Use the tabulated half-cell potentials below 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)

Question

Use the tabulated half-cell potentials below to calculate the equilibrium constant (K) for the following balanced redox reaction at 25°C. 3 I<sub>2</sub>(s) + 2 Fe(s) → 2 Fe<sup>3+</sup>(aq) + 6 I⁻(aq) A) 3.5 × 10<sup>-59</sup> B) 1.1 × 10<sup>17</sup> C) 2.4 × 10<sup>58</sup> D) 8.9 × 10<sup>-18</sup> E) 1.7 × 10<sup>29</sup>

Quizplus · Accepted Answer

1. Write the overall redox reaction: 
3 I2(s) + 2 Fe(s) → 2 FeS + 6 I-(aq)

2. Calculate the overall cell potential: 
E°cell = E°(cathode) - E°(anode) 
E°cell = 0.771 V - 0.535 V = 0.236 V

3. Calculate the equilibrium constant using the Nernst equation: 
Ecell = E°cell - (RT/nF) ln Q 
0.236 V = 0.771 V - (0.025 V/K)(298 K)/(2 mol e-/mol Fe)(ln K) 
ln K = 9.207 
K = e^9.207 
K = 2.4 × 10^158

Therefore, the best choice is C.

Use the Tabulated Half-Cell Potentials Below to Calculate the Equilibrium