Chapter 21: Electrochemistry: Chemical Change and
Electrical Work
1. Which one of the following is not a redox reaction?
A) Al(OH)4–(aq) + 4H+(aq) → Al3+(aq) + 4H2O(l)
B) C6H12O6(s) + 6O2(g) → 6CO2(g) + 6H2O(l)
C) Na6FeCl8(s) + 2Na(l) → 8NaCl(s) + Fe(s)
D) 2H2O2(aq) → 2H2O(l) + O2(g)
E) CO2(g) + H2(g) → CO(g) + H2O(g)
2. Consider the reaction
CuO(s) + H2(g) → Cu(s) + H2O(l)
In this reaction, which substances are the oxidant and reductant, respectively?
A) CuO and H2 D) H2O and H2
B) H2 and CuO E) none of the above
C) CuO and Cu
3. Consider the following balanced redox reaction
Mn2+(aq) + S2O82–(aq) + 2H2O(l) → MnO2(s) + 4H+(aq) + 2SO42–(aq)
Which of the following statements is true?
A) Mn2+(aq) is the oxidizing agent and is reduced.
B) Mn2+(aq) is the oxidizing agent and is oxidized.
C) Mn2+(aq) is the reducing agent and is oxidized.
D) Mn2+(aq) is the reducing agent and is reduced.
E) Manganese does not change its oxidation number in this reaction.
4. Consider the following balanced redox reaction
3CuO(s) + 2NH3(aq) → N2(g) + 3H2O(l) + 3Cu(s)
Which of the following statements is true?
A) CuO(s) is the oxidizing agent and copper is reduced.
B) CuO(s) is the oxidizing agent and copper is oxidized.
C) CuO(s) is the reducing agent and copper is oxidized.
D) CuO(s) is the reducing agent and copper is reduced.
E) CuO(s) is the oxidizing agent and N2(g) is the reducing agent.
Chapter 21: Electrochemistry: Chemical Change and Electrical Work
5. When the following redox equation is balanced with smallest whole number coefficients,
the coefficient for nitrogen dioxide will be _____.
I2(s) + HNO3(aq) → HIO3(aq) + NO2(g) + H2O(l)
A) 1 B) 2 C) 4 D) 10 E) none of the above
6. When the following redox equation is balanced with smallest whole number coefficients,
the coefficient for the hydrogen sulfate ion will be ______.
Al(s) + HSO4–(aq) + OH–(aq) → Al2O3(s) + S2–(aq) + H2O(l)
A) 1 B) 3 C) 4 D) 8 E) none of the above
7. When the following redox equation is balanced with smallest whole number coefficients,
the coefficient for zinc will be _____.
Zn(s) + ReO4–(aq) → Re(s) + Zn2+(aq) (acidic solution)
A) 2 B) 7 C) 8 D) 16 E) none of the above
8. When the following redox equation is balanced with smallest whole number coefficients,
the coefficient for the iodide ion will be _____.
I–(aq) + NO3–(aq) → NO(g) + I2(s) (acidic solution)
A) 2 B) 3 C) 6 D) 8 E) none of the above
9. When the following redox equation is balanced with smallest whole number coefficients,
the coefficient for Sn(OH)3– will be _____.
Bi(OH)3(s) + Sn(OH)3–(aq) → Sn(OH)62–(aq) + Bi(s) (basic solution)
A) 1 B) 2 C) 3 D) 6 E) none of the above
10. Consider the following redox equation
Mn(OH)2(s) + MnO4–(aq) → MnO42–(aq) (basic solution)
When the equation is balanced with smallest whole number coefficients, what is the coefficient
for OH–(aq) and on which side of the equation is OH–(aq) present?
A) 4, reactant side D) 6, product side
B) 4, product side E) none of the above
C) 6, reactant side
Chapter 21: Electrochemistry: Chemical Change and Electrical Work
11. Which of the following statements about voltaic and electrolytic cells is correct?
A) The electrons in the external wire flow from cathode to anode in both types of cell.
B) Oxidation occurs at the cathode only in a voltaic cell.
C) The free energy change, G, is negative for an electrolytic cell.
D) The cathode is labeled as positive (+) in a voltaic cell but negative (–) in an
electrolytic cell.
E) Reduction occurs at the anode in an electrolytic cell.
12. Which of the following statements about voltaic and electrolytic cells is correct?
A) The anode will definitely gain weight in a voltaic cell.
B) Oxidation occurs at the cathode of both cells.
C) The free energy change, G, is negative for the voltaic cell.
D) The electrons in the external wire flow from cathode to anode in an electrolytic
cell.
E) None of the above statements is correct.
13. Which one of the following pairs of substances could be used to construct a single redox
electrode (i.e., they have an element in common, but in different oxidation states)?
A) HCl and Cl– D) Fe3+ and Fe2O3
B) H+ and OH– E) MnO2 and Mn2+
C) H2O and H+
14. Which one of the following statements about electrochemical cells is correct?
A) In a salt bridge, current is carried by cations moving toward the anode, and anions
toward the cathode.
B) In the external wire, electrons travel from cathode to anode.
C) The anode of a voltaic cell is labeled minus (–).
D) Oxidation occurs at the cathode, in an electrolytic cell.
E) None of the above statements is correct.
15. Which of the following solids is commonly used as an inactive electrode in
electrochemical cells?
A) zinc B) graphite C) copper D) iron E) sodium
16. Which component of the following cell notation is the anode?
P | Q || R | S
A) P B) Q C) R D) S E) One of the | symbols is the anode.
Chapter 21: Electrochemistry: Chemical Change and Electrical Work
Page 376
17. A voltaic cell is prepared using copper and silver. Its cell notation is shown below.
Cu(s) | Cu2+(aq) || Ag+(aq) | Ag(s)
Which of the following processes occurs at the cathode?
A) Cu(s) → Cu2+(aq) + 2e–
B) Cu2+(aq) + 2e– → Cu(s)
C) Ag(s) → Ag+(aq) + e–
D) Ag+(aq) + e– → Ag(s)
E) Cu(s) + 2Ag+(aq) → Cu2+(aq) + 2Ag(s)
18. A voltaic cell prepared using aluminum and nickel has the following cell notation.
Al(s) | Al3+(aq) || Ni2+(aq) | Ni(s)
Which of the following reactions occurs at the anode?
A) Al(s) → Al3+(aq) + 3e– D) Ni2+(aq) + 2e– → Ni(s)
B) Al3+(aq) + 3e → Al(s) E) none of the above
C) Ni(s) → Ni2+(aq) + 2e–
19. A voltaic cell prepared using aluminum and nickel has the following cell notation.
Al(s) | Al3+(aq) || Ni2+(aq) | Ni(s)
Which of the following represents the correctly balanced spontaneous reaction equation for the
cell? A) Ni2+(aq) + Al(s) → Al3+(aq) + Ni(s)
B) 3Ni2+(aq) + 2Al(s) → 2Al3+(aq) + 3Ni(s)
C) Ni(s) + Al3+(aq) → Ni2+(aq) + Al(s)
D) 3Ni(s) + 2Al3+(aq) → 3Ni2+(aq) + 2Al(s)
E) none of the above
20. A voltaic cell prepared using zinc and iodine has the following cell notation.
Zn(s) | Zn2+(aq) || I–(aq) | I2(s) | C(graphite)
Which of the following equations correctly represents the balanced, spontaneous, cell reaction?
A) 2I–(aq) + Zn2+(aq) → I2(s) + Zn(s)
B) I2(s) + Zn(s) → 2I–(aq) + Zn2+(aq)
C) 2I–(aq) + Zn(s) → I2(s) + Zn2+(aq)
D) I2(s) + Zn2+(aq) → 2I–(aq) + Zn(s)
21. The line notation, Al(s) | Al3+(aq) || Co2+(aq) | Co(s), indicates that
A) Co is the reducing agent. D) Al3+ is the reducing agent.
B) Co2+ ions are oxidized. E) aluminum metal is the cathode.
C) Al is the reducing agent.
Chapter 21: Electrochemistry: Chemical Change and Electrical Work
Page 377
22. The line notation, Pt | H2(g) | H+(aq) || Cu2+(aq) | Cu(s), indicates that
A) copper metal is a product of the cell reaction.
B) hydrogen gas (H2) is a product of the cell reaction.
C) Cu is the anode.
D) Pt is the cathode.
E) Cu2+ is the reducing agent.
23. A cell can be prepared from copper and tin. What is the E°cell for the cell that forms from
the following half-reactions?
Cu2+(aq) + 2e– Cu(s) E° = 0.34 V
Sn4+(aq) + 2e– Sn2+(aq) E° = 0.13 V
A) 0.47 V B) 0.21 V C) –0.21 V D) –0.47 V E) 0.42 V
24. What is the E°cell for the cell represented by the combination of the following half-
reactions?
2Hg2+(aq) + 2e– Hg22+(aq) E°= 0.92 V
Cr3+ (aq) + 3e– Cr(s) E°= –0.74 V
A) –0.18 V B) 0.18 V C) 1.28 V D) 1.66 V E) 2.12 V
25. What is the E°cell for the cell represented by the combination of the following half-
reactions?
ClO4–(aq) + 8H+(aq) + 8e– Cl–(aq) + 4H2O(l) E° = 1.389 V
VO2+(aq) + 2H+(aq) + e– VO+(aq) + H2O(l) E° = 0.991 V
A) –0.398 V B) –2.380 V C) 0.398 V D) 2.380 V E) none of the above
26. The redox reaction of peroxydisulfate with iodide has been used for many years as part of
the iodine clock reaction which introduces students to kinetics. If E°cell = 1.587 V and E° of the
cathode half-cell is 0.536 V, what is E° of the anode half-cell?
S2O82–(aq) + 2H+ + 2I–(aq) → 2HSO4–(aq) + I2(aq)
A) –1.051 V B) –2.123 V C) 1.051 V D) 2.123 V E) none of the above
27. The voltaic cell made up of cobalt, copper, and their M2+ ions, has E°cell = 0.62 V. If E°
of the cathode half-cell is 0.34 V, what is E° of the anode half-cell?
Cu2+(aq) + Co(s) → Cu(s) + Co2+(aq)
Chapter 21: Electrochemistry: Chemical Change and Electrical Work
28. A voltaic cell has a standard cell potential equal to 0.74 V. If the standard electrode
(reduction) potential for the anode is – 0.22 V, what is the standard electrode potential for the
cathode?
A) 0.96 V
B) 0.52 V
C) -0.52 V
D) -0.96 V
E) Need to know the cell reaction in order to calculate the answer.
29. Given that E° for X + e– → Y is greater than E° for A + 2e– → B, it is correct to say
that, under standard conditions
A) X will oxidize A. D) B will oxidize X.
B) Y will oxidize A. E) B will reduce X.
C) Y will reduce A.
30. Examine the following half-reactions and select the strongest oxidizing agent among the
substances.
[PtCl4]2–(aq) + 2e– Pt(s) + 4Cl–(aq) E° = 0.755 V
RuO4(s) + 8H+(aq) + 8e– Ru(s) + 4H2O(l) E° = 1.038 V
FeO42–(aq) + 8H+(aq) + 3e– Fe3+(aq) + 4H2O(l) E° = 2.07 V
H4XeO6(aq) + 2H+(aq) + 2e– XeO3(aq) + 3H2O(l) E° = 2.42 V
A) [PtCl4]2–(aq) B) RuO4(s) C) HFeO4– (aq) D) H4XeO6(aq) E) Cl–(aq)
31. Examine the following half-reactions and select the strongest oxidizing agent among the
species listed.
Cr2+(aq) + 2e– Cr(s) E° = –0.913 V
Fe2+(aq) + 2e– Fe(s) E° = –0.447 V
Sr2+(aq) + 2e– Sr(s) E° = –2.89 V
Co2+(aq) + 2e– Co(s) E° = –0.28 V
A) Cr2+(aq) B) Fe(s) C) Fe2+(aq) D) Sr2+(aq) E) Co2+(aq)
Chapter 21: Electrochemistry: Chemical Change and Electrical Work
32. Examine the following half-reactions and select the weakest oxidizing agent among the
species listed.
AuBr4–(aq) + 3e– Au(s) + 4Br–(aq) E° = 0.854 V
Mn2+(aq) + 2e– Mn(s) E° = –1.185 V
K+(aq) + e– K(s) E° = –2.931 V
F2O(aq) + 2H+(aq) + 4e– 2F–(aq) + H2O(l) E° = 2.153 V
A) AuBr4–(aq) B) Mn2+(aq) C) K+(aq) D) F2O(aq) E) H+(aq)
33. Examine the following half-reactions and select the strongest reducing agent among the
species listed.
PbI2(s) + 2e– Pb(s) + 2I–(aq) E° = –0.365 V
Ca2+(aq) + 2e– Ca(s) E° = –2.868 V
Pt2+(aq) + 2e– Pt(s) E° = 1.18 V
Br2(l) + 2e– 2Br–(aq) E° = 1.066 V
A) Pb(s) B) Ca(s) C) Pt(s) D) Br–(aq) E) Pt2+(aq)
34. Examine the following half-reactions and select the strongest reducing agent among the
species listed.
HgO(s) + H2O(l) + 2e– Hg(l) + 2OH–(aq) E° = 0.0977 V
Zn(OH)2(s) + 2e– Zn(s) + 2OH–(aq) E° = –1.25 V
Ag2O(s) + H2O(l) + 2e– Ag(s) + 2OH–(aq) E° = 0.342 V
B(OH)3(aq) + 7H+(aq) + 8e– BH4–(aq) + 3H2O(l) E° = –0.481 V
A) Hg(l) B) Zn(s) C) Ag(s) D) BH4–(aq) E) Zn(OH)2(s)
35. Examine the following half-reactions and select the weakest reducing agent among the
substances.
Cr(OH)3(s) + 3e– Cr(s) + 3OH–(aq) E° = –1.48 V
SnO2(s) + 2H2O(l) + 4e– Sn(s) + 4OH–(aq) E° = –0.945 V
MnO2(s) + 4H+(aq) + 2e– Mn2+(aq) + 2H2O(l) E° = 1.224 V
Hg2SO4(s) + 2e– 2Hg(l) + SO42–(aq) E° = 0.613 V
A) Cr(s) B) Sn(s) C) Mn2+(aq) D) Hg(l) E) OH–(aq)
Chapter 21: Electrochemistry: Chemical Change and Electrical Work
36. Calculate E°cell and indicate whether the overall reaction shown is spontaneous or
nonspontaneous.
I2(s) + 2e– 2I–(aq) E° = 0.53 V
Cr3+(aq) + 3e– Cr(s) E° = –0.74 V
Overall reaction:
2Cr(s) + 3I2(s) → 2Cr3+(aq) + (aq) + 6I–(aq)
A) E°cell = –1.27 V, spontaneous D) E°cell = 1.27 V, nonspontaneous
B) E°cell = –1.27 V, nonspontaneous E) E°cell = 1.54 V, spontaneous
C) E°cell = 1.27 V, spontaneous
37. Calculate E°cell and indicate whether the overall reaction shown is spontaneous or
nonspontaneous.
Co3+(aq) + e– Co2+(aq) E° = 1.82 V
MnO4–(aq) + 2H2O(l) + 3e– MnO2(s) + 4OH–(aq) E° = 0.59 V
Overall reaction:
MnO4–(aq) + 2H2O(l) + 3Co2+(aq) → MnO2(s) + 3Co3+(aq) + 4OH–(aq)
A) E°cell = –1.23 V, spontaneous D) E°cell = 1.23 V, nonspontaneous
B) E°cell = –1.23 V, nonspontaneous E) E°cell = –0.05 V, nonspontaneous
C) E°cell = 1.23 V, spontaneous
38. Calculate E°cell and indicate whether the overall reaction shown is spontaneous or
nonspontaneous.
O2(g) + 4H+(aq) + 4e– 2H2O(l) E° = 1.229 V
Al3+(aq) + 3e– Al(s) E° = –1.662 V
Overall reaction:
4Al(s) + 3O2(g) + 12H+(aq) → 4Al3+(aq) + 6H2O(l)
A) E°cell = –2.891 V, nonspontaneous
B) E°cell = –2.891 V, spontaneous
C) E°cell = 2.891 V, nonspontaneous
D) E°cell = 2.891 V, spontaneous
E) Spontaneous, but none of the above values of E°cell is correct.
Chapter 21: Electrochemistry: Chemical Change and Electrical Work
39. Calculate E°cell and indicate whether the overall reaction shown is spontaneous or
nonspontaneous.
H2O2(aq) + 2H+(aq) + 2e– 2H2O(l) E° = 1.77 V
Fe3+ (aq) + e– Fe2+(aq) E° = 0.77 V
Overall reaction:
2Fe3+(aq) + 2H2O(l) → H2O2(aq) + 2H+(aq) + 2Fe2+(aq)
A) E°cell = –1.00 V, nonspontaneous D) E°cell = 1.00 V, spontaneous
B) E°cell = –1.00 V, spontaneous E) E°cell = –0.23 V, nonspontaneous
C) E°cell = 1.00 V, nonspontaneous
40. When metal A is placed in a solution of metal ions B2+, a reaction occurs between A and
B2+, and metal ions A2+ appear in the solution. When metal B is placed in acid solution, gas
bubbles form on its surface. When metal A is placed in a solution of metal ions C2+, no reaction
occurs. Which of the following reactions would not occur spontaneously?
A) C(s) + 2H+(aq) → H2(g) + C2+(aq)
B) C(s) + A2+(aq) → A(s) + C2+(aq)
C) B(s) + C2+(aq) → C(s) + B2+(aq)
D) A(s) + 2H+(aq) → H2(g) + A2+(aq)
E) B(s) + 2H+(aq) → H2(g) + B2+(aq)
41. Which of the following conditions is most likely to apply to a fully-charged secondary
cell?
A) Ecell = E°cell B) E°cell = 0 C) Q = 1 D) Q < K E) Q = K
42. A battery is considered “dead” when
A) Q < 1. B) Q = 1. C) Q > 1. D) Q = K. E) Q/K = 0.
43. Calculate the potential of a voltaic cell (E°cell) if it is required to do 5.43 × 10-3 kJ of
work when a charge of 2.50 C is transferred.
A) 2.17 × 103 V
B) 2.17 × 10-3 V
C) 2.17 V
D) 13.6 V
E) 1.36 × 10-2 V
Chapter 21: Electrochemistry: Chemical Change and Electrical Work
44. What is the value of the equilibrium constant for the cell reaction below at 25°C? E°cell =
0.30 V
Sn2+(aq) + Fe(s) Sn(s) + Fe2+(aq)
A) 1.2 × 105 B) 1.4 × 1010 C) 8.6 × 10–6 D) 7.1 × 10–11 E) 2.3 × 1023
45. What is the value of the equilibrium constant for the cell reaction below at 25°C? E°cell =
0.61 V
2Cr(s) + 3Pb2+(aq) 3Pb(s) + 2Cr3+(aq)
A) 4.1 × 1020
B) 8.2 × 1030
C) 3.3 × 1051
D) 7.4 × 1061
E) > 9.9 × 1099
46. The following half-reactions occur in the mercury battery used in calculators. If E°cell =
1.357 V, calculate the equilibrium constant for the cell reaction at 25°C. (Assume the
stoichiometric coefficients in the cell reaction are all equal to 1.)
HgO(s) + H2O(l) + 2e– Hg(l) + 2OH–(aq)
ZnO(s) + H2O(l) + 2e– Zn(s) + 2OH–(aq)
A) 9.4 × 1022
B) 7.5 × 1045
C) 6.4 × 1063
D) 7.8 × 1091
E) > 9.9 × 1099
47. Consider the non-aqueous cell reaction
2Na(l) + FeCl2(s) 2NaCl(s) + Fe(s)
for which E°cell = 2.35 V at 200°C. G° at this temperature is
A) 453 kJ. B) –453 kJ. C) 907 kJ. D) –907 kJ. E) none of the above.
48. Consider the reaction in the lead-acid cell
Pb(s) + PbO2(s) + 2H2SO4(aq) → 2PbSO4(aq) + 2H2O(l)
for which E°cell = 2.04 V at 298 K. G° for this reaction is
A) –3.94 × 105 kJ. D) –7.87 × 102 kJ.
B) –3.94 × 102 kJ. E) none of the above.
C) –1.97 × 105 kJ.
Chapter 21: Electrochemistry: Chemical Change and Electrical Work
49. The value of E°cell for the reaction
2Cr3+(aq) + 6Hg(l) → 2Cr(s) + 3Hg22+(aq)
is 1.59 V. Calculate G° for the reaction.
A) –921 kJ B) –767 kJ C) –460 kJ D) –307 kJ E) none of the above
50. Calculate G° for the reaction of iron(II) ions with one mole of permanganate ions.
MnO4–(aq) + 8H+(aq) + 5e– Mn2+(aq) + 4H2O(l) E° = 1.51 V
Fe3+(aq) + e– Fe2+(aq) E°= 0.77 V
A) –71.4 kJ B) –286 kJ C) –357 kJ D) –428 kJ E) none of the above
51. Calculate G° for the oxidation of 3 moles of copper by nitric acid.
Cu2+(aq) + 2e– Cu(s) E° = 0.34 V
NO3–(aq) + 4H+(aq) + 3e– NO(g) + 2H2O(l) E° = 0.957 V
A) –120 kJ B) –180 kJ C) –240 kJ D) –300 kJ E) –360 kJ
52. The value of the equilibrium constant for the reaction of nickel(II) ions with cadmium
metal is 1.17 × 105. Calculate G° for the reaction at 25°C.
A) –12.6 kJ B) –28.9 kJ C) 12.6 kJ D) 28.9 kJ E) none of the above
53. Calculate E°cell for the reaction of nickel(II) ions with cadmium metal at 25°C.
K = 1.17 × 105
Ni2+(aq) + Cd(s) → Cd2+(aq) + Ni(s)
A) 0.075 V B) 0.10 V C) 0.12 V D) 0.15 V E) 0.30 V
54. Consider the reaction of iodine with manganese dioxide
3I2(s) + 2MnO2(s) + 8OH–(aq) 6I–(aq) + 2MnO4–(aq) + 4H2O(l)
The equilibrium constant for the overall reaction is 8.30 × 10–7. Calculate G° for the reaction at
25°C.
A) –15.1 kJ B) –34.7 kJ C) 15.1 kJ D) 34.7 kJ E) none of the above
55. Consider the reaction of iodine with manganese dioxide
3I2(s) + 2MnO2(s) + 8OH–(aq) 6I–(aq) + 2MnO4–(aq) + 4H2O(l)
The equilibrium constant for the overall reaction is 8.30 × 10–7. Calculate E°cell for the reaction
at 25°C.
A) –0.36 V B) –0.18 V C) –0.12 V D) –0.060 V E) none of the above
Chapter 21: Electrochemistry: Chemical Change and Electrical Work
Page 384
56. A voltaic cell consists of a Mn/Mn2+ electrode (E° = –1.18 V) and a Fe/Fe2+ electrode
(E° = –0.44 V). Calculate [Fe2+] if [Mn2+] = 0.050 M and Ecell = 0.78 V at 25°C.
A) 0.040 M B) 0.24 M C) 1.1 M D) 1.8 M E) none of the above
57. A voltaic cell consists of an Au/Au3+ electrode (E° = 1.50 V) and a Cu/Cu2+ electrode
(E° = 0.34 V). Calculate [Au3+] if [Cu2+] = 1.20 M and Ecell = 1.13 V at 25°C.
A) 0.001 M B) 0.002 M C) 0.01 M D) 0.02 M E) 0.04 M
58. A voltaic cell consists of a Hg/Hg22+ electrode (E° = 0.85 V) and a Sn/Sn2+ electrode (E°
= –0.14 V). Calculate [Sn2+] if [Hg22+] = 0.24 M and Ecell = 1.04 V at 25°C.
A) 0.0001 M B) 0.0007 M C) 0.005 M D) 0.03 M E) 0.05 M
59. A voltaic cell consists of a Cd/Cd2+ electrode (E° = –0.40 V) and a Fe/Fe2+ electrode
(E° = –0.44 V). If Ecell = 0 and the temperature is 25°C, what is the ratio [Fe2+]/[Cd2+]?
A) 2 × 101 B) 1 × 101 C) 1 D) 1 × 10–1 E) 5 × 10–2
60. A voltaic cell consists of a Ag/Ag+ electrode (E° = 0.80 V) and a Fe2+/Fe3+ electrode
(E° = 0.77 V) with the following initial molar concentrations: [Fe2+] = 0.30 M;
[Fe3+] = 0.10 M; [Ag+] = 0.30 M. What is the equilibrium concentration of Fe3+? (Assume the
anode and cathode solutions are of equal volume, and a temperature of 25°C.)
A) 0.030 M B) 0.043 M C) 0.085 M D) 0.11 M E) 0.17 M
61. A concentration cell consists of two Zn/Zn2+ electrodes. The electrolyte in compartment
A is 0.10 M Zn(NO3)2 and in compartment B is 0.60 M Zn(NO3)2. What is the voltage of the cell
at 25°C?
A) 0.010 V
B) 0.020 V
C) 0.023 V
D) 0.046 V
E) none of the above
62. A concentration cell consists of two Al/Al3+electrodes. The electrolyte in compartment A
is 0.050 M Al(NO3)3 and in compartment B is 1.25 M Al(NO3)3. What is the voltage of the cell
at 25°C?
A) 0.083 V B) 0.062 V C) 0.041V D) 0.028 V E) none of the above
Chapter 21: Electrochemistry: Chemical Change and Electrical Work
63. Which one of the following statements relating to the glass electrode is correct?
A) The glass electrode detects hydrogen gas.
B) The glass of a glass electrode serves to conduct electrons.
C) When pH is measured, only a single electrode, the glass electrode, need be used.
D) The potential of the glass electrode varies linearly with the pH of the solution.
E) None of the above statements is correct.
64. A battery that cannot be recharged is a
A) fuel cell. D) simple battery.
B) primary battery. E) flow battery.
C) secondary battery.
65. Which, if any, of the following metals would be capable of acting as a sacrificial anode
when used with iron pipe? E°Fe = –0.44 V; all E° values refer to the M2+/M half-cell reactions.
A) copper, Cu, E° = 0.15 V
B) cobalt, Co, E° = –0.28 V
C) chromium, Cr, E° = –0.74 V
D) tin, Sn, E° = –0.14 V
E) None of these metals would be capable of acting as a sacrificial anode with iron.
66. Which, if any, of the following metals would not be capable of acting as a sacrificial
anode when used with iron E°Fe = –0.44 V; all E° values refer to the M2+/M half-cell reactions.
A) manganese, Mn, E° = –1.18 V
B) cadmium, Cd, E° = –0.40 V
C) magnesium, Mg, E° = –2.37 V
D) zinc, Zn, E° = –0.76 V
E) All of these metals are capable of acting as sacrificial anodes with iron.
67. What product forms at the cathode during the electrolysis of molten lithium iodide?
A) Li+(l) B) Li(l) C) I–(l) D) I2(g) E) I3–(l)
68. What product forms at the anode during the electrolysis of molten NaBr?
A) Na+(l) B) Na(l) C) Br–(l) D) Br3–(l) E) Br2(g)
Chapter 21: Electrochemistry: Chemical Change and Electrical Work
69. Which of the following elements could be prepared by electrolysis of the aqueous
solution shown?
A) sodium from Na3PO4(aq) D) potassium from KCl(aq)
B) sulfur from K2SO4(aq) E) nitrogen from AgNO3(aq)
C) oxygen from H2SO4(aq)
70. Which of the following elements can be isolated by electrolysis of the aqueous salt
shown?
A) phosphorus from K3PO4(aq) D) fluorine from KF(aq)
B) sodium from NaBr(aq) E) iodine from NaI(aq)
C) aluminum from AlCl3(aq)
71. In the electrolysis of aqueous potassium nitrate using inert electrodes, which one of the
following species is oxidized?
A) potassium ion
B) nitrate ion
C) water
D) oxygen
E) hydronium ion
72. In the electrolysis of aqueous sodium sulfate at electrodes of platinum, predict the
products of the cell reaction.
A) sodium and sulfur D) oxygen and sulfuric acid
B) hydrogen and sulfur E) hydrogen and oxygen
C) oxygen and sulfur
73. Predict the products of the cell reaction when a molten salt mixture of sodium bromide
and calcium fluoride is electrolyzed (spectator ions are not considered to be products).
A) calcium and bromine D) calcium and fluorine
B) sodium and fluorine E) sodium and bromine
C) calcium bromide
74. What mass of silver will be formed when 15.0 A are passed through molten AgCl for
25.0 minutes?
A) 0.419 g B) 6.29 g C) 12.6 g D) 25.2 g E) 33.4 g
Chapter 21: Electrochemistry: Chemical Change and Electrical Work
75. What mass of copper will be deposited when 18.2 A are passed through a CuSO4
solution for 45.0 minutes?
A) 16.2 g B) 33.4 g C) 40.6 g D) 81.3 g E) 163 g
76. A solution is prepared by dissolving 32.0 g of NiSO4 in water. What current would be
needed to deposit all of the nickel in 5.0 hours?
A) 1.1 A B) 2.2 A C) 3.3 A D) 4.4 A E) 5.5 A
77. How many grams of oxygen gas will be produced in the electrolysis of water, for every
gram of hydrogen gas formed?
Reaction: 2H2O(l) → 2H2(g) + O2(g)
A) 31.7 g B) 15.9 g C) 7.94 g D) 3.97 g E) 1.98 g
78. Two cells are connected in series, so that the same current flows through two electrodes
where the following half-reactions occur
Cu2+(aq) + 2e– → Cu(s) and Ag+(aq) + e– → Ag(s)
For every 1.00 g of copper produced in the first process, how many grams of silver will be
produced in the second one?
A) 0.294 g B) 0.588 g C) 0.850 g D) 1.70 g E) 3.40 g
79. A current of 250. A flows for 24.0 hours at an anode where the reaction occurring is
Mn2+(aq) + 2H2O(l) → MnO2(s) + 4H+(aq) + 2e–
What mass of MnO2 is deposited at this anode?
A) 19.5 kg B) 12.9 kg C) 4.87 kg D) 2.43 kg E) none of the above
80. Chromium metal is electroplated from acidic aqueous solutions containing the
dichromate ion, Cr2O72-. What is the minimum time needed to plate out 10.0 g of chromium
metal from such a solution, if the current is 50.0 A?
A) 6.2 minutes D) 24.7 minute
B) 12.4 minutes E) 37.1 minutes
C) 18.6 minutes
Chapter 21: Electrochemistry: Chemical Change and Electrical Work
81. In one or two short sentences each, explain what is meant by the following terms.
a. galvanic or voltaic cell
b. electrolytic cell
c. salt bridge
d. secondary battery or cell
e. primary battery or cell
f. glass electrode
82. A concentration cell is based on the aqueous reaction
Cu2+(1.00 M) → Cu2+(0.0100 M)
The cell consists of copper electrodes dipping into solutions of Cu2+ ions. The anions present are
sulfate ions. Draw a neat diagram to represent this cell, showing and labeling all necessary
components including: anode, cathode, electron flow, cation flow, and anion flow.
Chapter 21: Electrochemistry: Chemical Change and Electrical Work
83. A concentration cell is based on the aqueous reaction
Cu2+(1.00 M) → Cu2+(0.0100 M)
The cell consists of copper electrodes dipping into solutions of Cu2+ ions. The anions present are
sulfate ions. Write the shorthand cell notation for this cell.
84. A much-studied cell in electrochemistry has the following cell notation:
Ag(s) AgCl(s) HCl(aq) H2(g) Pt(s)
Bearing in mind that HCl(aq) consists of H+(aq) and Cl–(aq), and that this solution is in contact
with both electrodes (there is no salt bridge), write down balanced equations for
a. the anode half-reaction.
b. the cathode half-reaction.
c. the cell reaction.
85. A concentration cell is based on the aqueous reaction
Cu2+(1.00 M) → Cu2+(0.0100 M)
Calculate the potential of this cell if it operates at 25.0°C.
86. A galvanic cell is constructed using the two hypothetical half-reactions
A + e– → B E° = 1.50 V
and C + 2e– → D E° = –0.50 V
a. Write down the balanced equation representing the cell reaction.
b. Calculate the standard potential of this cell, E°cell.
c. Calculate G° for the cell reaction.
87. Write down equations representing the anode half-reaction, the cathode half-reaction,
and the overall cell reaction for the lead-acid battery.
Chapter 21: Electrochemistry: Chemical Change and Electrical Work
88. Explain what is meant by a fuel cell. Provide a balanced equation to represent the
reaction in any fuel cell of your choice.
89. A current of 1000. A flows for exactly 1 hour, through a cell in which the following
reaction occurs at one of the electrodes.
Mg2+ + 2e– → Mg
a. Calculate the charge, in coulombs, which passes through the circuit in this time.
b. Calculate the theoretical mass of Mg (magnesium metal) which is produced in this time.
90. Manganese dioxide (MnO2) for use in dry cells is made by electrolyzing solutions of
Mn2+, where the following reaction occurs at the anode:
Mn2+(aq) + 2H2O(l) → MnO2(s) + 4H+(aq) + 2e–
Using a current of 100. amperes (A), how many hours will it take to produce 5.00 kg of MnO2,
according to the above reaction?
91. a. Write a balanced equation to represent the overall reaction you would expect to occur
in the electrolysis of molten KCl.
b. Write a balanced equation to represent the overall reaction you would expect to occur in the
electrolysis of aqueous KCl.
c. Clearly explain why the products of the two processes are not the same.
92. Electrolytic cells utilize electrical energy to drive non-spontaneous redox reactions.
93. For the reaction occurring in a voltaic (galvanic) cell, G > 0.
94. Oxidation occurs at the cathode of a galvanic cell, but at the anode of an electrolytic cell.
Chapter 21: Electrochemistry: Chemical Change and Electrical Work
95. Electrons are produced at the cathode of a voltaic cell.
96. If the electrodes of a voltaic cell are connected with an external wire, electrons will flow
in this wire from the cathode to the anode.
97. In the electrolyte of an electrochemical cell, current is carried by anions moving toward
the anode and cations moving in the opposite direction.
98. In the electrolyte of an electrochemical cell, current is carried by electrons moving from
the anode to the cathode.
99. A salt bridge provides a path for electrons to move between the anode and cathode
compartments of a voltaic cell.
100. In the shorthand notation for cells, a double vertical line is used to separate the reduced
and oxidized forms of a redox couple.
101. In the shorthand notation for cells, a single vertical line represents a salt bridge.
102. A secondary cell (battery) can operate either as a galvanic or an electrolytic cell.
103. A primary battery is one which can be recharged.
104. The lead-acid battery is an example of a secondary battery.
105. In a fuel cell, an external source of electrical power is used to drive a non-spontaneous
reaction in which a fuel is produced.
106. In the absence of oxygen, iron will rust as long as moisture is present.
107. A buried iron pipe can be protected against corrosion by connecting it to a rod of
magnesium.
Chapter 21: Electrochemistry: Chemical Change and Electrical Work
108. A buried iron pipe can be protected against corrosion by connecting it to a rod of copper.