Nernst Equation

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By Expert Faculty of Sri Chaitanya

Question

For the cell reaction:
$2 {Fe}_{(aq)}^{3 +} + 2 I_{(aq)}^{-} ⟶ 2 {Fe}_{(aq)}^{2 +} + I_{2 (aq)}$
${E^{o}}_{cell} = 0.24 V at 298 K$ . The standard Gibbs energy $(Δ_{r} G °)$ of the cell reaction is
$[Given that Faraday constant F = 96500 {Cmol}^{- 1}]$

Moderate

Solution

$The standard Gibbs' energy (ΔG °) = - nFE °_{cell} Value of n = 2 ∴ ΔG ° = - 2 \times 96500 \times 0.24 = - 46320 J = - 46.32 kJ / mol$

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Similar Questions

The reaction which occur in the Galvanic cell is
$\begin{matrix} {MnO}_{4}^{-} + 8 H^{+} + 5 {Fe}^{2 +} ⟶ {Mn}^{2 +} + 5 {Fe}^{3 +} + 4 H_{2} O \\ E^{\circ} ({MnO}_{4}^{-}, {Mn}^{2 +}, H^{+} ∣ Pt) = 1 .51 V and \\ E^{\circ} ({Fe}^{3 +}, {Fe}^{2 +} ∣ Pt) = 0 .77 V \end{matrix}$
(Q). If $[H_{3} O^{+}] = 2 .0 M, [{Fe}^{3 +}] = 2 .0 M, [{Fe}^{2 +}] = 0 .5 M$ $[{Mn}^{2 +}] = 0 .5 M, [{MnO}_{4}^{-}] = 1 .0 M$ Calculate the emf of the cell.