An electrochemical device is assembled by coupling two half-cells under standard conditions at 298 K.
Half-cell A: Ce4+(aq) + e⁻ → Ce3+(aq) E° = +1.70 V
Half-cell B: I2(s) + 2 e⁻ → 2 I⁻(aq) E° = +0.54 V
The cell is first allowed to operate spontaneously, and then, in a separate trial, an external potential of 2.50 V is applied across the same electrodes so that current reverses direction.
(a) For the spontaneous operation, write the balanced overall chemical equation, including states of matter, for the reaction that proceeds.
(b) Calculate the standard cell potential and the corresponding standard Gibbs free-energy change, in kilojoules per mole of reaction, for the spontaneous process.
(c) Based on your answer to part (b), classify the spontaneous cell as galvanic or electrolytic and justify your classification.
(d) When a 2.50 V external source is connected so that current is forced to run opposite to the spontaneous direction, identify which species is reduced and which is oxidized. Explain how the applied voltage changes the thermodynamic favorability of the reaction.
(e) Predict qualitatively how the concentrations of Ce4+, Ce3+, I2, and I⁻ change over time during the forced (non-spontaneous) operation and relate these changes to the sign of ΔG under the applied conditions.