University of South Florida

CHM 2046

1)Calculate the number of moles of ammonia needed to dissolve 0.0010 mole of silver chloride in 4000. mL of solution at 25?C. Ksp for AgCl is 1.8 × 10-10 and the dissociation constant for [Ag(NH3)2] + is 6.3 × 10-8 .

1. 0.16 mole
2. 0.021 mole
3. 14 moles

4. 0.28 mole
5. 1.5 moles

2. In an electrolytic cell, the electrode that acts as a source of electrons to the solution is called the     ; the chemical change that occurs at this electrode is called                   .

1. anode; oxidation
2. anode; reduction
3. cathode; oxidation
4. cathode; reduction
5. Cannot answer unless we know the species being oxidized and reduced.

3. What product is formed at the anode when molten sodium chloride, NaCl, is electrolyzed using a Downs Cell?

1. O2
2. Cl2
3. NaOH
4. H2
5. Na metal

4. How long would a constant current of 4.5 amperes be required to flow in order to plate out 15 g of chromium from a chromium(III) sulfate solution?

1. 268 hr
2. 309 hr
3. 5.15 hr
4. 23.2 hr
5. 1.72 hr

5. Which of the following is not a feature of the lead storage battery?

1. The electrolyte is hydrochloric acid.
2. Lead is oxidized at the anode.
3. PbO2 is reduced at the cathode.
4. Lead (II) sulfate forms during discharge and sticks to the electrodes

6. Use the standard half-cell potentials listed below to calculate the standard cell potential for the following reaction occurring in an electrochemical cell at 25°C. (The equation is balanced.)

3 Cl2(g) + 2 Fe(s) → 6 Cl– (aq) + 2 Fe3+(aq) Cl2(g) + 2 e– → 2 Cl– (aq)                                 E° = +1.36 V

Fe3+(aq) + 3 e– → Fe(s)             E° = -0.04 V

A) +4.16 V

B) -1.40 V

C) -1.32 V

D) +1.32 V

E) +1.40 V

7. Use the standard half-cell potentials listed below to calculate the standard cell potential for the following reaction occurring in an electrochemical cell at 25°C. (The equation is balanced.)

Pb(s) + Br2(l) → Pb2+(aq) + 2 Br– (aq)

Pb2+(aq) + 2 e– → Pb(s)               E° = -0.13 V Br2(l) + 2 e– → 2 Br?(aq)                             E° = +1.07 V

A) +1.20 V

B) +0.94 V

C) -0.94 V

D) -1.20 V

E) -0.60 V

8. Use the standard half-cell potentials listed below to calculate the standard cell potential for the following reaction occurring in an electrochemical cell at 25°C. (The equation is balanced.)

Sn(s) + 2 Ag?(aq) → Sn2+(aq) + 2 Ag(s) Sn2+(aq) + 2 e– → Sn(s)                                   E° = -0.14 V

Ag?(aq) + e– → Ag(s)                  E° = +0.80 V

A) +1.74 V

B) +0.94 V

C) +1.08 V

D) -1.08 V

E) -1.74 V

9. Use the standard half-cell potentials listed below to calculate the standard cell potential for the following reaction occurring in an electrochemical cell at 25°C. (The equation is balanced.)

Mg(s) + Cu2+(aq) → Cu(s) + Mg2+(aq)

Mg2+(aq) + 2 e– → Mg(s)            E° = -2.38 V Cu2+(aq) + 2 e– → Cu(s)            E° = +0.34 V

A) +2.04 V

B) -2.04 V

C) +2.72 V

D) -1.36 V

E) +1.36 V

10. Use the standard half-cell potentials listed below to calculate the standard cell potential for the following reaction occurring in an electrochemical cell at 25 °C. (The equation is balanced.)

2 K(s) + I2(s) → 2 K+ (aq) + 2 I– (aq)

K+(aq) + e– → K (s)                         E° = -2.93 V I2(s) + 2 e– → 2 I?(aq)                                E° = +0.54 V

A) +6.40 V

B) +1.85 V

C) -5.32 V

D) +3.47 V

E) +5.32 V

11. Determine which of the following pairs of reactants will result in a spontaneous reaction at 25°C.
    1. Pb2+(aq) + Cu(s)
    2. Ag+(aq) + Br– (aq)

1. 3. Li+(aq) + Al(s)
   4. Fe3+(aq) + Ni(s)
   5. None of the above pairs will react.

12. Determine which of the following pairs of reactants will result in a spontaneous reaction at 25°C.
    1. Sn4+(aq) + Mg(s)
    2. Cr3+(aq) + Ni(s)
    3. Zn(s) + Na+(aq)
    4. Fe(s) + Ba2+(aq)
    5. None of the above pairs will react.

13. Determine which of the following pairs of reactants will result in a spontaneous reaction at 25°C.
    1. I-(aq) + Zn2+(aq)
    2. Ca(s) + Mg2+(aq)
    3. H2(g) + Cd2+(aq)
    4. Ag(s) + Sn2+(aq)
    5. All of the above pairs will react.

14. Which of the following metals will dissolve in nitric acid but not hydrochloric?
    1. Fe
    2. Pb
    3. Cu
    4. Sn
    5. Ni

15. Which of the following metals will dissolve in nitric acid but not hydrochloric?
    1. Cd
    2. Cr
    3. Mn
    4. Ag
    5. Al

16. Which of the following metals will dissolve in HCl?
    1. Ba
    2. Na
    3. Mg
    4. Al
    5. All of the above
17. Which of the following statements is true for the cell diagram below? Zn(s) ? Zn2+(aq)

    

     Cu2+(aq) ? Cu(s)

1. Zn is oxidized, Cu is reduced; the single vertical lines represent phase boundaries while the two vertical lines represent a salt bridge.
2. Zn2+is oxidized, Cu is reduced; the single vertical lines represent phase boundaries while the two vertical lines represent a salt bridge.
3. Zn is oxidized, Cu2+ is reduced; the single vertical lines represent phase boundaries while the two vertical lines represent a salt bridge.
4. Zn is oxidized, Cu2+ is reduced; the single vertical lines represent salt bridges while the two vertical lines represent a phase boundary.
5. Zn is reduced, Cu2+ is oxidized; the single vertical lines represent phase boundaries while the two vertical lines represent a salt bridge.

18. If the standard reduction potential of Zn is -0.76 V, which of the following statements about a cell whose half-cells are Zn2+/Zn and SHE is correct?
    1. SHE will be the cell's cathode, Zn(s) will be the cell's anode, and the measured cell potential will be 0.76 V.
    2. SHE will be the cell's anode, Zn(s) will be the cell's cathode, and the measured cell potential will be 0.76 V.
    3. SHE will be the cell's cathode, Zn(s) will be the cell's anode, and the measured cell potential will be -0.76 V.
    4. SHE will be the cell's cathode, Zn2+(aq) will be the cell's anode, and the measured cell potential will be 0.76 V.
    5. H+(aq) will be the cell's cathode, Zn2+(aq) will be the cell's anode, and the measured cell potential will be 0.76 V.

19. Identify the characteristics of a spontaneous reaction.
    1. ΔrG° > 0
    2. ΔE°cell < 0
    3. K < 0
    4. ΔE°cell > 0
    5. K = 0

20. A galvanic cell consists of a Ni2+/Ni half-cell and a standard hydrogen electrode. If the Ni2+/Ni half-cell standard cell functions as the anode, and the standard cell potential is 0.26 V, what is the standard reduction potential for the Ni2+/Ni half-cell?

A) -0.26 V

B) -0.13 V

C) +0.13 V

D) +0.26 V

21. A galvanic cell consists of one half-cell that contains Ag(s) and Ag+(aq), and one half-cell that contains Cu(s) and Cu2+(aq). What species are produced at the electrodes under standard conditions?

Ag+(aq) + e- → Ag(s)E° = +0.80 V Cu2+(aq) + 2 e- → Cu(s)                                                E° = +0.34 V

1. 1. Ag(aq) is formed at the cathode, and Cu(s) is formed at the anode.
   2. Ag(s) is formed at the cathode, and Cu2+(aq) is formed at the anode.
   3. Cu(s) is formed at the cathode, and Ag+(aq) is formed at the anode.
   4. Cu2+(aq) is formed at the cathode, and Cu(s) is formed at the anode.

22. Consider the following standard reduction potentials:

Ni2+(aq) + 2 e- → Ni(s)                  E° = -0.26 V I2(s) + 2 e- → 2 I-(aq)                                   E° = +0.54 V

Under standard conditions,

1. Ni2+(aq) is a stronger oxidizing agent than I2(s), and I-(aq) is a stronger reducing agent than Ni(s).
2. I2(s) is a stronger oxidizing agent than Ni2+(aq), and Ni(s) is a stronger reducing agent than I-(aq).
3. Ni(s) is a stronger oxidizing agent than I-(aq), and Ni2+(aq) is a stronger reducing agent than I2(s).
4. I-(aq) is a stronger oxidizing agent than Ni(s), and I2(s) is a stronger reducing agent than Ni2+(aq).

23. Consider the galvanic cell,

    

     Which one of the following changes to the cell would cause the cell potential to increase (i.e., become more positive)?
    1. increase the [Zn2+] concentration
    2. increase the [Pb2+] concentration
    3. increase the mass of Zn(s)
    4. decrease the mass of Zn(s)

24. Based on the following information,

Cl2(g) + 2 e- → 2Cl-(aq)                 E° = +1.36 V Mg2+(aq) + 2 e- → 2Mg(s)           E° = -2.37 V

which of the following chemical species is the strongest reducing agent?

1. Cl2(g)
2. Mg2+(aq)
3. Cl-(aq)
4. Mg(s)

25. Using the following standard reduction potentials,

Fe3+(aq) + e- →Fe2+(aq)             E° = +0.77 V Ni2+(aq) + 2 e- →Ni(s)                   E° = -0.23 V

calculate the standard cell potential for the galvanic cell reaction given below and determine whether or not this reaction is spontaneous under standard conditions.

Ni2+(aq) + 2Fe2+(aq) → 2Fe3+(aq) + Ni(s)

1. E° = -1.00 V, nonspontaneous
2. E° = -1.00 V, spontaneous
3. E° = +1.00 V, nonspontaneous
4. E° = +1.00 V, spontaneous

26. How many electrons are transferred in the following reaction? (The reaction is unbalanced.) Mg(s) + Al3+(aq) → Al(s) + Mg2+(aq)

1. 1. 6
   2. 2
   3. 3
   4. 1
   5. 4

27. How many electrons are transferred in the following reaction? (The reaction is unbalanced.) Fe2+(aq) + K(s) →  Fe(s) + K+(aq)
    1. 1
    2. 2
    3. 3
    4. 4
    5. 6

28. How many electrons are transferred in the following reaction? (The reaction is unbalanced.) I2(s) + Fe3+(aq) →  Fe(s) + I– (aq)
    1. 1
    2. 2
    3. 6
    4. 3
    5. 4

29. Which of the following reactions would have the smallest value of K at 298 K?
    1. A + B → C; E°cell = +1.22 V
    2. A + 2 B → C; E°cell = +0.98 V

C) A + B → 2 C; E°cell = -0.030 V

4. A + B → 3 C; E°cell = +0.15 V
5. More information is needed to determine.

30. Use the tabulated half-cell potentials to calculate ΔG°for the following balanced redox reaction.

Pb2+(aq) + Cu(s) → Pb(s) + Cu2+(aq)

1. 1. -41 kJ

B) -0.47 kJ

C) +46 kJ

D) +91 kJ

E) -21 kJ

31. Use the tabulated half-cell potentials to calculate ΔG° for the following balanced redox reaction.

3 I2(s) + 2 Fe(s) →  2 Fe3+(aq) + 6 I– (aq)

A) -1.1 x 102 kJ

B) +4.9 x 101 kJ

C) -9.7 x 101 kJ

D) +2.3 x 102 kJ

E) -3.3 x 102 kJ

32. Use the tabulated half-cell potentials to calculate ΔG° for the following redox reaction. 2 Al(s) + 3 Mg2+(aq) → 2 Al3+(aq) + 3 Mg(s)

A) +4.1 x 102 kJ

B) +1.4 x 102 kJ

C) -2.3 x 102 kJ

D) -7.8 x 102 kJ

E) +6.8 x 102 kJ

33. Which of the following reactions would be the most spontaneous at 298 K?
    1. A + 2 B → C; E°cell = +0.98 V

B) A + B → 2 C;  E°cell = -0.030 V

3. A + B → 3 C; E°cell = +0.15 V
4. A + B → C; E°cell = +1.22 V
5. More information is needed to determine.

34.  Use the provided reduction potentials to calculate the equilibrium constant (K) for the following balanced redox reaction at 25 °C:

2Al(s) + 3Mg2+(aq) → 2Al3+(aq) + 3Mg(s)

E°(Al3+/Al) = -1.66 V and E°(Mg2+/Mg) = -2.37 V A) 1.1 × 1072

B) 8.9 × 10-70

C) 9.7 × 10-73

D) 1.0 × 1024

E) 4.6 × 1031

35. Use the tabulated half-cell potentials to calculate the equilibrium constant (K) for the following balanced redox reaction at 25°C.

2 Al(s) + 3 Mg2+(aq) → 2 Al3+(aq) + 3 Mg(s)

A) 1.1 × 1072

B) 8.9 × 10-73

C) 1.1 x 10-72

D) 1.0 × 1024

E) 4.6 × 1031

36. Use the tabulated half-cell potentials to calculate the equilibrium constant (K) for the following balanced redox reaction at 25°C.

3 I2(s) + 2 Fe(s) →  2 Fe3+(aq) + 6 I– (aq)

A) 3.5 × 10-59

B) 1.1 × 1017

C) 2.4 × 1058

D) 8.9 × 10-18

E) 1.7 × 1029

37. Use the tabulated half-cell potentials to calculate the equilibrium constant (K) for the following balanced redox reaction at 25°C.

Pb2+(aq) + Cu(s) → Pb(s) + Cu2+(aq)

A) 7.9 × 10-8

B) 8.9 × 107

C) 7.9 × 1015

D) 1.3 × 10-16

E) 1.1 × 10-8

38.  The gas OF2 can be produced from the electrolysis of an aqueous solution of KF, as shown in the equation below:

OF2(g) + 2H+(aq) + 4 e- → H2O(l) + 2F-(aq)                           E° = +2.15 V

Using the given standard reduction potential, calculate the amount of OF2 that is produced, and the electrode at which the OF2 is produced, upon the passage of 0.480 faradays through an aqueous KF solution.

1. 6.48 g of OF2 at the anode
2. 26.0 g of OF2 at the anode
3. 6.48 g of OF2 at the cathode
4. 26.0 g of OF2 at the cathode

39. Calculate the cell potential for the following reaction that takes place in an electrochemical cell at 25°C.

Sn(s) ? Sn2+(aq, 1.8 M) || Ag+(aq, 0.055 M) ? Ag(s) A) -0.94 V

B) -0.85 V

C) +1.02 V

D) +0.98 V

E) +0.86 V

40. Calculate the cell potential for the following reaction that takes place in an electrochemical cell at 25°C.

Sn(s) ? Sn2+(aq, 0.022 M) || Ag+(aq, 2.7 M) ? Ag(s)

A) +1.01 V

B) -0.83 V

C) +1.31 V

D) +0.01 V

E) -0.66 V

41. Calculate the cell potential for the following reaction that takes place in an electrochemical

cell at 25°C.

Mg(s) ? Mg2+(aq, 2.74 M) || Cu2+(aq, 0.0033 M) ? Cu(s) A) -2.80 V

B) +2.62 V

C) +2.71 V

D) +2.12 V

E) -1.94 V

42. Calculate the cell potential for the following reaction that takes place in an electrochemical cell at 25°C.

Fe(s) ? Fe3+(aq, 0.0011 M) || Fe3+(aq, 2.33 M) ? Fe(s)

A) +0.066 V

B) -0.036 V

C) 0.00 V

D) -0.099 V

E) +0.20 V

43. Calculate the cell potential for the following reaction that takes place in an electrochemical cell at 25°C.

Cu(s) ? Cu2+(aq, 0.0032 M) || Cu2+(aq, 4.48 M) ? Cu(s) A) 0.00 V

B) +0.093 V

C) +0.34 V

D) +0.186 V

E) +0.052 V

44. Calculate the cell potential for the following reaction that takes place in an electrochemical cell at 25°C.

Al(s) ? Al3+(aq, 0.115 M) || Al3+(aq, 3.89 M) ? Al(s) A) +1.66 V

B) +0.060 V

C) 0.00 V

D) +0.090 V

E) +0.030 V

45. Given that E°red = -0.40 V for Cd2+/Cd at 25 °C, find E° and E for the concentration cell expressed using shorthand notation below:

Cd(s) ? Cd2+(1.0 × 10-5 mol L-1) ?? Cd2+(0.100 mol L-1) ? Cd(s) A) E° = 0.00 V and E = +0.24 V

B) E° = 0.00 V and E = +0.0789 V

C) E° = -0.40 V and E = -0.16 V

D) E° = -0.40 V and E = -0.28 V

46. The standard cell potential (E°) of a voltaic cell constructed using the cell reaction below is

0.76 V:

Zn(s) + 2H+(aq) → Zn2+(aq) + H2(g)

With

A) 2.0 × 10-2

B) 4.2 × 10-4

C) 1.4 × 10-1

D) 4.9 × 101

E) 1.0 × 10-12

47. A voltaic cell is constructed with two silver-silver chloride electrodes, where the half- reaction is

AgCl(s) + e- → Ag(s) + Cl- (aq) E° = +0.222 V

The concentrations of chloride ion in the two compartments are 0.0222 mol L-1 and 2.22 mol L- 1, respectively. The cell emf is                                  V.

A) 0.212

B) 0.118

C) 0.00222

D) 22.2

E) 0.232

48. A voltaic cell is constructed with two Zn2+-Zn electrodes, where the half-reaction is Zn2+ + 2e- → Zn (s)                               E° = -0.763 V

The concentrations of zinc ion in the two compartments are 5.50 mol L-1 and 1.11 × 10-2 mol L- 1, respectively. The cell emf is                                  V.

A) -1.54 × 10-3

B) -378

C) 0.0798

D) 0.160

E) -0.761

49. The standard emf for the cell using the overall cell reaction below is +2.20 V: 2Al(s) + 3I2(s) → 2Al3+(aq) + 6I-(aq)

The emf generated by the cell when [Al3+] = 4.5 × 10-3 mol L-1 and [I-] = 0.15 mol L-1 is

A) 2.20

B) 2.32

C) 2.10

D) 2.39

E) 2.23

50. The standard emf for the cell using the overall cell reaction below is +0.48 V:

Zn(s) + Ni2+ (aq) → Zn2+ (aq) + Ni(s)

The emf generated by the cell when [Ni2+] = 2.50 mol L-1 and [Zn2+] = 0.100 mol L-1 is

A) 0.40

B) 0.50

C) 0.52

D) 0.56

E) 0.44

51. Identify the battery that is in most automobiles.
    1. dry-cell battery
    2. lithium-ion battery
    3. lead-acid storage battery
    4. NiCad battery
    5. fuel cell

52. Identify the battery that is used as a common flashlight battery.
    1. dry-cell battery
    2. lithium-ion battery
    3. lead-acid storage battery
    4. NiCad battery
    5. fuel cell

53. Identify the components of a fuel cell.
    1. nickel-metal hydride
    2. lithium-ion
    3. hydrogen-oxygen
    4. nickel-cadmium
    5. zinc-manganese

54. Identify the battery type that has a high overcharge tolerance.
    1. NiCad battery
    2. lithium-ion battery
    3. nickel-metal hydride battery
    4. lead-acid storage battery
    5. zinc-manganese battery

55. What is the reaction at the anode in a breathalyzer?
    1. Ethanol is oxidized to acetic acid.
    2. Acetic acid is reduced to ethanol.
    3. Oxygen is reduced.
    4. Hydrogen is oxidized.
    5. Ethanol is oxidized to acetaldehyde.

56. What is the reaction at the cathode in a breathalyzer?
    1. Ethanol is oxidized to acetic acid.
    2. Acetic acid is reduced to ethanol.
    3. Oxygen is reduced.
    4. Hydrogen is oxidized.
    5. Ethanol is oxidized to acetaldehyde.

57. Describe how water can be made to be a good conductor of electrical current.
    1. use pure water
    2. heat the water
    3. add salt
    4. chill the water

1. 5. vaporize the water

58. What mass of silver can be plated onto an object in 33.5 minutes at 8.70 A of current? Ag+ (aq) + e–  → Ag(s)

A) 19.6 g

B) 0.326 g

C) 9.78 g

D) 3.07 g

E) 0.102 g

59. What mass of aluminum can be plated onto an object in 755 minutes at 5.80 A of current? A) 73.5 g

B) 24.5 g

C) 220. g

D) 147 g

E) 8.17 g

60. Nickel can be plated from aqueous solution according to the following half reaction. How long would it take (in min) to plate 29.6 g of nickel at 4.7 A?

Ni2+(aq) + 2 e– → Ni(s) A) 1.7 × 102 min

B) 5.9 × 102 min

C) 3.5 × 102 min

D) 4.8 × 102 min

E) 6.2 × 102 min

61. In the presence of which of the following would iron be under the highest risk of rusting?
    1. salt only
    2. moisture only
    3. acid only
    4. salt and moisture combined
    5. salt, moisture, and acid together

62. Predict the species that will be reduced first if a mixture of molten salts containing the following ions undergoes electrolysis:

Na+, Ca2+, Cl?, Br?, F?

1. 1. Na?
   2. Cl?
   3. Ca2+
   4. Br?
   5. F?

63. Predict the species that will be reduced first if a mixture of molten salts containing the following ions undergoes electrolysis:

Zn2+, Fe3+, Mg2+, Br-, I-

1. 1. Zn2+
   2. Mg2+
   3. Br-
   4. Fe3+
   5. I-

64. Predict the species that will be oxidized first if a mixture of molten salts containing the following ions undergoes electrolysis:

Cu2+, Mg2+, Cl?, Br?, F?

1. 1. Cl?
   2. F?
   3. Cu2+
   4. Mg2+
   5. Br?

65. The electrolysis of molten AlCl3 for 3.25 hr with an electrical current of 15.0 A produces

                      g of aluminum metal. A) 147

B) 0.606

C) 4.55 × 10-3

D) 16.4

E) 49.1

66. For a galvanic cell that uses the following two half-reactions,

Cr2O72-(aq) + 14H+(aq) + 6 e- → 2Cr3+(aq) + 7H2O(l)

Pb(s) → Pb2+(aq) + 2 e-

how many moles of Pb(s) are oxidized by three moles of Cr2O72-?

1. 3
2. 6
3. 9
4. 18

67. How many seconds are required to produce 4.00 g of aluminum metal from the electrolysis of molten AlCl3 with an electrical current of 12.0 A?

A) 27.0

B) 9.00

C) 1.19 × 103

D) 2.90 × 105

E) 3.57 × 103

68. How many grams of chromium metal are plated out when a constant current of 8.00 A is passed through an aqueous solution containing Cr3+ ions for 320. minutes?

A) 27.6 g

B) 49.2 g

C) 82.4 g

D) 248 g

69. How many grams of nickel metal are plated out when a constant current of 15.0 A is passed