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Chapter 19 - Electrochemistry
1. When balancing oxidation–reduction reactions in acidic solution by the half-reaction
method, the addition of the reactant H+ is required to balance the product
A)
O2.
B)
OH–.
C)
O.
D)
H2O.
E)
none of these
2. When the following oxidation–reduction reaction in acidic solution is balanced, what is the
lowest whole-number coefficient for Rb+(aq)?
Rb(s) + Cu2+(aq) → Rb+(aq) + Cu(s)
A)
5
B)
4
C)
1
D)
3
E)
2
3. When the following oxidation–reduction reaction in acidic solution is balanced, what is the
lowest whole-number coefficient for H+, and on which side of the balanced equation should
it appear?
MnO4–(aq) + Br–(aq) → Mn2+(aq) + Br2(l)
A)
1, reactant side
B)
2, product side
C)
8, reactant side
D)
16, reactant side
E)
4, product side
4. Balance the following half-reaction occurring in acidic solution.
NO3–(aq) → NO(g)
A)
NO3–(aq) + 4H+(aq) + 3e− → NO(g) + 2H2O(l)
B)
NO3–(aq) + 2H2O(l) + 3e− → NO(g) + 4H+(aq)
C)
NO3–(aq) + 4H+(aq) → NO(g) + 2H2O(l) + 3e−
D)
NO3–(aq) + 3e− → NO(g) + 4H+(aq) + 2H2O(l)
E)
NO3–(aq) + 4H+(aq)→ NO(g) + 2H2O(l)
5. When the following oxidation–reduction reaction in acidic solution is balanced, what is the
lowest whole-number coefficient for H+, and on which side of the balanced equation should
it appear?
S2O82–(aq) + NO(g) → SO42–(aq) + NO3–(aq)
A)
4, product side
B)
8, reactant side
C)
12, reactant side
D)
8, product side
E)
4, reactant side
6. Balance the following half-reaction occurring in basic solution.
MnO2(s) → Mn(OH)2(s)
A)
MnO2(s) + 2H2O(l) + 2e−→ Mn(OH)2(s)+ 2OH−(aq)
B)
MnO2(s) + 2H2O(l) + 4e−→ Mn(OH)2(s)+ (OH)2−(aq)
C)
MnO2(s) + H22+(aq) + 2e−→ Mn(OH)2(s)
D)
MnO2(s) + H2(g) → Mn(OH)2(s) + 2e−
E)
MnO2(s) + 2H2O(l) → Mn(OH)2(s)+ 2OH−(aq)
7. Balance the following oxidation-reduction occurring in acidic solution.
MnO4–(aq) + Co2+(aq)→ Mn2+(aq) + Co3+(aq)
A)
MnO4–(aq) + 8H+(aq) + 5Co2+(aq) → Mn2+(aq) + 4H2O(aq) + 5Co3+(aq)
B)
MnO4–(aq) + 8H+(aq) + Co2+(aq) → Mn2+(aq) + 4H2O(aq) + Co3+(aq)
C)
MnO4–(aq) + 4H2(g) + 5Co2+(aq)→ Mn2+(aq) + 4H2O(aq) + 5Co3+(aq)
D)
MnO4–(aq) + 8H+(aq) + 2Co2+(aq) → Mn2+(aq) + 4H2O(aq) + 2Co3+(aq)
E)
MnO4–(aq) + Co2+(aq) → Mn2+(aq) + 2O2(g) + Co3+(aq)
8. When the following oxidation–reduction reaction in basic solution is balanced, what is the
lowest whole-number coefficient for OH–, and on which side of the balanced equation
should it appear?
Cr2O3(s) → Cr(OH)2(aq) + CrO42–(aq)
A)
2, reactant side
B)
10, product side
C)
4, product side
D)
4, reactant side
E)
2, product side
9. Which of the following statements is true for a voltaic (galvanic) cell?
A)
The electron flow is from the anode to the cathode.
B)
The electron flow is from the positive electrode to the negative electrode.
C)
The electron flow is from the negative cathode to the positive anode.
D)
The electron flow is through the salt bridge.
E)
The electron flow is from the oxidizing agent to the reducing agent through an
external circuit.
10. Which of the following statements concerning a voltaic cell is/are correct?
1.
Reduction occurs at the cathode.
2.
A spontaneous reaction generates an electric current in a voltaic cell.
3.
Without a salt bridge charge buildup will cause the cell reaction to stop.
A)
1 only
B)
2 only
C)
3 only
D)
1 and 2
E)
1, 2, and 3
11. The electrochemical reaction which powers a lead-acid storage battery is as follows:
Pb(s) + PbO2(s) + 4H+(aq) + 2SO42–(aq) → 2PbSO4(s) + 2H2O(l)
A single cell of this battery consists of a Pb electrode and a PbO2 electrode, each submerged
in sulfuric acid. What reaction occurs at the cathode during discharge?
A)
Pb(s) is reduced to PbSO4(s).
B)
PbO2(s) is reduced to PbSO4(s).
C)
PbO2(s) is oxidized to PbSO4(s).
D)
Pb(s) is oxidized to PbSO4(s).
E)
H+ is oxidized to H2O(l) .
12. A lead storage battery involves the following two half-reactions:
PbSO4(s) + 2e– → Pb(s) + SO42–(aq); E° = –0.36 V
PbO2(s) + 4H+(aq) + SO42–(aq) + 2e– → PbSO4(s) + 2H2O(l); E° = 1.69 V
During the discharge reaction of the lead storage battery at 1.0 M concentrations, the cell
potential and the reducing agent are, respectively,
A)
2.05 V and PbO2.
B)
–2.05 V and PbO2.
C)
1.33 V and Pb.
D)
–2.05 V and Pb.
E)
2.05 V and Pb.
Test Bank General Chemistry, 10th edition 5
13. The anode in a voltaic cell and in an electrolytic cell is
A)
the site of oxidation and of reduction, respectively.
B)
the site of reduction and of oxidation, respectively.
C)
positive in both cells.
D)
the site of reduction in both cells.
E)
the site of oxidation in both cells.
14. The cell reaction for a certain alkaline dry cell is
Zn(s) + 2MnO2(s) + H2O(l) → Zn(OH)2(s) + Mn2O3(s)
Which of the following statements concerning this cell is/are correct?
1.
MnO2(s) is oxidized to Mn2O3(s) during cell discharge.
2.
Water is consumed during the discharge of this cell.
3.
Zn(OH)2(s) forms at the anode during discharge of this cell.
A)
1 only
B)
2 only
C)
3 only
D)
2 and 3
E)
1, 2 and 3
15. Which statement concerning the cathode in an electrochemical cell is correct?
A)
Metal ions may be deposited as metal atoms on the cathode during cell discharge.
B)
When connected to an external circuit, the cathode develops a positive charge.
C)
Reduction occurs at the cathode during cell discharge.
D)
Positive ions flow towards the cathode during cell discharge.
E)
All of the above.
16. The following reactions take place in a lead storage battery.
Discharging: Pb(s) + PbO2(s) + 4H+(aq) + 2SO42–(aq) → 2PbSO4(s) + 2H2O(l)
Charging: 2PbSO4(s) + 2H2O(l) → Pb(s) + PbO2(s) + 4H+(aq) + 2SO42–(aq)
Which of the following statements is true?
A)
The concentration of H2SO4 increases as the battery discharges.
B)
Pb is formed at the anode during discharge.
C)
Pb is formed at the cathode during charging.
D)
The mass of Pb decreases during charging.
E)
The mass of PbSO4 remains constant during charging and discharging.
17. A strip of iron is placed in a 1 M solution of iron(II) sulfate, and a strip of copper is placed
in a 1 M solution of copper(II) chloride. The two solutions are connected with a salt bridge,
and the two metals are connected by a wire.
Reduction Half-Reaction
E° (V)
Fe2+(aq) + 2e– Fe(s)
–0.41
Cu2+(aq) + 2e– Cu(s)
0.34
Which of the following takes place?
A)
Sulfur deposits at the iron electrode.
B)
The Fe(II) concentration of the iron half-cell decreases.
C)
Copper atoms deposit at the cathode.
D)
Chlorine is produced at the copper electrode.
E)
Chlorine is produced at the iron electrode.
18. Which reaction would be most likely to occur at the anode of a voltaic cell?
A)
PbSO4(s) + 2e– → Pb(s) + SO42–(aq)
B)
2H2O(l) + 2e– → H2(g) + 2OH–(aq)
C)
2H2O(l) → O2(g) + 4H+(aq) + 4e–
D)
PbSO4(s) → Pb2+(aq) + SO42–(aq)
E)
2H2O(l) → 2H2(g) + O2(g)
19. Which of the following statements is true concerning the voltaic cell shown below?
A)
The Cu cathode mass increases as the cell discharges.
B)
The Cu anode mass decreases as the cell discharges.
C)
The Cu anode mass increases as the cell discharges.
D)
The Cu cathode mass decreases as the cell discharges.
E)
The mass of the Cu electrode neither increases nor decreases as the cell discharges.
20. Which of the following statements is true concerning half-cell I as the voltaic cell shown
below discharges?
A)
[Zn2+] increases with time, and [Cl−] increases with time.
B)
[Zn2+] decreases with time, and [Cl−] increases with time.
C)
[Zn2+] decreases with time, and [Cl−] decreases with time.
D)
[Zn2+] decreases with time, and [NO3−] increases with time.
E)
[Zn2+] increases with time, and [NO3−] increases with time.
Test Bank General Chemistry, 10th edition 8
21. Which of the following cell reactions would require the use of an inert electrode?
A)
Fe(s) + 2Ag+(aq) → 2Ag(s) + Fe2+(aq)
B)
3Cu(s) + 2Au3+(aq) → 3Cu2+ + 2Au(s)
C)
Fe(s) + 2Ag+(aq) → Fe2+(aq) + 2Ag(s)
D)
Fe(s) + 2MnO2(s) + 2NH4+(aq) → Fe2+(aq) + Mn2O3(s) + 2NH3(aq) + H2O(l)
E)
3Zn2+(aq) + 2Al(s) → 3Zn(s) + 2Al3+(aq)
22. What is the effect on the cell when a salt bridge in an electrochemical cell is completely
clogged during cell discharge?
1.
Each half-cell reaction stops.
2.
The flow of ions to and from the salt bridge is disrupted.
3.
The flow of current through the external circuit slows but does not stop.
A)
1 only
B)
2 only
C)
3 only
D)
1 and 2
E)
1, 2 and 3
23. According to the following cell notation, which species is undergoing reduction?
Sn | Sn2+(aq) || Mn2+(aq) | MnO2(s) | Pt(s)
A)
Mn2+(aq)
B)
Sn2+(aq)
C)
Sn(s)
D)
MnO2(s)
E)
Pt(s)
24. In the following electrochemical cell, what is the role of the platinum?
Cu(s) | Cu2+(aq) || Fe3+(aq), Fe2+(aq) | Pt(s)
A)
The platinum serves as the anode.
B)
The oxidation of Fe2+ takes place at the surface of the platinum as the cell
discharges.
Test Bank General Chemistry, 10th edition 9
C)
The reduction of Fe3+ takes place at the surface of the platinum as the cell
discharges.
D)
A and C.
E)
A and B.
25. In the following electrochemical cell, what is the reduction half reaction?
Mn(s) | Mn2+(aq) || Fe3+(aq), Fe2+(aq) | Pt(s)
A)
Fe3+(aq) + e− → Fe2+(aq)
B)
Fe2+(aq) + e− → Fe3+(aq)
C)
Fe2+(aq) + Pt(s) → Fe3+(aq) + e−
D)
Mn2+(aq) → Mn(s) + 2e−
E)
Mn(s) → Mn2+(aq) + 2e−
26. A zinc–copper voltaic cell is represented as follows:
Zn(s) | Zn2+(1.0 M) || Cu2+(1.0 M) | Cu(s)
Which of the following statements is false?
A)
The copper electrode is the anode.
B)
Reduction occurs at the copper electrode during discharge.
C)
The mass of the zinc electrode decreases during discharge.
D)
Electrons flow through the external circuit from the zinc electrode to the copper
electrode.
E)
The concentration of Cu2+ decreases during discharge.
27. What is the cell reaction for the following voltaic cell?
Al(s) | Al3+(aq) || Br–(aq) | Br2(g) | Pt(s)
A)
2Al(s) + 3Br2(g) 2Al3+(aq) + 6Br–(aq)
B)
Al(s) + Al3+(aq) Br–(aq) + Br2(g)
C)
2Al3+(aq) + 6Br–(aq) 2Al(s) + 3Br2(g)
D)
Al(s) + 3Br2(g) Al3+(s) + 2Br–(aq)
E)
Al(s) + 2Br–(aq) Br2(g) + Al3+(aq)
Test Bank General Chemistry, 10th edition 10
28. What is the cell notation for the voltaic cell shown below?
A)
Zn2+(aq) | Zn(s) || Cu2+(aq) | Cu(s)
B)
Zn(s) | Zn2+(aq) || Cu(s) | Cu2+(aq)
C)
Zn(s) | Cu(s) || Zn2+(aq) | Cu2+(aq)
D)
Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)
E)
Zn2+(aq) | Zn(s) || Cu(s) | Cu2+(aq)
29. What is the cell reaction for the following electrochemical cell?
Cu | Cu2+(aq) || Ni2+(aq) | Ni
A)
2Cu(s) + Ni2+(aq) → Ni(s) + 2Cu2+(aq)
B)
Cu(s) + Cu2+(aq) → Ni(s) + Ni2+(aq)
C)
Cu(s) + Ni2+(aq) → Ni(s) + Cu2+(aq)
D)
2Ni(s) + Cu2+(aq) → Cu(s) + 2Ni2+(aq)
E)
Ni(s) + Cu2+(aq) → Cu(s) + Ni2+(aq)
30. What is the correct cell notation for a cell in which the hydrogen electrode is the anode and
the cathode half-reaction is Co3+(aq) + e− → Co2+(aq).
A)
Pt(s) | H2(g) | H+(aq) || Co3+(aq), Co2+(aq) | Pt(s)
B)
Pt(s) | H2(g) | H+(aq) || Co3+(aq), Co2+(aq)
C)
Co2+(aq), Co3+(aq) || H+(aq) | H2(g) | Pt(s)
D)
Pt(s)| Co2+(aq), Co3+(aq) || H+(aq) | H2(g) | Pt(s)
E)
H2(g) | H+(aq) || Co2+(aq), Co3+(aq)
31. Consider the following standard electrode potentials:
Ag+(aq) + e– Ag(s); E° = 0.80 V
Mn2+(aq) + 2e– Mn(s); E° = –1.18 V
Which of the following statements is false concerning the electrochemical cell given below?
Mn(s) | Mn2+(aq) || Ag+(aq) | Ag(s)
A)
The anode half-cell reaction is Mn(s) → Mn2+(aq) + 2e–.
B)
The reducing agent is Ag(s).
C)
Under standard-state conditions, the cell potential is 1.98 V.
D)
The cell potential decreases with time.
E)
The oxidizing agent is Ag+(aq).
32. For a galvanic cell using Fe | Fe2+(1.0 M) and Pb | Pb2+(1.0 M) half-cells, which of the
following statements is correct?
Fe2+(aq) + 2e– → Fe(s); E° = –0.41 V
Pb2+(aq) + 2e– → Pb(s); E° = –0.13 V
A)
The iron electrode is the cathode.
B)
When the cell has completely discharged, the concentration of Pb2+ is zero.
C)
The mass of the iron electrode increases during discharge.
D)
The concentration of Pb2+ decreases during discharge.
E)
Electrons leave the lead electrode to pass through the external circuit during
discharge.
33. What is the cell reaction for the following electrochemical cell?
Ni | Ni2+(aq) || Y3+(aq) | Y
A)
3Ni(s) + 2Y3+(aq) → 2Y(s) + 3Ni2+(aq)
B)
2Y(s) + 3Ni2+(aq) → 3Ni(s) + 2Y3+(aq)
C)
Ni(s) + Ni2+(aq) → Y(s) + Y3+(aq)
D)
Ni(s) + Y3+(aq) → Y(s) + Ni2+(aq)
E)
Y(s) + Ni2+(aq) → Ni(s) + Y3+(aq)
34. What is the cell reaction for the following electrochemical cell?
Pt | Fe2+(aq),Fe3+(aq) || Al3+(aq) | Al
A)
3Fe2+(aq) + Al3+(s) → Al(s) + 3Fe3+(aq)
B)
3Fe2+(aq) + 3Fe3+(aq) → Al(s) + Al3+(aq)
C)
Al(s) + 3Fe3+(aq) → Al3+(aq) + 3Fe2+(aq)
D)
Pt(s) + Fe2+(aq) + Al3+(aq) → Al(s) + Pt2+(aq) + Fe3+(aq)
E)
Al(s) + Al3+(aq) → 3Fe2+(aq) + 3Fe3+(aq)
35. What is the SI unit of potential difference?
A)
coulomb
B)
farad
C)
volt
D)
joule
E)
ampere
36. The Faraday constant is defined as
A)
The charge per mole of electrons.
B)
The charge on a single electron.
C)
The maximum work obtainable from an electrochemical cell.
D)
The electromotive force of the cell.
E)
The amount of charge moved between electrodes.
37. Which of the following statments concerning electrochemical cell potentials is/are correct?
1.
Electrons flow from a point of high electrical potential to a point of low
electrical potential.
2.
The potential difference between two points in a circuit is measured in
coulombs.
3.
The cell voltage is at its maximum when no current is flowing.
A)
1 only
B)
2 only
C)
3 only
D)
2 and 3
E)
1 and 3
38. 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 that this cell can achieve?
A)
1.14 V
B)
2.28 V
C)
2.01 V
D)
0.760 V
E)
13.7 V
39. The cell potential of an electrochemical cell with the cell reaction
Al(s) + Cr3+(aq) → Cr(s) + Al3+(aq)
is 1.63 V. What is the maximum electrical work obtainable from this cell when 0.50 g of Al
is consumed?
A)
−8.7 103
J
B)
−2.9 103
J
C)
−1.7 104
J
D)
−6.4 106
J
E)
−2.4 105
J
40. The cell potential of an electrochemical cell with the cell reaction
Al(s) + Cr3+(aq) → Cr(s) + Al3+(aq)
is 1.63 V. What is the maximum electrical work obtainable from this cell when 1.0 mol of
Al is consumed?
A)
−4.7 105
J
B)
−1.6 105
J
C)
−1.8 105
J
D)
−5.2 104
J
E)
−3.1 105
J
41. What is the balanced spontaneous reaction and standard cell potential of an electrochemical
cell constructed from half cells with the following half reactions?
E° = –0.763 V
E° = –0.130 V
A)
Pb2+(aq) + Zn(s) → Pb(s) + Zn2+(aq); 0.633 V
B)
Pb(s) + Zn2+(aq) → Pb2+(aq) + Zn(s); −0.633 V
C)
Pb2+(aq) + Zn2+(aq) → Pb(s) + Zn(s); –0.893 V
D)
Pb2+(aq) + Zn(s) → Pb(s) + Zn2+(aq); 0.317 V
E)
Pb(s) + Zn2+(aq) → Pb2+(aq) + Zn(s); −0.317 V
42. In a table of standard reduction potentials, the strongest reducing agents are the _______
species in the half-reactions with the _______ E° values.
A)
reduced, most negative
B)
oxidized, most positive
C)
reduced, most positive
D)
oxidized, most negative
E)
none of these
43. Given:
Mn2+(aq) + 2e– Mn(s); E° = –1.18 V
Cu2+(aq) + 2e– Cu(s); E° = 0.34 V
Cr2O72–(aq) + 14H+(aq) + 6e– 2Cr3+(aq) + 7H2O(l); E° = 1.33 V
Which of the following species is the strongest reducing agent?
A)
Mn
B)
Cu
C)
Cr3+
D)
Mn2+
E)
Cr2O72–
Test Bank General Chemistry, 10th edition 15
44. Given:
Ni2+(aq) + 2e– Ni(s); E° = –0.23 V
Cu2+(aq) + 2e– Cu(s); E° = 0.34 V
Cr2O72–(aq) + 14H+(aq) + 6e– 2Cr3+(aq) + 7H2O(l); E° = 1.33 V
Which of the following species is the strongest oxidizing agent?
A)
Ni
B)
Cr3+
C)
Cr2O72–
D)
Ni2+
E)
Cu
45. Given:
Ni2+(aq) + 2e– Ni(s); E° = –0.23 V
2H+(aq) + 2e– H2(g); E° = 0.00 V
Ag+(aq) + e– Ag(s); E° = 0.80 V
NO3–(aq) + 4H+(aq) + 3e– NO(g) + 2H2O(l); E° = 0.96 V
Which of the following statements is true?
A)
Ni2+ reacts spontaneously with 1 M H+(aq) to form H2.
B)
Ni2+ reacts spontaneously with H2(g).
C)
Ag(s) reacts spontaneously with Ni2+.
D)
Ag(s) reacts spontaneously with 1 M NO3– in 1 M H+(aq).
E)
Ag(s) reacts spontaneously with 1 M H+(aq).
46. Given:
Pb2+(aq) + 2e– Pb(s); E° = –0.13 V
Mg2+(aq) + 2e– Mg(s); E° = –2.38 V
Ag+(aq) + e– Ag(s); E° = 0.80 V
2H+(aq) + 2e– H2(g); E° = 0.00 V
Test Bank General Chemistry, 10th edition 16
Under standard-state conditions, which of the following species is the best reducing agent?
A)
H2
B)
Mg2+
C)
Ag+
D)
Pb
E)
Ag
47. Given:
Zn2+(aq) + 2e– Zn(s); E° = –0.76 V
Co2+(aq) + 2e– Co(s); E° = –0.28 V
Sn2+(aq) + 2e– Sn(s); E° = –0.15 V
Pb2+(aq) + 2e– Pb(s); E° = –0.13 V
Fe3+(aq) + e– Fe2+(aq); E° = 0.77 V
Under standard-state conditions, which of the following pairs of elements or ions is capable
of reducing Sn2+(aq) to Sn(s)?
A)
Zn(s) or Co(s)
B)
Pb(s) or Fe2+(aq)
C)
Co(s) or Pb(s)
D)
Fe2+(aq) or Zn(s)
E)
Zn2+(aq) or Co2+(aq)
48. Given:
Pb2+(aq) + 2e– Pb(s); E° = –0.13 V
2H+(aq) + 2e– H2(g); E° = 0.00 V
NO3–(aq) + 4H+(aq) + 3e– NO(g) + 2H2O(l); E° = 0.96 V
O2(g) + 4H+(aq) + 4e– 2H2O(l); E° = 1.23 V
PbO2(s) + SO42–(aq) + 4H+(aq) + 2e– PbSO4(s) + 2H2O(l); E° = 1.69 V
Under standard-state conditions, which of the following is the best oxidizing agent?
A)
H+
B)
Pb2+
C)
PbO2
D)
O2
E)
NO3–
49. Given:
2H+(aq) + 2e– H2(g); E° = 0.00 V
Li+(aq) + e– Li(s); E° = –3.04 V
F2(g) + 2e– 2F–(aq); E° = 2.87 V
Al3+(aq) + 3e– Al(s); E° = –1.66 V
Pb2+(aq) + 2e– Pb(s); E° = –0.13 V
Under standard-state conditions, which is the strongest reducing agent?
A)
Pb2+
B)
Al3+
C)
F–
D)
Li
E)
H+
50. Given:
W3+(aq) + 3e– W(s); E° = 0.10 V
Pb2+(aq) + 2e– Pb(s); E° = –0.13 V
Ni2+(aq) + 2e– Ni(s); E° = –0.23 V
Cd2+(aq) + 2e– Cd(s); E° = –0.40 V
Zn2+(aq) + 2e– Zn(s); E° = –0.76 V
Al3+(aq) + 3e– Al(s); E° = –1.66 V
Mg2+(aq) + 2e– Mg(s); E° = –2.38 V
Under standard-state conditions, which of the following metals will reduce W3+ to W but
will not reduce Ni2+ to Ni?
A)
Cd
B)
Pb
C)
Al
D)
Zn
E)
Mg
51. Given:
Zn2+(aq) + 2e– Zn(s); E° = –0.76 V
Cr3+(aq) + 3e– Cr(s); E° = –0.74 V
Fe2+(aq) + 2e– Fe(s); E° = –0.41 V
Cd2+(aq) + 2e– Cd(s); E° = –0.40 V
Sn2+(aq) + 2e– Sn(s); E° = –0.15 V
Hg2+(aq) + 2e– Hg(s); E° = 0.85 V
Au+(aq) + e– Au(s); E° = 1.69 V
Under standard-state conditions, which of the following metals will reduce Hg2+ to Hg but
will not reduce Cd2+ to Cd?
A)
Cr
B)
Zn
C)
Fe
D)
Sn
E)
Au
52. Given:
Zn2+(aq) + 2e– Zn(s); E° = –0.76 V
2H+(aq) + 2e– H2(g); E° = 0.00 V
I2(s) + 2e– 2I–(aq); E° = 0.54 V
Br2(l) + 2e– 2Br–(aq); E° = 1.07 V
Ni2+(aq) + 2e– Ni(s); E° = –0.23 V
Cu2+(aq) + 2e– Cu(s); E° = 0.34 V
Which of the following species will oxidize Ni but not Cu?
A)
Zn2+
B)
Br–
C)
H+
D)
I2
E)
Zn
53. Given:
Pb2+(aq) + 2e– Pb(s); E° = –0.13 V
Zn2+(aq) + 2e– Zn(s); E° = –0.76 V
Al3+(aq) + 3e– Al(s); E° = –1.66 V
Mg2+(aq) + 2e– Mg(s); E° = –2.38 V
V2+(aq) + 2e– V(s); E° = –1.18 V
Cu2+(aq) + 2e– Cu(s); E° = 0.34 V
Which of the following cations is capable of oxidizing Pb to Pb2+ under standard-state
conditions?
A)
Al3+
B)
V2+
C)
Cu2+
D)
Mg2+
E)
Zn2+
54. Given:
Li+(aq) + e– Li(s); E° = –3.04 V
Mg2+(aq) + 2e– Mg(s); E° = –2.38 V
Fe2+(aq) + 2e– Fe(s); E° = –0.41 V
Ag+(aq) + e– Ag(s); E° = 0.80 V
Br2(l) + 2e– 2Br–(aq); E° = 1.07 V
Which of the following species is the best oxidizing agent?
A)
Mg2+
B)
Br–
C)
Fe2+
D)
Li
E)
Ag+
55. Given:
Hg2+(aq) + 2e– Hg(s); E° = 0.85 V
Li+(aq) + e– Li(s); E° = –3.04 V
Br2(l) + 2e– 2Br–(aq); E° = 1.07 V
Au+(aq) + e– Au(s); E° = 1.69 V
Cl2(g) + 2e– 2Cl–(aq); E° = 1.36 V
Which of the following species is the best reducing agent?
A)
Hg
B)
Li+
C)
Au
D)
Cl–
E)
Br–
56. Which of the following statements is true about a voltaic cell for which E°cell = –1.00 V?
A)
The cathode is at a higher energy than the anode.
B)
It has G° < 0.
C)
The reaction is nonspontaneous.
D)
The system is at equilibrium.
E)
It has K = 1.
57. Which of the following is true for a reaction that is nonspontaneous as written?
A)
G > 0; Ecell < 0
B)
G < 0; Ecell > 0
C)
G < 0; Ecell < 0
D)
G > 0; Ecell > 0
E)
G > 0; Ecell = 0
58. Consider the following standard reduction potentials:
Mg2+(aq) + 2e– Mg(s); E° = –2.38 V
V2+(aq) + 2e– V(s); E° = –1.18 V
Cu2+(aq) + e– Cu+(aq); E° = 0.15 V
