Use the tabulated half-cell potentials below to calculate the equilibrium constant (K) for the following balanced redox reaction at 25°C.
Pb²⁺(aq) + Cu(s) → Pb(s) + Cu²⁺(aq)

Question

Use the tabulated half-cell potentials below to calculate the equilibrium constant (K) for the following balanced redox reaction at 25°C. Pb<sup>2+</sup>(aq) + Cu(s) → Pb(s) + Cu<sup>2+</sup>(aq) A) 7.9 × 10<sup>-8</sup> B) 8.9 × 10<sup>7</sup> C) 7.9 × 10<sup>15</sup> D) 1.3 × 10<sup>-16</sup> E) 1.1 × 10<sup>-8</sup>

Quizplus · Accepted Answer

We can use the Nernst equation to find the standard potential for the redox reaction:

E°cell = E°reduction (reduction) - E°reduction (oxidation)
E°cell = (-0.14 V) - (-0.13 V + 0.34 V)
E°cell = -0.15 V

Using the equation ∆G° = -nFE°cell and the ideal gas constant R = 8.31 J/mol*K, we can find ∆G°:

∆G° = -nFE°cell
∆G° = -(2 mol)(F)(-0.15 V)
∆G° = 0.4 kJ/mol

Then, we can use the equation ∆G° = -RTlnK to find the equilibrium constant K:

∆G° = -RTlnK
0.4 kJ/mol = -(8.31 J/mol*K)(298 K)lnK
lnK = -149.6
K = e^-149.6
K = 1.3 × 10^-65

Therefore, the best choice is D, which is the closest to the calculated value.

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