AP Chemistry
Chapter 19 Chemical Thermodynamics
Chapter 19. Chemical Thermodynamics
Sample Exercise 19.1 (p. 803)
Predict whether the following processes are spontaneous as described, are spontaneous in the reverse direction,
or are in equilibrium:
a)
When a piece of metal heated to 150oC is added to water at 40oC, the water gets hotter.
b)
Water at room temperature decomposes into H2(g) and O2(g).
c)
Benzene vapor, C6H6(g), at a pressure of 1 atm condenses to liquid benzene at the normal boiling point
of benzene, 80.1oC.
Practice Exercise 19.1
Under 1 atm pressure CO2(s) (“dry ice”) sublimes at -78oC. Is the transformation of CO2(s) to CO2(g) a
spontaneous process at -100oC and 1 atm pressure?
Sample Exercise 19.2 (p. 807)
The element, Hg, is a silvery liquid at room temperature. The normal freezing point of mercury is -38.9oC, and
its molar enthalpy of fusion is ∆Hfusion = 2.29 kJ/mol. What is the entropy change of the system when 50.0 g of
Hg(l) freezes at the normal freezing point?
(∆Ssys = -2.44 J/K)
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AP Chemistry
Chapter 19 Chemical Thermodynamics
Practice Exercise 19.2
The normal boiling point of ethanol, C2H5OH, is 78.3oC, and its molar enthalpy of vaporization is 38.56 kJ/mol.
What is the change in entropy in the system when 68.3 g of C2H5OH(g) at 1 atm condenses to liquid at the
normal boiling point?
(-163 J/K)
Sample Exercise 19.3 (p. 814)
Predict whether ∆S is positive or negative for each of the following processes, assuming each occurs at constant
temperature:
a)
H2O(l)  H2O(g)
b)
Ag+(aq) + Cl-(aq)  AgCl(s)
c)
4 Fe(s) + 3 O2(g)  2 Fe2O3(s)
d)
N2(g) + O2(g)  2 NO(g)
Practice Exercise 19.3
Indicate whether each of the following reactions produces an increase or decrease in the entropy of the system:
a)
CO2(s)  CO2(g)
b)
CaO(s) + CO2(g)  CaCO3(s)
c)
HCl(g) + NH3(g)  NH4Cl(s)
d)
2 SO2(g) + O2(g)  2 SO3(g)
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AP Chemistry
Chapter 19 Chemical Thermodynamics
Sample Exercise 19.4 (p. 815)
Choose the sample of matter that has greater entropy in each pair, and explain your choice:
a) 1 mol of NaCl(s) or 1 mol of HCl(g) at 25oC.
b) 2 mol of HCl(g) or 1 mol of HCl(g) at 25oC
c) 1 mol of HCl(g) or 1 mol or Ar(g) at 298 K.
Practice Exercise 19.4
Choose the substance with the greater entropy in each case:
a) 1 mol of H2(g) at STP or 1 mol of H2(g) at 100oC and 0.5 atm
b) 1 mol of H2O(s) at 0oC or 1 mol of H2O(l) at 25oC
c) 1 mol of H2(g) at STP or 1 mol of SO2(g) at STP
d) 1 mol of N2O4(g) at STP or 2 mol of NO2(g) at STP.
Sample Exercise 19.5 (p. 818)
Calculate ∆So for the synthesis of ammonia from N2(g) and H2(g) at 298 K.
N2(g) + 3 H2(g)  2 NH3(g)
(-198.3 J/K)
Practice Exercise 19.5
Using the standard entropies in Appendix C, calculate the standard entropy change, ∆So for the following
reaction at 298 K:
Al2O3(s) + 3 H2(g)  2 Al(s) + 3 H2O(g)
(180.39 J/K)
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AP Chemistry
Chapter 19 Chemical Thermodynamics
Sample Exercise 19.6 (p. 821)
Calculate the standard free energy change for the formation of NO(g) from N2(g) and O2(g) at 298 K:
N2(g) + O2(g)  2 NO(g)
Given that ∆Ho = 180.7 kJ and ∆So = 24.7 J/K. Is the reaction spontaneous under these circumstances?
Practice Exercise 19.6
A particular reaction has ∆Ho = 24.6 kJ and ∆So = 132 J/K at 298 K. Calculate ∆Go. Is the reaction spontaneous
under these conditions?
Sample Exercise 19.7 (p. 823)
a)
Use data from Appendix C to calculate the standard free-energy change for the following reaction at
298 K:
P4(g) + 6 Cl2(g)  4 PCl3(g)
(-1102.8 kJ)
b)
What is ∆Go for the reverse of the above reaction?
(+1102.8 kJ)
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AP Chemistry
Chapter 19 Chemical Thermodynamics
Practice Exercise 19.7
By using the data from Appendix C, calculate ∆Go at 298 K for the combustion of methane:
CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(g)
(-800.7 kJ)
Sample Exercise 19.8 (p. 823)
In Section 5.7 we used Hess’s law to calculate ∆Ho for the combustion of propane gas at 298 K:
C3H8(g) + 5 O2(g)  3 CO2(g) + 4 H2O(l)
∆Ho = -2220 kJ
a)
Without using data from Appendix C, predict whether ∆Go for this reaction is more negative or less
negative than ∆Ho.
b)
Use data from Appendix C to calculate the standard free-energy change for the reaction at 298 K. Is
your prediction from part (a) correct?
(-2108 kJ)
Practice Exercise 19.8
Consider the combustion of propane to form CO2(g) and H2O(g) at 298 K:
C3H8(g) + 5 O2(g)  3 CO2(g) + 4 H2O(g)
∆Ho = -2220 kJ
Would you expect ∆Go to be more negative or less negative than ∆Ho?
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AP Chemistry
Chapter 19 Chemical Thermodynamics
Sample Exercise 19.9 (p. 826)
The Haber process for the production of ammonia involves the following equilibrium:
N2(g) + 3 H2(g)  2 NH3(g)
Assume that ∆Ho and ∆So for this reaction do not change with temperature.
a)
b)
Predict the direction in which ∆Go for this reaction changes with increasing temperature.
Calculate the values of ∆Go for the reaction at 25oC and 500oC.
(-33.3 kJ, 61 kJ)
Practice Exercise 19.9
a)
Using standard enthalpies of formation and standard entropies in Appendix C, calculate ∆Ho and ∆So at
298 K for the following reaction:
2 SO2(g) + O2(g)  2 SO3(g).
(∆Ho = -196.6 kJ, ∆So = -189.6 J/K)
b)
Using the values obtained in part (a), estimate ∆Go at 400 K.
(∆Go = -120.8 kJ)
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AP Chemistry
Chapter 19 Chemical Thermodynamics
∆G = ∆Go + RTlnQ
Sample Exercise 19.10 (p. 827)
As we saw in Section 11.5, the normal boiling point is the temperature at which a pure liquid is in equilibrium
with its vapor at a pressure of 1 atm.
a)
Write the chemical equation that defines the normal boiling point of liquid carbon tetrachloride, CCl4(l).
b)
What is the value of ∆Go for the equilibrium in part (a)?
c)
Use thermodynamic data in Appendix C and ∆Go = ∆Ho – T∆So to estimate the normal boiling point
of CCl4. (70oC)
Practice Exercise 19.10
Use data in Appendix C to estimate the normal boiling point, in K, for elemental bromine, Br2(l).
(The experimental value is given in Table 11.3). (330 K)
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AP Chemistry
Chapter 19 Chemical Thermodynamics
Sample Exercise 19.11 (p. 828)
We will continue to explore the Haber process for the synthesis of ammonia:
N2(g) + 3 H2(g)  2 NH3(g)
Calculate ∆G at 298 K for a reaction mixture that consists of 1.0 atm N2, 3.0 atm H2, and 0.50 atm NH3.
(-44.9 kJ/mol)
Practice Exercise 19.11
Calculate ∆G at 298 K for the reaction of nitrogen and hydrogen to form ammonia if the reaction mixture
consists of 0.50 atm N2, 0.75 atm H2, and 2.0 atm NH3.
(-26.0 kJ/mol)
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AP Chemistry
Chapter 19 Chemical Thermodynamics
Sample Exercise 19.12 (p. 829)
Use standard free energies of formation to calculate the equilibrium constant Keq at 25oC for the reaction
involved in the Haber process:
N2(g) + 3 H2(g)  2 NH3(g)
(7 x 105)
Practice Exercise 19.12
Use data from Appendix C to calculate the standard free-energy change, ∆Go, and the equilibrium constant, K,
at 298 K for the following reaction:
H2(g) + Br2(l)  2 HBr(g)
(-106.4 kJ/mol; 4 x 1018)
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AP Chemistry
Chapter 19 Chemical Thermodynamics
Sample Integrative Exercise 19 (p. 831)
Consider the simple salts NaCl(s) and AgCl(s). We will examine the equilibria in which these salts dissolve in
water to form aqueous solutions of ions:
NaCl(s)  Na+(aq) + Cl-(aq)
AgCl(s)  Ag+(aq) + Cl-(aq)
a)
Calculate the value of ∆Go at 298 K for each of the preceding reactions.
b)
The two values from part (a) are very different. Is this difference primarily due to the enthalpy term or
the entropy term of the standard free-energy change?
c)
Use the values of ∆Go to calculate Ksp values for the two salts at 298 K.
d)
Sodium chloride is considered a soluble salt, whereas silver chloride is considered insoluble. Are these
descriptions consistent with the answers to part (c)?
e)
How will ∆Go for the solution process of these salts change with increasing T? What effect should this
change have on the solubility of the salts?
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