Emf And Equilibrium Constant

The link between standard cell potential and the equilibrium constant arises from combining ΔG° = -nFE°_cell with ΔG° = -RT ln K, giving log K = (n × E°_cell)/0.0591 at 298 K. Here n = 2 and E°_cell = 0.295 V, so log K = (2 × 0.295)/0.0591 = 0.590/0.0591 ≈ 9.98, which is essentially 10. Therefore K ≈ 10^10. The large value confirms the reaction proceeds almost completely to products, consistent with the positive standard potential. Option 10 underestimates by treating the exponent as the value itself. Option 10^5 would require E°_cell near 0.15 V. Option 10^-10 wrongly takes the potential as negative, which would describe a non-spontaneous reaction. This connection between thermodynamics and electrochemistry is a recurring JEE Advanced theme. This concept also bridges to Chemical Thermodynamics and Equilibrium, so mastering it strengthens performance on linked questions from those topics as well. Such questions reward conceptual clarity, since a student who truly grasps equilibrium constant can solve many superficially different variants with the same approach. Plausibility check: a positive E°_cell must give K greater than one, and the magnitude near 10^10 matches the moderate two-electron potential supplied.