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CHEMISTRY - SILBERBERG 8E

CH.6 - THERMOCHEMISTRY

CONCEPT: ENERGY CHANGES AND ENERGY CONSERVATION

_______________________ is the branch of physical science concerned with heat and its transformations to and from

other forms of energy.

_______________________ is the branch of chemistry that deals with the heat involved in chemical and physical changes.

Energy Changes and Energy Conservation

• The ______________________ is the specific part of the universe that we are focused on.

• The _________________________ deals with everything outside of it.

When talking about the movement of energy or heat between the ____________________ & ____________________ we

use the terms: Endothermic & Exothermic.

2 H2 (g) + O2 (g) 2 H2O (l) +

HEAT

HEAT

+ 2 HgO (s) 2 Hg(l) + O2 (g)

EXAMPLE: Classify each of the following process as either exothermic or endothermic: a) Fusion of Ice.

b) Sublimation of CO2.

c) Vaporization of aqueous water.

d) Deposition of chlorine gas.

e) Condensation of water vapor.



CONCEPT: ENERGY FLOW TO AND FROM A SYSTEM

The _______ Law of Thermodynamics states that energy cannot be created nor destroyed, but only converted from one

form to another.

In chemistry, we are normally concerned with the energy changes associated with the system, not with its surroundings.

∆E = q + w q = ∆H (enthalpy) w = - P∆V

∆E =

q = * For q: (+) when system __________, __________, __________, heat or energy,

(-) when system __________, __________, __________, __________ heat or energy.

w = * For w : (+) when work done _____ system _____ the surroundings. Key word: volume ______________

(-) when work done _____ system _____ the surroundings. Key word: volume _______________

EXAMPLE: Which of the following signs on q and w represent a system that is doing work on the surroundings, as well as

losing heat to the surroundings?

a) q = - , w = - b) q = +, w = + c) q = -, w = + d) q = +, w = -



PRACTICE: ENERGY FLOW TO AND FROM A SYSTEM

EXAMPLE: An unknown gas expands in a container increasing the volume from 4.3 L to 8.2 L at a constant pressure of 931

mmHg.

a. Calculate the work done (in kJ) by the gas as it expands. (1 L · atm = 101.3 J)

b. Using part A, calculate the internal energy of the system if the system absorbs 2.3 kJ of energy.

c. Using part B, calculate the internal energy of the system if the system does work against a vacuum.

PRACTICE: The reaction of nitrogen with hydrogen to make ammonia has an enthalpy, ∆H = - 92.2 kJ:

N2 (g) + 3 H2 (g) 2 NH3 (g)

What is in the internal energy of the system if the reaction is done at a constant pressure of 20.0 atm and the volume compresses from 10 L to 5 L?



CONCEPT: CONSTANT-VOLUME CALORIMETRY

Every object has its own _________________________ (C), the quantity of heat required to change its temperature by 1 K.

C = qΔT

[in units of JK ]

_________________________________ (c), the quantity of heat required to change 1 gram of a substance by 1 degree K.

c = [in units of J

g•K ]

If we know c of a substance, we can algebraically solve the amount of heat absorbed or released:

q =

EXAMPLE: At constant volume, the heat of combustion of a particular compound is – 4621.0 kJ/mol. When 2.319 grams of

this compound (molar mass = 192.75 g/mol) was burned in a bomb calorimeter, the temperature of the calorimeter

(including its contents) rose by 3.138oC. What is the heat capacity of the calorimeter in J/K?



PRACTICE: CONSTANT-VOLUME CALORIMETRY

EXAMPLE: In an experiment a 9.87 carat (1 carat = 0.200g) diamond is heated to 72.25oC and immersed in 22.08 g of

water in a calorimeter. If the initial temperature of the water was 31.0oC what is the final temperature of the water? (cdiamond =

0.519J

g• oC) (cwater = 4.184

Jg• oC

).

PRACTICE 1: A sample of copper absorbs 35.3 kJ of heat, which increases the temperature by 25oC, determine the mass

(in kg) of the copper sample if the specific heat capacity of copper is 0.385 J

g• oC.

PRACTICE 2: 50.00 g of heated metal ore is placed into an insulated beaker containing 822.5 g of water. Once the metal

heats up the final temperature of the water is 32.08oC. If the metal gains 14.55 kJ of energy, what is the initial temperature

of the water?



CONCEPT: CONSTANT-PRESSURE CALORIMETRY The ______________ of a reaction can be calculated through the use of

a coffee-cup calorimeter.

EXAMPLE: You place 50.0 mL of 0.100 M NaOH into a coffee-cup calorimeter at 50.00oC and carefully add 75.0 mL of 0.100 M HCl,

also at 50.00oC. After stirring, the final temperature is 76.12oC. (Heat capacity and density of water: 4.184 J

g• oC and 1.00

gmL ).

HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l)

a) Calculate qsoln (in J)

b) Calculate the enthalpy, ∆Hrxn (in J/mol), for the formation of water.



CONCEPT: HEAT SUMMATION

Many reactions are difficult, even impossible, to carry out in a single chemical step.

• They may often times require multiple steps to get to the final products.

__________ Law states that the enthalpy change (∆H) of an overall process is the sum of the enthalpy changes of its

individual steps.

EXAMPLE: For the following example calculate the unknown ∆H from the given ∆H values of the other equations.

Calculate the ∆Hrxn for

S(s) + 32

O2 (g) SO3 (g) ∆H = ?

Given the following set of reactions:

12

S (s) + 12

O2 (g) 12

SO2 (g) ∆H1 = – 296.8 kJ

2 SO3 (g) 2 SO2(g) + O2 (g) ∆H2 = 198.4 kJ



PRACTICE: HEAT SUMMATION

A. Calculate the ∆Hrxn for

CO(g) + NO (g) CO2 (g) + 12

N2 (g) ∆H = ?

Given the following set of reactions:

CO2 (g) CO (g) + 12

O2 (g) ∆H1 = 283.0 kJ

N2 (g) + O2 (g) 2 NO (g) ∆H2 = 180.6 kJ

B. Calculate the ∆Hrxn for

ClF (g) + F2 (g) ClF3 (g) ∆Hrxn = ?

Given the following reactions:

Cl2O (g) + F2O 2 ClF (g) + O2 (g) ∆Hrxn = - 167.4 kJ

4 ClF3 (g) + 4 O2 (g) 2 Cl2O (g) + 6 F2O (g) ∆Hrxn = 682.8 kJ

2 F2 (g) + O2 (g) 2 F2O (g) ∆Hrxn = -181.7 kJ



CONCEPT: STANDARD HEATS OF FORMATION (∆HRXN)

In a _______________ equation, 1 mole of a compound forms from its elements. The ______________________________

(∆Hof) is the enthalpy change for the chemical equation when all the substances are in their standard states.

C (graphite) + 2 H2 (g) CH4 (g) ΔHfo = −74.9kJ

When calculating ∆Hof remember:

1) An element in its standard state (elemental state) is given an ∆Hof of zero.

Ex: Na (s) P4 (s) Cl2 (g) S8 (g)

2) Most compounds have a negative ∆Hof. 3) To find the ∆Hrxn use the following formula:

ΔHrxno = ΔHf(products)

o −ΔHf(reac tan ts)o

EXAMPLE: The oxidation of ammonia is given by the following reaction:

4 NH3 (g) + 5 O2 (g) 4 NO (g) + 6 H2O (g)

Calculate the ∆Horxn if the ΔHfo

value for NH3 , NO and H2O are – 45.9 kJ/mol, 90.3 kJ/mol and – 241.8 kJ/mol

respectively.

PRACTICE: Ibuprofen is used as an anti-inflammatory agent used to deal with pain and bring down fevers. If it has a

molecular formula of C13H18O2 determine the balanced chemical equation that would give you directly the enthalpy of

formation for ibuprofen.



PRACTICE: STANDARD HEATS OF FORMATION (∆HRXN)

EXAMPLE: Use the following bond strength values (kJ/mol):

C–H 412 C–O 360 C=O 743 C–C 348 H–H 436

C=C 611 C≡C 837 C≡O 1072 O–H 464 O=O 498

Calculate the enthalpy of the reaction shown in the formula below:

H–C≡C–H + H–H + C=O

C

C C

O

HH

H H



4. An unknown gas expands in a container increasing the volume from 8.7 L to 18.9 L at a constant pressure of 1380 mmHg. (a) Calculate the work done (in J) by the gas as it expands. (1 L· atm = 101.3 J). (b) Calculate the internal energy of the system if the system absorbs 235.5 J of energy.

(c) Calculate the internal energy of the system if work was done against a vacuum. (1 L · atm = 101.3 J).



8. Calculate the amount of heat absorbed when 12.0 g of water is heated from 20oC to 100oC. (c = 4.184 J/g· oC).



9. 101.3 g of an unknown metal has an initial temperature of 25oC. If it absorbs 639.1 J of energy to obtain a final

temperature of 32.01oC identify the unknown metal.

Metal Specific Heat Capacity (J/g·oC)

Au 0.129

Fe 0.444

Al 0.900

Hg 0.139



10. Which substance has the highest molar heat capacity?

a) Copper (specific heat Cu (s): 0.39J

g ⋅ oC )

b) Silver (specific heat Ag (s): 0.23J

g ⋅ oC )

c) Iron (specific heat Fe(s): 0.46J

g ⋅ oC )

d) Lead (specific heat Pb (s): 0.13J

g ⋅ oC )



11. 25.00 g of heated metal ore is placed into an insulated beaker containing 615.5 g of water at 42.18oC. If the metal

gains 19.11 kJ of energy, what is the final temperature of the water? (cwater = 4.184 J/g · oC).



12. If 53.2 g Al at 25.0 oC is placed in 110.0 g H2O at 90 oC, what is the final temperature of the mixture? The specific

heat capacities of water and aluminum are 4.184 J/g · oC and 0.897 J/g · oC, respectively.



13. A 20.0 g sample of iron (specific heat Fe (s) = 0.46 J

g ⋅ oC ) has an initial temperature of 30.2 oC. If 0.310 kJ are

applied to the iron sample, calculate its final temperature.



14. A sample of H2O (l) containing 2.50 moles has a final temperature of 45.0 oC. If the sample absorbs 3.00 kJ of

heat, what is the initial temperature of the H2O (l)? The specific heat of H2O (l) is 4.184 J

g ⋅ oC .



Determine the heat released when 80.0 g H2O (l) at 90 oC is cooled to ice at – 10.0 oC. Specific Heat of H2O (l) = 4.184

Jg ⋅ oC . Specific Heat of H2O (s) = 2.09

Jg ⋅ oC . Heat of Fusion of water = 333

Jg .



If 1050 g of aluminum metal with a specific heat capacity of 0.902 J

g ⋅ oC at – 20 oC is placed in liquid water at 0.00 oC,

how many grams of liquid water are frozen by the time that the aluminum metal has warmed to – 10 oC? Heat of Fusion of

water = 333 Jg .



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