Chapter 11: Electrochemistry

1. Ammonium metavanadate reacts with sulfur dioxide in acidic solution as follows (hydrogen
ions and H 2 O omitted):
xVO 3 – + ySO 2 → xVO2+ + ySO 4 2–
The ratio x : y is
A) 1 : 1
B) 1 : 2
C) 2 : 1
D) 1 : 3
E) 3 : 1
ANS: C PTS: 1 DIF: easy TOP: 11.1
KEY: general chemistry | electrochemistry | half-reaction | Balancing Oxidation-Reduction
Reactions | Half-Reaction Method

2. What is the oxidation state of Mn in MnO 2 ?

A) +2
B) +4
C) +9
D) –1
E) +3
ANS: B PTS: 1 DIF: easy TOP: 11.1
KEY: general chemistry | chemical reactions | types of chemical reactions | oxidation-
reduction reaction | assigning oxidation-number

3. How many electrons are transferred in the following reaction?

2Cr 2 O 7 2– + 14H+ + 6Cl– → 2Cr3+ + 3Cl 2 + 7H 2 O
A) 2
B) 4
C) 6
D) 8
E) none of these
ANS: C PTS: 1 DIF: easy TOP: 11.1
KEY: general chemistry | electrochemistry | half-reaction | Balancing Oxidation-Reduction
Reactions | half-reaction method

4. How many electrons are transferred in the following reaction?

SO 3 2–(aq) + MnO 4 –(aq) → SO 4 2–(aq) + Mn2+(aq)
A) 6
B) 2
C) 10
D) 4
E) 3
ANS: C PTS: 1 DIF: easy TOP: 11.1
KEY: general chemistry | electrochemistry | half-reaction | Balancing Oxidation-Reduction
Reactions | half-reaction method
5. Refer to the galvanic cell below (the contents of each half-cell are written beneath each
compartment).

The standard reduction potentials are as follows:

MnO 4 – + 8H+ + 5e– → Mn2+ + 4H 2 O E° = 1.51 V
Cr 2 O 7 2– +14H+ + 6e– → 2 Cr3+ +7H 2 O E° = 1.33 V

What is the oxidation state of Cr in Cr 2 O 7 2–?

A) +7
B) +6
C) +12
D) –1
E) –2
ANS: B PTS: 1 DIF: easy TOP: 11.1
KEY: general chemistry | chemical reactions | types of chemical reactions | oxidation-
reduction reaction | assigning oxidation-number

6. The following reaction occurs in basic solution:

Ag+ + Cu → Ag + Cu2+
When the equation is balanced, what is the sum of the coefficients?
A) 4
B) 5
C) 6
D) 7
E) 8
ANS: C PTS: 1 DIF: easy TOP: 11.1
KEY: general chemistry | electrochemistry | half-reaction | Balancing Oxidation-Reduction
Reactions | half-reaction method

7. When the equation for the following reaction in basic solution is balanced, what is the sum
of the coefficients?
MnO 2 + HO 2 – → MnO 4 –
A) 11
B) 31
C) 14
D) 9
E) 18
ANS: D PTS: 1 DIF: moderate TOP: 11.1
 KEY: general chemistry | electrochemistry | half-reaction | Balancing Oxidation-Reduction
Reactions | half-reaction method

8. When the equation for the following reaction in basic solution is balanced, what is the sum
of the coefficients?
MnO 4 –(aq) + CN–(aq) → MnO 2 (s) + CNO–(aq)
A) 13
B) 8
C) 10
D) 20
E) 11
ANS: A PTS: 1 DIF: moderate TOP: 11.1
KEY: general chemistry | acids and bases | half-reaction | Balancing Oxidation-Reduction
Reactions | half-reaction method

9. The reaction below occurs in basic solution. In the balanced equation, what is the sum of the
coefficients?
Zn + NO 3 – → Zn(OH) 4 2– + NH 3
A) 12
B) 15
C) 19
D) 23
E) 27
ANS: D PTS: 1 DIF: moderate TOP: 11.1
KEY: general chemistry | electrochemistry | half-reaction | Balancing Oxidation-Reduction
Reactions | half-reaction method

10. A strip of copper is placed in a 1 M solution of copper nitrate, and a strip of silver is placed
in a 1 M solution of silver nitrate. The two metal strips are connected to a voltmeter by
wires, and a salt bridge connects the solutions. The following standard reduction potentials
apply:
Ag+(aq) + e– → Ag(s) E° = +0.80 V
Cu2+(aq) + 2e– → Cu(s) E° = +0.34 V
When the voltmeter is removed and the two electrodes are connected by a wire, which of the
following does not take place?
A) Electrons flow in the external circuit from the copper electrode to the silver
electrode.
B) The silver electrode increases in mass as the cell operates.
C) There is a net general movement of silver ions through the salt bridge to the copper
half-cell.
D) Negative ions pass through the salt bridge from the silver half-cell to the copper
half-cell.
E) Some positive copper ions pass through the salt bridge from the copper half-cell to
the silver half-cell.
ANS: C PTS: 1 DIF: moderate TOP: 11.1 | 11.2
KEY: general chemistry | electrochemistry | voltaic cells | construction of voltaic cells
11. Which metal, Al or Ni, could reduce Zn2+ to Zn(s) if placed in a Zn2+(aq) solution?
Zn2+ + 2e– → Zn E° = –0.76 V
Al + 3e → Al
3+ –
E° = –1.66 V
Ni + 2e → Ni
2+ –
E° = –0.23 V

A) Al
B) Ni
C) Both Al and Ni would work.
D) Neither Al nor Ni would work.
E) This cannot be determined.
ANS: A PTS: 1 DIF: easy TOP: 11.2
KEY: general chemistry | electrochemistry | voltaic cells | standard cell emfs and standard
electrode potentials | strength of oxidizing and reducing agents

12. Which of the following is the best reducing agent?

Cl2 + 2e– → 2Cl– E° = 1.36 V
Mg + 2e → Mg
2+ –
E° = –2.37 V
2H + 2e → H2
+ –
E° = 0.00 V

A) Cl2
B) H2
C) Mg
D) Mg2+
E) Cl-
ANS: C PTS: 1 DIF: easy TOP: 11.2
KEY: general chemistry | electrochemistry | voltaic cells | standard cell emfs and standard
electrode potentials | strength of oxidizing and reducing agents

13. The following two half-reactions take place in a galvanic cell. At standard conditions, what
species are produced at each electrode?
Sn2+ + 2e– → Sn E° = –0.14 V
Cu + 2e → Cu
2+ –
E° = 0.34 V

A) Sn is produced at the anode, and Cu2+ is produced at the cathode.

B) Sn is produced at the anode, and Cu is produced at the cathode.
C) Sn is produced at the cathode, and Cu2+ is produced at the anode.
D) Cu is produced at the cathode, and Sn2+ is produced at the anode.
E) Cu is produced at the anode, and Sn2+ is produced at the cathode.
ANS: D PTS: 1 DIF: easy TOP: 11.2
KEY: general chemistry | electrochemistry | voltaic cells | construction of voltaic cells

14. Which of the following is the strongest oxidizing agent?

MnO4– + 4H+ + 3e– → MnO2 + 2H2O E° = 1.68 V
I2 + 2e → 2I
– –
E° = 0.54 V
Zn + 2e → Zn
2+ –
E° = –0.76 V

A) MnO4–
 B) I2
C) Zn2+
D) Zn
E) MnO2
ANS: A PTS: 1 DIF: easy TOP: 11.2
KEY: general chemistry | electrochemistry | voltaic cells | standard cell emfs and standard
electrode potentials | strength of oxidizing and reducing agents

Reaction E° (volts)
Na+ + e– → Na –2.71
Al3+ + 3e– → Al –1.66
Fe2+ + 2e– → Fe –0.44
Co2+ + 2e– → Co –0.28
Cu2+ + 2e– → Cu +0.34
Ag+ + e– → Ag +0.80
Cl2 + 2e– → 2Cl– +1.36
F2 + 2e– → 2F– +2.87

15. Which of the following would be the best reducing agent?

A) Cl2
B) F2
C) Na
D) Na+
E) F–
ANS: C PTS: 1 DIF: easy TOP: 11.2
KEY: general chemistry | electrochemistry | voltaic cells | standard cell emfs and standard
electrode potentials | strength of oxidizing and reducing agents

16. Silver will spontaneously reduce which of the following?

A) Fe2+
B) Co2+
C) Na+
D) Al3+
E) none of these
ANS: E PTS: 1 DIF: easy TOP: 11.2
KEY: general chemistry | electrochemistry | voltaic cells | standard cell emfs and standard
electrode potentials | strength of oxidizing and reducing agents

17. Determine the standard potential, E°, of a cell that employs the reaction
Co + Ag+ → Co2+ +Ag
A) 1.08 V
B) 1.88 V
C) –1.08 V
D) –0.52 V
E) none of these
ANS: A PTS: 1 DIF: easy TOP: 11.2
 KEY: general chemistry | electrochemistry | voltaic cells | standard cell emfs and standard
electrode potentials | calculating cell emfs

18. The standard potential for the reaction Zn + 2Ag+ → Zn2+ + 2Ag is 1.56 V. Given that the
standard reduction potential for Ag+ + e– → Ag is 0.80 V, determine the standard reduction
potential for Zn2+ + 2e– → Zn.
A) –0.76 V
B) 0.04 V
C) 0.76 V
D) –0.38 V
E) none of these
ANS: A PTS: 1 DIF: easy TOP: 11.2
KEY: general chemistry | electrochemistry | voltaic cells | standard cell emfs and standard
electrode potentials | calculating cell emfs

The following questions refer to a galvanic cell that utilizes the following reaction
(unbalanced):

(AuCl4)–(aq) + Cu(s) → Au(s) + Cl–(aq) + Cu2+(aq)

19. Given the following information, determine the standard cell potential.
Species Standard Reduction Potential (V)
Au3+(aq) 1.498
2+
Cu (aq) 0.339

A) 1.159 V
B) 1.837 V
C) 1.979 V
D) 1.462 V
E) 2.102 V
ANS: A PTS: 1 DIF: easy TOP: 11.2
KEY: general chemistry | electrochemistry | voltaic cells | standard cell emfs and standard
electrode potentials | calculating cell emfs

20. Determine the number of electrons transferred during the reaction (when balanced).
A) 2
B) 3
C) 4
D) 6
E) 9
ANS: D PTS: 1 DIF: easy TOP: 11.2
KEY: general chemistry | electrochemistry | half-reaction | Balancing Oxidation-Reduction
Reactions | half-reaction method

21. Choose the correct statement(s) given the following information:

Fe3+(aq) + e– → Fe2+(aq) E° = 0.77 V

 Fe(CN)63– + e– → Fe(CN)64– E° = 0.36 V

I. Fe2+(aq) is more likely to be oxidized than Fe2+ complexed to CN–.

II. Fe3+(aq) is more likely to be reduced than Fe3+ complexed to CN–.
III. Complexation of Fe ions with CN– has no effect on their tendencies to become
oxidized or reduced.

A) I only
B) II only
C) I and II
D) III only
E) None of these is true.
ANS: B PTS: 1 DIF: easy TOP: 11.2
KEY: general chemistry | electrochemistry | voltaic cells | standard cell emfs and standard
electrode potentials | strength of oxidizing and reducing agents

Refer to the galvanic cell below (the contents of each half-cell are written beneath each
compartment).

The standard reduction potentials are as follows:

MnO4– + 8H+ + 5e– → Mn2+ + 4H2O E° = 1.51 V
Cr2O72– +14H+ + 6e– → 2Cr3+ +7H2O E° = 1.33 V

22. What is the value of E°cell?

A) –0.18 V
B) 2.84 V
C) 0.18 V
D) 1.79 V
E) 2.29 V
ANS: C PTS: 1 DIF: easy TOP: 11.2
KEY: general chemistry | electrochemistry | voltaic cells | standard cell emfs and standard
electrode potentials | calculating cell emfs

23. In the balanced cell reaction, what is the stoichiometric coefficient for H+?
A) 5
B) 6
C) 30
D) 22
E) 2
 ANS: D PTS: 1 DIF: easy TOP: 11.2
KEY: general chemistry | electrochemistry | half-reaction | Balancing Oxidation-Reduction
Reactions | half-reaction method

Consider an electrochemical cell with a copper electrode immersed in 1.0 M Cu2+ and a
silver electrode immersed in 1.0 M Ag+.
Cu2+ + 2e– → Cu E° = 0.34 V
Ag + e → Ag
+ –
E° = 0.80 V

24. Calculate E° for this cell.

A) 1.48 V
B) 1.26 V
C) 1.14 V
D) 0.46 V
E) none of these
ANS: D PTS: 1 DIF: easy TOP: 11.2
KEY: general chemistry | electrochemistry | voltaic cells | standard cell emfs and standard
electrode potentials | calculating cell emfs

25. Which of the electrodes is the anode?

A) the copper electrode
B) the silver electrode
ANS: A PTS: 1 DIF: easy TOP: 11.2
KEY: general chemistry | electrochemistry | voltaic cells | construction of voltaic cells

Consider the galvanic cell shown below (the contents of each half-cell are written beneath
each compartment).

The standard reduction potentials are as follows:

Cr3+ + 3e– → Cr(s) E° = –0.73 V
Br2(aq) + 2e → 2Br
– –
E° = +1.09 V

26. What is E° for this cell?

A) 1.82 V
B) 0.36 V
C) 4.75 V
D) 1.79 V
E) 4.40 V
ANS: A PTS: 1 DIF: easy TOP: 11.2
 KEY: general chemistry | electrochemistry | voltaic cells | standard cell emfs and standard
electrode potentials | calculating cell emfs

27. Which of the following statements about this cell is false?

A) This is a galvanic cell.
B) Electrons flow from the Pt electrode to the Cr electrode.
C) Reduction occurs at the Pt electrode.
D) The cell is not at standard conditions.
E) To complete the circuit, cations migrate into the left half-cell and anions migrate
into the right half-cell from the salt bridge.
ANS: B PTS: 1 DIF: moderate TOP: 11.2
KEY: general chemistry | electrochemistry | voltaic cells | construction of voltaic cells

The reaction Cr(s) + NO3–(aq) → Cr3+(aq) + NO(g) takes place in acidic solution.

28. How many electrons are transferred per mole of NO(g)?

A) 1
B) 2
C) 3
D) 4
E) 5
ANS: C PTS: 1 DIF: easy TOP: 11.2
KEY: general chemistry | electrochemistry | half-reaction | Balancing Oxidation-Reduction
Reactions | half-reaction method

29. What is the coefficient of H2O(l) in the balanced equation?

A) 1
B) 2
C) 3
D) 4
E) 5
ANS: B PTS: 1 DIF: moderate TOP: 11.2
KEY: general chemistry | electrochemistry | half-reaction | Balancing Oxidation-Reduction
Reactions | half-reaction method

Refer to the galvanic cell below (the contents of each half-cell are written beneath each
compartment).

The standard reduction potentials are as follows:

 MnO4– + 8H+ + 5e– → Mn2+ + 4H2O E° = 1.51 V
Cr2O72– +14H+ + 6e– → 2 Cr3+ +7H2O E° = 1.33 V

30. When current is allowed to flow, which species is oxidized?

A) Cr2O72–
B) Cr3+
C) MnO4–
D) Mn2+
E) H+
ANS: B PTS: 1 DIF: easy TOP: 11.2 | 11.4
KEY: general chemistry | electrochemistry | voltaic cells | standard cell emfs and standard
electrode potentials | strength of oxidizing and reducing agents

31. When current is allowed to flow, which species is reduced?

A) Cr2O72–
B) Cr3+
C) MnO4–
D) Mn2+
E) H+
ANS: C PTS: 1 DIF: easy TOP: 11.2 | 11.4
KEY: general chemistry | electrochemistry | voltaic cells | standard cell emfs and standard
electrode potentials | strength of oxidizing and reducing agents

32. In which direction do electrons flow in the external circuit?

A) left to right
B) right to left
C) No current flows; the cell is at equilibrium.
ANS: B PTS: 1 DIF: easy TOP: 11.2 | 11.4
KEY: general chemistry | electrochemistry | voltaic cells | construction of voltaic cells

33. How many electrons are transferred in the balanced reaction (that is, what will be the value
of n in the Nernst equation)?
A) 5
B) 6
C) 30
D) 22
E) 2
ANS: C PTS: 1 DIF: easy TOP: 11.2 | 11.4
KEY: general chemistry | electrochemistry | half-reaction | Balancing Oxidation-Reduction
Reactions | half-reaction method

34. In a common car battery, six identical cells each carry out the following reaction:
Pb + PbO2 + 2HSO4– + 2H+ → 2PbSO4 + 2H2O

For such a cell, E° is 2.04 V. Calculate ∆G° at 25°C.

A) –787 kJ
 B) –98 kJ
C) –394 kJ
D) –197 kJ
E) –0.121 kJ
ANS: C PTS: 1 DIF: moderate TOP: 11.3
KEY: general chemistry | electrochemistry | voltaic cells | electromotive force

35. The reduction potentials for Au3+ and Ni2+ are as follows:
Au3+ + 3e– → Au E° = +1.50 V
Ni2+ + 2e– → Ni E° = –0.23 V
Calculate ∆G° (at 25°C) for the reaction
2Au3+ + 3Ni → 3Ni2+ + 2Au
A) –5.00 × 102 kJ
B) +5.00 × 102 kJ
C) –2140 kJ
D) +1.00 × 103 kJ
E) –1.00 × 103 kJ
ANS: E PTS: 1 DIF: moderate TOP: 11.3
KEY: general chemistry | electrochemistry | voltaic cells | electromotive force

36. Consider the following reduction potentials:

Cu2+ + 2e– → Cu E° = +0.34 V
Pb2+ + 2e– → Pb E° = –0.13 V
For a galvanic cell employing the Cu, Cu2+ and Pb, Pb2+ couples, calculate the maximum
amount of work that would accompany the reaction of 1 mol of lead under standard
conditions.
A) –40.5 kJ
B) –45.3 kJ
C) –90.7 kJ
D) No work can be done. The system is at equilibrium.
E) none of these
ANS: C PTS: 1 DIF: moderate TOP: 11.3
KEY: general chemistry | electrochemistry | voltaic cells | electromotive force

37. The standard free energies of formation of several aqueous species are
kJ/mol
+
H (aq) 0
H2O(l) –237
CH3OH(aq) –163
HCOOH(aq) –351
–
e 0
What is the standard reduction potential of methanoic acid in aqueous solution (that is, for
HCOOH + 4H+ + 4e– → CH3OH + H2O)?
A) +0.13 V
B) +0.17 V
C) +0.25 V
D) –0.13 V
 E) –0.25 V
ANS: A PTS: 1 DIF: moderate TOP: 11.3
KEY: general chemistry | electrochemistry | voltaic cells | electromotive force

38. In which of the following cases can E° be equal to zero?

I. In any cell at equilibrium
II. In a concentration cell
III. E° can never be equal to zero.

A) I only
B) II only
C) III
D) I and II
ANS: B PTS: 1 DIF: easy TOP: 11.3 | 11.4
KEY: general chemistry | electrochemistry | voltaic cells | dependence of emf on
concentration

39. In which of the following cases must E be equal to zero?

I. In any cell at equilibrium
II. In a concentration cell
III. E° can never be equal to zero.

A) I only
B) II only
C) III
D) I and II
ANS: A PTS: 1 DIF: easy TOP: 11.3 | 11.4
KEY: general chemistry | electrochemistry | voltaic cells | dependence of emf on
concentration | Nernst equation

40. Consider the hydrogen–oxygen fuel cell where

H2(g) + O2(g) H2O(l) ∆G° = –237.18 kJ/mol H2
Which of the following statements is(are) true?
I. At standard conditions, the maximum work the fuel cell could do on the surroundings
is 237.18 kJ/mol.
II. In the real world, the actual amount of useful work the cell can do is less than 237.18
kJ.
III. More energy is dissipated as waste heat in the fuel cell than in the reversible pathway.

A) I
B) II
C) III
D) I, II, and III
E) None of the statements is true.
ANS: D PTS: 1 DIF: easy TOP: 11.3 | 11.5
KEY: general chemistry | electrochemistry | voltaic cells | commercial voltaic cells
41. A fuel cell designed to react grain alcohol with oxygen has the following net reaction:
C2H5OH(l) + 3O2(g) → 2CO2(g) + 3H2O(l)
The maximum work that 1 mol of alcohol can yield by this process is 1320 kJ. What is the
theoretical maximum voltage this cell can achieve?
A) 0.760 V
B) 1.14 V
C) 2.01 V
D) 2.28 V
E) 13.7 V
ANS: B PTS: 1 DIF: moderate TOP: 11.3 | 11.5
KEY: general chemistry | electrochemistry | voltaic cells | electromotive force

42. An excess of finely divided iron is stirred up with a solution that contains Cu2+ ion, and the
system is allowed to come to equilibrium. The solid materials are then filtered off, and
electrodes of solid copper and solid iron are inserted into the remaining solution. What
potential develops between these two electrodes at 25°C?
A) 0
B) –0.78 V
C) 0.592 V
D) 0.296 V
E) Not enough information is given.
ANS: A PTS: 1 DIF: easy TOP: 11.4
KEY: general chemistry | electrochemistry | voltaic cells | equilibrium constants from emfs

You make a cell with a copper electrode in a solution of copper nitrate and a silver electrode
in a solution of silver nitrate.
Cu2+ + 2e– → Cu E° = 0.34 V
Ag+ + e– → Ag E° = 0.80 V

43. If you could increase the concentration of Ag+, which of the following would be true about
the cell potential?
A) It would increase.
B) It would decrease.
C) It would remain constant.
D) This cannot be determined.
ANS: A PTS: 1 DIF: easy TOP: 11.4
KEY: general chemistry | electrochemistry | voltaic cells | dependence of emf on
concentration

44. If you could increase the concentration of Cu2+, which of the following would be true about
the cell potential?
A) It would increase.
B) It would decrease.
C) It would remain constant.
D) This cannot be determined.
ANS: B PTS: 1 DIF: easy TOP: 11.4
 KEY: general chemistry | electrochemistry | voltaic cells | dependence of emf on
concentration

45. A cell is set up with copper and lead electrodes in contact with CuSO4(aq) and
Pb(NO3)2(aq), respectively, at 25°C. The standard reduction potentials are
Pb2+ + 2e– → Pb E° = –0.13 V
Cu2+ + 2e– → Cu E° = +0.34 V
If sulfuric acid is added to the Pb(NO3)2 solution, forming a precipitate of PbSO4, the cell
potential
A) increases.
B) decreases.
C) is unchanged.
D) It is impossible to tell what will happen.
ANS: A PTS: 1 DIF: easy TOP: 11.4
KEY: general chemistry | electrochemistry | voltaic cells | dependence of emf on
concentration

46. A concentration cell is constructed using two Ni electrodes with Ni2+ concentrations of
1.70 M and 4.40 × 10–4 M in the two half-cells. The reduction potential of Ni2+ is –0.23 V.
Calculate the potential of the cell at 25°C.
A) +0.336 V
B) –0.212 V
C) –0.244 V
D) +0.106 V
E) –0.0180 V
ANS: D PTS: 1 DIF: moderate TOP: 11.4
KEY: general chemistry | electrochemistry | voltaic cells | dependence of emf on
concentration | Nernst equation

47. The standard potential for the reaction A(s) + B3+(aq) A3+(aq) + B(s) is 0.51 V. What
is the equilibrium constant K for this reaction at 25°C?
A) 6.3 × 1025
B) 1.6 × 10–26
C) 17
D) –17
E) 0.015
ANS: E PTS: 1 DIF: moderate TOP: 11.4
KEY: general chemistry | electrochemistry | voltaic cells | equilibrium constants from emfs

48. A concentration cell is constructed using two Co electrodes with Co2+ concentrations of
0.14 M and 1.8 × 10–5 M in the two half-cells. Determine the standard reduction potential of
Co2+, given that the potential of the cell at 25°C is 0.115 V.
A) 0 V
B) +0.115 V
C) –0.115 V
D) +0.230 V
E) This cannot be determined with the information given.
 ANS: E PTS: 1 DIF: moderate TOP: 11.4
KEY: general chemistry | electrochemistry | voltaic cells | dependence of emf on
concentration | Nernst equation

49. Consider the galvanic cell shown below (the contents of each half-cell are written beneath
each compartment).

The standard reduction potentials are as follows:

Cr3+ + 3e– → Cr(s) E° = –0.73 V
Br2(aq) + 2e → 2Br
– –
E° = +1.09 V

What is the value of E for this cell at 25°C?

A) 2.21 V
B) 1.76 V
C) 2.12 V
D) 1.88 V
E) 0.59 V
ANS: D PTS: 1 DIF: moderate TOP: 11.4
KEY: general chemistry | electrochemistry | voltaic cells | dependence of emf on
concentration | Nernst equation

50. Consider an electrochemical cell with a copper electrode immersed in 1.0 M Cu2+ and a
silver electrode immersed in 1.0 M Ag+.
Cu2+ + 2e– → Cu E° = 0.34 V
Ag+ + e– → Ag E° = 0.80 V

If [Cu2+]0 is 0.0010 M and [Ag+]0 is 0.24 M, calculate E.

A) 0.53 V
B) 0.62 V
C) 0.51 V
D) 0.41 V
E) 0.39 V
ANS: C PTS: 1 DIF: moderate TOP: 11.4
KEY: general chemistry | electrochemistry | voltaic cells | dependence of emf on
concentration | Nernst equation

51. For a reaction in a voltaic cell, both ∆H° and ∆S° are positive. Which of the following
statements is true?
A) E°cell will increase with an increase in temperature.
 B) E°cell will decrease with an increase in temperature.
C) E°cell will not change when the temperature increases.
D) ∆G° > 0 for all temperatures.
E) None of the above statements is true.
ANS: A PTS: 1 DIF: moderate TOP: 11.4
KEY: general chemistry | electrochemistry | voltaic cells | electromotive force

52. Determine the equilibrium constant at 25°C for the reaction

Zn + Ni2+ → Zn2+ + Ni
Zn2+ + 2e– → Zn E° = –0.76 V
Ni + 2e → Ni
-2+ –
E° = –0.23 V

A) 8.6 × 1017
B) 9.3 × 108
C) 4.1 × 103
D) 6.4 × 10–6
E) none of these
ANS: A PTS: 1 DIF: moderate TOP: 11.4
KEY: general chemistry | electrochemistry | voltaic cells | equilibrium constants from emfs

53. If a reducing agent M reacts with an oxidizing agent Q2+ to give M2+ and Q, and the
equilibrium constant for the reaction is 1.0, then what is the E° value for the oxidation–
reduction reaction at 25°C?
A) 0.0 V
B) –1.0 V
C) 1.0 V
D) 0.030 V
E) 0.41 V
ANS: A PTS: 1 DIF: moderate TOP: 11.4
KEY: general chemistry | electrochemistry | voltaic cells

54. A galvanic cell is constructed with copper electrodes and Cu2+ in each compartment. In one
compartment, [Cu2+] = 3.4 × 10–3 M, and in the other compartment, [Cu2+] = 2.5 M.
Calculate the potential for this cell at 25°C. The standard reduction potential for Cu2+ is
+0.34 V.
A) 0.42 V
B) –0.42 V
C) 0.085 V
D) –0.085 V
E) 0.76 V
ANS: C PTS: 1 DIF: moderate TOP: 11.4
KEY: general chemistry | electrochemistry | voltaic cells | dependence of emf on
concentration | Nernst equation
55. An excess of finely divided iron is stirred up with a solution that contains Cu2+ ion, and the
system is allowed to come to equilibrium. The solid materials are then filtered off, and
electrodes of solid copper and solid iron are inserted into the remaining solution. What is the
value of the ratio [Fe2+]/[Cu2+] at 25°C?
A) 1
B) 0
C) 2.5 × 1026
D) 4.0 × 10–27
E) none of these
ANS: C PTS: 1 DIF: moderate TOP: 11.4
KEY: general chemistry | electrochemistry | voltaic cells | equilibrium constants from emfs

Consider an electrochemical cell that has a nickel electrode immersed in 1.0 M Ni2+ and a
platinum electrode immersed in 0.010 M Co2+.
Co2+ + 2e– → Co E° = –0.28 V
Ni + 2e → Ni
2+ –
E° = –0.23 V

56. Calculate E for this cell.

A) 0.05 V
B) 0.11 V
C) 0.51 V
D) 0.56 V
E) none of these
ANS: B PTS: 1 DIF: moderate TOP: 11.4
KEY: general chemistry | electrochemistry | voltaic cells | dependence of emf on
concentration | Nernst equation

57. Calculate the concentration of Ni2+ if the cell is allowed to run to equilibrium at 25°C.
A) 1.1 M
B) 1.0 M
C) 0.020 M
D) 1.9 × 10–4 M
E) none of these
ANS: C PTS: 1 DIF: difficult TOP: 11.4
KEY: general chemistry | electrochemistry | voltaic cells | dependence of emf on
concentration | Nernst equation

Refer to the galvanic cell below (the contents of each half-cell are written beneath each
compartment).
 The standard reduction potentials are as follows:
MnO4– + 8H+ + 5e– → Mn2+ + 4H2O E° = 1.51 V
Cr2O72– +14H+ + 6e– → 2 Cr3+ +7H2O E° = 1.33 V

58. What is the value of Q, the reaction quotient, for this cell reaction?
A) 6.7 × 1040
B) 1.5 × 10–41
C) 1.5 × 10–4
D) 6.7 × 103
ANS: B PTS: 1 DIF: moderate TOP: 11.4
KEY: general chemistry | electrochemistry | voltaic cells | dependence of emf on
concentration

59. What is the cell potential at 25°C as read on the digital voltmeter?
A) 0.18 V
B) 2.58 V
C) 0.10 V
D) 0.59 V
E) 0.26 V
ANS: E PTS: 1 DIF: difficult TOP: 11.4
KEY: general chemistry | electrochemistry | voltaic cells | dependence of emf on
concentration | Nernst equation

60. What is the value of the equilibrium constant at 25°C for the net spontaneous cell reaction?
A) 7.3 × 10–11
B) 4.3 × 10–92
C) 91
D) 1.1 × 103
E) 2.3 × 1091
ANS: E PTS: 1 DIF: difficult TOP: 11.4
KEY: general chemistry | electrochemistry | voltaic cells | equilibrium constants from emfs

61. Calculate E at 25°C for this cell, given the following data:

Ag+ + e– → Ag(s) E° = 0.80 V

Ni2+ + 2e– → Ni(s) E° = –0.23 V
Ksp for AgCl = 1.6 × 10–10
 A) 0.83 V
B) 0.54 V
C) 1.01 V
D) 2.98 V
E) This cannot be determined from the data given
ANS: B PTS: 1 DIF: difficult TOP: 11.4
KEY: general chemistry | electrochemistry | voltaic cells | dependence of emf on
concentration | Nernst equation

62. Use the following data to calculate the Ksp value at 25°C for PbSO4(s).
E°
PbO2 + 4H+ + SO42– + 2e– → PbSO4(s) + 2H2O +1.69
PbO2 + 4H+ + 2e– → Pb2+ + 2H2O +1.46

A) 2.57 × 10105
B) 3.89 × 10–105
C) 5.9 × 107
D) 1.7 × 10–8
E) None of these is within 5% of the correct answer.
ANS: D PTS: 1 DIF: difficult TOP: 11.4
KEY: general chemistry | electrochemistry | voltaic cells | dependence of emf on
concentration | Nernst equation

63. Calculate the solubility product of silver iodide at 25°C, given the following data:
E° (V)
AgI(s) + e → Ag(s) + I
– –
–0.15
I2(s) + 2e → 2I
– –
+0.54
Ag + e → Ag(s)
+ –
+0.80

A) 2.9 × 10–3
B) 1.9 × 10–4
C) 2.1 × 10–12
D) 8.4 × 10–17
E) 3.5 × 10–20
ANS: D PTS: 1 DIF: difficult TOP: 11.4
KEY: general chemistry | electrochemistry | voltaic cells | dependence of emf on
concentration | Nernst equation

64. Why is aluminum protected from corrosion? (Note: The standard reduction potential for
Al3+ is –1.66 V.)
A) Aluminum forms a protective oxide coating.
B) The oxidation of aluminum is not a favored process, as seen by the standard
reduction potential for Al3+.
C) Oxygen and aluminum have no affinity for one another.
D) At least two of these are correct.
E) Aluminum is not protected from corrosion.
 ANS: A PTS: 1 DIF: easy TOP: 11.6
KEY: general chemistry | electrochemistry | voltaic cells | corrosion

65. If a constant current of 4.9 amperes is passed through a cell containing Cr3+ for 2.3 hour,
how many grams of Cr will plate out onto the cathode? (The atomic mass of Cr is
51.996 g/mol.)
A) 22 g
B) 0.12 g
C) 7.3 g
D) 66 g
E) 1.1 g
ANS: C PTS: 1 DIF: easy TOP: 11.7
KEY: general chemistry | electrochemistry | electrolytic cells | aqueous electrolysis |
electroplating of metals

66. In a common car battery, six identical cells each carry out the following reaction:
Pb + PbO2 + 2HSO4– + 2H+ → 2PbSO4 + 2H2O

Suppose that to start a car on a cold morning, 126 amperes is drawn for 13.0 seconds from
such a cell. How many grams of Pb are consumed? (The atomic mass of Pb is
207.19 g/mol.)
A) 7.03 g
B) 1.76 g
C) 3.52 g
D) 0.01041 g
E) 0.569 g
ANS: B PTS: 1 DIF: moderate TOP: 11.7
KEY: general chemistry | electrochemistry | voltaic cells

An antique automobile bumper is to be chrome plated. The bumper, which is dipped into an
acidic Cr2O72– solution, serves as a cathode of an electrolytic cell. The atomic mass of Cr is
51.996; 1 faraday = 96,485 coulombs.

67. If oxidation of H2O occurs at the anode, how many moles of oxygen gas will evolve for
every 2.98 × 102 g of Cr(s) deposited?
A) 8.60 mol
B) 5.73 mol
C) 3.82 mol
D) 7.64 mol
E) 2.87 mol
ANS: A PTS: 1 DIF: moderate TOP: 11.7
KEY: general chemistry | electrochemistry | electrolytic cells | aqueous electrolysis |
electroplating of metals

68. If the current is 34.4 amperes, how long will it take to deposit 1.58 × 102 g of Cr(s) onto the
bumper?
A) 28.4 h
 B) 14.2 h
C) 23.7 min
D) 23.4 h
E) 9.5 h
ANS: B PTS: 1 DIF: moderate TOP: 11.7
KEY: general chemistry | electrochemistry | electrolytic cells | aqueous electrolysis |
electroplating of metals

69. Copper is electroplated from an aqueous CuSO4 solution. A constant current of 5.42 amp is
applied by an external power supply. How long will it take to deposit 4.06 × 102 g of Cu?
The atomic mass of copper is 63.546 g/mol.
A) 63.2 h
B) 31.6 h
C) 2.58 h
D) 126.4 h
E) 94.8 h
ANS: A PTS: 1 DIF: moderate TOP: 11.7
KEY: general chemistry | electrochemistry | electrolytic cells | aqueous electrolysis |
electroplating of metals

70. What quantity of charge is required to reduce 31.2 g of CrCl3 to chromium metal?
(1 faraday = 96,485 coulombs)
A) 1.90 × 104 C
B) 5.70 × 104 C
C) 1.74 × 105 C
D) 3.80 × 104 C
E) 7.60 × 104 C
ANS: B PTS: 1 DIF: moderate TOP: 11.7
KEY: general chemistry | electrochemistry | electrolytic cells | aqueous electrolysis |
electroplating of metals

71. Electrolysis of a molten salt with the formula MCl, using a current of 3.86 amp for
16.2 min, deposits 1.52 g of metal. Identify the metal. (1 faraday = 96,485 coulombs)
A) Li
B) Na
C) K
D) Rb
E) Ca
ANS: C PTS: 1 DIF: moderate TOP: 11.7
KEY: general chemistry | electrochemistry | electrolytic cells | electrolysis of molten salts

72. If an electrolysis plant operates its electrolytic cells at a total current of 1.0 × 106 amp, how
long will it take to produce one metric ton (one million grams) of Mg(s) from seawater
containing Mg2+? (1 faraday = 96,485 coulombs)
A) 2.2 h
B) 2.4 days
 C) 55 min
D) 3.7 h
E) 1 year
ANS: A PTS: 1 DIF: moderate TOP: 11.7
KEY: general chemistry | electrochemistry | electrolytic cells | aqueous electrolysis

73. Nickel is electroplated from a NiSO4 solution. A constant current of 4.74 amp is applied by
an external power supply. How long will it take to deposit 1.66 × 102 g of Ni? (The atomic
mass of Ni is 58.69 g/mol.)
A) 32.0 h
B) 16.0 h
C) 1.00 h
D) 48.0 h
E) 19.2 min
ANS: A PTS: 1 DIF: moderate TOP: 11.7
KEY: general chemistry | electrochemistry | electrolytic cells | aqueous electrolysis |
electroplating of metals

74. A solution of MnO42– is electrolytically reduced to Mn3+. A current of 8.14 amp is passed
through the solution for 12.4 minutes. What is the number of moles of Mn3+ produced in this
process? (1 faraday = 96,485 coulombs)
A) 0.0628 mol
B) 0.000349 mol
C) 0.0157 mol
D) 0.0209 mol
E) 0.0314 mol
ANS: D PTS: 1 DIF: moderate TOP: 11.7
KEY: general chemistry | electrochemistry | electrolytic cells | aqueous electrolysis |
electroplating of metals

75. How many seconds would it take to deposit 21.40 g of Ag (atomic mass = 107.87) from a
solution of AgNO3 using a current of 10.00 amp?
A) 9649 s
B) 4825 s
C) 3828 s
D) 1914 s
E) none of these
ANS: D PTS: 1 DIF: moderate TOP: 11.7
KEY: general chemistry | electrochemistry | electrolytic cells | aqueous electrolysis

76. Gold (atomic mass = 197 g/mol) is plated from a solution of chlorauric acid, HAuCl4; it
deposits on the cathode. Calculate the time it takes to deposit 0.65 g of gold, passing a
current of 0.14 amperes. (1 faraday = 96,485 coulombs)
A) 0.63 h
B) 1.9 h
C) 2.5 h
D) 0.0025 days
 E) 1.3 h
ANS: B PTS: 1 DIF: moderate TOP: 11.7
KEY: general chemistry | electrochemistry | electrolytic cells | aqueous electrolysis |
electroplating of metals

77. An electrolytic cell process involves plating Zr(s) from a solution containing Zr4+. If 5.80
amp is run through this mixture for 1.86 h, what mass of Zr is plated?
A) 36.7 g
B) 0.153 g
C) 0.101 g
D) 9.18 g
E) none of these
ANS: D PTS: 1 DIF: moderate TOP: 11.7
KEY: general chemistry | electrochemistry | electrolytic cells | aqueous electrolysis

78. Gold is produced electrochemically from an aqueous solution of Au(CN)2– containing an

excess of CN–. Gold metal and oxygen gas are produced at the electrodes. How many moles
of O2 will be produced during the production of 1.00 mol of gold?
A) 0.25 mol
B) 0.50 mol
C) 1.00 mol
D) 3.56 mol
E) 4.00 mol
ANS: A PTS: 1 DIF: difficult TOP: 11.7
KEY: general chemistry | electrochemistry | electrolytic cells | aqueous electrolysis

79. Fe2+ + 2e– → Fe(s) E° = –0.440 V

2H+ + 2e– → H2(g) E° = 0.000 V
In a galvanic cell, the iron compartment contains an iron electrode, and [Fe2+] = 1.00 × 10–3
M. The hydrogen compartment contains a platinum electrode (PH2 = 1.00 atm) and a weak
acid HA at an initial concentration of 1.00 M. If the observed cell potential is 0.333 V at
25°C, calculate Ka for the weak acid HA at 25°C.

ANS:
Ka = 2.43 × 10–7

PTS: 1 DIF: difficult TOP: 11.4

KEY: general chemistry | electrochemistry | voltaic cells | dependence of emf on
concentration | Nernst equation

80. Ag+ + e– → Ag(s) E° = 0.80 V

Cu2+ + 2e– → Cu(s) E° = 0.34 V
In a galvanic cell, the silver compartment contains a silver electrode and excess AgCl(s)
(Ksp = 1.6 × 10–10). The copper compartment contains a copper electrode, and [Cu2+] = 2.0
M.

A. Calculate the potential for this cell at 25°C.

 B. Assuming 1.0 L of 2.0 M Cu2+ in the copper compartment, calculate how many moles of
NH3 would have to be added to establish the cell potential at 0.52 V at 25°C (assume no
volume change on addition of NH3).

Cu2+ + 4NH3 Cu(NH3)42+ Kf = 1.0 × 1013

ANS:
A. 0.16 V; B. 8.6 mol of NH3

PTS: 1 DIF: difficult TOP: 11.4

KEY: general chemistry | electrochemistry | voltaic cells | dependence of emf on
concentration | Nernst equation

81. Cu2+ + 2e– → Cu(s) E° = 0.34 V

V2+ + 2e– → V(s) E° = -1.20 V

In a galvanic cell, the copper compartment contains a copper electrode and [Cu2+] = 1.00 M.
The vanadium compartment contains a vanadium electrode and V2+ at an unknown
concentration. The compartment containing the vanadium (1.00 L of solution) was titrated
with 0.0800 M H2EDTA2–, resulting in the reaction:

H2EDTA2– + V2+ VEDTA2– + 2H+

The potential of the cell was monitored to determine the stoichiometric point for the
process, which occurred when 500.0 mL of H2EDTA2– solution was added. At the
stoichiometric point, Ecell was 1.98 V. The solution was buffered at pH 10.00.

A. Calculate Ecell before titration was carried out.

B. Calculate the equilibrium constant K for the titration reaction.
C. Calculate Ecell at the halfway point of titration.

ANS:
A. 1.58 V; B. 1.4; C. 1.59 V

PTS: 1 DIF: difficult TOP: 11.4

KEY: general chemistry | electrochemistry | voltaic cells | dependence of emf on
concentration | Nernst equation

82. In the following equation

DG = DG° + RT nK

A) DG° = 0
B) DG = 0
C) K=1
D) DG = DG°
E) All of the above
 ANS: B PTS: 1

83. Which of the following would be a good sacrificial electrode to protect an iron boat from corrosion?

Half Reaction E °(V)

-0.23
-0.44
-0.50
-1.66

A) Nickel
B) Chromium
C) Aluminum
D) B and C
E) More information needed to answer the question

ANS: A PTS: 1

84. When a battery dies, which of the following is true?

A) DG° = 0
B) DG = 0
C) E° = 0
D) Q=1
E) A and C
F) B and D

ANS: B PTS: 1

85. The SI unit for current is the __________ and the SI unit for charge is the _________.

A) volt, coloumb
B) coulomb, ampere
C) coulomb, faraday
D) ampere, coulomb
E) ampere, faraday
F) ampere, volt

ANS: D PTS: 1