA2] Andreeva, M. N. [68AND/KHA2] Angell, C. A. [66ANG/GRU] Angnes, L. [82NEV/ANG] Antill, J. E. [67ANT/WAR]


LIST OF CITED AUTHORS

586

Author Reference Aprile, F. [68LIN/APR] Arai, Y. [86TOD/HAS] Archer, D. G. [90ARC/WAN], [98ARC/RAR], [99ARC] Archibald, R. C. [66STO/ARC] Arculus, R. J. [86HOL/NEI] Ariya, S. M. [71ARI/MOR] Arnek, R. [68ARN] Arnold, R. G. [75ARN/MAL] Artamonova, E. A. [89ART/KAS] Aruga, R. [74ARU], [75ARU] Arvia, A. J. [63BOL/JAU] Asao, H. [69MOR/SAT] Ashcroft, S. J. [70ASH] Ashurst, K. G. [77ASH/HAN] Asper, R. [77OSW/ASP] Atkinson, G. [84ANA/ATK] Auffredic, J.-P. [72AUF/CAR], [72AUF/CAR2] Avramenko, N. I. [83AVR/BLO] Baes, Jr., C. F. [74MES/BAE], [76BAE/MES] Baeyens, B. [2003BAE/BRA] Bahta, A. [97BAH/PAR] Bailey, S. M. [82WAG/EVA] Bain, R. [63PEN/BAI] Balarev, K. [89LIL/TEP] Bale, C. W. [83GAU/BAL] Balej, J. [89BAL], [93BAL/DIV], [93BAL/DIV2],

[97BAL/DIV] Ball, R. G. J. [92BAL/DIC] Balz, O. [21WOH/BAL] Banerjea, D. [81BAN/KAD] Banwart, S. A. [97ALL/BAN] Bányai, I. [92BAN/BLI] Bard, A. J. [85BAR/PAR] Barin, I. [73BAR/KNA], [77BAR/KNA] Barnard, R. [80BAR/RAN], [81BAR/RAN] Bartholin, H. [81ADA/BIL] Bartholomew, C. H. [69IZA/EAT] Bartovská, A. [74BAR/CER] Bartovský, A. [74BAR/CER] Barvinok, M. S. [75BAR/MAS] Basavalingu [91TAR/FAZ] Bates, R. G. [73BAT] Bath, G. A. [71BHA/SUB] Batrakov, V. V. [93MAK/VOE] Batyaev, I. M. [74NET/BAT]



587

Author Reference Baucke, F. G. K. [2002BUC/RON] Beaudouin, B. [84BRA/DEL] Bechtold, D. B. [78BEC/LIU] Beck, M. T. [62BEC/BJE], [87BEC], [90BEC/NAG] Becker, K. [67SIG/BEC] Becraft, K. A. [98ALM/NOV] Behr, M. [24MOS/BEH] Behrens, R. G. [72BEH/ROS] Belford, T. N. [64ALC/BEL] Bell, J. [37BEL], [40BEL] Belokurova, N. A. [86VAS/DMI] Bender, J. A. [23VIL/BEN] Benrath, A. [34BEN/THI] Berger, E. [21BER/CRU] Berglund, S. [76BER2] Bergman, G. A. [82GLU/GUR] Berry, L. G. [81NIC/BER] Bestul, A. B. [71CHA/BES] Bevington, P. R. [69BEV] Bezborodov, A. A. [76ZHO/BEZ] Bhandage, G. T. [91TAR/FAZ] Bhatia, S. N. [77BHA/CAR] Bhattacharya, S. [86BHA/MUK] Biader Ceipidor, U. [76BIA/CAR] Bichowsky, F. R. [36BIC/ROS] Bidoglio, G. [95SIL/BID] Bidwell, L. R. [67RIZ/BID] Biedermann, G. [71OHT/BIE], [75BIE] Bigliardi, G. [63FER/BRA] Bijvoet, J. M. [33BIJ/NIE] Bilinski, H. [73NOV/COS] Billerey, D. [77BIL/TER], [80ADA/BIL], [81ADA/BIL] Biltz, W. [36BIL/VOI] Binford Jr., J. S. [67BIN/STR], [70BIN/HEB] Bish, D. L. [82LIV/BIS] Bishop, J. J. [61MAR/BYD] Bixler, J. W. [74BIX/LAR], [83SOL/BON] Bjerrum, J. [62BEC/BJE], [86BJE/SKI], [87BJE], [88BJE],

[88BJE2], [89BJE], [90BJE], [90BJE2] Blackie, M. S. [59BLA/GOL], [59BLA/GOL2] Blaszkiewicz, G. [67BLA] Blixt, J. [92BAN/BLI] Blokhin, V. V. [74BLO/RAZ], [76KUL/BLO], [81KUL/BLO],

[83AVR/BLO] Bode, H. [66BOD/DEH]



588

Author Reference Boerstoel, B. M. [64CHA/BOE] Bogatzki, D. P. [38BOG] Boldamdaeva, O. K. [93UVA/TIM] Bolzan, J. A. [63BOL/JAU] Bond, A. M. [70BON], [71BON], [72BON/HEF], [83SOL/BON] Bonilla, C. F. [52CAR/BON] Bonnell, D. G. R. [35BON/BUR] Borchardt G. [84ROG/BOR] Bornemann, K. [08BOR], [10BOR] Boström, D. [87BOS] Bouet, G. [95KAH/CRO] Boye, E. [33BOY] Boyle, B. J. [54BOY/KIN] Brackett, T. E. [56HOR/BRA] Braconnier, J. J. [84BRA/DEL] Bradbury, M. H. [2003BAE/BRA] Bradley, D. J. [79BRA/PIT] Braibanti, A. [63FER/BRA] Branica, M. [73NOV/COS] Bratsch, S. G. [89BRA] Braun, M. [66VOL/KOH] Breckpot, R. [69RAN/BRE] Brett, C. M. A. [2002BUC/RON] Britton, H. T. S. [25BRI] Brodale, G. E. [66STO/ARC] Broers, G. H. J. [75WEE/KOE] Bronson, H. L. [36BRO/WIL] Brønsted, J. N. [22BRO], [22BRO2] Brooker, E. J. [52ROB/BRO] Brown, B. R. [2004BRO/MER] Brown, D. S. [93ALL/BRO] Brown, G. R. [65JAM/BRO] Brown, Jr., G. E. [82BRO] Brown, P. G. M. [55BRO/PRU] Brown, T. L. [77CHE/BRO] Bruner, L. [09BRU/ZAW] Brunetti, B. [96BRU/PIA] Brungs, M. P. [80TRA/BRU] Bruno, J. [86BRU], [97ALL/BAN] Buck, R. P. [2002BUC/RON] Bucko, M. [2000ROG/KOZ] Bugarčiç, Z. D. [80MIL/BUG] Burdese, A. [63BUR/ABB] Burke, K. A. [20BUR]



589

Author Reference Burkov, K. A. [65BUR/LIL], [65BUR/LIL2], [66BUR/IVA],

[71BUR/ZIN], [78BUR/KAM] Burridge, L. W. [35BON/BUR] Busey, R. H. [52BUS/GIA], [52BUS/GIA2], [53BUS/GIA] Bydalek, T. J. [61MAR/BYD] Caillère, S. [62CAI/POB] Caldero, J. M. [62TRI/CAL], [64TRI/CAL] Cali, R. [73GUR/CAL] Calk, L. C. [73WAA/CAL] Caminiti, R. [80CAM/LIC], [82CAM] Camoes, M. F. [2002BUC/RON] Campbell, A. B. [96LEM/CAM], [97PAN/CAM] Campbell, F. E. [68CAM/ROE] Cantor, S. [65CAN] Capdevila, H. [92CAP], [93GIF/VIT2], [94CAP/VIT], [95CAP/VIT] Carbonaro, N. [99MAL/CAR] Carel, C. [72AUF/CAR], [72AUF/CAR2] Carlin, R. L. [77BHA/CAR] Carney, J. H. [70CAR/LAI] Carpenter, S. A. [98ALM/NOV] Carr, D. S. [52CAR/BON] Carunchio, V. [76BIA/CAR] Castet, S. [91AND/CAS] Catalano, E. [54STO/CAT], [55CAT/STO], [55STO/CAT] Cemic, L. [86CEM/KLE], [87CEM/KLE] Černý, C. [74BAR/CER] Cerutti, P. [68LAR/CER] Chaberek, Jr., S. [52CHA/COU] Chakravorty, A. [86BHA/MUK] Chandrasekharaiah, M. S. [75CHA/KAR] Chang, S. S. [71CHA/BES] Chang, Y. A. [77LIN/IPS], [78LIN/HU], [79SHA/CHA],

[80SHA/CHA], [86HSI/CHA] Charette, G. G. [68CHA/FLE] Charnock, J. M. [2001HEN/RED] Charykov, N. A. [85FIL/CHA], [86FIL/CHA] Chase, Jr., M. W. [98CHA] Chattopadhyay, G. [75CHA/KAR] Chen, A. [90PAS/TAY] Cheng, C. P. [77CHE/BRO] Chialvo, A. A. [2002CHI/SIM] Chichirova, N. D. [85SAP/CHI] Chickerur, N. S. [80CHI/SAB], [82CHI/SAB] Chierice, G. O. [82NEV/ANG] Childs, C. W. [70CHI]



590

Author Reference Christ, C. L. [65GAR/CHR] Christensen, J. J. [63CHR/IZA], [69IZA/EAT], [71IZA/JOH] Christensen, Jr., A. U. [58KIN/CHR] Chukhlantsev, V. G. [56CHU3], [56CHU4], [57CHU], [61CHU/ALY2] Chukurov, P. M. [73CHU/DRA] Chunaeva, V. D. [95MUL/CHU] Churney, K. L. [82WAG/EVA] Chyzy, R. [2001MIK/MIG] Ciavatta, L. [80CIA], [87CIA/IUL], [88CIA], [90CIA] Ciret, N. [77BIL/TER] Claire, Y. [80DUB/CLA] Clare, B. W. [75CLA/KEP] Clark, C. W. [35KEE/CLA] Coffetti, G. [05COF/FOR] Coleman, J. S. [65COL/PET] Colombier, L. [34COL2] Combs, K. E. S. [89ZIE/JON] Comert, H. [84COM/PRA] Conard, B. B. [77CON/SRI] Connelly, D. L. [71CON/LOO] Conway, B. E. [62CON/GIL] Conway, K. C. [54BOY/KIN] Cook, R. S. [26PEA/COO] Cordes, S. M. [69SUB/COR] Cordfunke, E. H. P. [73COR], [90COR/KON2], [92BAL/DIC] Corlis, C. [81COR/SUG] Cosovič, B. [73NOV/COS] Coughlin, J. P. [51COU] Courtney, R. C. [52CHA/COU] Covington, A. K. [82SAR/COV], [2002BUC/RON] Cowell, J. C. [74WOR/COW] Cox, J. D. [89COX/WAG] Craig, J. R. [76CRA/SCO] Crego, A. [62FAU/CRE] Criaud, A. [84FOU/CRI] Criss, C. M. [96CRI/MIL] Cronier, D. [95KAH/CRO] Crut, G. [21BER/CRU], [24CRU], [24CRU2] Csokán, P. [38CSO] Cucinotta, V. [78SIR/SEM] Cudennec, Y. [87CUD/LEC] Cummiskey, C. J. [71LAN/CUM] Čuprová [59KIS/CUP] Curti, E. [2003HUM/CUR] Cuta, F. [56CUT/KSA]



591

Author Reference Cvitaš, T. [93MIL/CVI] Dabrowski, J. [32SKA/DAB] Dabrowski, W. [93PRZ/WIS] Damiani, R. [63BUR/ABB] D'Aprano, A. [71APR] Darab, J. G. [99HOF/DAR] Das, R. C. [61MOH/DAS], [71DAS/DAS] Dash, A. C. [71DAS/DAS] Dash, U. N. [71DAS/DAS] Date, M. [83KAT/SUG] Davies, A. G. [61DAV/SMI] Davies, C. W. [32MON/DAV], [62DAV] Davis, C. E. S. [77NIC/ROB] Day, P. [76DAY/DIN] De Jong, W. F. [27JON/WIL] De Jonge, W. J. M. [72POL/HER] De Moras, M. [85MAG/MOR] De Nobel, J. [64CHA/BOE] De Ranter, C. [69RAN/BRE] De Souza, B. C. [63SHA/DES] De Stefano, C. [94STE/FOT] De Waal, S. A. [73WAA/CAL] De Wijs, H. J. [25WIJ] De Zoubov, N. [63DEL/ZOU] DeFord, D. D. [51DEF/HUM] Dehmelt, K. [66BOD/DEH] Delmas, C. [84BRA/DEL] Deltombe, E. [63DEL/ZOU] Derby, I. H. [16DER/YNG] Desai, P. D. [73HUL/DES] Deshpande, D. A. [79NAN/DES] Desnoyers, J. E. [81PER/ROU] DeWitt, B. J. [40SEL/DEW] Dickinson, S. [92BAL/DIC] Dickson, A. G. [81TUR/WHI] Diebler, H. [76TAY/DIE] Dinsdale, A. T. [76DAY/DIN], [91DIN] Divisek, J. [93BAL/DIV], [93BAL/DIV2], [97BAL/DIV] Dmitrieva, N. G. [84VAS/VAS], [86VAS/DMI] Dobrokhtov, G. N. [54DOB] Dodgen, H. W. [68LIN/APR], [78BEC/LIU], [83GOW/DOD] Dojčiloviç, J. [96DOJ/NOV] Dokoupil, Z. [72KLA/DOK] Domange, L. [37DOM] Domash, L. [55TOM/DOM]



592

Author Reference Domiciano, J. B. [90JUR/DOM] Donaldson, R. H. [65DON/EDM] Dönges, E. [47DON] Drakin, S. I. [73CHU/DRA], [85DRA/MAD] Drever, J. I. [82DRE], [96POU/DRE] Droll, H. A. [63NET/DRO] Drowart, J. [92BAL/DIC] Du Chatenier, F. J. [64CHA/BOE] DuBose, J. B. [79DUB/NAU] Dubusc, M. [80DUB/CLA] Dunn, L. [73EDW/FAR] Dyrssen, D. [85DYR], [88DYR], [90DYR/KRE] Dziembaj, R. [2001MIK/MIG] Eatough, D. J. [69IZA/EAT], [71IZA/JOH] Eckstrom, H. C. [37PEA/ECK] Economou, M. [81MAR/ECO], [83ECO/MAR] Edmonds, D. T. [65DON/EDM] Edwards, O. W. [73EDW/FAR] Efimov, A. I. [70EFI/KUD] Efimov, M. E. [88EFI/EVD], [89EVD/EFI], [90EFI/EVD],

[90EFI/FUR] Egorov, A. M. [65EGO/SMI] Ehler, T. C. [64EHL/KEN] Ehrlich, P. [36BIL/VOI] Einerhand, J. [50GLE/EIN] Eitel, W. [54EIT] Ekström, C. G. [76PER/EKS] Elliott, R. P. [65ELL] Elmasry, M. A. A. [98ELM/GAB] Elmore, A. L. [98JAI/ELM] Emara, M. M. [87EMA/FAR] Enderby, J. E. [83NEI/END] Enoki, T. [78ENO/TSU] Ephraim, J. H. [97ALL/BAN] Eriksson, G. [95KON/ERI] Erstigneeva, T. L. [69KUL/ERS] Espelund, A. W [73SKE/ESP] Euler, H. [04EUL] Evans, W. H. [52ROS/WAG], [82WAG/EVA] Evdokimova, V. I. [88EFI/EVD], [90EFI/EVD] Evdokimova, V. P. [89EVD/EFI] Falicov, L. M. [90PAS/TAY], [92FRE/FAL] Fan, G. [83RAM/FAN] Fanelli, N. [99MAL/CAR] Fanghänel, T. [97KON/FAN], [2003GUI/FAN]



593

Author Reference Farber, M. [58FAR/MEY] Farhi, R. [84RAY/PET] Farid, N. A. [87EMA/FAR] Farr, T. D. [73EDW/FAR] Faucherre, J. [62FAU/CRE] Fazeli, A. R. [91TAR/FAZ] Feber, R. C. [68HER/FEB] Fedorov, V. A. [73FED/SHM] Fedorov, Y. A. [86FIL/CHA] Fedotov, N. V. [74MAK/MAS] Feitknecht, W. [36LOT/FEI], [54FEI/HAR], [63FEI/SCH] Felmy, A. R. [97MAT/RAI], [98RAI/FEL] Ferguson, R. B. [74GRI/FER] Ferier, A. [63GEO/FER] Fernelius, W. C. [77FER] Ferrante, M. J. [78STU/FER], [82FER/GOK] Ferrari, A. [63FER/BRA] Ferri, D. [83FER/GRE], [85FER/GRE] Ferrini, S. [99MAL/CAR] Fialkov, Y. A. [50AKS/FIA], [56FIA/SHE] Figlarz, M. [84BRA/DEL] Filho, A. P. [77BHA/CAR] Filippov, V. K. [85FIL/CHA], [86FIL/CHA] Filippova, E. P. [74MAK/MAS] Fisher, F. H. [79FIS/FOX] Fitzner, K. [75MOS/FIT], [75MOS/FIT2] Fjellvåg, H. [94STO/FJE], [96SEI/FJE] Fleet, M. E. [72FLE], [77FLE], [87FLE] Flengas, S. N. [68CHA/FLE], [73FLE] Florence, T. M. [66FLO] Flowers, G. C. [81HEL/KIR] Fontaine, A. [80LAG/FON] Foote, H. W. [23FOO] Forsling, W. [78FOR] Forstat, H. [59SPE/FOR] Förster, F. [05COF/FOR] Foti, C. [94STE/FOT] Fouassier, C. [84BRA/DEL] Fouillac, C. [84FOU/CRI] Fox, A. P. [79FIS/FOX] Frank, D. F. [67RIZ/BID] Franke, E. [1895FRA] Fraunberger, F. [04MUT/FRA] Fray, D. J. [94KAL/FRA] Freeman, R. D. [84FRE]



594

Author Reference Freericks, J. K. [92FRE/FAL] Freiser, H. [69SUB/COR] Freund, H. [59FRE/SCH] Frey, C. M. [72FRE/STU], [78FRE/STU] Fricke, R. [42FRI/WEI] Friedberg, S. A. [60ROB/FRI], [73HAM/FRI] Fronaeus, S. [53FRO2], [62FRO/LAR] Fuger, J. [92FUG/KHO], [92GRE/FUG], [2001LEM/FUG],

[2003GUI/FAN] Fujimura, K. [78IWA/FUJ] Fujinaga, T. [66SHI/FUJ] Fujisawa, H. [65AKI/FUJ] Fulton, J. L. [99HOF/DAR] Funahashi, S. [69FUN/TAN] Funk, R. [1899FUN] Fuoss, R. M. [59FUO/ACC] Furkalyuk, M. Y. [90EFI/FUR] Furumi, Y. [76SHI/TSU] Gaber, A. [98ELM/GAB] Galchenko, G. L. [84LAV/TIM] Galus, Z. [74KUT/GAL] Gamsjäger, H. [82GAM/REI], [98GAM/KON], [99KON/KON],

[2001GAM/PRE], [2002GAM/WAL], [2002WAL/PRE], [2003GAM/KON]

Garber, H. K. [68LAR/CER] Garrels, R. M. [65GAR/CHR] Garrett, A. B. [49GAY/GAR] Garvin, D. [87GAR/PAR] Gatterer, K. [2002WAL/GAT] Gauthier, M. [83GAU/BAL] Gayer, K. H. [49GAY/GAR], [52GAY/WOO] Gee, R. [76GEE/SHE] Geier, G. [71HOH/GEI] Geiger, G. H. [73VEN/GEI] Geoffray, H. [63GEO/FER] Gerault, Y. [87CUD/LEC] Gessner, H. [14THI/GES] Giacomelli, A. [99MAL/CAR] Gianguzza, A. [94STE/FOT] Giauque, W. F. [41STO/GIA], [52BUS/GIA], [52BUS/GIA2],

[53BUS/GIA], [56HOR/BRA], [66STO/ARC] Giera, E. [75KAK/GIE] Giesekus, A. [90PAS/TAY] Giffaut, E. [93GIF/VIT2], [94GIF] Gil, N. S. [59GIL/NYH]



595

Author Reference Gilbert, G. [63PEN/BAI] Gileadi, E. [62CON/GIL] Glaser, J. [92BAN/BLI] Glasstone, S. [26GLA] Glebov, V. A. [71KLY/GLE] Gleiser, M. [73HUL/DES] Glemser, O. [50GLE/EIN] Glowacz, E. [2000ROG/KOZ] Glushko, V. P. [82GLU/GUR] Gokcen, N. A. [82FER/GOK], [84PAN/STU] Gold, V. [59BLA/GOL], [59BLA/GOL2] Goldberg, R. N. [66GOL/RID], [79GOL/NUT], [82SAR/COV] Goldschmidt, V. M. [26GOL] Golovanova, J. G. [75LEV/GOL] Golub, A. M. [79GOL/KOH] Goncharov, V. I. [73CHU/DRA] Goryozono, Y. [96UCH/GOR] Goto, S. [77SAD/KAW] Gower, S. A. [83GOW/DOD] Grad, B. [99MIK/MIG], [2001MIK/MIG] Graham, R. D. [74GRA/WIL] Grauer, R. [97ALL/BAN] Grayson-Smith, H. [50VAS/GRA] Gregán, F. [96SAH/SAH] Grenthe, I. [83FER/GRE], [85FER/GRE], [87RIG/VIT],

[92GRE/FUG], [94PLY/GRE], [95GRE/PUI], [97ALL/BAN], [97GRE/PLY], [97GRE/PLY2], [97PUI/RAR], [2000GRE/WAN], [2003GUI/FAN]

Gricaenko, G. S. [53GRI/SLU] Grice, J. D. [74GRI/FER] Griffiths, T. R. [69GRI/SCA] Grimaldi, M. [64GRI/LIB] Grønvold, F. [91STO/GRO], [94STO/FJE], [95GRO/STO],

[96SEI/FJE] Gruen, D. M. [66ANG/GRU] Gruener, D. W. [65JAM/BRO] Grzybkowski, W. [90GRZ] Guedes de Carvalho, J. R. F. [73RAU/GUE] Guggenheim, E. A. [35GUG], [66GUG] Guillaumont, R. [2003GUI/FAN] Gurrieri, S. [73GUR/CAL] Gurvich, L. V. [82GLU/GUR] Haas, C. [81KUI/SAN] Hadermann, J. [97ALL/BAN] Hadley, W. B. [64STO/HAD]



596

Author Reference Hagenmuller, P. [84BRA/DEL] Hakem, N. [98ALM/NOV] Hale, J. D. [63CHR/IZA] Hallenbeck, D. R. [83SOL/BON] Halloff, E. [70HAL/VAN] Halow, I. [82WAG/EVA] Hamann, S. D. [82HAM] Hamburger, A. I. [73HAM/FRI] Hamer, W. J. [59HAM] Hammes, G. G. [64HAM/MOR] Hancock, R. D. [77ASH/HAN] Hansen, J. W. [71IZA/JOH] Haring, M. M. [29HAR/BOS] Harned, H. S. [58HAR/OWE] Hartmann, L. [54FEI/HAR] Hashimoto, K. [86TOD/HAS], [86TOD/HAS], [94HAS/TOD],

[94HAS/TOD] Hatfield, J. D. [73EDW/FAR] Haury, G. L. [74RUT/HAU] Hawkins, D. T. [73HUL/DES] Hay, R. G. [55TOM/DOM] Heath, G. A. [77HEA/HEF] Hebert, T. H. [70BIN/HEB] Hedlund, T. [88HED] Hefter, G. [72BON/HEF], [74HEF], [77HEA/HEF] Heindl, R. [82GAM/REI] Hejtmánek, M. [56CUT/KSA] Helgeson, H. C. [74HEL/KIR], [74HEL/KIR2], [81HEL/KIR],

[88SHO/HEL], [88TAN/HEL], [89SHO/HEL], [89SHO/HEL2], [90OEL/HEL], [92SHO/OEL]

Helm, L. [2002MON/HEL] Hemingway, B. S. [84ROB/HEM], [90HEM], [95ROB/HEM] Henderson, C. M. B. [2001HEN/RED] Henry, W. F. [30ROD/HEN] Hepler, L. G. [63KO/HEP], [66GOL/RID], [68LAR/CER],

[78SPI/SIN], [78SPI/SIN2], [79SPI/OLO], [89HOV/HEP], [96HEP/HOV]

Herbert, T. H. [67BIN/STR] Herrick, C. C. [68HER/FEB] Hertz, H. G. [79WEI/HER], [79WEI/MUL] Herweijer, A. [72POL/HER] Hess, N. J. [98RAI/FEL] Heyding, R. D. [58TAY/HEY] Hietanen, S. [85FER/GRE] Higginson, W. C. E. [73HUT/HIG]



597

Author Reference Hofer, F. [65SCH/ALT] Hoffmann, H. [74DIC/HOF] Hoffmann, M. M. [99HOF/DAR] Hohl, H. [71HOH/GEI] Holmes, H. F. [83HOL/MES] Holmes, R. D. [86HOL/NEI] Homann, K. [93MIL/CVI] Hood, G. C. [51HOO/WOY] Hopkins, H. P. [66GOL/RID] Hornung, E. W. [56HOR/BRA] Hostettler, J. D. [84HOS] Hounslow, A. W. [59WIL/THR] Hovey, J. K. [89HOV/HEP], [96HEP/HOV] Howell, I. [97HOW/NEI] Hsieh, K. [86HSI/CHA] Hu, D. C. [78LIN/HU] Hubbard, W. N. [67RUD/DEV] Huber, C. O. [78KAR/HUB] Huebner, J. S. [70HUE/SAT] Hugel, R. [71PIE/HUG] Hultgren, R. [73HUL/DES] Hume, D. N. [50HUM/KOL], [51DEF/HUM], [64GEE/HUM],

[72RET/HUM] Hummel, W. [97ALL/BAN], [2003BAE/BRA], [2003HUM/CUR] Hummelstedt, L. [71PAA/HUM] Hunt, J. P. [68LIN/APR], [78BEC/LIU], [83GOW/DOD] Hutchinson, M. H. [73HUT/HIG] Hutta, J. J. [63PHI/HUT] Hyams, D. [97HYA] Ichihashi, M. [82WAK/ICH] Imanakunov, B. I. [83KIM/IMA] In Hefter, G. T. [2003GAM/KON] Ingraham, T. R. [66ING], [70NAG/ING] Ipser, H. [77LIN/IPS] Isaacs, T. [63ISA] Isabaev, A. S. [93MUL/ISA] Isabaev, S. M. [93MUL/ISA], [95MUL/CHU] Isono, T. [71ISO] Ito, J. [84ROB/HEM] Itoh, C. T. [88NIS/ITO] Iuliano, M. [87CIA/IUL], [89IUL/POR] Ivanova, L. V. [66BUR/IVA] Iwase, M. [78IWA/FUJ] Izatt, R. M. [63CHR/IZA], [69IZA/EAT], [71IZA/JOH] Jablonski, M. [93PRZ/WIS]



598

Author Reference Jacob, K. T. [86JAC/KAL] Jaffe, I. [52ROS/WAG] Jäger, H. [93OBE/KAS] Jain, M. [98JAI/ELM] Jakob, A. [97ALL/BAN] Jamieson, J. W. S. [65JAM/BRO] Janjic, T. [74DIC/HOF] Jauregui, E. A. [63BOL/JAU] Jeanloz, R. [90PAS/TAY] Jellinek, K. [25JEL/ZAK], [26JEL/ULO], [26JEL/ULO2],

[28JEL/RUD] Jena, P. K. [56JEN/PRA] Jensen, K. A. [71JEN] Ji, L.-N. [96SAH/SAH] Job, A. [23JOB/SAM] Job, P. [36JOB] Johnson, F. M. G. [04STE/JOH] Johnson, J. W. [92SHO/OEL] Johnson, K. S. [79JOH/PYT] Johnston, H. D. [71IZA/JOH] Jones, G. P. [74SLO/JON] Jones, L. H. [60MCC/JON], [63PEN/BAI] Jones, M. E. [89ZIE/JON] Jones, T. S. [54ORI/JON] Jordan, J. [85BAR/PAR] Juraitis, K. R. [90JUR/DOM] Kabir-Ud-Din [74KAB] Kaden, T. A. [81BAN/KAD] Kahn, M. A. [95KAH/CRO] Kakokowicz, W. [75KAK/GIE] Kale, G. M. [86JAC/KAL], [94KAL/FRA] Kalinichenko, I. I. [66KAL/PUR] Kalinkina, I. N. [64KOS/KAL] Kallay, N. [93MIL/CVI] Kamenetskaya, N. V. [78BUR/KAM] Kamimura, T. [97NOZ/KOB] Kanchiku, Y. [66SHI/FUJ] Kapustinsky, A. F. [36KAP/SIL] Karapet'yants, M. K. [73CHU/DRA] Karapiperis, T. [97ALL/BAN] Karato, S. [2001LAW/ZHA] Karkhanavala, M. D. [75CHA/KAR] Karlina, O. [85DRA/MAD] Karlsson, R. [76KAR/KUL] Karpinskaya, N. M. [68AND/KHA]



599

Author Reference Karwan, T. [78KAR/MAL] Karweik, D. H. [78KAR/HUB] Kasaoka, S. [77KAS/SAK] Kaschnitz, E. [93OBE/KAS] Kasenov, B. K. [89ART/KAS] Kasper, R. B. [76NAV/KAS] Katayama, I. [79KEM/KAT] Katayama, S. [73KAT], [77KAT] Katsumata, K. [83KAT/SUG] Kawai, T. [73KAW/OTS] Kawakami, M. [77SAD/KAW] Kazybaev, S. A. [83KIM/IMA] Keesom, W. H. [35KEE/CLA] Kehiaian, H. V. [96KEH] Kelley, K. K. [35KEL/AND], [60KEL], [61KEL/KIN],

[64WEL/KEL], [73HUL/DES] Kellogg, H. H. [69KEL], [75MEY/WAR] Kemori, N. [79KEM/KAT] Kennedy, B. [69KEN2] Kennedy, M. B. [66KEN/LIS] Kent, R. A. [64EHL/KEN] Kenttämaa, J. [56KEN], [58KEN], [59KEN] Kepert, D. L. [75CLA/KEP] Kerr, P. F. [45KER] Ketclaar, J. A. A. [34KET] Khachkuruzov, G. A. [82GLU/GUR] Khaldin, V. G. [68AND/KHA], [68AND/KHA2] Khan, G. A. [59SPE/FOR] Khater, E. M. H. [98ELM/GAB] Kher, V. G. [79NAN/DES] Khodakovsky, I. L. [92FUG/KHO] Kildahl, N. K. [69KOL/KIL] Kim, J. I. [97KON/FAN] Kim, J. J. [74PIT/KIM] Kim, T. P. [83KIM/IMA] King, E. G. [54BOY/KIN], [57KIN], [58KIN/CHR], [61KEL/KIN],

[65ADA/KIN] Kipp, R. L. [57MOR/ZEL] Kirkham, D. H. [74HEL/KIR], [74HEL/KIR2], [81HEL/KIR] Kišova, L. [59KIS/CUP] Kitayama, K. [89KIT] Kiukkola, K. [57KIU/WAG] Kivalo, P. [57KIV/LUO] Klaaijsen, F. W. [72KLA/DOK] Kleinclauss, J. [77BIL/TER]



600

Author Reference Klepp, K. O. [72KLE/KOM] Kleppa, O. J. [68NAV/KLE], [86CEM/KLE], [87CEM/KLE] Kloger, W. [82KLO] Klygin, A. E. [71KLY/GLE] Knacke, O. [73BAR/KNA], [77BAR/KNA] Knauer, H. [39SIM/KNA] Knight, K. S. [2001HEN/RED] Ko, H. C. [63KO/HEP], [78STU/FER] Kobayashi, H. [97NOZ/KOB] Koekoek, F. R. J. [75WEE/KOE] Köhler, H. [79GOL/KOH] Kohler, K. [64KOH/ZAS] Kohlhaas, R. [66VOL/KOH] Kohls, D. W. [66KOH/ROD] Kokurin, N. I. [84VAS/VAS] Kolonin, G. R. [91STO/GRO] Kolski, G. B. [68KOL/MAR], [69KOL/KIL] Kolthoff, I. M. [31KOL], [50HUM/KOL] Kolyada, N. S. [71KLY/GLE] Komarek, K. L. [72KLE/KOM] Königsberger, E. [95KON/ERI], [98GAM/KON], [99KON/KON],

[2003GAM/KON] Königsberger, L.-C. [99KON/KON] Konings, R. J. M. [90COR/KON2], [92BAL/DIC], [92GRE/FUG] Könnecke, T. [97KON/FAN] Korableva, V. D. [58YAT/KOR] Korolev, V. P. [95MAN/KOR] Korotkevich, I. I. [92VID/KOR] Kosaki, A. [69SOR/KOS] Kostryukov, V. N. [64KOS/KAL] Kostryukova, M. O. [68KOS], [69KOS], [75KOS] Köther, W. [78KOT/MUL] Kowalewska, G. [78LIB/KOW], [78LIB/MAC] Kozlowska-Róg, A. [98ROG/KOZ], [2000ROG/KOZ] Kozuka, Z. [69MOR/SAT], [79KEM/KAT] Krauss, F. [55KRA/WAR] Krausz E. R. [76DAY/DIN] Kremling, K. [90DYR/KRE] Krupka, K. M. [84ROB/HEM] Kryzhanovskii, M. M. [70MIR/MAK] Ksandr, Z. [56CUT/KSA] Ksenzhek, O. S. [70TYS/KSE] Kubaschewski, O. [72KUB], [77BAR/KNA], [93KUB/ALC] Kuchitsu, K. [93MIL/CVI] Kudasheva, T. V. [93MAK/VOE]



601

Author Reference Kudryashova, Z. P. [70EFI/KUD] Kuindersma, S. R. [81KUI/SAN] Kukhtina, V. A. [73FED/SHM] Kulagov, E. A. [69KUL/ERS] Kullberg, L. [74KUL3], [76KAR/KUL] Kullerud, G. [62KUL/YUN] Kul'vinova, L. A. [76KUL/BLO], [81KUL/BLO], [83AVR/BLO] Kurakane, K. [76MUR/KUR] Kutner, W. [74KUT/GAL] Laffitte, M. [59LAF], [63LIN/LAF], [82LAF] Lagarde, P. [80LAG/FON] Laitinen, H. A. [70CAR/LAI] Lambert, S. M. [57LAM/WAT] LaMontagne, R. A. [65JAM/BRO] Landers, L. [71LAN/CUM] Lange, E. [56LAN/MIE], [59LAN] Lange, N. [97PRA/LAN] Langmuir, D. [99LAN/MAH] Larson, J. W. [68LAR/CER] Larson, T. M. [74BIX/LAR] Larsson, R. [56AHR/LAR], [62FRO/LAR], [64LAR] Latimer, W. M. [51LAT], [52LAT] Lavut, E. G. [84LAV/TIM] Lawlis, J. D. [2001LAW/ZHA] Lebedev, B. G. [61LEB/LEV] Lebedev, S. V. [71LEB/SAV] LeBlanc, J. C. [80TRE/LEB2] Lee, S. Y. [90LEE/NAS] Lee, W. H. [79LEE/WHE] Leegard, T. [64LEE/ROS] Leigh, G. J. [90LEI] Lekae, V. A. [71KLY/GLE] Lemire, R. J. [92GRE/FUG], [95LEM], [96LEM/CAM],

[2001LEM/FUG] Leserf, A. [87CUD/LEC] Levin, E. M. [69LEV/ROB] Levine, S. [52ROS/WAG] Levitskii, V. A. [61LEB/LEV], [75LEV/GOL] Lewis, G. N. [61LEW/RAN] Liberti, A. [64GRI/LIB] Libowitz, G. G. [69LIG/LIB] Libus, W. [80LIB/SAD], [82SAD/LIB] Libus, Z. [69LIB/SAD], [73LIB], [75LIB/TIA], [78LIB/KOW],

[78LIB/MAC], [80LIB/SAD] Licheri, G. [80CAM/LIC], [83LIC/PAS], [85MAG/MOR]



602

Author Reference Licht, S. [88LIC2] Lightstone, J. B. [69LIG/LIB] Lileev, A. S. [89LIL/TEP] Lilic, L. S. [65BUR/LIL], [65BUR/LIL2] Lilich, L. S. [71BUR/ZIN] Lin, R. Y. [77LIN/IPS], [78LIN/HU] Lincoln, S. F. [68LIN/APR] Liné, G. [63LIN/LAF] Linke, W. F. [65LIN/SEI] Lister, M. W. [60LIS/ROS], [61LIS/WIL], [66KEN/LIS] Liu, G. [78BEC/LIU] Livingstone, A. [82LIV/BIS] Lobo, V. M. M. [89LOB/QUA], [89LOB] Lobry de Bruyn, C. A. [18SMI/LOB] Lofgren, N. L. [66LOF/MCI] Logsdon, K. M. [83SOL/BON] Long, F. A. [51LON] Loomis, J. S. [71CON/LOO] Lopez-Lopez, A. [62LOP] Lotmar, W. [36LOT/FEI] Luber, A. [30MUL/LUB] Luoto, R. [57KIV/LUO] Luther III, G. W. [96LUT/RIC] Lutz, H. D. [97PRA/LAN] Lyashchenko, A. K. [89LIL/TEP] Ma, C.-B. [74MA] Mac Arthur, D. M. [70MAC] MacDonald, A. [99LAN/MAH] MacDonald, R. D. [77NIC/ROB] Maciejewski, W. [78LIB/MAC] Madanat, A. T. [85DRA/MAD] Magini, M. [82MAG/PAS], [85MAG/MOR] Mah, A. D. [76MAH/PAN] Mahapatra, P. P. [80CHI/SAB], [82CHI/SAB] Mahoney, J. [99LAN/MAH] Mainard, R. [80ADA/BIL] Mainwaring, P. R. [85SKE/MAI] Major, E. [42MAJ] Makarov, G. V. [93MAK/VOE] Makarow, L. L. [79OJK/MAK] Makashev, Y. A. [70MIR/MAK], [74BLO/RAZ], [81KUL/BLO] Makovskaya, G. V. [69MAK/SPI] Maksimova, I. N. [74MAK/MAS] Malatesta, F. [97MAL/ZAM], [99MAL/CAR] Malik, O. P. [75ARN/MAL]



603

Author Reference Malinowski, C. [78KAR/MAL] Mal'kov, I. V. [79VOR/VAS] Malyavinskaya, O. N. [68MAL/TUR], [71TUR/MAL] Manin, N. G. [95MAN/KOR] Mapother, D. E. [71CON/LOO] Marchal, G. [25MAR] Marcopoulos, T. [81MAR/ECO], [83ECO/MAR] Margerum, D. W. [61MAR/BYD], [68KOL/MAR], [69KOL/KIL] Margrave, J. L. [58FAR/MEY], [64EHL/KEN] Marshall, W. L. [84MAR/MES] Martell, A. E. [52CHA/COU], [64SIL/MAR], [71SIL/MAR],

[76SMI/MAR] Marumo, F. [74YAG/MAR] Masaki, K. [32MAS] Mashkov, L. V. [75BAR/MAS] Mashovets, V. P. [74MAK/MAS] Masuda, I. [82WAK/ICH] Matarrese, L. M. [54MAT/STO] Matthews, M. A. [98JAI/ELM] Mattigod, S. V. [77MAT/SPO], [79MAT/SPO], [97MAT/RAI] Mavrina, I. Y. [70MIR/MAK] Mayorga, G. [73PIT/MAY], [74PIT/MAY] Mazo, L. H. [82NEV/ANG] McBreen, J. [99MCB] McCormick, D. B. [67SIG/BEC] McCullough, R. L. [60MCC/JON] McCurdy, K. G. [78SPI/SIN] McDonald, H. J. [40SEL/DEW] McIver, E. J. [66LOF/MCI] McMurdie, H. F. [69LEV/ROB] McRae, B. R. [2004BRO/MER] Meade, C. [90PAS/TAY] Medvedev, V. A. [82GLU/GUR], [88EFI/EVD], [89COX/WAG],

[92FUG/KHO] Mehrotra, G. M. [85MEH/TAR] Meisel, K. [36BIL/VOI] Merbach, A. E. [2002MON/HEL] Mériel, P. [80ADA/BIL] Merkley, E. D. [2004BRO/MER] Merrill, P. W. [68MOO/MER] Mesmer, R. E. [74MES/BAE], [76BAE/MES], [83HOL/MES],

[84MAR/MES], [91AND/CAS] Meusser, A. [01MEU] Meyer, G. A. [75MEY/WAR] Meyer, R. T. [58FAR/MEY]



604

Author Reference Mibuchi, T. [82WAK/ICH] Miederer, W. [56LAN/MIE] Migdał-Mikuli, A. [99MIK/MIG], [2001MIK/MIG] Migliardo, P. [80LAG/FON] Mikhailin, B. V. [79VOR/VAS] Mikuli, E. [99MIK/MIG], [2001MIK/MIG] Miliç, N. B. [80MIL/BUG] Miljutin G. [36TRA/SCH] Miller, A. R. [78SCH/MIL] Miller, M. L. [56MIL/SHE] Millero, F. J. [79MIL], [94ZHA/MIL], [96CRI/MIL] Mills, I. [93MIL/CVI] Mills, K. C. [74MIL] Milton, M. J. T. [2002BUC/RON] Minder, W. [65SCH/ALT] Mironov, L. A. [83AVR/BLO] Mironov, V. E. [70MIR/MAK], [73FED/SHM], [74BLO/RAZ],

[76KUL/BLO], [81KUL/BLO] Miura, K. [94HAS/TOD] Mohapatra, P. P. [61MOH/DAS] Molinou, I. E. [2000TSI/MOL], [2001MOL/TSI] Money, R. W. [32MON/DAV] Monk, C. B. [49MON] Monlien, F. J. [2002MON/HEL] Monnin, C. [90MON] Montgomery, C. W. [55TOM/DOM] Moore, C. E. [68MOO/MER] Moore, D. A. [98RAI/FEL] Morgan, J. J. [96STU/MOR2] Mori, T. [78IWA/FUJ] Moriyama, J. [69MOR/SAT] Morozova, M. P. [71ARI/MOR] Morrell, M. L. [64HAM/MOR] Morris, D. F. C. [65MOR/REE] Morris, J. P. [57MOR/ZEL] Moser, L. [24MOS/BEH] Moser, Z. [75MOS/FIT], [75MOS/FIT2] Mu, J. [82MU/PER] Muhammed, M. [98PLY/ZHA] Mukherjee, R. [86BHA/MUK] Muldagalieva, R. A. [95MUL/CHU] Muldagalieva, S. A. [93MUL/ISA] Muldrow, C. N. [58MUL] Muller, A. B. [85MUL], [92GRE/FUG] Müller, C. [79WEI/MUL]



605

Author Reference Müller, E. [30MUL/LUB] Müller, F. [78KOT/MUL] Mumme, W. G. [93NIC/ROB] Munoz, J. L. [86NOR/MUN] Murai, R. [76MUR/KUR] Murata, K. [28MUR] Murray, J. J. [74MUR/POT] Mussini, T. [2002BUC/RON] Muthmann, W. [04MUT/FRA] Nàer-Neumann, E. [85FER/GRE] Nagamori, M. [70NAG/ING] Nagypál, I. [90BEC/NAG], [91ZEK/NAG] Nair, V. S. K. [59NAI/NAN] Naito, M. [96UCH/GOR] Nancollas, G. H. [59NAI/NAN], [67NAN/TOR] Nandi, P. N. [79NAN/DES] Napijalo, M. L. [96DOJ/NOV] Napijalo, M. M. [96DOJ/NOV] Näsänen, R. [43NAS] Nash, P. [87SIN/NAS], [90LEE/NAS] Nassau, K. [73NAS/SHI] Naugle, D. G. [79DUB/NAU] Naumann, R. [2002BUC/RON] Navratil, J. D. [92FUG/KHO] Navrotsky, A. [68NAV/KLE], [71NAV], [73NAV], [73NAV2],

[76NAV/KAS] Neck, V. [2003GUI/FAN] Neilson, G. W. [83NEI/END], [97HOW/NEI] Netsvetaeva, V. S. [74NET/BAT] Netzel, D. A. [63NET/DRO] Neumann, B. [1894NEU] Neuschütz, D. [67GES/NEU] Neves, E. A. [82NEV/ANG] Newmann, A. A. [75NEW2] Nguyen-Trung, C. [92GRE/FUG] Nickel, E. H. [77NIC/ROB], [81NIC/BER], [93NIC/ROB] Nieuwenkamp, W. [33BIJ/NIE] Nikol'skaya, N. A. [71KLY/GLE] Nishimura, T. [88NIS/ITO] Nitsche, H. [2001LEM/FUG] Noläng, B. [2005OLI/NOL] Nordstrom, D. K. [86NOR/MUN] Novak, C. F. [98ALM/NOV] Novak-Adamič, D. M. [73NOV/COS] Novakoviç, L. [96DOJ/NOV]



606

Author Reference Novo-Gradac, K. H. [93ALL/BRO] Novotný, P. [85SOH/NOV] Nozue, T. [97NOZ/KOB] Nuttall, R. L. [79GOL/NUT], [82SAR/COV], [82WAG/EVA] Nyholm, R. S. [59GIL/NYH], [63PEN/BAI] Oakes, C. S. [2001OAK/RAI] Obendrauf, W. [93OBE/KAS] Obozova, L. A. [75BAR/MAS] Oelkers, E. H. [90OEL/HEL], [92SHO/OEL] Ogawa, S. [77OGA] Ogino, H. [61TAN/OGI], [63TAN/SAI] Öhman, L-O. [2005OLI/NOL] Ohtaki, H. [71OHT/BIE], [73KAW/OTS] Oikova, T. [79OIK] Ojkova, T. [79OJK/MAK] Oka, Y. [38OKA] Okaji, M. [76OKA/WAT] Olette, M. [63GEO/FER] Olin, Å. [2005OLI/NOL] Olofsson, I. V. [78SPI/SIN2], [79SPI/OLO] Olroyd, A. [96LUT/RIC] Omarova, F. M. [80OMA/SHA] O'Neill, H. S. C. [86HOL/NEI], [87NEI], [93NEI/POW] Oppermann, H. [80OPP/TOS] Oranskaya, M. A. [54SHC/TOL] Oriani, R. A. [54ORI/JON] Origlia-Luster, M. L. [2004BRO/MER] Osadchii, E. [90OSA/ROS] Osadchii, E. G. [2005OLI/NOL] Osswald, H. [34BRI/OSS] Östhols, E. [99WAN/OST], [2000GRE/WAN], [2000OST/WAN],

[2000WAN/OST] Oswald, H. R. [77OSW/ASP] Otsuka, H. [73KAW/OTS] Ouweltjes, W. [90COR/KON2] Owen, B. B. [58HAR/OWE] Paatero, J. [71PAA/HUM] Pabalan, R. T. [91PAP/PIT] Pack, R. T. [63CHR/IZA] Palmer, B. J. [99HOF/DAR] Palmer, D. A. [2003GUI/FAN] Pan, P. [97PAN/CAM] Pankratz, L. B. [76MAH/PAN], [84PAN/STU], [84PAN] Pant, G. N. [76PAN/PAT] Paoletti, O. [65PAO/SAB], [65PAO]



607

Author Reference Parker, G. A. [97BAH/PAR] Parker, V. B. [82WAG/EVA], [87GAR/PAR] Parsons, R. [74PAR], [85BAR/PAR] Paschina, G. [80CAM/LIC], [82MAG/PAS], [83LIC/PAS] Pascucci, G. [76BIA/CAR] Pasternak, M. P. [90PAS/TAY] Pathak, D. N. [76PAN/PAT] Patil, S. K. [76PAT/RAM] Patterson, B. A. [2002PAT/WOO] Pauling, L. [40PAU] Pavlinova, L. A. [71ARI/MOR] Pawel, R. E. [65PAW/STA] Pawlowski, T. [91SWI/PAW] Payne, S. L. [57MOR/ZEL] Peacock, M. A. [47PEA] Pearce, J. N. [37PEA/ECK] Pease, R. N. [26PEA/COO] Peiluck, R. V. [65JAM/BRO] Penneman, R. A. [60MCC/JON], [63PEN/BAI], [65COL/PET] Perlmutter, D. D. [82MU/PER] Perlmutter-Hayman, B. [73PER/SEC] Perrin, D. D. [64PER] Perron, G. [81PER/ROU] Persson, H. [71PER], [74PER], [76PER/EKS], [76PER] Petersen, Jr. H. [65COL/PET] Peterson, J. R. [78PIT/PET] Pethybridge, A. D. [80PET/TAB] Petot, C. [84RAY/PET] Petot-Ervas, G. [84RAY/PET] Petrov, G. V. [79PET/SHE], [79PET/SHE2], [80PET/SHE] Pfanhauser, W. [01PFA] Phillips, B. [63PHI/HUT] Piacente, V. [96BRU/PIA] Piccaluga, G. [80CAM/LIC], [82MAG/PAS], [83LIC/PAS],

[85MAG/MOR] Pierrard, J-C. [71PIE/HUG] Pinna, G. [80CAM/LIC], [83LIC/PAS] Pitzer, K. S. [73PIT/MAY], [73PIT], [74PIT/KIM], [74PIT/MAY],

[75PIT], [76PIT/SIL], [78PIT/PET], [79BRA/PIT], [79PIT], [91PAP/PIT], [91PIT]

Plyasunov, A. V. [94PLY/GRE], [97ALL/BAN], [97GRE/PLY], [97GRE/PLY2], [97PUI/RAR]

Plyasunova, N. V. [98PLY/ZHA] Pobegun, T. [62CAI/POB] Polgar, L. G. [72POL/HER]



608

Author Reference Polydoropoulos, K. [53SCH/POL] Ponomareva, V. L. [71ARI/MOR] Popov, C. G. [75LEV/GOL] Popov, N. I. [76ZHO/BEZ] Porto, R. [87CIA/IUL], [89IUL/POR] Pottel, R. [65POT] Potter, P. E. [2001LEM/FUG] Pottie, R. F. [74MUR/POT] Pottlacher, G. [93OBE/KAS] Poulson, S. R. [96POU/DRE] Pourbaix, M. [63DEL/ZOU] Powell, H. K. J. [73POW] Pownceby, M. I. [93NEI/POW] Pracht, G. [97PRA/LAN] Prasad, B. [56JEN/PRA], [68PRA] Pratt, J. N. [84COM/PRA] Pratt, K. W. [2002BUC/RON] Predel, B. [72PRE/RUG], [72PRE/RUG2] Preis, W. [98GAM/KON], [2001GAM/PRE], [2002GAM/WAL],

[2002WAL/PRE] Prescott, B. E. [73NAS/SHI] Prewitt, C. T. [74RAJ/PRE] Prue, J. E. [55BRO/PRU] Przepiera, A. [93PRZ/WIS] Ptak, W. [78KAR/MAL] Puigdomènech, I. [95GRE/PUI], [95SIL/BID], [97ALL/BAN],

[97PUI/RAR] Pupp, C. [74MUR/POT] Purtov, A. I. [66KAL/PUR] Pytkowicz, R. M. [79JOH/PYT], [79PYT] Quaresma, J. L. [89LOB/QUA] Rabbering, G. [75RAB/WAN] Rai, D. [97MAT/RAI], [98RAI/FEL], [2001OAK/RAI] Rajamani, V. [74RAJ/PRE] Rama Char, T. L. [58VAI/RAM] Ramachandran, A. [86JAC/KAL] Ramakrishna, V. V. [76PAT/RAM] Ramette, R. W. [83RAM/FAN], [86RAM] Rand, M. H. [95GRE/PUI], [95SIL/BID], [99RAR/RAN],

[2001LEM/FUG], [2003GUI/FAN] Randall, M. [61LEW/RAN] Randell, C. F. [80BAR/RAN], [81BAR/RAN] Rao, L. [97MAT/RAI] Rao, Y. K. [75MEY/WAR] Raoux, D. [80LAG/FON]



609

Author Reference Rapp, R. A. [63RAP] Rard, J. A. [87RAR3], [92RAR2], [97ALL/BAN], [97PUI/RAR],

[98ARC/RAR], [99RAR/RAN] Rau, H. [73RAU/GUE], [75RAU2], [76RAU] Ray, D. K. [84RAY/PET] Razmyslova, L. I. [74BLO/RAZ] Reddy, G. K. N. [63PEN/BAI] Redfern, S. A. T. [2001HEN/RED] Reed, G. L. [65MOR/REE] Regnault, L. P. [80ADA/BIL], [81ADA/BIL] Regnault, V. [1841REG] Reiterer, F. [80REI], [82GAM/REI] Retajczyk, T. F. [72RET/HUM] Rey, J. [80DUB/CLA] Richens, D. T. [97RIC] Richter, J. [79RIC/VRE] Rickard, D. T. [96LUT/RIC] Riddell, R. G. [66GOL/RID] Riglet, C. [87RIG/VIT], [89RIG/ROB], [90RIG] Ringwood, A. E. [62RIN] Riou, A. [87CUD/LEC] Rizzo, F. E. [67RIZ/BID] Robbins, C. R. [69LEV/ROB] Robbins, D. J. [76DAY/DIN] Robie, R. A. [84ROB/HEM], [95ROB/HEM] Robinson, B. W. [77NIC/ROB], [93NIC/ROB] Robinson, R. A. [49ROB/STO], [59ROB/STO] Robinson, S. C. [52ROB/BRO] Robinson, W. K. [60ROB/FRI] Robouch, P. [89RIG/ROB], [89ROB], [95SIL/BID] Robov, A. M. [73FED/SHM] Rodda, J. L. [66KOH/ROD] Rodebush, W. H. [30ROD/HEN] Roeder, P. [68CAM/ROE] Róg, G. [84ROG/BOR], [98ROG/KOZ], [2000ROG/KOZ] Rohmer, R. [39ROH] Rondini, S. [2002BUC/RON] Rosén, E. [90OSA/ROS], [2005OLI/NOL] Rosenblatt, G. M. [72BEH/ROS] Rosenblum, P. [60LIS/ROS] Rosengren, K. [56AHR/LAR], [56AHR/ROS] Rosenqvist, T. [54ROS], [64LEE/ROS] Rossat-Mignod, J. [80ADA/BIL], [81ADA/BIL] Rossetti-François, J. [52ROS] Rossini, F. D. [36BIC/ROS], [52ROS/WAG]



610

Author Reference Rossotti, F. J. C. [59ROS/ROS], [61ROS/ROS] Rossotti, H. [59ROS/ROS], [61ROS/ROS], [69ROS] Roth, C. [70ROT] Roux, A. [81PER/ROU] Rowson, J. [99LAN/MAH] Rubin, J. J. [65UIT/WIL] Rudat, A. [28JEL/RUD] Rudzitis, E. [67RUD/DEV] Ruge, H. [72PRE/RUG], [72PRE/RUG2] Rutner, E. [74RUT/HAU] Ruvinskii, O. E. [70TUR/RUV] Rydberg, J. [2001LEM/FUG] Sabat, B. B. [80CHI/SAB], [82CHI/SAB] Sabatini, A. [65PAO/SAB] Sadakane, K. [77SAD/KAW] Sadoc, A. [80LAG/FON] Sadowska, T. [69LIB/SAD], [80LIB/SAD], [82SAD/LIB] Sage, M. W. [08THO/SAG] Saha, A. [96SAH/SAH] Saha, N. [96SAH/SAH] Saito, Y. [63TAN/SAI] Saitton, B. [90OSA/ROS] Sajadi, S. A. A. [96SAH/SAH] Sakata, Y. [77KAS/SAK] Salvatore, F. [83FER/GRE], [85FER/GRE] Samuel, A. [23JOB/SAM], [43SAM] Sanchez, J. P. [81KUI/SAN] Sandino, M. C. A. [95GRE/PUI] Sano, K. [37SAN], [37SAN2], [53SAN] Sano, W. [90JUR/DOM] Saprykova, Z. A. [85SAP/CHI] Sarbar, M. [82SAR/COV] Sato, M. [70HUE/SAT], [97NOZ/KOB] Sato, N. [69MOR/SAT] Satyanarayan, D. [71DAS/DAS] Savvatimskii, A. I. [71LEB/SAV] Saxena, S. [97ALL/BAN] Scarrow, R. K. [69GRI/SCA] Scatchard, G. [36SCA], [76SCA] Schaefer, S. C. [81SCH] Schenck, R. [39SCH/FOR] Schiffman, R. A. [78SCH/MIL] Schildbach, R. [10SCH] Schindler, P. [63FEI/SCH], [63SCH], [65SCH/ALT] Schlesinger, M. E. [2002SCH]



611

Author Reference Schmalzried, H. [64TAY/SCH], [65TRE/SCH] Schneider, C. R. [59FRE/SCH] Schoch, E. P. [09SCH] Schott, J. [91AND/CAS] Schreiner, L. [34SIE/SCH] Schubnikow, L. [36TRA/SCH] Schuijff, A. [75RAB/WAN] Schumb, W. C. [23SCH] Schumm, R. H. [82WAG/EVA] Schwab, G.-M. [53SCH/POL] Schweitzer, A. [09SCH2] Scott, S. D. [76CRA/SCO] Secco, F. [73PER/SEC] Seidell, A. [65LIN/SEI] Seim, H. [94STO/FJE], [96SEI/FJE] Seki, S. [69SOR/KOS] Sekine, T. [76MUR/KUR] Seltz, H. [40SEL/DEW] Seminara, A. [78SIR/SEM] Sergeyeva, E. I. [92FUG/KHO] Shankar, J. [63SHA/DES] Shannon, R. D. [76SHA] Sharipov, M. S. [80OMA/SHA] Sharma, R. C. [79SHA/CHA], [80SHA/CHA] Shchigol, M. B. [61SHC] Shchukarev, S. A. [54SHC/TOL] Shehata, H. A. [87EMA/FAR], [93SHE] Sheka, Z. A. [56FIA/SHE] Shelton, R. A. J. [76GEE/SHE] Sheridan, X. [56MIL/SHE] Sherwood, R. C. [65UIT/WIL] Shevchuk, V. G. [79PET/SHE], [79PET/SHE2], [80PET/SHE] Shibata, N. [66SHI/FUJ] Shiever, J. W. [73NAS/SHI] Shimizu, K. [76SHI/TSU] Shirata, M. [77KAS/SAK] Shmyd'ko, I. I. [73FED/SHM] Shock, E. L. [88SHO/HEL], [89SHO/HEL], [89SHO/HEL2],

[92SHO/OEL] Short, E. L. [65MOR/REE] Shukla, A. K. [86JAC/KAL] Siemens, A. [04SIE] Sigel, H. [67SIG/BEC], [81BAN/KAD], [96SAH/SAH] Silberman, A. [36KAP/SIL] Sillén, L. G. [64SIL/MAR], [65BUR/LIL], [71SIL/MAR]



612

Author Reference Silva, R. J. [95SIL/BID] Silverman, D. C. [81SIL] Silvester, L. F. [76PIT/SIL], [78PIT/PET] Simaeva, L. S. [73FED/SHM] Simon, A. [39SIM/KNA] Simonson, J. M. [2002CHI/SIM] Singh, P. P. [78SPI/SIN], [78SPI/SIN2], [79SPI/OLO] Singleton, M. [87SIN/NAS] Siracusa, G. [73GUR/CAL], [78SIR/SEM] Skapsi, A. [32SKA/DAB] Skeaff, J. M. [73SKE/ESP], [85SKE/MAI] Skibsted, L. H. [86BJE/SKI] Slater, D. N. [65MOR/REE] Slough, W. [74SLO/JON] Sludskaja, N. P. [53GRI/SLU] Smirnov, Y. B. [71LEB/SAV] Smirnova, I. D. [71KLY/GLE] Smirnova, M. F. [68AND/KHA], [68AND/KHA2] Smirnova, M. N. [65EGO/SMI] Smith, MacF. W. [61DAV/SMI] Smith, R. I. [2001HEN/RED] Smith, R. M. [76SMI/MAR] Smith-Magowan, D. [82SMI/WOO] Smits, A. [18SMI/LOB] Smoes, S. [92BAL/DIC] Smurov, A. A. [38SMU] Söhnel, O. [85SOH/NOV] Sokolova, M. A. [56SOK] Solechnik, N. D. [85FIL/CHA] Solomon, L. R. [83SOL/BON] Song, B. [96SAH/SAH] Sorai, M. [69SOR/KOS] Spahiu, K. [83SPA], [97ALL/BAN], [97GRE/PLY2],

[2001LEM/FUG] Speelman, J. L. [85SKE/MAI] Spence, R. D. [59SPE/FOR] Spencer, P. J. [93KUB/ALC] Spitzer, J. J. [78SPI/SIN], [78SPI/SIN2], [79SPI/OLO] Spitzer, P. [2002BUC/RON] Spivakovskii, V. B. [69MAK/SPI], [91SPI/TRI] Sposito, G. [77MAT/SPO], [79MAT/SPO] Sridhar, R. [77CON/SRI] Stansbury, E. E. [65PAW/STA] Staples, B. R. [79GOL/NUT] Staveley, L. A. K. [82WHI/STA]



613

Author Reference Staveley, L. A. W. [74WOR/COW] Steele, B. C. H. [65STE/ALC] Steele, B. D. [04STE/JOH] Stokes, R. H. [49ROB/STO], [59ROB/STO], [91STO] Stølen, S. [91STO/GRO], [94STO/FJE], [95GRO/STO],

[96SEI/FJE] Stolyrova, T. A. [81STO], [82STO] Stout, J. W. [41STO/GIA], [54MAT/STO], [54STO/CAT],

[55CAT/STO], [55STO/CAT], [64STO/HAD], [66STO/ARC]

Strohmenger, J. M. [67BIN/STR] Stuehr, J. E. [72FRE/STU], [78FRE/STU] Stumm, W. [96STU/MOR2] Stuve, J. M. [78STU/FER], [84PAN/STU] Subbaraman, P. R. [69SUB/COR] Subrahmanya, R. S. [71BHA/SUB] Sudo, K. [52SUD], [64ALC/SUD] Suga, H. [69SOR/KOS] Sugar, J. [81COR/SUG] Sugiyama, K. [83KAT/SUG] Sukiennik, M. [78KAR/MAL] Sullivan, J. C. [2001LEM/FUG] Suzuki T. [97NOZ/KOB] Sverjensky, D. A. [89SHO/HEL2], [92SHO/OEL] Swanson, H. E. [53SWA/TAT] Swiatek, J. [91SWI/PAW] Sykes, C. [38SYK/WIL] Taba, S. S. [80PET/TAB] Tamamushi, R. [75TAM] Tanaka, M. [69FUN/TAN] Tanaka, N. [61TAN/OGI], [63TAN/SAI] Tanger, IV, J. C. [88TAN/HEL] Taniewska-Osinska, S. [61TAN2] Tantzov, N. V. [24TAN] Tare, V. B. [85MEH/TAR] Tareen, J. A. K. [91TAR/FAZ] Tatge, E. [53SWA/TAT] Taylor, G. [59SPE/FOR] Taylor, J. B. [58TAY/HEY] Taylor, J. R. [82TAY] Taylor, R. D. [90PAS/TAY] Taylor, R. S. [76TAY/DIE] Taylor, R. W. [64TAY/SCH] Tepavitcharova, S. [89LIL/TEP] Terrier, C. [77BIL/TER], [80ADA/BIL], [81ADA/BIL]



614

Author Reference Theberge, S. [96LUT/RIC] Thiel, A. [14THI/GES] Thiemann, W. [34BEN/THI] Thoenen, T. [99THO] Thompson, D. M. [08THO/SAG] Thomsen, J. [1883THO], [06THO] Threadgold, I. M. [59WIL/THR], [63THR] Tialowska, H. [75LIB/TIA] Tilden, W. A. [02TIL], [02TIL2] Tillett, D. M. [82SMI/WOO] Timofeeva, T. G. [93UVA/TIM] Timofeyev, B. I. [84LAV/TIM] Toda, Y. [86TOD/HAS], [94HAS/TOD] Tolmacheva, T. A. [54SHC/TOL] Tomlinson, J. R. [55TOM/DOM] Torrance, K. [67NAN/TOR] Toschew, A. [80OPP/TOS] Tóth, I. [92BAN/BLI] Tozawa, K. [88NIS/ITO] Tran, T. [80TRA/BRU] Trapeznikowa, O. [36TRA/SCH] Tremaine, P. R. [80TRE/LEB2] Trémillon, B. [58TRE] Tretjakow, J. D. [65TRE/SCH] Tribalat, S. [62TRI/CAL], [64TRI/CAL] Trizna, L. G. [91SPI/TRI] Trojan, M. [90TRO] Tsertsov, V. N. [75LEV/GOL] Tsierkezos, N. G. [2000TSI/MOL], [2001MOL/TSI] Tsuchihashi, N. [76SHI/TSU] Tsujikawa, I. [78ENO/TSU] Tuck, D. G. [97BAH/PAR] Turner, D. R. [81TUR/WHI] Tur'yan, Y. I. [68MAL/TUR], [70TUR/RUV], [71TUR/MAL] Tye, F. L. [80BAR/RAN], [81BAR/RAN] Tysyachnyi, V. P. [70TYS/KSE] Uchida, E. [96UCH/GOR] Uesawa, A. [97NOZ/KOB] Ullman, W. J. [2001LEM/FUG] Uloth, R. [26JEL/ULO], [26JEL/ULO2] Uvaliev, Y. K. [93UVA/TIM] Uvarov, L. [62VAG/UVA] Vacca, A. [65PAO/SAB] Vagramyan, A. [62VAG/UVA] Vaid, J. [58VAI/RAM]



615

Author Reference Valla, E. [11AGE/VAL] Van Deventer, E. H. [67RUD/DEV] Van Geet, A. L. [64GEE/HUM] Van Uitert, L. G. [65UIT/WIL] Vanden Bosche, E. G. [29HAR/BOS] Vannerberg, N.-G. [70HAL/VAN] Varet, R. [1896VAR] Vasca, E. [89IUL/POR] Vasiç, M. V. [80MIL/BUG] Vasileff, H. D. [50VAS/GRA] Vasilev, V. A. [79VOR/VAS] Vasil'ev, V. P. [56YAT/VAS2], [62VAS], [84VAS/VAS],

[86VAS/DMI] Vasil'eva, V. N. [84VAS/VAS], [86VAS/DMI] Vasilyev, V. P. [56YAT/VAS] Veits, I. V. [82GLU/GUR] Venal, W. V. [73VEN/GEI] Vidavskii, L. M. [92VID/KOR] Vierling, F. [95KAH/CRO] Vilbrandt, F. C. [23VIL/BEN] Vitorge, P. [87RIG/VIT], [89RIG/ROB], [93GIF/VIT2],

[94CAP/VIT], [95CAP/VIT], [2001LEM/FUG] Voevoda, N. Y. [93MAK/VOE] Vogel, W. M. [68VOG] Vogt, N. [2001VOG] Voigt, A. [36BIL/VOI] Vollmer, O. [66VOL/KOH] Von Brintzinger, H. [34BRI/OSS] Von der Forst, P. [39SCH/FOR] Von Dickert, F. [74DIC/HOF] Von Gesell, H. [67GES/NEU] Von Kiss, A. V. [41KIS/ACS] Von Sieverts, A. [34SIE/SCH] Vorob'ev, A. F. [79VOR/VAS] Vreuls, W. [79RIC/VRE] Wagman, D. D. [52ROS/WAG], [73HUL/DES], [82WAG/EVA],

[89COX/WAG] Wagner, C. [57KIU/WAG] Wagner, Jr., J. B. [85MEH/TAR] Wakita, H. [82WAK/ICH] Wallner, H. [2001GAM/PRE], [2002GAM/WAL],

[2002WAL/GAT], [2002WAL/PRE] Wang, P. [90ARC/WAN]



616

Author Reference Wanner, H. [88WAN], [91WAN], [92GRE/FUG], [95SIL/BID],

[99RAR/RAN], [99WAN/OST], [2000GRE/WAN], [2000OST/WAN], [2000WAN/OST], [2001LEM/FUG]

Wanrooy, J. [75RAB/WAN] Warburton, J. B. [67ANT/WAR] Warncke, H. [55KRA/WAR] Warner, J. S. [75MEY/WAR], [77CON/SRI] Warshaw, I. [63PHI/HUT] Watanabe, H. [82WAT], [93WAT] Watanabe, M. [33WAT] Watanabe, T. [76OKA/WAT] Waters, D. N. [65MOR/REE] Watt, G. D. [63CHR/IZA] Watters, J. I. [57LAM/WAT] Weenk, J. W. [75WEE/KOE] Weibke, F. [36BIL/VOI] Weidner, J. W. [98JAI/ELM] Weigel, D. [72AUF/CAR], [72AUF/CAR2] Weigel, O. [07WEI] Weingärtner, H. [79WEI/HER], [79WEI/MUL] Weitbrecht, G. [42FRI/WEI] Weller, W. W. [64WEL/KEL], [65WEL] Wendlandt, W. W. [74WEN] Westgren, A. [38WES] Westrum, Jr., E. F. [91STO/GRO], [94STO/FJE] Wheaton, R. J. [79LEE/WHE] Whiffen, D. H. [79WHI2] White, H. W. [74WHI] White, Jr., H. J. [87GAR/PAR] White, M. A. [82WHI/STA] Whitfield, M. [79WHI], [81TUR/WHI] Whiting, K. S. [92WHI] Wilkinson, H. [38SYK/WIL] Willems, H. W. V. [27JON/WIL] Williams, D. R. [74GRA/WIL] Williams, H. J. [65UIT/WIL] Williams, K. L. [59WIL/THR], [60WIL2] Williams, T. [62WIL2] Wilson, A. J. C. [36BRO/WIL] Wilson, D. W. [61LIS/WIL] Wilson, G. S. [2002BUC/RON] Wingard, M. R. [66GOL/RID] Wisniewski, M. [93PRZ/WIS] Witte, J. [66BOD/DEH] Wöhler, L. [21WOH/BAL]



617

Author Reference Wood, R. H. [82SMI/WOO] Woolley, E. M. [2002PAT/WOO], [2004BRO/MER] Woontner, L. [52GAY/WOO] Worswick, R. D. [74WOR/COW] Woyski, M. M. [51HOO/WOY] Wróbel, S. [99MIK/MIG] Wulff, C. A. [66GOL/RID] Yagi, T. [74YAG/MAR] Yamatera, H. [81YOK/YAM] Yashkova, V. I. [86VAS/DMI] Yashko-Zakharova, O. E. [69KUL/ERS] Yatsimirskii, K. B. [56YAT/VAS], [56YAT/VAS2], [58YAT/KOR] Yngve, V. [16DER/YNG] Yokoyama, H. [81YOK/YAM] Yoshihara, K. [66SHI/FUJ] Yui, M. [98RAI/FEL] Yuldasheva, V. M. [84LAV/TIM] Yund, R. A. [61YUN], [62KUL/YUN] Yungman, V. S. [82GLU/GUR] Zador, S. [64ALC/SUD] Zakowski, J. [25JEL/ZAK] Zakulski, W. [75MOS/FIT] Zamboni, R. [97MAL/ZAM] Zäske, P. [64KOH/ZAS] Zavrazhnova, D. M. [71KLY/GLE] Zawadzki, J. [09BRU/ZAW] Zékány, L. [91ZEK/NAG] Zellars, G. R. [57MOR/ZEL] Zhambekov, M. I. [93MUL/ISA] Zhang, J.-Z. [94ZHA/MIL] Zhang, Y. [98PLY/ZHA] Zhao, J. [96SAH/SAH] Zhao, Y.-H. [2001LAW/ZHA] Zhorov, V. A. [76ZHO/BEZ] Ziemniak, S. E. [89ZIE/JON] Zinevich, N. I. [71BUR/ZIN] Zirel'nikov, V. I. [88EFI/EVD]

CHEMICAL THERMODYNAMICS OF NICKEL · 2013-09-09 · of Inorganic and Analytical Chemistry,...

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