The standard e.m.f of a cell, involving one electron change is found to be 0.591 V at `25^(@) C`. The equilibrium constant of the reaction is : `(F=96,500C mol^(-1)`: R=8.314 `Jk^(-1)mol^(-1)`

Options

A. `(1.0xx10^(1)`
B. `(1.0xx10^(30)`
C. `1.0xx10^(10)`
D. `(1.0xx10^(5)`

Explanation

To find the equilibrium constant (Kc) of the reaction given the standard e.m.f (E°) of a cell, we can use the Nernst equation in its logarithmic form. Here’s a step-by-step solution:

### Step 1: Write the Nernst Equation
The Nernst equation in terms of standard cell potential (E°) and the equilibrium constant (Kc) is given by:

\[ E^\circ = \frac{RT}{nF} \ln K_c \]
...