AP Chemistry Chapter 18 & 19: Thermodynamics & Electrochemistry
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AP Chemistry Chapter 18 & 19: Thermodynamics & Electrochemistry
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AP Chemistry Chapter 18 & 19: Thermodynamics & Electrochemistry
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•Warm-ups and problems will be collected before you take the test.
•Unit 13. Read Chapter 18: Entropy, Free Energy, and Equilibrium & Read Chapter 19: Electrochemistry
Answer the following problems in the space provided. For problems involving an equation, carry out the
following steps: 1. Write the equation. 2. Substitute numbers and units. 3. Show the final answer with units.
There is no credit without showing work.
Spontaneous Processes and Entropy
1. Explain what is meant by a “spontaneous” process (thermodynamically favored process). What two factors
determine whether a process is thermodynamically favored?
2. Which of the following processes are spontaneous and which are nonspontaneous?
a. dissolving table salt (NaCl) in hot soup
b. climbing Mt. Everest
c. spreading fragrance in a room by removing the cap from a perfume bottle
d. separating helium and neon from a mixture of the gases
2. Which of the following processes are spontaneous and which are nonspontaneous at a given temperature?
a. NaNO3(s)
→
NaNO3(aq) (saturated solution)
b. NaNO3(s)
→
NaNO3(aq) (unsaturated solution)
c. NaNO3(s)
HO
→
NaNO3(aq) (supersaturated solution)
H2O
H2O
2
3. How does the entropy of a system change (increase (I), decrease (D), or stay the same (S)) for each of the
following processes?
a. A solid melts
b. A liquid freezes
c. A vapor is converted to a solid
d. A vapor condenses to a liquid
e. A solid sublimes
f.
Urea dissolves in water.
Second Law of Thermodynamics
4. Arrange the following substances (1 mole each) in order of increasing entropy at 25°C:
Ne(g), SO2(g), Na(s), NaCl(s), NH3(g). Give the reasons for your arrangement.
AP Chemistry Chapter 18 & 19: Thermodynamics & Electrochemistry
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5. Using the data in Appendix 3, calculate the standard entropy changes for the following reactions at 25°C:
a. 2Al(s) + 3ZnO(s)  Al2O3(s) + 3Zn(s)
b. CH4(g) + 2O2(g)  CO2(g) + 2H2O(l)
6. State whether the sign of the entropy change expected for each of the following processes will be positive or
negative, and explain your predictions.
a. PC13(l) + C12(g)  PCl5(s)
b. 2HgO(s)  2Hg(l) + O2(g)
c. H2(g)  2H(g)
Gibbs Free Energy
7. Calculate G° for the following reactions at 25°C, and state whether each is spontaneous:
a. 2SO2(g) + O2(g)  2SO3(g)
b. 2C2H6(g) + 7O2(g)  4CO2(g) + 6H2O(l)
8. From the signs and/or values of H and S, predict which of the following reactions would be spontaneous at
25oC. If any reaction is nonspontaneous at 25oC, at what temperature would it become spontaneous?
a. H = 10.5 kJ and S = 30 J/K
b. H = 1.8 kJ and S = -113 J/K
c. H = -126 kJ and S = 84 J/K
d. H = -11.7 kJ and S = -105 J/K
AP Chemistry Chapter 18 & 19: Thermodynamics & Electrochemistry
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9. Predict the signs of , S, and G of the system for the following processes at 1 atm, given the normal
melting point of ammonia is -77.7°C.
H
S
G
a. ammonia melts at -60°C
b. ammonia melts at -77.7°C
c. ammonia melts at -100°C
10. Ammonium nitrate (NH4NO3) readily dissolves in water, and the solution temperature decreases during
dissolution. What can you deduce about the sign of S, and G for the solution process?
11. Crystallization of sodium acetate from a supersaturated solution occurs spontaneously. What can you deduce
about the signs of S, and G?
12. The reaction below proceeds spontaneously at 25 °C even though there is a decrease in disorder in the system
(gases are converted to a solid). Explain.
NH3(g) + HCl(g)  NH4C1(s)
13. Heating copper(II) oxide at 400oC does not produce any appreciable amount of Cu.
CuO(s)  Cu(s) + ½O2(g)
Go = 127.2 kJ
However, if this reaction is coupled with the conversion of graphite to carbon dioxide, it becomes spontaneous.
Write an equation for the coupled process, and calculate the free energy change for the coupled reaction.
14. In the human body, glucose is burned to provide energy according to the reaction:
C6H12O6(s) + 6O2(g)  6CO2(g) + 6H2O(l)
Go = -2880 kJ/mol
Rather than using this energy immediately, much of this energy is stored in cells by converting adenosine
diphosphate (ADP) to adenosine triphosphate (ATP) in an endothermic reaction (shown below) coupled with
the burning of glucose. The energy thus stored in ATP in the cell can be used at a later time.
ADP(aq) + H2PO4-(aq)  ATP(aq) +H2O(l)
Go = 30.5 kJ/mol
If all the free energy from burning one molecule of glucose goes into forming ATP from ADP, how many ATP
molecules can be produced?
AP Chemistry Chapter 18 & 19: Thermodynamics & Electrochemistry
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15. Explain the difference between a thermodynamically controlled reaction and a kinetically controlled reaction.
16. Calculate G° for the process: C(diamond)  C(graphite).
Is the reaction spontaneous at 25°C? If so, why is it that diamonds do not become graphite on standing?
17. The heat of vaporization of water as given in the N. Y. State reference tables is 2260 J/g.
a. Write a chemical equation for this process.
b. Calculate the enthalpy change in kJ/mol at the phase transition temperature.
c. Calculate the entropy change for this process at the phase transition temperature.
d.
What is the value of G at the phase transition temperature?
Free Energy and Chemical Equilibrium
18. Explain the difference between G and Go.
19. For the autoionization of water at 25°C:
H2O(l)  H+(aq) + OH-(aq)
Kw is 1.0 x 10-14. What is G° for the process? Does your answer make sense?
20. Calculate G for the reaction below at 25 °C for the following conditions:
H2O(l)  H+(aq) + OH-(aq)
a. [H+] = 1.0 x 10-7 M, [OH-] = 1.0 x 10-7 M
b. [H+] = 1.0 x 10-3 M, [OH-] = 1.0 x 10-4 M
c. [H+] = 3.5 M, [OH-] = 4.8 x 10-4 M
AP Chemistry Chapter 18 & 19: Thermodynamics & Electrochemistry
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21. The equilibrium constant (KP) for the reaction for the reaction below is 4.40 at 2000 K.
H2(g) + CO2(g)  H2O(g) + CO(g)
a. Calculate G° for the reaction.
b. Calculate G for the reaction when the partial pressures are PH2 = 0.25 atm, PCO2 = 0.78 atm, PH2O = 0.66
atm, and PCO = 1.20 atm.
22. The equilibrium constant KP for the reaction below is 5.62 x 1035 at 25°C. Calculate Gfo for COC12 at 25°C.
CO(g) + Cl2(g)  COCl2(g)
23. Calculate the pressure of O2 (in atm) over a sample of NiO at 25°C if G° = 212 kJ for the reaction
NiO(s)  Ni(s) + ½O2(g)
24. Carbon monoxide (CO) and nitric oxide (NO) are polluting gases contained in automobile exhaust. Using a
catalytic converter in the exhaust pipe, these gases can be made to react to form nitrogen (N2) and the less
harmful carbon dioxide (CO2).
a. Write a balanced equation for this reaction.
b. Identify the oxidizing and reducing agents.
c. Calculate the KP for the reaction at 25 °C.
AP Chemistry Chapter 18 & 19: Thermodynamics & Electrochemistry
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d. Under normal atmospheric conditions, the partial pressures are PN2 = 0.80 atm, PCO2 = 3.0 x 10-4 atm, PCO =
5.0 x 10-5 atm, and PNO = 5.0 x 10-7 atm . Calculate QP and predict the direction toward which the reaction
will proceed.
e. Will raising the temperature favor the formation of N2 and CO2?
25. One of the steps in the extraction of iron from its ore (FeO) is the reduction of iron(II) oxide by carbon
monoxide at 900°C:
FeO(s) + CO(g)  Fe(s) + CO2(g)
Calculate the equilibrium constant for this reaction. State any assumption.
Balancing Redox Reactions
26. Balance the following redox equations.
a. Mn2+ + H2O2  MnO2 + H2O (in basic solution)
b. Bi(OH)3 + SnO22-  SnO32- + Bi (in basic solution)
c. Cr2O72- + C2O42-  Cr3+ + CO2 (in acidic solution)
AP Chemistry Chapter 18 & 19: Thermodynamics & Electrochemistry
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27. Calcium oxalate (CaC2O4) is insoluble in water. This property has been used to determine the amount of Ca2+
ions in blood. The calcium oxalate isolated from blood is dissolved in acid and titrated against a standardized
KMnO4 solution using the unbalanced equation below.
MnO4- + C2O42-  Mn2+ + CO2
(in acidic solution)
In one test it is found that the calcium oxalate isolated from a 10.0-mL sample of blood requires 24.2 mL of
9.56 x 10-4 M KMnO4 for titration. First balance the equation, and then calculate the number of milligrams of
calcium per milliliter of blood.
Electrochemical Cells and Standard EMF
28. What is the function of the salt bridge in a galvanic cell? What are the requirements of the salt used in the salt
bridge?
29. Calculate the standard emf of a cell that uses Ag/Ag+ and A1/A13+ half-cell reactions, and write the cell
reaction that occurs under standard-state conditions.
30. Which of the following reagents can oxidize H2O to O2(g) under standard-state conditions?
a. H+(aq)
b. Cl-(aq)
c. Cl2(g)
d. Cu2+(aq)
e. Pb2+(aq)
31. Predict whether the following reactions occur spontaneously in aqueous solution at 25°C. Assume that the
initial concentrations of dissolved species are all 1.0 M.
a. Ca(s) + Cd2+(aq)  Ca2+(aq) + Cd(s)
b. 2Br-(aq) + Sn2+(aq)  Br2(l) + Sn(s)
c. 2Ag(s) + Ni2+(aq)  2Ag+(aq) + Ni(s)
AP Chemistry Chapter 18 & 19: Thermodynamics & Electrochemistry
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32. Which species in each pair is a better reducing agent under standard-state conditions?
a. Na or Li
b. H2 or I2
c. Fe2+ or Ag
d. Br- or Co2+
Spontaneity of Redox Reactions
33. The equilibrium constant for the reaction below is 2.69 x 1012 at 25°C.
Sr(s) + Mg2+(aq)  Sr2+(aq) + Mg(s)
Calculate E° for a cell made up of Sr/Sr2+ and Mg/Mg2+ half-cells.
34. Calculate Eocell, G°, and Kc for the following reactions at 25°C:
a. Mg(s) + Pb2+(aq)  Mg2+(aq) + Pb(s)
b. O2(g) + 4H+(aq) + 4Fe2+(aq)  2H2O(l) + 4Fe3+(aq)
35. Comment on whether F2 will become a stronger oxidizing agent with increasing H+ concentration.
The Effect of Concentration on Cell EMF
36. Calculate E°, and predict whether the voltage would be greater, the same, or less than Eo for the following cell
reactions.
a. Mg(s) + Sn2+(aq)  Mg2+(aq) + Sn(s)
with [Mg2+] = 0.045 M and [Sn2+] = 1.00 M
b. 3Zn(s) + 2Cr3+(aq)  3Zn2+(aq) + 2Cr(s)
with [Cr3+] = 0.010 M and [Zn2+] = 1.00 M
AP Chemistry Chapter 18 & 19: Thermodynamics & Electrochemistry
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Electrolysis
37. Consider the electrolysis of molten barium chloride, BaCl2.
a. Write the half-reactions and identify the anode and cathode.
b. How many grams of barium metal can be produced by supplying 0.50 A for 30 min?
38. If the cost of electricity to produce magnesium by the electrolysis of molten magnesium chloride is $155 per
ton of metal, what is the dollar cost of the electricity necessary to produce 10.0 tons of aluminum? (1 ton =
2000 lb, 1 lb = 453.6 g)
39. How many faradays of electricity are required to produce:
a. 0.84 L of O2 at exactly 1 atm and 25°C from aqueous H2SO4 solution
b. 1.5 L of C12 at 750 mmHg and 20°C from molten NaCl
40. A constant electric current flows for 3.75 h through two electrolytic cells connected in series. One contains a
solution of AgNO3 and the second a solution of CuCl2. During this time 2.00 g of silver are deposited in the
first cell.
a. How many grams of copper are deposited in the second cell?
b. What is the current flowing, in amperes?
AP Chemistry Chapter 18 & 19: Thermodynamics & Electrochemistry
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41. In the electrolysis of acidified water, 0.845 L of H2 is collected at 25°C and 782 mmHg. How many faradays of
electricity had to pass through the solution?
Review
42. This year we studied three types of spectroscopy. In each type, the energy of the photon absorbed matches
some physical process taking place within the atom or molecule. For each type of spectroscopy, describe the
process taking place within the atom or molecule that gives rise to the absorption of a photon, and state how
each type of spectroscopy is used in a practical sense.
a. UV/visible spectroscopy (UV/vis)
Process:
Use:
b. Photoelectron spectroscopy (PES)
Process:
Use:
c. Infrared spectroscopy (IR)
Process:
Use: