Q: Use the tabulated half-cell potentials to calculate $$\Delta G^\circ$$ for the following balanced redox reaction: $$\mathrm{I_2$ (s) + 2 Fe$$^{2+}$$ (aq) \rightarrow 2 Fe$$^{3+}$$ (aq) + 2 I^- (aq)}$$ $$E^\circ_{\mathrm{I_2$/I^-}} = +0.54\ \mathrm{V}$$ $$E^\circ_{\mathrm{Fe$$^{3+}$$/Fe^{2+}$$}} = +0.77\ \mathrm{V}$

$$\Delta G^\circ = -n F E^\circ_{cell} = -2 \times 96485 \times (0.77 - 0.54) = -44.3\ \mathrm{kJ}$$

Q: Use the tabulated half-cell potentials to calculate the equilibrium constant ($$K) $$for the following balanced redox reaction at 25$$^\circC$$: $$\mathrm{I_2$ (s) + 2 Fe$$^{2+}$$ (aq) \rightarrow 2 Fe$$^{3+}$$ (aq) + 2 I^- (aq)}$$ $$E^\circ_{cell} = 0.23\ \mathrm{V}$$

$$K = \exp\left(\frac{n F E^\circ_{cell}}{RT}\right) = \exp\left(\frac{2 \times 96485 \times 0.23}{8.314 \times 298}\right) \approx 1.1 \times 10^7$

Question 1

Which element is being reduced in the following redox reaction?

$$\mathrm{$$MnO_4$$^- (aq) + H_2C_2O_4 (aq) \rightarrow Mn$$^{2+}$$ (aq) + CO_2$ (g)}$$

Accepted Answer

C

Question 2

What element is being oxidized in the following redox reaction?

$$\mathrm{$$MnO_4$$^- (aq) + H_2C_2O_4 (aq) \rightarrow Mn$$^{2+}$$ (aq) + CO_2$ (g)}$$

Accepted Answer

C

Question 3

In the reaction $$\mathrm{Cr_2O_7$$^{2-}$$ + HBr \rightarrow Cr$$^{3+}$$ + Br_2$}$$, which element is being oxidized?

Accepted Answer

Br

Question 4

Balance the following redox reaction in acidic solution. What are the coefficients in front of $$\mathrm{H_2C_2O_4}$$ and $$\mathrm{H^+} in $$the balanced reaction?

$$\mathrm{$$MnO_4$$^- (aq) + H_2C_2O_4 (aq) \rightarrow Mn$$^{2+}$$ (aq) + CO_2$ (g)}$$

Accepted Answer

$$\mathrm{H_2C_2O_4 = 5, H^+ = 16}$$

Question 5

Balance the following redox reaction in basic solution. What are the coefficients in front of $$\mathrm{Cl_2$}$$ and $$\mathrm{OH^-} in $$the balanced reaction?

$$\mathrm{Cl_2$ (g) + OH^- (aq) \rightarrow $$ClO_2$$^- (aq) + Cl^- (aq)}$$

Accepted Answer

$$\mathrm{Cl_2$ = 1, OH^- = 6}$$

Question 6

A voltaic cell:

Accepted Answer

produces electrical current from a spontaneous chemical reaction.

Question 7

Identify the location of oxidation in an electrochemical cell.

Accepted Answer

the anode

Question 8

Identify the location of reduction in an electrochemical cell.

Accepted Answer

the cathode

Question 9

Determine the cell notation for the redox reaction below:

$$\mathrm{3 Br_2$ (l) + 2 Fe$$^{2+}$$ (aq) \rightarrow 6 Br^- (aq) + 2 Fe$$^{3+}$$ (aq)}$$

Accepted Answer

$$\mathrm{Br_2$ (l) | Br^- (aq) || Fe$$^{2+}$$ (aq) | Fe$$^{3+}$$ (aq)}$$

Question 10

Use line notation to represent the following electrochemical cell:

$$\mathrm{O_2$ (g) + 4 H^+ (aq) + 2 Zn (s) \rightarrow 2 H_2O (l) + 2 Zn$$^{2+}$$ (aq)}$$

Accepted Answer

$$\mathrm{Zn (s) | Zn$$^{2+}$$ (aq) || O_2$ (g) | H^+ (aq) | H_2O (l)}$$

Question 11

Given the following redox reaction, write out the half reaction that occurs at the anode:

$$\mathrm{O_2$ (g) + 4 H^+ (aq) + 2 Zn (s) \rightarrow 2 H_2O (l) + 2 Zn$$^{2+}$$ (aq)}$$

Accepted Answer

$$\mathrm{Zn (s) \rightarrow Zn$$^{2+}$$ (aq) + 2 e^-}$$

Question 12

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$$^\circC. ($$The equation is balanced.)

$$\mathrm{Cl_2$ (g) + 2 Fe$$^{2+}$$ (aq) \rightarrow 2 Cl^- (aq) + 2 Fe$$^{3+}$$ (aq)}$$

$$E^\circ_{\mathrm{Cl_2$/Cl^-}} = +1.36\ \mathrm{V}$$
$$E^\circ_{\mathrm{Fe$$^{3+}$$/Fe^{2+}$$}} = +0.77\ \mathrm{V}$

Accepted Answer

$$1.36\ \mathrm{V} - 0.77\ \mathrm{V} = 0.59\ \mathrm{V}$$

Question 13

Identify the characteristics of a nonspontaneous reaction.

Accepted Answer

$$\Delta G \u003e 0$$

Question 14

Use the tabulated half-cell potentials to calculate $$\Delta G^\circ$$ for the following balanced redox reaction:

$$\mathrm{I_2$ (s) + 2 Fe$$^{2+}$$ (aq) \rightarrow 2 Fe$$^{3+}$$ (aq) + 2 I^- (aq)}$$

$$E^\circ_{\mathrm{I_2$/I^-}} = +0.54\ \mathrm{V}$$
$$E^\circ_{\mathrm{Fe$$^{3+}$$/Fe^{2+}$$}} = +0.77\ \mathrm{V}$

Accepted Answer

$$\Delta G^\circ = -n F E^\circ_{cell} = -2 	imes 96485 	imes (0.77 - 0.54) = -44.3\ \mathrm{kJ}$$

Question 15

Use the tabulated half-cell potentials to calculate the equilibrium constant ($$K) $$for the following balanced redox reaction at 25$$^\circC$$:

$$\mathrm{I_2$ (s) + 2 Fe$$^{2+}$$ (aq) \rightarrow 2 Fe$$^{3+}$$ (aq) + 2 I^- (aq)}$$

$$E^\circ_{cell} = 0.23\ \mathrm{V}$$

Accepted Answer

$$K = \exp\left(\frac{n F E^\circ_{cell}}{RT}ight) = \exp\left(\frac{2 	imes 96485 	imes 0.23}{8.314 	imes 298}ight) \approx 1.1 	imes 10^7$

Question 16

Calculate the cell potential for the following reaction that takes place in an electrochemical cell at 25$$^\circC$$:

$$\mathrm{Mg(s) | Mg$$^{2+}$$ (aq, 2.74 M) || Cu$$^{2+}$$ (aq, 0.0033 M) | Cu(s)}$$

$$E^\circ_{\mathrm{Cu$$^{2+}$$/Cu}} = +0.34\ \mathrm{V}$$
$$E^\circ_{\mathrm{Mg$$^{2+}$$/Mg}} = -2.37\ \mathrm{V}$$

Accepted Answer

$$E_{cell} = E^\circ_{cell} - \frac{0.0592}{2} \log \left( \frac{[Mg$$^{2+}$$]}{[Cu$$^{2+}$$]} \right) = 2.71\ \mathrm{V}$$

Question 17

What mass of silver can be plated onto an object in 67 minutes at 8.70 A of current?

$$\mathrm{Ag^+ (aq) + e^- \rightarrow Ag (s)}$$

Accepted Answer

56.9 g

Question 18

Copper can be electroplated at the cathode of an electrolysis cell by the half-reaction:

$$\mathrm{Cu$$^{2+}$$ (aq) + 2e^- \rightarrow Cu (s)}$$

How much time would it take for 235 mg of copper to be plated at a current of 7.5 A?

Accepted Answer

51.5 s

Question 19

A sacrificial electrode is composed of a metal that:

Accepted Answer

oxidizes more easily than iron.

Question 20

The standard cell potential for the cell using the overall cell reaction below is +2.20 V.

$$\mathrm{2 Al(s) + 3 I_2$ (s) ightarrow 2 Al$$^{3+}$$ (aq) + 6 I^- (aq)}$$

The cell generated by the cell when $$[Al$$^{3+}$$] = 1.0 	imes 10$$^{-3}$$ M and $$[I^-] = 0.20$$ M is ... V.

Accepted Answer

2.11

Electrochemistry: Redox Reactions, Electrochemical Cells, and Applications

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