Electrochemistry Chapter 19 Verified Test Bank

Chemistry: Molecular Nature of Matter, 8e (Jespersen)

Chapter 19 Electrochemistry

1) An electrolyte is

A) a solid that conducts electrical energy.

B) an inert electrode that conducts electrical energy.

C) a metal that conducts electrical energy through a solution.

D) a compound that conducts electricity either in solution or in the molten state.

E) a solvent that conducts electricity.

Diff: 1

Section: 19.1

2) Anions

A) are charged ions that move toward the anode of a galvanic or electrolytic cell.

B) are charged ions that move toward the negative electrode of an electrolytic cell.

C) are charged ions that move toward the north pole of a magnetic field.

D) are positively charged ions that result from electrical discharge in a liquid solution.

E) are ions that attach themselves to any electrode to react chemically during electrolytic.

Diff: 1

Section: 19.1

3) Cations

A) are negatively charged ions that result from electrical discharge in a liquid solution.

B) are charged ions that move toward the cathode of a galvanic or electrolytic cell.

C) are charged ions that move toward the positive electrode of an electrolytic cell.

D) are charged ions that move toward the south pole of a magnetic field.

E) are ions that attach themselves to any electrode to react chemically during electrolysis.

Diff: 1

Section: 19.1

4) A galvanic cell has two electrodes. Which statement is correct?

A) Oxidation takes place at the anode, which is positively charged.

B) Oxidation takes place at the anode, which is negatively charged.

C) Oxidation takes place at the cathode, which is positively charged.

D) Oxidation takes place at the cathode, which is negatively charged.

E) Oxidation take place at the dynode, which is uncharged.

Diff: 1

Section: 19.1

5) A galvanic cell has two electrodes. Which statement is correct?

A) Reduction takes place at the anode, which is positively charged.

B) Reduction takes place at the anode, which is negatively charged.

C) Reduction takes place at the cathode, which is positively charged.

D) Reduction takes place at the cathode, which is negatively charged.

E) Reduction takes place at the dynode, which is uncharged.

Diff: 1

Section: 19.1

6) A galvanic cell consists of a Cu(s)|Cu2+(aq) half-cell and a Zn(s)|Zn2+(aq) half-cell, connected by a salt bridge. Oxidation occurs in the zinc half-cell. The cell can be represented in standard notation as:

A) Cu(s)|Cu2+(aq)|Zn(s)|Zn2+(aq)

B) Zn(s)|Zn2+(aq)||Cu(s)|Cu2+(aq)

C) Cu2+(aq)|Cu(s)||Zn(s)|Zn2+(aq)

D) Zn(s)|Zn2+(aq)||Cu2+(aq)|Cu(s)

E) Zn2+(aq)|Zn(s)||Cu(s)|Cu2+(aq)

Diff: 1

Section: 19.1

7) A galvanic cell consists of an Ag(s)|Ag+(aq) half-cell and a Zn(s)|Zn2+(aq) half-cell connected by a salt bridge. Oxidation occurs in the zinc half-cell. The cell can be represented in standard notation as

A) Ag(s)|Ag+(aq)|Zn(s)|Zn2+(aq)

B) Zn(s)|Zn2+(aq)||Ag(s)|Ag+(aq)

C) Ag+(aq)|Ag(s)||Zn(s)|Zn2+(aq)

D) Zn2+(aq)|Zn(s)||Ag(s)|Ag+(aq)

E) Zn(s)|Zn2+(aq)||Ag+(aq)|Ag(s)

Diff: 1

Section: 19.1

8) A galvanic cell consists of a Cd(s)|Cd2+(aq) half-cell and a Zn(s)|Zn2+(aq) half-cell connected by a salt bridge. Reduction occurs in the cadmium half-cell. The cell can be represented in standard notation as:

A) Cd(s)|Cd2+(aq)|Zn(s)|Zn2+(aq)

B) Zn(s)|Zn2+(aq)||Cd(s)|Cd2+(aq)

C) Zn(s)|Zn2+(aq)||Cd2+(aq)|Cd(s)

D) Zn2+(aq)|Zn(s)||Cd(s)|Cd2+(aq)

E) Cd2+(aq)|Cd(s)||Zn(s)|Zn2+(aq)

Diff: 1

Section: 19.1

9) A galvanic cell consists of a Cu(s)|Cu2+(aq) half-cell and a Cd(s)|Cd2+(aq) half-cell connected by a salt bridge. Oxidation occurs in the cadmium half-cell. The cell can be represented in standard notation as:

A) Cu(s)|Cu2+(aq)|Cd(s)|Cd2+(aq)

B) Cd(s)|Cd2+(aq)||Cu2+(aq)|Cu(s)

C) Cd2+(aq)|Cd(s)||Cu(s)|Cu2+(aq)

D) Cd(s)|Cd2+(aq)||Cu(s)|Cu2+(aq)

E) Cu2+(aq)|Cu(s)||Cd(s)|Cd2+(aq)

Diff: 1

Section: 19.1

10) A galvanic cell consists of an Ag(s)|Ag+(aq) half-cell and a Cu(s)|Cu2+(aq) half-cell connected by a salt bridge. The cell can be represented in standard notation as:

A) Ag(s)|Ag+(aq)|Cu(s)|Cu2+(aq)

B) Cu(s)|Cu2+(aq)||Ag(s)|Ag+(aq)

C) Ag+(aq)|Ag(s)||Cu(s)|Cu2+(aq)

D) Cu2+(aq)|Cu(s)||Ag(s)|Ag+(aq)

E) Cu(s)|Cu2+(aq)||Ag+(aq)|Ag(s)

Diff: 1

Section: 19.1

11) A galvanic cell consists of an Ag(s)|Ag+(aq) half-cell and a Cu(s)|Cu2+(aq) half-cell connected by a salt bridge. The cell can be represented in standard notation as: Cu(s)|Cu2+(aq)||Ag+(aq)|Ag(s).

Which species is being reduced?

A) Ag+(aq)

B) Cu2+(aq)

C) Ag(s)

D) Cu(s)

E) Cu+(aq)

Diff: 1

Section: 19.1

12) The electrode for which the standard reduction potential of 0.00 V is assigned uses the half-reaction:

A) Zn2+(aq) + 2e- Zn(s)

B) Cu2+(aq) + 2e- Cu(s)

C) Ag+(aq) + e- Ag(s)

D) 2 H+(aq) + 2e- H2(g)

E) 2 NH4+(aq) + 2e- H2(g) + 2 NH3(g)

Diff: 2

Section: 19.2

13) Using these metal ion/metal standard reduction potentials:

Cd2+(aq)|Cd(s) Zn2+(aq)|Zn(s) Ni2+(aq)|Ni(s) Cu2+(aq)|Cu(s)

-0.40 V -0.76 V -0.25 V +0.34 V

Calculate the standard cell potential for the cell whose reaction is:

Cu2+(aq) + Cd(s) → Cd2+(aq) + Cu(s)

A) +0.76 V

B) +0.06 V

C) -0.06 V

D) +0.74 V

E) +0.20 V

Diff: 1

Section: 19.2

14) Using these metal ion/metal standard reduction potentials:

Cd2+(aq)|Cd(s) Zn2+(aq)|Zn(s) Ni2+(aq)|Ni(s) Cu2+(aq)/Cu(s)

-0.40 V -0.76 V -0.25 V +0.34 V

Calculate the standard cell potential for the cell whose reaction is:

Ni2+(aq) + Zn(s) → Zn2+(aq) + Ni(s)

A) +0.51 V

B) -1.02 V

C) -1.01 V

D) +1.01 V

E) -0.51 V

Diff: 2

Section: 19.2

15) Using these metal ion/metal standard reduction potentials:

Fe2+(aq)|Fe(s) Zn2+(aq)|Zn(s) Cr3+(aq)|Cr(s) Cu2+(aq)/Cu(s)

-0.44 V -0.76 V -0.74 V +0.34 V

Calculate the standard cell potential for the cell whose reaction is:

Fe2+(aq) + Cr(s) → Fe(s) + Cr3+(aq)

A) -0.30 V

B) -1.18 V

C) +0.30 V

D) +0.16 V

E) -0.16 V

Diff: 2

Section: 19.2

16) Using these metal ion/metal reaction potentials:

Cu2+(aq)|Cu(s) Ag+(aq)|Ag(s) Co2+(aq)|Co(s) Zn2+(aq)|Zn(s)

+0.34 V +0.80 V -0.28 V -0.76 V

Calculate the standard cell potential for the cell whose reaction is:

Co(s) + Cu2+(aq) → Co2+(aq) + Cu(s)

A) -0.06 V

B) +0.06 V

C) -0.62 V

D) +0.62 V

E) +0.68 V

Diff: 2

Section: 19.2

17) Using the standard reduction potentials:

Au3+(aq) + 3 e- Au(s) +1.42 V

Ca2+(aq) + 2 e- Ca(s) -2.76 V

Calculate the value of E°cell for the reaction:

2Au(s) + 3Ca2+(aq) → 2Au3+(aq) + 3Ca(s)

A) -1.43 V

B) +1.34 V

C) -4.18 V

D) +4.18 V

E) -1.34 V

Diff: 2

Section: 19.2

18) Using the standard reduction potentials:

Cr3+(aq) + 3e- Cr(s) -0.74 V

Cl2(g) + 2e- 2Cl-(aq) +1.36 V

Calculate the value of E°cell for the cell with the reaction:

2Cr(s) + 3Cl2(g) → 2Cr3+(aq) + 6Cl-(aq)

A) -0.96 V

B) +0.96 V

C) +2.10 V

D) -2.10 V

E) +0.98 V

Diff: 2

Section: 19.2

19) For the reaction, 2 Cr2+(aq) + Cl2(g) 2 Cr3+(aq) + 2 Cl-(aq), the value of E°cell is

1.78 V. What is the value of E°cell for the following reaction?

Cr3+(aq) + Cl-(aq) → Cr2+(aq) + ½ Cl2(g)

A) -1.78 V

B) +0.89 V

C) +1.78 V

D) -0.89 V

E) -3.56 V

Diff: 2

Section: 19.2

20) The cell described by the net reaction:

2U(s) + 3Cl2(g) 6 Cl-(aq) + 2U3+(aq)

has a standard cell potential of 3.16 Vs. Using the standard reduction potential value shown for:

Cl2(g) + 2 e- 2 Cl-(aq) E° = +1.36 V

determine the standard reduction potential of the U3+(aq)|U(s) half-cell

A) -1.80 V

B) +1.80 V

C) -1.96 V

D) -4.52 V

E) +4.52 V

Diff: 2

Section: 19.2

21) Consider this electrochemical cell:

Pt | Pu3+(aq), Pu4+(aq) || Cl2(g), Cl-(aq) | Pt

Given that the standard cell potential is 0.35 V, what is the standard reduction potential E°(Pu4+/Pu3+)?

A) 2.37 V

B) 1.71 V

C) 1.01 V

D) -1.71 V

E) -1.01 V

Diff: 2

Section: 19.2

22) The cell described by the reaction,

2 Co3+(aq) + 2 Cl-(aq) 2 Co2+(aq) + Cl2(g)

has a standard potential of 0.46 V. Using the standard reduction potential value shown for

Cl2(g) + 2 e- 2 Cl-(aq) E° = +1.36 V

determine a value for the standard reduction potential of the cathode half-cell.

A) -0.90 V

B) +0.90 V

C) +0.91 V

D) -1.82 V

E) +1.82 V

Diff: 2

Section: 19.2

23) Consider these metal ion/metal standard reduction potentials

Cu2+(aq)|Cu(s) Ag+(aq)|Ag(s) Co2+(aq)|Co(s) Zn2+(aq)|Zn(s)

+0.34 V +0.80 V -0.28 V -0.76 V

Based on the data above, which one of the species below is the best reducing agent?

A) Co(s)

B) Zn(s)

C) Cu2+(aq)

D) Cu(s)

E) Ag(s)

Diff: 2

Section: 19.2

24) Using the standard reduction potentials

Mg2+(aq) + 2 e- ] Mg(s) -2.37 V

NO3(aq) + 4 H+(aq) + 3 e- NO(g) + 2 H2O +0.96 V

Calculate the value of E°cell for the cell with the reaction:

3 Mg(s) + 2 NO3-(aq) + 8 H+(aq) → 3 Mg2+(aq) + 2 NO(g) + 4 H2O

A) +1.41 V

B) -1.41 V

C) +3.33 V

D) +8.46 V

E) -8.46 V

Diff: 2

Section: 19.2

25) Using the standard reduction potentials

Ni2+(aq) + 2 e- Ni(s) -0.25 V

Fe3+(aq) + e- Fe2+(aq) +0.77 V

Calculate the value of E°cell for the cell with the following reaction.

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

A) +0.52 V

B) -1.02 V

C) +2.81 V

D) +1.02 V

E) -2.81 V

Diff: 2

Section: 19.2

26) Consider these metal ion/metal standard reduction potentials

Cu2+(aq)|Cu(s) Ag+(aq)|Ag(s) Co2+(aq)|Co(s) Zn2+(aq)|Zn(s) Cd2+(aq)|Cd(s)

+0.34 V +0.80 V -0.28 V -0.76 V -0.40 V

Based on the data above, which species is the best oxidizing agent?

A) Co2+(aq)

B) Zn2+(aq)

C) Cu2+(aq)

D) Cd2+(aq)

E) Ag+(aq)

Diff: 2

Section: 19.2

27) Consider these metal ion/metal standard reduction potentials

Cd2+(aq)|Cd(s) Zn2+(aq)|Zn(s) Ni2+(aq)|Ni(s) Cu2+(aq)|Cu(s) Ag+(aq)|Ag(s)

-0.40 V -0.76 V -0.25 V +0.34 V +0.80 V

Based on the data above, which species is the best reducing agent?

A) Cd(s)

B) Ag(s)

C) Ni(s)

D) Zn(s)

E) Cu(s)

Diff: 2

Section: 19.2

28) Consider these metal ion/metal standard reduction potentials

Cd2+(aq)|Cd(s) Zn2+(aq)|Zn(s) Ni2+(aq)|Ni(s) Cu2+(aq)|Cu(s) Co2+(aq)|Co(s)

-0.40 V -0.76 V -0.25 V +0.34 V -0.28 V

Based on the data above, which species is the best oxidizing agent?

A) Cd2+(aq)

B) Zn2+(aq)

C) Co2+(aq)

D) Cu2+(aq)

E) Ni2+(aq)

Diff: 2

Section: 19.2

29) Consider these metal ion/metal standard reduction potentials

Cd2+(aq)|Cd(s) Zn2+(aq)|Zn(s) Ni2+(aq)|Ni(s) Cu2+(aq)|Cu(s) Co2+(aq)|Co(s)

-0.40 V -0.76 V -0.25 V +0.34 V -0.28 V

Based on the data above, which species is the best reducing agent?

A) Co(s)

B) Cu(s)

C) Cd2+(aq)

D) Zn2+(aq)

E) Zn(s)

Diff: 2

Section: 19.2

30) Consider these metal ion/metal standard reduction potentials

Al3+(aq)|Al(s) Na+(aq)|Na(s) Au3+(aq)|Au(s) Ni2+(aq)|Ni(s) Cu2+(aq)|Cu(s)

-1.66 V -2.71 V +1.42 V -0.25 V +0.34 V

Based on the data above, which species is the best reducing agent?

A) Ni(s)

B) Na(s)

C) Au(s)

D) Cu(s)

E) Al(s)

Diff: 2

Section: 19.2

31) Which statement is true concerning a galvanic cell?

A) E° for the cell is always positive.

B) E° for the cell is always negative.

C) The standard reduction potential for the anode reaction is always positive.

D) The standard reduction potential for the anode reaction is always negative.

E) The standard reduction potential for the cathode reaction is always positive.

Diff: 1

Section: 19.2

32) A certain electrochemical cell has a cell potential of +0.34 V. Which of the following is a true statement about the electrochemical reaction?

A) The reaction favors the formation of reactants and would be considered an electrolytic cell.

B) The reaction favors the formation of reactants and would be considered a galvanic cell.

C) The reaction favors the formation of products and would be considered an electrolytic cell.

D) The reaction is at equilibrium and is a galvanic cell.

E) The reaction favors the formation of products and would be considered a galvanic cell.

Diff: 1

Section: 19.3

33) A unit of electrical energy is the

A) ampere.

B) coulomb.

C) joule.

D) volt.

E) watt.

Diff: 1

Section: 19.3

34) Consider the following reaction: 2Fe2+(aq) + Cu2+ → 2Fe3+(aq) + Cu.

When the ion concentrations change to the point where the reaction comes to equilibrium, what would be the cell voltage?

A) 1.11 V

B) -0.43 V

C) 0.0 V

D) 0.43 V

E) 0.78 V

Diff: 1

Section: 19.3

35) A unit of electrical charge used is the

A) ampere.

B) coulomb.

C) V.

D) joule.

E) watt.

Diff: 1

Section: 19.4

36) The Faraday constant is equal to the ________ on 1 mole of electrons.

A) capacitance

B) current

C) power

D) pressure

E) electrical charge

Diff: 1

Section: 19.4

37) One mole of electrical charge contains

A) 4.184 joules.

B) 3,600 coulombs.

C) 23,060 joules.

D) 96,485 coulombs.

E) 3.47 × 108 coulombs.

Diff: 1

Section: 19.4

38) A galvanic cell is composed of these two half-cells, with the standard reduction potentials shown:

Co2+(aq) + 2 e- Co(s) -0.28 V

Cd2+(aq) + 2 e- Cd(s) -0.40 V

What is the standard free energy change for the cell reaction of this galvanic cell?

A) -12 kJ

B) +12 kJ

C) -23 kJ

D) +23 kJ

E) -46 kJ

Diff: 1

Section: 19.4

39) A galvanic cell is composed of these two half-cells, with the standard reduction potentials shown:

Cu2+(aq) + 2 e- Cu(s) +0.34 V

Cd2+(aq) + 2 e- Cd(s) -0.40 V

What is the standard free energy (ΔG°)change for the cell reaction of this galvanic cell?

A) +12 kJ

B) -12 kJ

C) +143 kJ

D) -143 kJ

E) -71 kJ

Diff: 2

Section: 19.4

40) A galvanic cell is composed of these two half-cells, with the standard reduction potentials shown:

Zn2+(aq) + 2 e- Zn(s) -0.76 V

Cd2+(aq) + 2 e- Cd(s) -0.40 V

What is the standard free energy change (ΔG°) for the cell reaction of this galvanic cell?

A) 69 kJ

B) +69 kJ

C) -224 kJ

D) +224 kJ

E) -35 kJ

Diff: 2

Section: 19.4

41) A galvanic cell is composed of these two half-cells, with the standard reduction potentials shown:

Co2+(aq) + 2 e- Co(s) -0.28 V

Cr3+(aq) + 3 e- Cr(s) -0.74 V

What is the standard free energy (ΔG°) change for the cell reaction of this galvanic cell?

A) -88.8 kJ

B) -178 kJ

C) -266 kJ

D) -295 kJ

E) -590 kJ

Diff: 2

Section: 19.4

42) A galvanic cell is composed of these two half-cells, with the standard reduction potentials shown:

Fe2+(aq) + 2e- Fe(s) -0.44V

Al3+(aq) + 3e- Al(s) -1.66V

What is the standard free energy (ΔG°) change for the cell reaction of this galvanic cell?

A) -806 kJ

B) -1.22 × 103 kJ

C) -706 kJ

D) -540 kJ

E) -600 kJ

Diff: 2

Section: 19.4

43) A galvanic cell is composed of these two half-cells, with the standard reduction potentials shown:

Sn2+(aq) + 2e- Sn(s) -0.14V

Zn2+(aq) + 2e- Zn(s) -0.76V

What is the standard free energy (ΔG°) change for the cell reaction of this galvanic cell?

A) -2.22 × 102 kJ

B) -3.14 × 102 kJ

C) -1.74 × 102 kJ

D) -6.02 × 102 kJ

E) -1.20 × 102 kJ

Diff: 2

Section: 19.4

44) A galvanic cell is composed of these two half-cells, with the standard reduction potentials shown:

Cu2+(aq) + 2e- Cu(s) +0.34V

Fe3+(aq) + 3e- Fe(s) +0.77V

What is the standard free energy (ΔG°) change for the cell reaction of this galvanic cell?

A) -2.49 × 102 kJ

B) -3.21 × 102 kJ

C) -6.43 × 102 kJ

D) -5.32 × 102 kJ

E) -4.31 × 102 kJ

Diff: 2

Section: 19.4

45) Using the standard reduction potentials

Au3+(aq) + 3 e- Au(s) +1.42 V

Ca2+(aq) + 2 e- Ca(s) -2.76 V

Calculate the standard free energy (ΔG°) change for the cell reaction:

2 Au(s) + 3 Ca2+(aq) → 2 Au3+(aq) + 3 Ca(s)

A) 2420 kJ

B) 388 kJ

C) -766 kJ

D) 766 kJ

E) -1210 kJ

Diff: 2

Section: 19.4

46) Using these metal ion/metal standard reduction potentials

Fe2+(aq)|Fe(s) Zn2+(aq)|Zn(s) Cr3+(aq)|Cr(s) Cu2+(aq)/Cu(s)

-0.44 V -0.76 V -0.74 V +0.34 V

Calculate the standard free energy (ΔG°) change for the cell reaction:

Fe2+(aq) + Cr(s) → Fe(s) + Cr3+(aq)

A) -92.6 kJ

B) -86.8 kJ

C) 683.1 kJ

D) -57.9 kJ

E) -173.7 kJ

Diff: 2

Section: 19.4

47) Using the reduction potentials given, calculate the equilibrium constant, K, at 25°C for the reaction,

Ag+(aq) + Fe2+(aq) Ag(s) + Fe3+(aq)

Fe3+(aq) + e- Fe2+(aq) +0.77 V

Ag+(aq) + e- Ag(s) +0.80 V

A) 1.66

B) 6.4

C) 3.2

D) 6.1 × 10-4

E) 1.6 × 104

Diff: 2

Section: 19.4

48) The equilibrium constant, Kc, was found to be 1.2 × 103 at 25°C for the reaction,

2X(s) + Cu2+(aq) 2X+(aq) + Cu(s)

Using the following reduction potential for copper, what is the reduction potential for the other half reaction involving the substance X?

Cu2+(aq) + 2e- Cu(s) +0.34 V

A) -0.16 V

B) 0.091 V

C) 0.52 V

D) 0.18 V

E) -0.25 V

Diff: 2

Section: 19.4

49) Using the reduction potentials given, calculate the equilibrium constant, K, at 25°C for the reaction,

2I-(aq) + Br2(s) I2(s) + 2Br-(aq)

Br2(s) + 2e- 2Br-(aq) +1.07 V

A) 8.5 × 1017

B) 6.8

C) 2.4 × 104

D) 1.02

E) 1.2 × 10-18

Diff: 2

Section: 19.4

50) The equilibrium constant, Kc, was found to be 2.4 × 108 at 25°C for the following reaction,

2X(s) + 3Y2+(aq) 2X3+(aq) + 3Y(s)

Using this information, what is the standard reduction potential for this reaction?

A) 0.25 V

B) 0.083 V

C) 0.17 V

D) 0.50 V

E) 0.21 V

Diff: 2

Section: 19.4

51) A galvanic cell is composed of these two half-cells, with the standard reduction potentials shown:

Co2+(aq) + 2 e- Co(s) -0.28 V

Cd2+(aq) + 2 e- Cd(s) -0.40 V

The actual concentrations are: [Co2+(aq)] = 0.00100 M, [Cd2+] = 0.100 M. What is the potential of this galvanic cell?

A) +0.18 V

B) +0.12 V

C) +0.24 V

D) +0.060 V

E) +0.68 V

Hint: First calculate E°, then apply the solution concentrations of the galvanic cell using the Nernst equation.

Diff: 3

Section: 19.5

52) A galvanic cell is composed of these two half-cells, with the standard reduction potentials shown:

Co2+(aq) + 2 e- Co(s) -0.28 V

Cd2+(aq) + 2 e- Cd(s) -0.40 V

The actual concentrations in the cell are: [Co2+](aq) = 0.100 M, [Cd2+] = 0.0100 M. What is the potential of this galvanic cell?

Hint: First calculate E°, then apply the solution concentrations of the galvanic cell using the Nernst equation.

A) +0.06 V

B) +0.09 V

C) +0.15 V

D) +0.18 V

E) +0.24 V

Diff: 3

Section: 19.5

53) A galvanic cell is composed of these two half-cells, with the standard reduction potentials shown:

Co2+(aq) + 2 e- Co(s) -0.28 V

Cr3+(aq) + 3 e- Cr(s) -0.74 V

The actual concentrations in the cell are: [Co2+](aq) = 0.0100 M, [Cr3+] = 0.00100 M. What is the potential of this galvanic cell?

Hint: First calculate E°, then apply the solution concentrations of the galvanic cell using the Nernst equation.

A) +0.40 V

B) +0.46 V

C) +0.52 V

D) +0.54 V

E) +1.02 V

Diff: 3

Section: 19.5

54) A galvanic cell is composed of these two half-cells, with the standard reduction potentials shown:

Co2+(aq) + 2 e- Co(s) -0.28 V

Cr3+(aq) + 3 e- Cr(s) -0.74 V

The actual concentrations in the cell are: [Co2+] = 0.00100 M, [Cr3+]= 0.100 M. What is the potential of this galvanic cell?

Hint: First calculate E°, then apply the solution concentrations of the galvanic cell using the Nernst equation.

A) +0.33 V

B) +0.39 V

C) +0.45 V

D) +0.94 V

E) +1.61 V

Diff: 3

Section: 19.5

55) A galvanic cell is composed of these two half-cells:

Cr3+(aq) | Cr(s)

Cu2+(aq) | Cu(s)

The actual concentrations in the cell are: [Cu2+] = 0.00350 M, [Cr3+] = 0.360 M. What is the potential of this galvanic cell at 25°C?

Hint: First calculate E°, then apply the solution concentrations of the galvanic cell using the Nernst equation.

A) +1.06 V

B) -0.16 V

C) +1.02 V

D) +1.14 V

E) +1.98 V

Diff: 3

Section: 19.5

56) The standard reduction potentials of Cu2+(aq)|Cu(s) and Ag+(aq)|Ag(s) are +0.34 and

+0.80 V, respectively. Determine the value of the actual cell potential, Ecell, (in V) for the following cell at 25.0 °C.

Cu(s)|Cu2+(0.250 M)||Ag+(0.0010 M)|Ag(s)

Hint: First calculate E°, then apply the solution concentrations of the galvanic cell using the Nernst equation.

A) +0.30 V

B) +0.14 V

C) +0.62 V

D) +0.78 V

E) +0.39 V

Diff: 3

Section: 19.5

57) Fuel cells are different from other traditional batteries because

A) they require a constant supply of reactants to produce voltage.

B) they are only used in space.

C) they have a solid medium.

D) they require voltage to work.

E) they utilize heat from combustion of gases.

Diff: 1

Section: 19.6

58) When fused (molten) sodium chloride is electrolyzed what occurs?

A) Gaseous chlorine is formed at the cathode.

B) Hydrogen gas is formed at the cathode.

C) Liquid sodium is formed at the cathode.

D) Liquid chlorine is formed at the anode.

E) Solid sodium is formed at the anode.

Diff: 1

Section: 19.7

59) In doping semiconductor materials,

A) the energy gap between the valence band and the conduction band is completely removed and a covalent bond is formed.

B) impurities are added that either provide extra electrons, or 'holes' for electrons to move through.

C) a large flow of electricity is added to the material kicking electrons out of the material and creating 'holes'.

D) the material is destroyed using an acid.

E) the material is dissolved in an organic solvent.

Diff: 2

Section: 19.6

60) When an aqueous solution of AgNO3 is electrolyzed, a gas is formed at the anode. The gas is

A) dinitrogen tetroxide.

B) hydrogen.

C) mononitrogen monoxide.

D) nitrogen dioxide.

E) oxygen.

Diff: 2

Section: 19.7

61) Electrolysis is

A) the splitting of atomic nuclei by electrical energy.

B) the splitting of atoms by electrical energy.

C) the passage of electrical energy through a split-field armature.

D) the chemical reaction which results when electrical energy is passed through a liquid electrolyte.

E) the chemical reaction which results when electrical energy is passed through a metallic liquid.

Diff: 2

Section: 19.7

62) Which statement below is true?

A) Electrolysis cells generate alternating current when their terminals are reversed.

B) Electrolysis was discovered by Lewis Latimer.

C) Galvanic cells generate electrical energy rather than consuming it.

D) Galvanic cells were invented by Thomas Edison.

E) The Laws of Electrolysis were discovered by Alberta Nernst.

Diff: 2

Section: 19.7

63) When an aqueous solution of sodium chloride is electrolyzed, hydrogen gas is evolved at the cathode. The solution near the cathode becomes

A) acidic.

B) basic.

C) bubbly.

D) colored.

E) viscous.

Diff: 2

Section: 19.7

64) When an aqueous solution of sodium sulfate is electrolyzed, a gas is evolved at the anode. The solution near the anode becomes

A) acidic.

B) basic.

C) bubbly.

D) colored.

E) viscous.

Diff: 2

Section: 19.7

65) An electrolytic cell has two electrodes. Which statement is correct?

A) Oxidation takes place at the anode, which is positively charged.

B) Oxidation takes place at the anode, which is negatively charged.

C) Oxidation takes place at the cathode, which is positively charged.

D) Oxidation takes place at the cathode, which is negatively charged.

E) Oxidation take place at the dynode, which is uncharged.

Diff: 2

Section: 19.7

66) An electrolysis cell has two electrodes. Which statement is correct?

A) Reduction takes place at the anode, which is positively charged.

B) Reduction takes place at the anode, which is negatively charged.

C) Reduction takes place at the cathode, which is positively charged.

D) Reduction takes place at the cathode, which is negatively charged.

E) Reduction takes place at the dynode, which is uncharged.

Diff: 2

Section: 19.7

67) Which metal can be prepared by electrolysis of an aqueous solution of one of its salts?

A) aluminum

B) copper

C) magnesium

D) potassium

E) sodium

Diff: 2

Section: 19.7

68) The products of the electrolysis of molten magnesium chloride using platinum electrodes are

A) hydrogen gas and chlorine gas.

B) magnesium metal and chlorine gas.

C) magnesium metal and oxygen gas.

D) magnesium metal and hydroxide ions.

E) chlorine gas and platinum-magnesium alloy.

Diff: 2

Section: 19.7

69) The products of the electrolysis of aqueous magnesium chloride using platinum electrodes are

A) magnesium metal and chlorine gas.

B) magnesium metal and oxygen gas.

C) magnesium metal and hydroxide ions.

D) hydrogen gas and chlorine gas.

E) chlorine gas and platinum-magnesium alloy.

Diff: 2

Section: 19.7

70) When molten sodium chloride is electrolyzed, a gas is observed to form at the anode. The gas is

A) chlorine.

B) hydrogen.

C) hydrogen peroxide.

D) oxygen.

E) sodium.

Diff: 2

Section: 19.7

71) When an aqueous solution of copper sulfate is electrolyzed, a gas is observed to form at the anode. The gas is

A) hydrogen.

B) hydrogen sulfide.

C) hydrogen peroxide.

D) oxygen.

E) sulfur dioxide.

Diff: 2

Section: 19.7

72) When an aqueous solution of sodium sulfate is electrolyzed, a gas is observed to form at the anode. The gas is

A) hydrogen.

B) hydrogen sulfide.

C) hydrogen peroxide.

D) oxygen.

E) sulfur dioxide.

Diff: 2

Section: 19.7

73) When an aqueous solution of sodium sulfate is electrolyzed, a gas is observed to form at the cathode. The gas is

A) hydrogen.

B) hydrogen sulfide.

C) hydrogen peroxide.

D) oxygen.

E) sulfur dioxide.

Diff: 2

Section: 19.7

74) When an aqueous solution of magnesium sulfate is electrolyzed, what product is formed at the cathode?

A) hydrogen

B) hydrogen sulfide

C) magnesium

D) oxygen

E) sulfur dioxide

Diff: 2

Section: 19.7

75) When an aqueous solution of nickel sulfate is electrolyzed, what product is formed at the anode?

A) hydrogen

B) hydrogen sulfide

C) nickel

D) oxygen

E) sulfur dioxide

Diff: 2

Section: 19.7

76) The half-reaction that occurs at the cathode during electrolysis of aqueous sodium iodide solution is:

A) 2 H2O(l) + 2 e- → H2(g) + 2 OH-(aq)

B) I2(aq) + 2 e- → 2 I-(aq)

C) 2 I-(aq) → I2(aq) + 2 e-

D) Na+(aq) + e- → Na(s)

E) Na(s) → Na+(aq) + e-

Diff: 2

Section: 19.7

77) The half-reaction that occurs at the cathode during electrolysis of aqueous CuCl2 solution is:

A) Cl2(g) + 2 e- → 2 Cl-(aq)

B) 2 Cl(aq) → Cl2(g) + 2 e-

C) Cu2+(aq) + 2 e- → Cu(s)

D) Cu+(aq) + e- → Cu(s)

E) 2 H2O + 2 e- → H2(g) + 2 OH-(aq)

Diff: 2

Section: 19.7

78) The half-reaction that should occur at the anode during electrolysis of aqueous potassium bromide solution is:

A) Br2(g) + 2 e- → 2 Br-(aq)

B) 2 Br-(aq) → Br2(l) + 2 e-

C) 2 H2O → O2(g) + 4 H+(aq) + 4 e-

D) 2 H+(aq) + e- → H2(g)

E) Na+(aq) + e- → Na(s)

Diff: 2

Section: 19.7

79) The SI unit for electric current is the

A) ampere.

B) coulomb.

C) volt.

D) joule.

E) watt.

Diff: 2

Section: 19.7

80) Using the same current and similar conditions, which will require the shorter length of time?

A) Depositing 0.10 mol Ag from a Ag+ solution

B) Depositing 0.10 mol Cr from a Cr3+ solution

C) Depositing 0.10 mol Cu from a Cu2+ solution

D) Depositing 0.20 mol Cu from a Cu2+ solution

E) They should all take the same time.

Diff: 1

Section: 19.8

81) Using the same current and similar conditions, which will require the longest length of time?

A) Depositing 0.20 mol Ag from a Ag+ solution

B) Depositing 0.10 mol Cr from a Cr3+ solution

C) Depositing 0.10 mol Cu from a Cu2+ solution

D) Depositing 0.20 mol Cu from a Cu2+ solution

E) They should all take the same time.

Diff: 1

Section: 19.8

82) Using the same current and similar conditions, which will require the longest length of time?

A) Depositing 0.20 mol Ag from a Ag+ solution

B) Depositing 0.20 mol Cr from a Cr3+ solution

C) Depositing 0.20 mol Cu from a Cu2+ solution

D) Depositing 0.10 mol Cu from a Al3+ solution

E) They should all take the same time.

Diff: 1

Section: 19.8

83) Using the same current and similar conditions, which will require the shortest length of time?

A) Depositing 21.6 g Ag from a Ag+ solution

B) Depositing 10.4 g Cr from a Cr3+ solution

C) Depositing 12.7 g Cu from a Cu2+ solution

D) Depositing 8.8 g Ni from a Ni2+ solution

E) Depositing 2.7 g Al from a Al3+ solution

Diff: 2

Section: 19.8

84) Using the same current and similar conditions, which will require the longest length of time?

A) Depositing 21.6 g Ag from a Ag+ solution

B) Depositing 10.4 g Cr from a Cr3+ solution

C) Depositing 12.7 g Cu from a Cu2+ solution

D) Depositing 8.8 g Ni from a Ni2+ solution

E) Depositing 2.7 g Al from a Al3+ solution

Diff: 2

Section: 19.8

85) The electric charge can be calculated as

A) the product of current in amperes by time in seconds.

B) the product of potential in V by time in seconds.

C) the product of power in watts by time in seconds.

D) the product of energy in joules by time in seconds.

E) the product of V times coulombs.

Diff: 1

Section: 19.8

86) How many coulombs of charge are required for the reduction of 0.20 mol of Cr3+ ions to Cr metal?

A) 0.60 coulombs

B) 3.0 coulombs

C) 2.9 × 104 coulombs

D) 5.8 × 104 coulombs

E) 9.65 × 104 coulombs

Diff: 1

Section: 19.8

87) How many coulombs of charge are required for the reduction of 0.30 mol Co2+ to Co metal?

A) 0.60 coulombs

B) 0.30 coulombs

C) 5.8 × 104 coulombs

D) 2.9 × 104 coulombs

E) 9.6 × 103 coulombs

Diff: 1

Section: 19.8

88) A metal object is to be gold-plated by an electrolytic procedure using aqueous Au(CN)4, a gold(III) complex ion, as the electrolyte. Calculate the number of grams of gold deposited in 3.00 minutes by a constant current of 10.0 amperes. (Au: 196.97 g/mol)

A) 0.00689 gram

B) 1.22 gram

C) 1.32 gram

D) 1.77 gram

E) 3.55 gram

Diff: 2

Section: 19.8

89) How long will it take to produce 78.0 g of Al metal by the reduction of Al3+ ions in an electrolytic cell with a current of 2.00 amperes? (Al: 26.98 g/mol)

A) 419 seconds

B) 4.34 hours

C) 13.0 hours

D) 116 hours

E) 6.98 × 103 hours

Diff: 2

Section: 19.8

90) How many grams of nickel would be electroplated by passing a constant current of 7.20 amperes through a solution of NiSO4 for 90.0 minutes? (Ni: 58.69 g/mol)

A) 0.200 g

B) 0.400 g

C) 11.8 g

D) 23.7 g

E) 47.3 g

Diff: 2

Section: 19.8

91) How many coulombs of electrical charge must pass through an electrolytic cell to reduce 0.44 mol of Ca2+ ion to calcium metal? (Ca: 40.08 g/mol)

A) 0.88 coulomb

B) 2.1 × 104 coulomb

C) 4.25 × 104 coulomb

D) 8.5 × 104 coulomb

E) 1.93 × 104 coulomb

Diff: 2

Section: 19.8

92) How many grams of chromium would be electroplated by passing a constant current of 5.2 amperes through a solution containing chromium (III) sulfate for 45.0 minutes?

(Cr: 52.00 g/mol)

A) 9.3 × 10-4 g

B) 0.042 g

C) 24 g

D) 2.5 g

E) 2.3 × 1010 g

Diff: 2

Section: 19.8

93) How many coulombs would be required to electroplate 35.0 grams of chromium by passing an electrical current through a solution containing aqueous CrCl3?

(Cr: 52.00 g/mol, Cl: 35.453 g/mol)

A) 1.60 × 103 coulomb

B) 6.40 × 104 coulomb

C) 6.50 × 104 coulomb

D) 1.95 × 105 coulomb

E) 1.01 × 107 coulomb

Diff: 2

Section: 19.8

94) How many minutes would be required to electroplate 25.0 grams of chromium by passing a constant current of 4.80 amperes through a solution containing CrCl3?

(Cr: 52.00 g/mol, Cl: 35.453 g/mol)

A) 161 minutes

B) 322 minutes

C) 4.80 × 102 minutes

D) 1.11 × 104 minutes

E) 2.01 × 104 minutes

Diff: 2

Section: 19.8

95) During electrolysis, a current of 3.20 amperes was used for a period of 2 hours and 49 minutes. How many moles of e- flowed through the electrolytic cell?

A) 9.06

B) 0.336

C) 0.326

D) 0.240

E) 0.105

Diff: 2

Section: 19.8

96) How long would it take to deposit 50.0 g of silver metal from a solution containing Ag+ ions, using a current of 3.00 amperes? (Ag: 107.87 g/mol)

A) 41.4 minutes

B) 124 minutes

C) 249 minutes

D) 497 minutes

E) 1490 minutes

Diff: 2

Section: 19.8

97) How many liters of chlorine gas, measured at STP, would be produced by the electrolysis of molten sodium chloride, using a current of 3.50 amperes for 46 minutes?

Hint: Remember that 1 mole of an ideal gas occupies a volume of 22. L at STP.

A) 1.12 liters

B) 2.24 liters

C) 11.2 liters

D) 22.4 liters

E) 44.8 liters

Diff: 3

Section: 19.8

98) How many liters of chlorine gas, measured at STP, would be produced by the electrolysis of molten sodium chloride, using a current of 2.50 amperes for 56 minutes?

Hint: Remember that 1 mole of an ideal gas occupies a volume of 22. L at STP.

A) 4.26 liters

B) 0.975 liters

C) 2.13 liters

D) 0.0177 liters

E) 0.0355 liters

Diff: 3

Section: 19.8

99) When copper is refined using electrolysis, which process occurs?

A) Impure copper goes into solution at the anode, and pure copper plates out on the cathode.

B) Impure copper goes into solution at the cathode, and pure copper plates out on the anode.

C) Pure copper goes into solution from the anode and falls to the bottom of the tank.

D) Pure copper goes into solution from the cathode and falls to the bottom of the tank.

E) Pure copper on the bottom of the tank goes into solution and plates out on the cathode.

Diff: 1

Section: 19.9

100) Why is cryolite, Na3AlF6, mixed with aluminum oxide prior to electrolysis to produce Al?

A) because sodium is produced which in turn displaces the aluminum.

B) the aluminum in cryolite is easier to reduce than that in aluminum oxide.

C) to provide more fluoride ions.

D) to reduce the temperature required to melt the aluminum oxide.

E) to shift the equilibrium to the right.

Diff: 1

Section: 19.9

101) When aqueous brine is electrolyzed, the products are:

A) chlorine gas, hydrogen gas, and sodium hydroxide solution

B) chlorine gas and sodium metal

C) Na2O and HCl

D) NaCl and sodium metal

E) NaClO3 and NaClO2

Diff: 1

Section: 19.9

102) The Downs cell is used in the electrolytic manufacture of

A) chlorine gas and sodium hydroxide.

B) hydrogen gas and chlorine gas.

C) hydrogen gas and sodium hydroxide.

D) sodium metal and chlorine gas.

E) sodium metal and sodium hydroxide.

Diff: 1

Section: 19.9

103) Iron objects, such as storage tanks and underground pipelines, can be protected from corrosion, by connecting them to a piece of

A) copper.

B) lead.

C) zinc.

D) silver.

E) tin.

Diff: 2

Section: Chemistry Outside the Classroom 19.1

104) Steel objects that are exposed to weather can be protected from corrosion by a coating of ________, a process called galvanizing.

A) copper

B) lead

C) tin

D) cobalt

E) zinc

Diff: 2

Section: Chemistry Outside the Classroom 19.1

105) The cathode in a galvanic cell has a ________ polarity.

Diff: 1

Section: 19.1

106) The anode in a galvanic cell has a ________ polarity.

Diff: 1

Section: 19.1

107) The salt bridge in a galvanic cell allows the solutions in the half-cells to remain electrically neutral by allowing the transfer of ________ between the half-cells.

Diff: 1

Section: 19.1

108) The Faraday constant is equal to the ________ on 1 mole of electrons.

Diff: 1

Section: 19.1

109) What is the purpose of an electrolyte in an electrochemical reaction?

Diff: 1

Section: 19.1

110) Given the following notation for an electrochemical cell

Pt(s) | H2(g) | H+(aq) || Ag+(aq) | Ag(s),

what is the balanced overall (net) cell reaction?

Diff: 1

Section: 19.1

111) Sketch a galvanic cell with metallic zinc (E° = -0.76 V) and metallic tin (E° = -0.14 V) electrodes. Make sure to label the anode and cathode, write the two half-reactions, and suggest a good electrolyte.

Zn2+(aq) + 2e- → Zn(s) E° = -0.76 V

Sn2+(aq) + 2e- → Sn(s) E° = -0.14 V

We need to utilize the equation E°cell = E°cathode - E°anode so that we end up with a positive E°cell. We arrive at the following equation:

Zn(s) + Sn2+(aq) Zn2+(aq) + Sn(s) Eocell = 0.62 V

This tells us the cathode must be the tin electrode, and the anode must be the zinc electrode. A sketch of the system is shown below. A suitable electrolyte would be Na2SO4.)

Diff: 2

Section: 19.1

112) Consider an electrochemical cell based on the following cell diagram:

In | In+(aq) || Br-(aq),Br2(g) | Pt

If the standard cell potential is 1.21 V, what is the standard reduction potential for Indium (In)?

Diff: 2

Section: 19.2

113) Using these metal ion/metal standard reduction potentials

Fe2+(aq)|Fe(s) Zn2+(aq)|Zn(s) Cr3+(aq)|Cr(s) Cu2+(aq)/Cu(s)

-0.44 V -0.76 V -0.74 V +0.34 V

A galvanic cell is composed of these two half-cells:

Cr3+(aq) | Cr(s)

Cu2+(aq) | Cu(s)

What is the standard reduction potential for the cell reaction of this galvanic cell?

Diff: 2

Section: 19.2

114) If the measured voltage of the cell Zn(s) | Zn2+(aq) || Ag+(aq) | Ag(s) is 1.37 V when the concentration of Zn2+ ion is 0.010 M, what is the Ag+ ion concentration?

Diff: 2

Section: 19.5

115) Consider the following reaction: 2Fe2+(aq) + Cu2+ → 2Fe3+(aq) + Cu. When the reaction comes to equilibrium, what is the cell voltage?

Diff: 2

Section: 19.4

116) Consider the following reaction: 2Cu+(aq) + Ni2+ → 2Cu2+(aq) + Ni(s). When the reaction comes to equilibrium, what is the cell voltage?

Diff: 2

Section: 19.4

117) In the lead storage battery, the electrolyte is a solution of ________.

Diff: 1

Section: 19.6

118) In lithium ion batteries, lithium ions are able to slip between layers in certain crystals, in a process called ________.

Diff: 1

Section: 19.6

119) Semiconductor material is ________ with another element, to increase the material's ability to conduct electricity.

Diff: 1

Section: 19.6

120) In the zinc-carbon dry cell, the electrode at which oxidation takes place has a ________ charge.

Diff: 2

Section: 19.6

121) How many grams of Ag are deposited on the cathode of an electrolytic cell, if a current of 2.00 A is applied to a solution of AgNO3 for 10.00 minutes?

Ag+ + e- → Ag(s)

2.00 A × 10 min × = 1.20 × 103 C

1.20 × 103C × = 0.0124 mol e-

0.0124 mol e- × × = 1.34 g Ag

Diff: 2

Section: 19.8

122) How long would it take to deposit 3.00 grams of Cu from a CuSO4 solution, if a current of 1.5 A is applied? (Give your answer in hours)

Cu2+ + 2e- → Cu(s)

3.00 g Cu × × = 0.0944 mol e-

0.0944 mol e- × = 9.11 × C or A∙ s

9.11 × A∙ s × = 1.69 hours

Diff: 2

Section: 19.8

123) In the Hall-Héroult process for manufacture of aluminum, the anodes are constructed of ________.

Diff: 1

Section: 19.9

124) For every oxidation in a chemical process, there must be an associated reduction.

Diff: 1

Section: 19.1

125) When adding half reactions and canceling out electrons, multiply the reduction potential of each half reaction by any coefficients.

Diff: 1

Section: 19.1

126) Lithium metal can be produced readily by the electrolysis of aqueous solutions of lithium nitrate, or lithium chloride.

Diff: 2

Section: 19.7

127) During the electrolysis of water, oxygen gas is formed at the anode.

Diff: 2

Section: 19.7

128) When an electrical current is passed through fused (molten) magnesium chloride, chlorine gas is produced at the negatively charged electrode.

Diff: 2

Section: 19.7

129) If a brine bath is stirred during its electrolysis, a different set of products are obtained than would be obtained by electrolysis of an unstirred bath.

Diff: 2

Section: 19.7

130) Using a table of standard electrode potentials and the information given below, decide which of the following statements is completely true.

(Note that some imaginary elements are being used in these questions).

(a) Nt2+(aq) + 2e- Nt(s) -0.422 V

(b) NtSO4(s) + 2e- Nt(s) + SO42-(aq) -0.715 V

(c) Nt(NH3)42+(aq) + 2e- Nt(s) + 4NH3(aq) -0.922 V

(d) Nt(CN)42-(aq) + 2e- Nt(s) + 4CN-(aq) -1.014 V

A. Nt2+ can oxidize Cu2+, and Al3+ can reduce H+.

B. Br2 can oxidize Ni, and H2 can reduce Cl-.

C. Cu2+ can oxidize H2, and Nt can reduce Cl2.

Additional information for each answer choice:

A: If Nt2+ can oxidize Cu2+, then the potential to reduce Ni2+ must be greater than the reduction potential for Cu2+. But the key to this problem is to realize that oxidation of Cu2+ would require the additional loss of electrons from Cu2+, which is difficult (there is no listed equation for this for that reason). Therefore Ni2+ will not oxidize Cu2+

A similar argument can be made for Al3+. For Al3+ to reduce H+, it must be oxidized, which would again require loss of electrons from a highly charged species.

B: Br2 can oxidize Nt, as the reduction potential of Br2 is higher than Ni in the reduction table. If you calculated the reduction potential for this, you would get a positive value.

If the second part was correct, H2 can reduce Cl-, then you would be reducing Cl-, which would result in a Cl-2, and that is not a favorable process (not on table).

C: Cu2+ can oxidize H2 as the reduction potential for Cu2+ is higher than that for H2 on the table (hydrogen is a value of zero). Ni can also reduce Cl2, because the reduction potential of Cl2 is larger than that of Nt.

One key for this type of problem is to write out the half reaction that refers to the oxidation or reduction process mentioned and see if you can compare the potentials for these equations if possible. If the reduction potential is not available, you need to logically ask if the process would be easy to do or not.

Diff: 2

Section: 19.2

131) Using a table of standard electrode potentials, decide which of the following statements is completely true.

A. Cl2 can oxidize Cd, and Al3+ can reduce Cu2+.

B. Cd2+ can oxidize Cu, and Cd2+ can reduce Cl2.

C. Cu2+ can oxidize Cd, and Cd can reduce Al3+.

Additional information for each answer choice:

A: If Cl2 can oxidize Cd, then Cl2 would be reduced and Cd would be oxidized. This results in the following reactions:

Cl2(g) + 2e- → 2Cl-(aq) E° = +1.36

Cd(s) → Cd2+(aq) + 2e- E° = -(-0.40) (Sign switched for oxidation)

Which gives a potential of +0.96 V, so this is spontaneous and should happen.

If Al can reduce Cu2+, then Al would be oxidized and Cu2+ would be reduced. This results in the following reactions:

Al(s) → Al3+(aq) + 3e- E° = -(-1.66) (Sign switched for oxidation)

Cu2+(aq) + 2e- → Cu(s) E° = +0.34

Which gives a potential of +2.00 V, so this is spontaneous and should happen.

B: If Cd2+ can oxidize Cu, then Cd2+ would be reduced and Cu would be oxidized. This results in the following reactions:

Cd2+(aq) + 2e- → Cd(s) E° = -0.40

Cu(s) → Cu2+(aq) + 2e- E° = -(0.34) (Sign switched for oxidation)

Which gives a potential of -0.74 V, so this is not spontaneous and should not happen.

If Cd2+ can reduce Cl2, then Cd2+ would be oxidized and Cl2 would be reduced. This results in the following reactions:

Cd(s) → Cd2+(aq) + 2e- E° = -(-0.40) (Sign switched for oxidation)

Cl2(g) + 2e- → 2Cl-(aq) E° = +1.36

Which gives a potential of +1.76 V, so this is spontaneous and should happen.

But, because the first is not true, then this option is not correct.

C: If Cu2+ can oxidize Cd, then Cu2+ would be reduced and Cd would be oxidized. This results in the following reactions:

Cu2+(aq) + 2e- → Cu(s) E° = -0.34

Cd(s) → Cd2+(aq) + 2e- E° = -(-0.40) (Sign switched for oxidation)

Which gives a potential of +0.06 V, so this is spontaneous and should happen.

If Cd can reduce Al3+, then Cd would be oxidized and Al3+ would be reduced. This results in the following reactions:

Cd(s) → Cd2+(aq) + 2e- E° = -(-0.40) (Sign switched for oxidation)

Al3+(aq) + 3e- → Al(s) E° = -1.66

Which gives a potential of -1.26 V, so this is not spontaneous and should not happen.

Because the second is not true, then this option is not correct.

Diff: 2

Section: 19.2

132) Which one of the following, when added to a piece of Cu(s) in solution, is capable of oxidizing Cu(s) to Cu2+(aq)?

A) I- (aq)

B) Ni2+ (s)

C) Ag+ (aq)

D) Al3+ (aq)

E) H2 (g)

Diff: 2

Section: 19.3

133) Use the standard reduction potentials below to determine the equilibrium constant for the following reaction:

Nt(NH3)42+(aq) + 4CN-(aq) Nt(CN)42-(aq) + 4NH3(aq)

Nt(NH3)42+(aq) + 2e- Nt(s) + NH3(aq) -0.922 V

Nt(CN)42-(aq) + 2e- Nt(s) + 4CN-(aq) -1.014 V

Note that the imaginary element Nt is being used in this reaction.

Diff: 2

Section: 19.4

134) The equation for the chemical reaction from which the expression for the solubility product of NtSO4(s) is derived is: NtSO4(s) Nt2+(aq) + SO42-(aq). Use this equation and the reduction potentials below to help you determine the solubility product of NtSO4 (s). (Note that the imaginary element Nt is being used).

(a) Nt2+(aq) + 2e- Nt(s) -0.422 V

(b) NtSO4(s) + 2e- Nt(s) + SO42-(aq) -0.715 V

(c) Nt(NH3)42+(aq) + 2e- Nt(s) + 4NH3(aq) -0.922 V

(d) Nt(CN)42-(aq) + 2e- Nt(s) + 4CN-(aq) -1.014 V

Diff: 2

Section: 19.4

135) Calculate the value of the equilibrium constant, Kc, for the reaction

Nt(CN)42-(aq) + SO42-(aq) NtSO4(s) + 4CN-(aq)

Hint: Use these standard potentials in answering this question. (Note that the imaginary element Nt is being used).

(a) Nt2+(aq) + 2e- Nt(s) -0.422 V

(b) NtSO4(s) + 2e- Nt(s) + SO42-(aq) -0.715 V

(c) Nt(NH3)42+(aq) + 2e- Nt(s) + 4NH3(aq) -0.922 V

(d) Nt(CN)42-(aq) + 2e- Nt(s) + 4CN-(aq) -1.014 V

Diff: 2

Section: 19.4

136) The solubility product constant for NtCrO4 is 4.50 × 10-12. Using the potentials provided in the listing above, calculate the value for the equilibrium constant, Kc, for the reaction

Nt(CN)42-(aq) + CrO42-(aq) NtCrO4(s) + 4CN-(aq)

Use these standard potentials in answering this question. (Note that imaginary elements are being used in these questions).

(a) Nt2+(aq) + 2e- Nt(s) -0.422 V

(b) NtSO4(s) + 2e- Nt(s) + SO42-(aq) -0.715 V

(c) Nt(NH3)42+(aq) + 2e- Nt(s) + 4NH3(aq) -0.922 V

(d) Nt(CN)42-(aq) + 2e- Nt(s) + 4CN-(aq) -1.014 V

Hint: Combine two reactions in the list to find a third reaction and an equilibrium constant that combines with the solubilization of NtCrO4 to get the target equation.

Diff: 3

Section: 19.4

137) What is value of the formation constant for the tetracyanonortonate ion, Nt(CN)42? The reaction is:

Nt2+(aq) + 4CN-(aq) Nt(CN)42-(aq)

Use these standard potentials in answering this question. (Note that an imaginary element is being used in this question).

(a) Nt2+(aq) + 2e- Nt(s) -0.422 V

(b) NtSO4(s) + 2e- Nt(s) + SO42-(aq) -0.715 V

(c) Nt(NH3)42+(aq) + 2e- Nt(s) + 4NH3(aq) -0.922 V

(d) Nt(CN)42-(aq) + 2e- Nt(s) + 4CN-(aq) -1.014 V

Diff: 2

Section: 19.4

138) Given the following standard reduction potentials,

Ag+(aq) + e- Ag(s) E° = 0.80 V

AgCN(s) + e- Ag(s) + CN-(aq) E° = -0.01 V

calculate the solubility product (Ksp) of AgCN at 25°C.

The first step in solving this problem is identify the equation that relates to the Ksp of AgCN(s).

AgCN(s) Ag+(aq) + CN-(aq)

This equation can be built from the two given equations by reversing equation 1. When reversing equation 1 you need to change the sign of E°.

Ag(s) Ag+(aq) + e- E° = -0.80 V

AgCN(s) + e- Ag(s) + CN-(aq) E° = -0.01 V

Adding these two equations then gives you the Ksp equation. If the two equations are additive, so are their reduction potentials.

So E°cell = -0.80 V + -0.01 V = -0.81 V

After E°cell is found, the equilibrium constant can be found using the equation

after solving for K.

This gives you the value of 2.1 × 10-14 for the Ksp of AgCN.)

Diff: 2

Section: 19.4

139) Calculate the cell emf for the following balanced reaction at 25°C:

Cu2+(0.10 M) + H2(1 atm) → Cu(s) + 2H+ (pH = 3.00)

Hint: Use the ideal gas law to find [H2], pH to find [H+], then apply the Nernst equation to find the cell emf.

Diff: 3

Section: 19.5

140) For the electrochemical cell

Ni(s) | Ni2+(1 M) || H+(1 M) | H2(1 atm) | Pt(s)

which one of the following changes will cause a decrease in the cell voltage?

A) Increase the pressure of H2 to 2.0 atm.

B) Decrease the mass of the nickel electrode.

C) Lower the pH of the cell electrolyte.

D) Decrease the concentration of Ni2+ ion.

E) None of these

Diff: 2

Section: 19.5

141) The measured voltage of the cell Pt(s) | H2 (1.0 atm) | H+(aq) (pH = ??) || Ag+(1.0 M) | Ag(s) is 1.02 V at 25°C. Calculate the pH of the solution.

Hint: Use the ideal gas law to find [H2], the apply the Nernst equation to find [H+] and then find pH.

Diff: 3

Section: 19.5

142) An old classmate who works for Eagle Metals Company on the night shift was supposed to do a gold plating job, in which he put 4.5000 troy ounces of gold (1 troy ounce = 31.103 gram) on a medallion using the reaction,

Au(CN)4-(aq) + 3 e- Au(s) + 4 CN-

He somehow got 4.8527 troy ounces of gold on the medallion instead of the correct amount. Calculate how many hours, minutes, and seconds he should run the cells with the leads reversed, using a current of 1.350 amperes, to fix the error and bring the medallions down to the correct weight.

Diff: 2

Section: 19.8

143) Note: Imaginary elements are used in the following question. An electrolysis study was carried to determine the atomic weight of nortonium from a study on the compound, nortonium nitrate. Another study had established that the formula was probably Nt(NO3)2. In the present study, electrolysis was carried on an aqueous sample of nortonium nitrate, and the quantities of the electrolysis products were carefully measured. At the anode, a gas was collected which, after correction for the vapor pressure of water, registered a volume of 94.25 mL when the temperature was 25.0 °C and the pressure was 748.5 torr. The mass of the nortonium that plated out on the cathode was measured as 836.2 mg. From this data, calculate a value for the atomic weight of nortonium.

Hint: Use the ideal gas law and balanced redox equation to find moles of nortonium, then find atomic weight.

Diff: 3

Section: 19.8

144) Note: Imaginary elements are used in the following question. An electrolysis study was carried to determine the atomic weight of nortonium from a study on the compound, nortonium chloride. Another study had established that the formula was probably NtCl2. In the present study, electrolysis was carried on an aqueous sample of nortonium chloride, and the quantity of electrical energy was carefully measured. A current measuring 1.550 amperes at a voltage of 12.25 V, was passed through the electrolysis with the aqueous nortonium chloride solution, starting at 2:00:00 P.M. and ending at 2:53:30 P.M., the same day. The amount of nortonium which plated out on the cathode measured 2.848 g. From this data, calculate a value for the atomic weight of nortonium.

Hint: Use dimensional analysis to find the moles of nortonium, then find atomic weight.

Diff: 3

Section: 19.8

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