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. The relationship between the standard e.m.f and the equilibrium constant is given by the equation:

\[ E^\circ = \frac{0.059}{n} \log K_c \]

Where:
- \( E^\circ \) is the standard e.m.f of the cell (in volts).
- \( n \) is the number of electrons transferred in the reaction.
- \( K_c \) is the equilibrium constant.
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