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AP Thermodynamics Packet Unit 9
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AP Thermodynamics Packet Unit 9
Thermodynamics Review
Students should be able to demonstrate an understanding of the following essential knowledge:
 3.C.2 Net changes in energy for a chemical reaction can be endothermic or exothermic.
 5.A.1 Temperature is a measure of the average kinetic energy of atoms and molecules.
 5.A.2 The process of kinetic energy transfer at the particulate scale is referred to in this course as heat transfer,
and the spontaneous direction of the transfer from a hot to a cold body.
 5.B.2 When two systems are ion contact with each other and are otherwise isolated, the energy that comes out of
one system is equal to the energy that goes into the other system. The combined energy of the two systems
remains fixed. Energy transfer can occur through either heat exchange or work.
 5.B.3 Chemical systems undergo three main processes that change their energy: heating/cooling, phase transitions,
and chemical reactions.
 5.B.4 Calorimetry is an experimental technique that is used to determine the heat exchanged/transferred in a
chemical system.
 5.C.1 Potential energy is associated with a particular geometric arrangement of atoms or ions and the electrostatic
interactions between them.
 5.D.1 Potential energy is associated with the interaction of molecules; as molecules draw near each other, they
experience an attractive force.
Answer the following questions using the heat formula. Show work with units and significant
figures.
1. How many Joules of energy are needed to change the temperature of 100.0 grams of water from
20.0C to 40.0C?
2. How many kilojoules of energy are needed to change the temperature of 15.0 grams of water from
35.0C to 75.0C?
3. If the temperature of water is changed from 10.0C to 35.0C by the addition of 350.0J, how many
grams were heated?
4. If the temperature of water is changed from 100.0C to 250.0C by the addition of 5000.0J, how
many grams were heated?
5. If 3500.0J of energy are applied to 150.0 grams of water at 50.0C, what is the final temperature?
6. If 425J of energy are released from 25.0 grams of water at 25.0C, what is the final temperature?
7. Look at the rearranged equation for heat, solved for specific heat. What are the units of specific
heat based on this rearranged equation?
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AP Thermodynamics Packet Unit 9
8. What is the specific heat of silver if an 80.0 gram sample is heated from 24.0C to 49.0C by adding
468.2J?
9. What is the specific heat of copper if a 75.0 gram sample of copper is heated from 20.0C to 24.0C
by adding 117J?
10. The specific heat of aluminum is 0.88J/gC. How many joules of heat does it take to heat 50.0 grams
of aluminum from 20.0 to 70.0C?
11. What is the change in temperature when 3.00 grams of Iron (specific heat = 0.45J/gC) is subjected
to 350.0.J of energy?
12. What mass of Aluminum (specific heat = 0.902 J/gC) can be heated from 25.0C to 90.0C with the
addition of 100.0J of heat?
13. How many joules of heat must be released in order to change the temperature of 50.0 grams of air
(specific heat 1.01 J/gC) from 35.0C to 25.0C?
Base answers to questions 14-17 on the diagram:
14. How much heat is added to change the
substance from the coldest to the warmest
pure liquid state?
15. If the specific heat of this substances in the liquid
phase is 3.56 J/gC calculate how many grams were
heated in question 14?
16. If the same sample is heated as a solid at 40C to 60C will it have the same specific heat? Show work
to support your answer.
17. How much heat will need to be added to the same sample in order to raise the temperature from
20.0C to 40.0C? (Notice this is not on the graph, you must calculate it based on your answers
above.)
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AP Thermodynamics Packet Unit 9
AP Calorimetric Calculations
Students should be able to demonstrate an understanding of the following essential knowledge:
 5.B.4 Calorimetry is an experimental technique that is used to determine the heat exchanged/transferred in a
chemical system.
The first law of thermodynamics: Energy can never be created nor destroyed. Therefore, the energy of
the universe is constant.
1. It takes 585J of energy to raise the temperature of 125.6g of Hg from 20.0 to 53.5C. Calculate the
specific heat of Hg.
2. A 46.2g sample of Cu is heated to 95.4C and placed in a calorimeter containing 75.0g of water at
19.6C. The final temperature inside the calorimeter equals 21.8C. Calculate the specific heat of
copper.
3. The specific heat of aluminum is 0.9000J/gC and the density is 2.71g/cm3.
a. Calculate the energy needed to raise the temperature of 4.36x105cm3 block from 22.8 to
94.6C.
b. Calculate the molar heat capacity of aluminum.
4. A piece of iron with a mass of 56.0g and specific heat of 0.45J/gC is placed in 155g of water at
21.0C. The final temperature is 33.5C. Calculate the original temperature of the iron.
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AP Thermodynamics Packet Unit 9
Enthalpy
Students should be able to demonstrate an understanding of the following essential knowledge:
 3.C.2 Net changes in energy for a chemical reaction can be endothermic or exothermic.
 5.C.2 The net change during a reaction is the sum of the energy required to break the bonds in the reactant
molecules and the energy released in forming the bond of the product molecules. The net change in energy may be
positive for endothermic reactions where energy is required or negative for exothermic reactions where energy is
released.
1. Consider the reaction:
2Mg + O2  2MgO
a. Is this reaction endothermic or exothermic?
ΔH=-1204kJ/mol
b. Calculate the heat transferred when 3.60g of Mg reacts with excess oxygen.
c. How many grams of MgO are produced during the enthalpy change of -96.0kJ?
2. Consider the reaction:
2AgBr + Cl2  2AgCl + Br2
ΔH=-55.2kJ/mol
a. Calculate the heat transferred by 37.56g of silver bromide reacting with excess chlorine.
b. How many grams of liquid bromine are produced when the enthalpy change equals -106kJ?
c. Calculate the heat of reaction when 60.0g of solid silver chloride are produced.
3. Given the following information: C2H6 + 7/2 O2  2CO2 + 3H2O
a. What is the enthalpy change for the reverse reaction?
ΔH = -91 kJ/mol
b. Which reaction (forward or reverse) is more favorable?
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AP Thermodynamics Packet Unit 9
Enthalpy of Formation
Students should be able to demonstrate an understanding of the following essential knowledge:
 3.C.2 Net changes in energy for a chemical reaction can be endothermic or exothermic.
 5.C.2 The net change during a reaction is the sum of the energy required to break the bonds in the reactant
molecules and the energy released in forming the bond of the product molecules. The net change in energy may be
positive for endothermic reactions where energy is required or negative for exothermic reactions where energy is
released.
1. The thermite reaction is highly exothermic and is used for welding :
2Al(s) + Fe2O3  2Fe(s) + Al2O3(s)
Calculate the heat of this reaction using enthalpies of formation.
2. Calculate the enthalpy of these reactions using enthalpies of formation:
a. 4FeO(s) + O2(g)  2Fe2O3(s)
b. SiCl4(l) + 2H2O(l)  SiO2(s) + 4HCl(aq)
c. NH3(g) + HCl(g)  NH4Cl(s)
d. MgO(s) + H2O(l)  Mg(OH) 2 (s)
e. C2H5OH(l) + 3O2 (g)  2CO2 (g) + 3H2O(g)
Reminder of an old Enthalpy Question:
3. Calculate the change in energy that accompanies the following reaction given the data below.
H2(g) + F2(g) → 2 HF(g) Bond Type Bond Energy H−H 432 kJ/mol F−F 154 kJ/mol H−F 565 kJ/mol
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AP Thermodynamics Packet Unit 9
Hess’s Law
Students should be able to demonstrate an understanding of the following essential knowledge:
 3.C.2 Net changes in energy for a chemical reaction can be endothermic or exothermic.
 5.C.1 Potential energy is associated with a particular geometric arrangement of atoms or ions and the electrostatic
interactions between them.
 5.C.2 The net change during a reaction is the sum of the energy required to break the bonds in the reactant
molecules and the energy released in forming the bond of the product molecules. The net change in energy may be
positive for endothermic reactions where energy is required or negative for exothermic reactions where energy is
released.
1. Given the following data, calculate the heat of S(s) + O2(g)  SO2(g)
S(s) + 3/2 O2(g)  SO3(g)
ΔH= -395.2kJ/mol
2SO2(g) + O2(g)  2SO3(g)
ΔH= -198.2kJ/mol
2. Given the following data, calculate the heat of C6H4(OH)2(aq) + H2O2 (aq)  C6H4O2 (aq) + 2H2O(l)
C6H4(OH)2(aq)  C6H4O2 (aq) + H2(g)
H= 177.4kJ/mol
H2 (g) + O2 (g)  H2O2 (aq)
ΔH= -191.2kJ/mol
H2 (g) + ½ O2 (g)  H2O(g)
ΔH= -241.8kJ/mol
H2O(g)  H2O(l)
ΔH= -43.8kJ/mol
3. Given the following data, calculate the heat of
2CO2 (g) + H2O(g)  C2H2 (g) + 5/2 O2 (g)
C2H6(g)  C2H2 (g) + 2H2 (g)
ΔH= 850.5kJ/mol
H2 (g) + ½ O2 (g)  H2O(g)
ΔH= -641.2kJ/mol
C2H6(g) + 7/2 O2 (g)  2CO2 (g) + 3H2O(g)
ΔH= -2547kJ/mol
4. Given the following data, calculate the heat of NO(g) + O(g)  NO2(g)
2O3 (g)  3O2(g)
ΔH= -427kJ/mol
O2 (g)  2O(g)
ΔH= 495kJ/mol
NO(g) + O3(g)  NO2 (g) + O2 (g)
ΔH= -199kJ/mol
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AP Thermodynamics Packet Unit 9
Thermodynamics Practice
1. Propane, C3H8, is a hydrocarbon that is commonly used as fuel for cooking.
(a) Write a balanced equation for the complete combustion of propane gas, which yields CO2(g) and H2O(l).
(b) Calculate the volume of air at 30C and 1.00 atmosphere that is needed to burn completely 10.0 grams of
propane. Assume that air is 21.0 percent O2 by volume.
(c) The heat of combustion of propane is -2,220.1 kJ/mol. Calculate the heat of formation, ΔHf, of propane given
that ΔHf of H2O(l) = -285.3 kJ/mol and ΔHf of CO2(g) = -393.5 kJ/mol.
(d) Assuming that all of the heat evolved in burning 30.0 grams of propane is transferred to 8.00 kilograms of
water (specific heat = 4.18 J/g.K), calculate the increase in temperature of water.
2. Pentane, C5H12, is a hydrocarbon used in the production of Styrofoam and is present in certain fuels.
(a) Write a balanced equation for the complete combustion of pentane gas, which yields CO2(g) and H2O(l).
(b) Calculate the volume of air at 25C and 1.00 atmosphere that is needed to burn completely 50.5 grams of
pentane. Assume that air is 21.0 percent O2 by volume.
(c) The heat of combustion of pentane is -3,285.3 kJ/mol. Calculate the heat of formation, ΔHf, of pentane given
that ΔHf of H2O(l) = -285.3 kJ/mol and ΔHf of CO2(g) = -393.5 kJ/mol.
(d) Assuming that all of the heat evolved in burning 50.5 grams of pentane is transferred to 10.0 kilograms of
water (specific heat = 4.18 J/g.K), calculate the increase in temperature of water.
3. (a) The specific heat of fluorine gas is 0.037 J/gK. Calculate the molar heat capacity (in J/molK) of fluorine
gas.
(b) The molar heat capacity of a compound with the formula C4H10SO is 43.6 J/molK. Calculate the specific
heat, c, of this substance.
4. Given the following data:
Calculate ∆Hºrxn for the reaction:
S(s) + 3/2 O2(g)  SO3(g)
2 SO2(g) + O2  2 SO3(g)
S(s) + O2(g)  SO2(g)
∆Hfº =-395.2 kJ
∆Hfº=-198.2 kJ
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AP Thermodynamics Packet Unit 9
Work
Students should be able to demonstrate an understanding of the following essential knowledge:
 5.B.2 When two systems are ion contact with each other and are otherwise isolated, the energy that comes out of
one system is equal to the energy that goes into the other system. The combined energy of the two systems
remains fixed. Energy transfer can occur through either heat exchange or work.
Energy is often defined as the “ability to do work”.
Algebraic signs of q:


∆E = q (heat) + w (work)
+q if heat absorbed
–q if heat released
Algebraic sign of w:



+ w if work done on the system (i.e., compression)
−w if work done by the system (i.e., expansion)
When related to gases, work is a function of pressure. Pressure is defined as force per unit of area,
so when the volume is changed work was either done on the gas or by the gas.
work = −P∆V
1. Calculate ∆E for a system undergoing an endothermic process in which 15.6 kJ of heat flows and where 1.4
kJ of work is done on the system.
2. Calculate the work associated with the expansion of a gas from 46 L to 64 L at a constant external pressure
of 15 atm.
3. A balloon is being inflated to its full extent by heating the air inside it. In the final stages of this process, the
volume of the balloon changes from 4.00 × 106 L to 4.50 × 106 L by the addition of 1.3 × 108 J of energy as
heat. Assuming that the balloon expands against a constant pressure of 1.0 atm, calculate ∆E for the
process. (To convert between L ⋅ atm and J, use 1 L ⋅ atm = 101.3 J.)
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AP Thermodynamics Packet Unit 9
Entropy
Students should be able to demonstrate an understanding of the following essential knowledge:
 5.E.1 Entropy is a measure of the dispersal of matter and energy.
Entropy is the degree of disorder or randomness in a substance. The symbol for change in entropy
is ΔS.
Solids are very ordered and have low entropy. Liquids and aqueous ions have more entropy because they
move about more freely, and gases have an even larger amount of entropy. According to the Second Law
of Thermodynamics, nature is always proceeding to a state of higher entropy.
Determine whether the following reactions show an increase or decrease in entropy.
1. 2KClO3(s) → 2KCl(s) + 3O2(g)
________________________
2. H2O(l) → H2O(s)
________________________
3. N2(g) + 3H2(g) → 2NH3(g)
________________________
4. 2NaCl(s) → 2Na(s) + Cl2 (g)
________________________
5. KCl(s) → KCl(l)
________________________
6. CO2(s) → CO2(g)
________________________
7. H+(aq) + C2H3O2-(aq) → HC2H3O3(l)
________________________
8. C(s) + O2(g) → CO2(g)
________________________
9. H2(g) + Cl2(g) → 2HCl(g)
________________________
10. Ag+(aq) + Cl-(aq) → AgCl(s)
________________________
11. 2N2O5(g) → 4NO2(g) + O2(g)
________________________
12. 2Al(s) + 3I2(s) → 2AlI3(s)
________________________
13. H+(aq) + OH-(aq) → H2O(l) ( str.acid-base rxn)
________________________
14. 2NO(g) → N2(g) + O2(g)
________________________
15. H2O(g) → H2O(l)
________________________
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AP Thermodynamics Packet Unit 9
Entropy Calculations
Students should be able to demonstrate an understanding of the following essential knowledge:
 5.E.1 Entropy is a measure of the dispersal of matter and energy.
 5.E.2 Some physical or chemical processes involve both a decrease in the internal energy of the components
(ΔH<0) under consideration and an increase in the entropy of those components (ΔS>0). These processes are
necessarily “thermodynamically favored” (ΔG<0).
The second law of thermodynamics: the universe is constantly increasing the dispersal of matter and
energy.
The third law of thermodynamics: the entropy of a perfect crystal at 0 K is zero. [Not a lot of perfect
crystals out there so, entropy values are RARELY ever zero—even elements] So what? This means the
absolute entropy of a substance can then be determined at any temp. higher than zero K. (Handy to know
if you ever need to defend why G & H for elements = 0. . . . BUT S does not!)
ΔS is + when dispersal/disorder increases (favored)
ΔS is – when dispersal/disorder decreases
NOTE: Units are usually J/(molrxn • K) (not kJ!)
1. Calculate the entropy change at 25°C, in J/(molrxn • K) for: 2 SO2(g) + O2(g) → 2 SO3(g) Given the
following data: SO2(g) 248.1 J/(mol• K) O2(g) 205.3 J/(mol• K) SO3(g) 256.6 J/(mol • K)
Gibbs Free Energy Calculations
2. 2 H2O + O2(g) → 2 H2O2 Calculate the free energy of formation for the oxidation of water to produce
hydrogen peroxide given the following information ∆G values: H2O ‒56.7 kcal/molrxn O2(g) 0 kcal/molrxn
H2O2 ‒27.2 kcal/molrxn
3. Cdiamond(s) + O2(g) → CO2(g) ∆G°= ‒397 kJ
Cgraphite(s) + O2(g) → CO2(g)
∆G°= ‒394 kJ
Calculate ∆G° for the reaction Cdiamond(s)→Cgraphite(s)
4. 2 SO2(g) + O2(g) → 2 SO3(g) The reaction above was carried out at 25°C and 1 atm. Calculate ∆H°, ∆S°, and
∆G° using the following data:
5. The overall reaction for the corrosion (rusting) of iron by oxygen is 4 Fe(s) + 3 O2(g) → 2 Fe2O3(s) Using the
following data, calculate the equilibrium constant for this reaction at 25°C
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AP Thermodynamics Packet Unit 9
Gibbs Free Energy
Students should be able to demonstrate an understanding of the following essential knowledge:
 5.E.2 Some physical or chemical processes involve both a decrease in the internal energy of the components
(ΔH<0) under consideration and an increase in the entropy of those components (ΔS>0). These processes are
necessarily “thermodynamically favored” (ΔG<0).
 5.E.3 If a chemical or physical process is not driven by both entropy and enthalpy changes, then the Gibbs free
energy change can be used to determine whether the process is thermodynamically favored.
 5.E.4 External sources of energy can be used to drive change in cases where the Gibbs free energy change is
positive.
 5.E.5 A thermodynamically favored process may not occur due to kinetic constraints (kinetic vs. thermodynamic
control).
For a physical or chemical reaction to be spontaneous, the sign of ΔG (Gibbs Free Energy) must be
negative. The mathematical formula for this value is:
ΔG = ΔH – TΔS
where ΔH = change in enthalpy or heat of reaction
ΔS = change in entropy or randomness
T = temperature in Kelvin
Complete the table for the sign of ΔG; +, - or undetermined. When
conditions allow for an undetermined sign of ΔG, temperature will
decide spontaneity. (temp. dependent)
ΔH
ΔS
-
+
+
-
ΔG
1. The conditions in which ΔG is always negative is when ΔH is
_________________ and ΔS is _________________.
+
+
2. The conditions in which ΔG is always positive is when ΔH is
_________________ and ΔS is _________________.
3. When the situation is indeterminate, a low temperature favors the ( entropy / enthalpy ) factor,
and a high temperature favors the ( entropy / enthalpy ) factor.
Answer Problems 4-6 with always, sometimes or never.
4. The reaction: 2KClO3(s) + heat → 2KCl(s) + 3O2g) will _________________ be spontaneous.
5. The reaction: 2H2(g) + O2(g) → 2H2O(l) + heat will _________________ be spontaneous.
6. The reaction: heat + H2O(s) → H2O(l) will _________________ be spontaneous.
7. What is the value of ΔG if ΔH = -32.0 kJ/mol, ΔS = +25 J/molK and T = 293 K? _________________
8. Is the reaction in Problem 7 spontaneous? ______ Explain ______________
9. What is the value of ΔG if ΔH = +12.0 kJ/mol, ΔS = - 5 J/molK and T = 290 K? _________________
10. Is the reaction is Problem 9 spontaneous? _________________
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AP Thermodynamics Packet Unit 9
AP Problems
PCl5(g)  PCl3(g) + Cl2(g)
1. For the reaction above, H = +22.1 kilocalories per mole at 25C
(a) Does the tendency of reactions to proceed to a state of minimum energy favor the formation of the products
of this reaction? Explain
(b) Does the tendency of reactions to proceed to a state of maximum entropy favor the formation of the
products of this reaction? Explain.
(c) State whether an increase in temperature drives this reaction to the right, to the left, or has no effect.
Explain.
(d) State whether a decrease in the volume of the system at constant temperature drives this reaction to the
right, to the left or has no effect. Explain.
2.
Standard Heat of
Absolute
Formation, Hf,
Entropy, S,
-1
Substance
in kJ mol
in J mol-1 K-1
C(s)
0.00
5.69
CO2(g)
-393.5
213.6
H2(g)
0.00
130.6
H2O(l)
-285.85
69.91
O2(g)
0.00
205.0
C3H7COOH(l)
?
226.3
The enthalpy change for the combustion of butyric acid at 25C, Hcomb, is -2,183.5 kilojoules per mole. The
combustion reaction is
C3H7COOH(l) + 5 O2(g)  4 CO2(g) + 4 H2O(l)
(a) From the above data, calculate the standard heat of formation, Hf, for butyric acid.
(b)Write a correctly balanced equation for the formation of butyric acid from its elements.
(c) Calculate the standard entropy change, Sf, for the formation of butyric acid at 25C. The entropy
change, S, for the combustion reaction above is -117.1 J K-1 at 25C.
(d)Calculate the standard free energy of formation, Gf, for butyric acid at 25C.
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AP Thermodynamics Packet Unit 9
3. When crystals of barium hydroxide, Ba(OH)2.8H2O, are mixed with crystals of ammonium thiocyanate,
NH4SCN, at room temperature in an open beaker, the mixture liquefies, the temperature drops dramatically, and
the odor of ammonia is detected. The reaction that occurs is the following:
Ba(OH)2.8H2O(s) + 2 NH4SCN(s)  Ba2+ + 2 SCN- + 2 NH3(g) + 10 H2O(l)
(a) Indicate how the enthalpy, the entropy, and the free energy of this system change as the reaction occurs.
Explain your predictions.
(b)If the beaker in which the reaction is taking place is put on a block of wet wood, the water on the wood
immediately freezes and the beaker adheres to the wood. Yet the water inside the beaker, formed as the
reaction proceeds, does not freeze even though the temperature of the reaction mixture drops to -15C.
Explain these observations.
4.
Br2(l)  Br2(g)
At 25C the equilibrium constant, Kp, for the reaction above is 0.281 atmosphere.
(a) What is the G298 for this reaction?
(b)It takes 193 joules to vaporize 1.00 gram of Br2(l) at 25C and 1.00 atmosphere pressure. What are the
values of H298 and S298 for this reaction?
(c) Calculate the normal boiling point of bromine. Assume that H and S remain constant as the
temperature is changed.
(d)What is the equilibrium vapor pressure of bromine at 25C?
5.
BCl3(g) + NH3(g)  Cl3BNH3(s)
The reaction represented above is a reversible reaction.
(a) Predict the sign of the entropy change, S, as the reaction proceeds to the right. Explain your prediction.
(b) If the reaction spontaneously proceeds to the right, predict the sign of the enthalpy change, H. Explain
your prediction.
(c) The direction in which the reaction spontaneously proceeds changes as the temperature is increased above a
specific temperature. Explain.
(d) What is the value of the equilibrium constant at the temperature referred to in (c); that is, the specific
temperature at which the direction of the spontaneous reaction changes? Explain.
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AP Thermodynamics Packet Unit 9
6.
Cl2(g) + 3 F2(g)  2 ClF3(g)
ClF3 can be prepared by the reaction represented by the equation above. For ClF3 the standard enthalpy
of formation, Hf, is -163.2 kilojoules/mole and the standard free energy of formation, Gf, is -123.0
kilojoules/mole.
(a) Calculate the value of the equilibrium constant for the reaction at 298K.
(b)Calculate the standard entropy change, S, for the reaction at 298K.
(c) If ClF3 were produced as a liquid rather than as a gas, how would the sign and the magnitude of S for
the reaction be affected? Explain.
(d)At 298K the absolute entropies of Cl2(g) and ClF3(g) are 222.96 joules per mole-Kelvin and 281.50 joules
per mole-Kelvin, respectively.
(i) Account for the larger entropy of ClF3(g) relative to that of Cl2(g).
(ii) Calculate the value of the absolute entropy of F2(g) at 298K.
7.
2 C4H10(g) + 13 O2(g)  8 CO2(g) + 10 H2O(l)
The reaction represented above is spontaneous at 25C. Assume that all reactants and products are in their
standard state.
(a) Predict the sign of S for the reaction and justify your prediction.
(b)What is the sign of G for the reaction? How would the sign and magnitude of G be affected by an
increase in temperature to 50C? Explain your answer.
(c) What must be the sign of H for the reaction at 25C? How does the total bond energy of the reactants
compare to that of the products?
(d)When the reactants are place together in a container, no change is observed even though the reaction is
known to be spontaneous. Explain this observation.
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AP Thermodynamics Packet Unit 9
8.
2H2S(g) + SO2(g)  3 S(s) + 2 H2O(g)
At 298 K, the standard enthalpy change, H for the reaction represented above is -145 kilojoules.
(a) Predict the sign of the standard entropy change, S, for the reaction. Explain the basis for your
prediction.
(b)At 298 K, the forward reaction (i.e., toward the right) is spontaneous. What change, if any, would occur
in the value of G for this reaction as the temperature is increased? Explain your reasoning using
thermodynamic principles.
(c) What change, if any, would occur in the value of the equilibrium constant, Keq, for the situation
described in (b)? Explain your reasoning.
(d)The absolute temperature at which the forward reaction becomes nonspontaneous can be predicted.
Write the equation that is used to make the prediction. Why does this equation predict only an approximate
value for the temperature?
9. Propane, C3H8, is a hydrocarbon that is commonly used as fuel for cooking.
(a) Write a balanced equation for the complete combustion of propane gas, which yields CO2(g) and H2O(l).
(b)Calculate the volume of air at 30C and 1.00 atmosphere that is needed to burn completely 10.0 grams
of propane. Assume that air is 21.0 percent O2 by volume.
(c) The heat of combustion of propane is -2,220.1 kJ/mol. Calculate the heat of formation, Hf, of propane
given that Hf of H2O(l) = -285.3 kJ/mol and Hf of CO2(g) = -393.5 kJ/mol.
(d)Assuming that all of the heat evolved in burning 30.0 grams of propane is transferred to 8.00 kilograms
of water (specific heat = 4.18 J/g.K), calculate the increase in temperature of water.
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AP Thermodynamics Packet Unit 9
10. Lead iodide is a dense, golden yellow, slightly soluble solid. At 25C, lead iodide dissolves in water
forming a system represented by the following equation.
PbI2(s)  Pb2+ + 2 IH = +46.5 kilojoules
(a) How does the entropy of the system PbI2(s) + H2O(l) change as PbI2(s) dissolves in water at 25C? Explain
(b)If the temperature of the system were lowered from 25C to 15C, what would be the effect on the value
of Ksp? Explain.
(c) If additional solid PbI2 were added to the system at equilibrium, what would be the effect on the
concentration of I- in the solution? Explain.
(d)At equilibrium, G = 0. What is the initial effect on the value of G of adding a small amount of
Pb(NO3)2 to the system at equilibrium? Explain.
C2H2(g) + 2 H2(g)  C2H6(g)
11.
Information about the substances involved in the reaction represented above is summarized in the following tables.
Substance
S (J/molK)
Hf (kJ/mol)
Bond
Bond Energy (kJ/mol)
C2H2(g)
200.9
226.7
C-C
347
H2(g)
130.7
0
C=C
611
-84.7
C-H
414
H-H
436
C2H6(g)
----
(a) If the value of the standard entropy change, S, for the reaction is -232.7 joules per moleKelvin,
calculate the standard molar entropy, S, of C2H6 gas.
(b)Calculate the value of the standard free-energy change, G, for the reaction. What does the sign of G
indicate about the reaction above?
(c) Calculate the value of the equilibrium constant, K, for the reaction at 298 K.
(d)Calculate the value of the CC bond energy in C2H2 in kilojoules per mole.
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AP Thermodynamics Packet Unit 9
12. For the gaseous equilibrium represented below, it is observed that greater amounts of PCl 3 and Cl2 are
produced as the temperature is increased.
PCl5(g)  PCl3(g) + Cl2(g)
(a) What is the sign of S for the reaction? Explain.
(b)What change, if any, will occur in G for the reaction as the temperature is increased? Explain your
reasoning in terms of thermodynamic principles.
(c) If He gas is added to the original reaction mixture at constant volume and temperature, what will happen
to the partial pressure of Cl2? Explain.
(d)If the volume of the reaction mixture is decreased at constant temperature to half the original volume,
what will happen to the number of moles of Cl2 in the reaction vessel? Explain.
13.
C6H5OH(s) + 7 O2(g)  6 CO2(g) + 3 H2O(l)
When a 2.000-gram sample of pure phenol, C6H5OH(s), is completely burned according to the equation above,
64.98 kilojoules of heat is released. Use the information in the table below to answer the questions that follow.
Standard Heat of Absolute Entropy,
Formation, Hâ;
S, at 25C
(J/molòK)
Substance
at 25C (kJ/mol)
C(graphite)
0.00
5.69
CO2(g)
-393.5
213.6
H2(g)
0.00
130.6
H2O(l)
-285.85
69.91
O2(g)
0.00
205.0
C6H5OH(s)
?
144.0
(a) Calculate the molar heat of combustion of phenol in kilojoules per mole at 25C.
(b)Calculate the standard heat of formation, Hâ, of phenol in kilojoules per mole at 25C.
(c) Calculate the value of the standard free-energy change, G, for the combustion of phenol at 25C.
(d)If the volume of the combustion container is 10.0 liters, calculate the final pressure in the container
when the temperature is changed to 110.C. (Assume no oxygen remains unreacted and that all products are
gaseous.)
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AP Thermodynamics Packet Unit 9
14. A student is asked to determine the molar enthalpy of neutralization, ∆Hneut, for the reaction represented
above. The student combines equal volumes of 1.0 M HCl and 1.0 M NaOH in an open polystyrene cup
calorimeter. The heat released by the reaction is determined by using the equation q = mc∆T.
Assume the following.
• Both solutions are at the same temperature before they are combined.
• The densities of all the solutions are the same as that of water.
• Any heat lost to the calorimeter or to the air is negligible.
• The specific heat capacity of the combined solutions is the same as that of water.
(a) Give appropriate units for each of the terms in the equation q = mc∆T.
(b)List the measurements that must be made in order to obtain the value of q.
(c) Explain how to calculate each of the following.
(i) The number of moles of water formed during the experiment
(ii) The value of the molar enthalpy of neutralization, ∆Hneut, for the reaction between HCl(aq) and
NaOH(aq)
(d)The student repeats the experiment with the same equal volumes as before, but this time uses 2.0 M HCl
and 2.0 M NaOH.
(i) Indicate whether the value of q increases, decreases, or stays the same when compared to the first
experiment. Justify your prediction.
(ii) Indicate whether the value of the molar enthalpy of neutralization, ∆Hneut, increases, decreases, or
stays the same when compared to the first experiment. Justify your prediction.
(e) Suppose that a significant amount of heat were lost to the air during the experiment. What effect would
this have on the calculated value of the molar enthalpy of neutralization, ∆Hneut? Justify your answer.
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AP Thermodynamics Packet Unit 9
15. Answer the following questions that relate to the chemistry of nitrogen.
(a) Two nitrogen atoms combine to form a nitrogen molecule, as represented by the following equation.
2 N(g) N2(g)
Using the table of average bond energies below, determine the enthalpy change, ∆H, for the reaction.
Bond
Average Bond
Energy (kJ mol–
1
)
N–N
160
N=N
420
NN
950
(b) The reaction between nitrogen and hydrogen to form ammonia is represented below.
N2(g) + 3 H2(g)  2 NH3(g) ∆H˚ = –92.2 kJ
Predict the sign of the standard entropy change, ∆S˚, for the reaction. Justify your answer.
(c) The value of ∆G˚ for the reaction represented in part (b) is negative at low temperatures but positive at
high temperatures. Explain.
(d)When N2(g) and H2(g) are placed in a sealed container at a low temperature, no measurable amount of
NH3(g) is produced. Explain.
3
2 Fe(s) + 2 O2(g)  Fe2O3(s) ∆Hf˚ = -824 kJ mol–1
16.
Iron reacts with oxygen to produce iron(III) oxide as represented above. A 75.0 g sample of Fe(s) is mixed
with 11.5 L of O2(g) at 2.66 atm and 298 K.
(a) Calculate the number of moles of each of the following before the reaction occurs.
(i) Fe(s)
(ii)
O2(g)
(b)Identify the limiting reactant when the mixture is heated to produce Fe2O3. Support your answer with
calculations.
(c) Calculate the number of moles of Fe2O3 produced when the reaction proceeds to completion.
(d) The standard free energy of formation, ∆Gf˚ of Fe2O3 is –740. kJ mol–1 at 298 K.
(i) Calculate the standard entropy of formation ∆Sf˚ of Fe2O3 at 298 K. Include units with your
answer.
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AP Thermodynamics Packet Unit 9
(ii) Which is more responsible for the spontaneity of the formation reaction at 298K, the standard enthalpy
or the standard entropy?
The reaction represented below also produces iron(III) oxide. The value of ∆H˚ for the reaction is –280 kJ
per mol.
1
2 FeO(s) + 2 O2(g)  Fe2O3(s)
(e) Calculate the standard enthalpy of formation, ∆Hf˚ of FeO(s).
17.
CO(g) +
1
2
O2(g)  CO2(g)
The combustion of carbon monoxide is represented by the equation above.
(a) Determine the value of the standard enthalpy change, ∆H˚rxn for the combustion of CO(g) at 298 K using
the following information.
1
C(s) + O2(g)  CO(g)
∆H˚298 = –110.5 kJ mol-1
2
C(s) + O2(g)  CO2(g)
∆H˚298 = –393.5 kJ mol-1
(b)Determine the value of the standard entropy change, ∆S˚rxn, for the combustion of CO(g) at 298 K using
the information in the following table.
S˚298
Substance
-1 -1
(J mol K )
CO(g)
197.7
CO2(g)
213.7
O2(g)
205.1
(c) Determine the standard free energy change, ∆G˚rxn, for the reaction at 298 K. Include units with your
answer.
(d)Is the reaction spontaneous under standard conditions at 298 K? Justify your answer.
(e) Calculate the value of the equilibrium constant, Keq, for the reaction at 298 K.
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AP Thermodynamics Packet Unit 9
18.
N2(g) + 3 F2(g)  2 NF3(g)
o
o
= – 264 kJ mol–1; S298
= – 278 J K–1 mol–1
H 298
The following questions relate to the synthesis reaction represented by the chemical equation in the box above.
o
(a) Calculate the value of the standard free energy change, G298
for the reaction.
(b) Determine the temperature at which the equilibrium constant, Keq, for the reaction is equal to 1.00.
(Assume that ∆H˚ and ∆S˚ are independent of temperature.)
(c) Calculate the standard enthalpy change, ∆H˚, that occurs when a 0.256 mol sample of NF3(g) is formed
from N2(g) and F2(g) at 1.00 atm and 298 K.
The enthalpy change in a chemical reaction is the difference between energy absorbed in breaking bonds in the
reactants and energy released by bond formation in the products.
(d) How many bonds are formed when two molecules of NF3 are produced according to the equation in
the box above?
(e) Use both the information in the box above and the table of average bond enthalpies below to calculate
the average enthalpy of the F–F bond.
Bond
Average Bond Enthalpy
(kJ mol-1)
NN
946
N–F
272
F–F
?
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AP Thermodynamics Packet Unit 9
Calculations Summary: For each factor, provide the ways to find their values and describe how you will know
to use the calculation. One example is provided for you. Use you reference table and your packet!
ΔH/q Enthalpy/heat
(J/mol or kJ/mol)
ΔS
Entropy/disorder
(J/mol K)
ΔG Gibbs Free Energy/ Spontaneity
(kJ/mol)
ΔG=-nFE
We didn’t cover this
yet (next unit) but it shows up on
the RT. Used when you know the
voltage of a battery.
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AP Thermodynamics Packet Unit 9
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