CHAPTER 13:
CHEMICAL THERMODYNAMICS
The following information is often presented on the first page of exams covering this chapter under the heading: “potentially useful information.”
E = E° - (0.0257/n) ln Q 1 F = 96,500 C R = 8.314 J/mol-K
ΔG = - nFE ΔG° = - RT ln K
Many of the following questions assume that a truncated table of standard-state thermodynamic data is attached to the exam. The table usually contains ΔHac° and S° data for 6-10 compounds, which pertain to the particular questions on that exam.
Internal Energy and Enthalpy
13-1. Which of the following changes would increase the internal energy of a system?
(a) the system gains heat and performs work.
(b) the system gains heat and has worked performed on it.
(c) the system loses heat and performs work.
(d) the system loses heat and has work performed on it.
(e) none of these changes would increase the internal energy of a system.
Answer: (b)
13-2. The heat transferred from a system to its surroundings (or vice versa) when a chemical reaction is run under conditions of constant pressure is equal to:
(a) the change in the enthalpy of the system (ΔH).
(b) the change in the energy of the system (ΔE).
(c) the change in the entropy of the system (ΔS).
(d) the change in the free energy of the system (ΔG).
(e) the work done by the system.
Answer: (a)
13-3. The standard enthalpy of atom combination of CO2 is -1608.53 kJ/molrxn and O2 is -498.34 kJ/molrxn and the standard enthalpy of reaction for the following reaction is -283.0 kJ/molrxn.
CO(g) + ½ O2(g) ➝ CO2(g)
What is the standard enthalpy of atom combination of CO(g) in kJ/mol?
(a) -827.21 (b) -1076.36 (c) -1359.36 (d) -1393.17 (e) -1642.34
Answer: (b)
13-4. A reaction in which carbon monoxide spontaneously disproportionates to graphite and carbon dioxide causes problems in the upper part of iron blast furnaces. The enthalpy of atom combination of CO is -1076.38 kJ/molrxn, graphite is -716.68 kJ/molrxn, and CO2 is –1608.53 kJ/molrxn. What is the enthalpy change in this disproportionation reaction at 298K?
(a) -613 kJ/molrxn (b) -503 kJ/molrxn (c) -283 kJ/molrxn
(d) -172 kJ/molrxn (e) 283 kJ/molrxn
Answer: (d)
13-5. For which of the following reactions is ΔH approximately equal to ΔE?
(a) 2 H2O2(g) ➝ 2 H2O(g) + O2(g)
(b) 2 H2(g) + O2(g) ➝ 2 H2O(g)
(c) 2 NH3(g) ➝ N2(g) + 3 H2(g)
(d) 2 NO(g) ➝ N2(g) + O2(g)
(e) NH4NO3(s) ➝ N2(g) + 1/2 O2 (g) + 2 H2O(g)
Answer: (d)
Entropy
13-6. Which of the following elements could be described as the most disorganized at room temperature?
(a) Cl2(g) S° = 223 J/mol-K
(b) diamond S° = 2.43 J/mol-K
(c) tin S° = 51.5 J/mol-K
(d) mercury S° = 77.4 J/mol-K
(e) N2(g) S° = 192 J/mol-K
Answer: (a)
13-7. Use your gut-level understanding of entropy to predict which of the following processes will have the largest entropy change.
(a) H2O(s) ➝ H2O(l) (b) H2O(s) ➝ H2O(g)
(c) H2O2(l) ➝ H2O(g) + ½ O2(g) (d) H2O(l) ➝ H2O(g)
(e) H2O2(g) ➝ H2O(g) + ½ O2(g)
Answer: (c)
13-8. Which of the following would have the most positive value for ΔS°?
(a) H2O(l) ➝ H2O(s) (b) NaNO3(s) ➝ Na+(aq) + NO3-(aq)
(c) 2 HCl(g) ➝ H2(g) + Cl2(g) (d) 2 H2(g) + O2(g) ➝ 2 H2O(g)
Answer: (b)
13-9. Which of the following processes would have a negative entropy change?
(a) decomposition of hydrogen peroxide.
(b) sublimation of dry ice (solid CO2).
(c) evaporation of water.
(d) formation of Al2O3 from its elements.
(e) dissolution of salt in water.
Answer: (d)
13-10. For which of the following highly exothermic processes would you expect ΔH°and ΔG° to be about the same?
(a) 2 Al(s) + 3/2 O2(g) ➝ Al2O3(s) (b) 2 H2(g) + O2(g) ➝ 2 H2O(g)
(c) 2 Na(s) + 2 H2O(l) ➝ 2 NaOH(aq) + H2(g) (d) 2 NO(g) ➝ N2O4(g)
(e) 2 Al(s) + Fe2O3(s) ➝ 2 Fe(s) + Al2O3(s)
Answer: (e)
13-11. An exothermic chemical reaction in which 2 moles of gaseous products are formed from 3 moles of gaseous reactants is favored by:
(a) increasing both temperature and pressure.
(b) increasing temperature and decreasing pressure.
(c) decreasing both temperature and pressure.
(d) decreasing temperature and increasing pressure.
(e) none of the above changes.
Answer: (d)
13-12. Which of the following reactions would have a positive value for ΔS?
(a) 3 NO(g) ➝ NO2(g) + N2O(g) (b) 2 CO2(g) ➝ 2 CO(g) + O2(g)
(c) 2 I(g) ➝ I2(g) (d) NH3(g) ➝ NH3(l)
(e) None of these reactions would have a positive ΔS.
Answer: (b)
13-13. Which of the following would have a positive value for ΔS?
(a) 2 H2O(g) ➝ 2 H2(g) + O2(g) (b) N2(g) + 3 H2(g) ➝ 2 NH3(g)
(c) CO2(g) + H2(g) ➝ CO(g) + H2O(g) (d) H2O(l) ➝ H2O(s)
(e) none of these reactions would have a positive ΔS.
Answer: (a)
13-14. Which of the following would have a negative entropy change?
(a) Formation of Al2O3 from its elements.
(b) Sublimation of dry ice — solid CO2 — at room temperature.
(c) Evaporation of water.
(d) Decomposition of hydrogen peroxide.
(e) Dissolution of NaCl in water.
Answer: (a)
13-15. Which of the following reactions would have a negative ΔS°?
(a) H2(g) ➝ 2 H(g) (b) H2(l) ➝ H2(g)
(c) H2(g) + Cl2(g) ➝ 2 HCl(g) (d) 2 H2(g) + CO(g) ➝ CH3OH(g)
(e) none of the above
Answer: (d)
Signs of ΔH°, ΔS°, and ΔG°
13-16. Use your gut-level understanding of thermodynamics to predict the signs of ΔH° and ΔS° for the following reaction.
2 H2(g) + O2(g) ➝ 2 H2O(g)
(a) ΔH° = - and ΔS° = - (b) ΔH° = - and ΔS° = +
(c) ΔH° = + and ΔS° = - (d) ΔH° = + and ΔS° = +
(e) ΔH° = - and ΔS° = 0
Answer: (a)
13-17. Which of the following reactions is unfavorable at low temperatures but becomes favorable as the temperature increases?
(a) 2 CO(g) + O2(g) ➝ 2 CO2(g) ΔH° = -566 kJ, ΔS° = -173 J/K
(b) 2 H2O(g) ➝ 2 H2(g) + O2(g) ΔH° = 484 kJ, ΔS°= 90.0 J/K
(c) 2 N2O(g) ➝ 2 N2(g) + O2(g) ΔH° = -164 kJ, ΔS° = 149 J/K
(d) PbCl2(s) ➝ Pb2+(aq) + 2 Cl-(aq) ΔH° = 23.4 kJ, ΔS° = -12.5 J/K
(e) none of the above
Answer: (b)
13-18. If the following reaction is spontaneous as written,
Zn(s) + Cu2+(aq) ➝ Zn2+(aq) + Cu(s)
which of the following statements is true?
(a) Kc > 1 and ΔG° > 0 (b) Kc > 1 and ΔG° < 0
(c) Kc < 1 and ΔG° < 0 (d) Kc < 1 and ΔG° < 0
(e) Kc < 1 and ΔG° = 0
Answer: (b)
13-19. At about what temperature does the equilibrium constant for the reaction in the previous question become larger than one?
(a) between -200°C and 0°C (b) between 0°C and 100°C
(c) between 100°C and 200°C (d) between 200°C and 400°C
(e) between 400°C and 600°C
Answer: (c)
13-20. Which statement is true if the following reaction is spontaneous as written?
Zn + Cu2+ ➝ Zn2+ + Cu
(a) Kc is larger than 1, ΔG° is negative, and E° is positive.
(b) Kc is larger than 1, ΔG° is positive, and E° is negative.
(c) Kc is larger than 1, ΔG° is negative, and E° is negative.
(d) Kc is smaller than 1, ΔG° is negative, and E° is positive.
(e) Kc is smaller than 1, ΔG° is positive, and E° is negative.
Answer: (a)
13-21. What happens to the magnitude of the equilibrium constant for the following reaction as the temperature of the system increases?
2 HI(g) ➝ H2(g) + I2(g) ΔH° = -9.48 kJ and ΔS°= -21.8 J/K.
(a) Kc first increases and then decreases.
(b) Kc first decreases and then increases.
(c) Kc increases.
(d) Kc decreases.
(e) it is impossible to determine what happens to Kc.
Answer: (d)
13-22. Which of the following always corresponds to a spontaneous reaction?
(a) ΔH° < 0, ΔS° < 0 (b) ΔH° > 0, ΔS° < 0
(c) ΔH° < 0, ΔS° > 0 (d) ΔH° > 0, ΔS° > 0
(e) ΔH° < 0, ΔS° = 0
Answer: (c)
13-23. Under what conditions will a reaction that isn’t spontaneous at room temperature become spontaneous as the temperature increases?
(a) ΔH = + and ΔS = + (b) ΔH = + and ΔS = -
(c) ΔH = - and ΔS = + (d) ΔH = - and ΔS = -
(e) a reaction cannot become more spontaneous as the temperature increases, regardless of the signs of ΔH and ΔS
Answer: (a)
Calculating ΔS° and ΔG°
13-24. Use the following data to calculate ΔS° for the reaction:
2 NH3(g) ➝ N2(g) + 3 H2(g)
Compound S° (J/mol-K)
H2(g) -98.74
N2(g) -114.99
NH3(g) -304.99
(a) -198.8 J/K (b) -91.26 J/K (c) -106.22 J/K (d) 91.26 J/K (e) 198.8 J/K
Answer: (e)
13-25. Assuming that ΔH0 for the reaction in the previous question is 92.2 kJ/molrxn, calculate ΔG° at 25°C.
(a) -106.5 kJ/molrxn (b) -33.0 kJ/molrxn
(c) 33.0 kJ/molrxn (d) 106.5 kJ/molrxn
(e) none of these
Answer: (c)
13-26. What is the equilibrium constant at 25°C for the reaction in the previous two questions?
(a) 1.7 x 10-6 (b) 1.3 x 10-3 (c) 7.8 x 102
(d) 6.1 x 105 (e) none of these
Answer: (a)
13-27. Use the S° values for tin metal (S° = -124.35 J/mol-K) , chlorine gas (S° = -107.33 J/mol-K) and tin tetrachloride (S° = -570.7 J/mol-K) to calculate ΔS° for the following reaction.
Sn(s) + 2 Cl2(g) ➝ SnCl4(l)
(a) -339.0 (b) -231.7 (c) 231.7 (d) 339.0 (e) none of the above
Answer: (b)
Questions 28-31 refer to the following reaction.
CO(g) + 2 H2(g) ➝ CH3OH(g)
for which the following data are available.
Compound ΔHac° (kJ/mol) S° (J/mol-K)
CO(g) -1076.38 -121.48
H2(g) -435.3 -98.74
CH3OH(g) -2037.11 -538.19
13-28. What is the value of ΔS° per mole of CH3OH produced in this reaction?
(a) -219.2 J/K (b) -164.1 J/K (c) -88.4 J/K
(d) 88.4 J/K (e) 219.2 J/K
Answer: (a)
13-29. What is ΔG° per mole of CH3OH produced in this reaction at 25°C?
(a) -155.5 kJ (b) -24.8 kJ (c) +24.8 kJ (d) 155.5 kJ (e) none of the above
Answer: (b)
13-30. Which of the following statements about this reaction is true?
(a) the reaction can never become favorable.
(b) the reaction can never become unfavorable.
(c) the reaction will become more favorable as the temperature is increased.
(d) the reaction will become less favorable as the temperature is increased.
(e) temperature has no effect on whether this reaction is favorable.
Answer: (d)
13-31. Calculate the equilibrium constant for this reaction at 25°C
(a) 5.6 x 10-28 (b) 4.5 x 10-5 (c) 2.2 x 104 (d) 1.8 x 1027
(e) none of the above
Answer: (c)
Questions 32-34 refer to the following oxidation-reduction reaction.
2 Fe3+(aq) + Cu(s) ➝ 2 Fe2+(aq) + Cu2+(aq)
Compound ΔHac° (kJ/mol) S° (J/mol-K)
Cu(s) -338.32 -133.23
Cu2+(aq) -273.55 -266.0
Fe2+(aq) -505.4 -318.2
Fe3+(aq) -464.8 -496.4
13-32. The value of ΔS° for this reaction is:
(a) between -400 and -200 J/K
(b) between -200 and 0 J/K
(c) between 0 and 200 J/K
(d) between 200 and 400 J/K
(e) more than 400 J/K
Answer: (d)
13-33. The equilibrium constant for this reaction at 25°C is:
(a) between 10-30 and 10-16 (b) between 10-16 and 10-10
(c) between 10-10 and 106 (d) between 106 and 1016
(e) larger than 1016
Answer: (d)
13-34. If ΔG° = -nFE°, then:
(a) E° is between -1.00 and -0.5 V, and the reaction is not spontaneous.
(b) E° is between -0.5 and 0 V, and the reaction is not spontaneous.
(c) E° is between 0 and 0.5 V, and the reaction is spontaneous.
(d) E° is between 0.5 and 1.0 V, and the reaction is spontaneous.
(e) none of the above are true.
Answer: (c)
13-35. Which statement correctly describes the following reaction?
PbO(s) + C(s) ➝ Pb(s) + CO(g) ΔS° = 188 J/K, ΔH° = 106 kJ/molrxn
(a) The reaction is spontaneous at all temperatures.
(b) The reaction is spontaneous only below about 300°C.
(c) The reaction becomes spontaneous at about 300°C.
(d) The reaction becomes spontaneous at about 560°C.
(e) The reaction is spontaneous only above 2000°C
Answer: (c)
ΔG versus ΔG°
13-36. The entropy change in the process: 2 H(g) ➝ H2(g) is -98.74 J/mol-K. The enthalpy of atom combination of the H2 molecule is 435 kJ/mol. Estimate the temperature above which the formation of H2 molecules from H atoms will no longer be a spontaneous process.
(a) 500 K (b) 1250 K (c) 2300 K (d) 4400 K (e) 6500 K
Answer: (d)
13-37. Which of the following statements is true?
(a) ΔG = ΔG° when the reaction is at equilibrium.
(b) ΔG is a measure of how far the reaction is from equilibrium.
(c) ΔG° is positive for reactions that have to much reactant in the standard state.
(d) The value of ΔG does not depend on the temperature of the reaction.
(e) none of these statements are true.
Answer: (b)
13-38. Which of the following statements is false?
(a) ΔG is equal to ΔG° when the system is at the standard state.
(b) ΔG is zero when the system is at equilibrium.
(c) ΔG measures how far the reaction is from equilibrium.
(d) When ΔG is positive, the reaction should proceed forward to form more product.
(e) al1 of the above statements are true.
Answer: (d)
ΔG° and Equilibrium Constants
13-39. When a reaction is at equilibrium, which of the following is true?
(a) ΔG = 0 (b) ΔG = ΔG° (c) ΔG° = 0 (d) Q = 0 (e) ln K = 0
Answer: (a)
13-40. ΔG° = -53 6 kJ/molrxn for the following reaction at 25° C.
Zn2+ + 4 NH3 ➝ Zn(NH3)42+
What is the value of the complex formation equilibrium constant for this reaction?
(a) 5 x 10-21 (b) 4 x 10-10 (c) 2.5 x 109 (d) 2.0 x 1020
Answer: (c)
13-41. Use the following information to answer this question.
H2(g) + F2(g) ➝ 2 HF(g) ΔG° = -190 kJ
H2(g) + Cl2(g) ➝ 2 HCl(g) ΔG° = -546 kJ
H2(g) + Br2(g) ➝ 2 HBr(g) ΔG° = -106 kJ
H2(g) + I2(g) ➝ 2 HI(g) ΔG° = 4 kJ
An excess of H2 is allowed to react with equal amounts of F2, Cl2, Br2, and I2 in a sealed flask. Which of the following products is present in the largest concentration at equilibrium?
(a) HF (b) HCl (c) HBr (d) HI
(e) all four products are present in equal concentrations.
Answer: (b)
13-42. What is the equilibrium constant for the following reaction if ΔG° at 298 K is -16.5 kJ/molrxn?
N2(g) + 3 H2(g) ➝ 2 NH3(g)
(a) 2.0 x 10-5 (b) 1.3 x 10-3 (c) 780 (d) 4.6 x 106
Answer: (c)
13-43. E° = 0.62 V for the following reaction.
Sn2+(aq) + 2 Fe3+(aq) ➝ Sn4+(aq) + 2 Fe2+(aq)
What is Δ G0 for this reaction at 25°C?
(a) -119.6 kJ (b) -59.8 kJ (c) 59.8 kJ (d) 119.6 kJ (e) none of these
Answer: (a)
13-44. Ka for benzoic acid (C6H5CO2H) is 3.0 x 10-5. What is the free energy change for the dissociation of this acid at 25°C?
(a) -25.8 kJ (b) -2.2 kJ (c) 2.2 kJ (d) 25.8 kJ (e) none of the above
Answer: (d)
13-45. Although the diatomic molecule N2 has an extremely strong N≡N bond it will break apart into nitrogen atoms — N2(g) ➝ 2 N(g) — if it is heated to a high enough temperature because:
(a) The value of ΔG becomes more positive as the temperature increases.
(b) the unfavorable enthalpy change for the reaction disappears at higher temperatures.
(c) the entropy change for the reaction is favorable, and this factor becomes more important as the temperature of the system increases.
(d) the equilibrium constant for the reaction decreases with temperature.
(e) none of the above statements is a correct explanation.
Answer: (c)
13-46. The following reaction is used to produce a mixture of CO and H2 known as synthesis gas.
C(s) + H2O(g) ➝ CO(g) + H2(g)
If ΔH° = 131.3 kJ/molrxn and ΔS° = 134.4 J/mol-K, at what temperature must the reaction be carried out so that the equilibrium constant is larger than 1?
(a) About 1 K. (b) About 1°C (c) About 700 K (d) About 700°C
(e) The reaction never has an equilibrium constant larger than 1.
Answer: (d)
13-47. For the reaction: 2 SO3(g) ➝ 2 SO2(g) + O2(g) the following thermodynamic data are available.
Compound ΔHac° (kJ/mol) S0 (J/mol-K)
SO3(g) -1422.04 -394.23
SO2(g) -1073.95 -241.21
O2(g) -498.34 -116.97
What is the equilibrium constant for this reaction at 25°C?
(a) between 10-50 and 10-30 (b) between 10-30 and 10-20
(c) roughly 1 (d) between 1020 and 1030
(e) between 1030 and 1050
Answer: (b)
13-48. The standard-state cell potential for the following reaction is 0.23 V.
2 Fe3+(aq) + 2 I-(aq) ➝ 2 Fe2+(aq) + I2(aq)
If ΔG° = - nFE° and F = 96,500 C, how large is the equilibrium constant for this reaction at 25°C?
(a) between 10-20 and 10-8 (b) between 10-8 and 10-5
(c) between 10-5 and 105 (d) between 105 and 108
(e) between 108 and 1020
Answer: (d)