The standard e.m.f. of a cell,involving one e | vedclass

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(C) The relationship between the standard $e.m.f.$ $(E^{\circ}_{cell})$ and the equilibrium constant $(K_{eq})$ is given by the Nernst equation at equ

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$1.0 	imes 10^{10}$

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(C) The relationship between the standard $e.m.f.$ $(E^{\circ}_{cell})$ and the equilibrium constant $(K_{eq})$ is given by the Nernst equation at equilibrium: $E^{\circ}_{cell} = \frac{2.303 RT}{nF} \log K_{eq}$ Given: $E^{\circ}_{cell} = 0.591 \ V$ $n = 1$ $T = 25^{\circ} C = 298 \ K$ At $298 \ K$,the value of $\frac{2.303 RT}{F} \approx 0.0591 \ V$. Substituting the values: $0.591 = \frac{0.0591}{1} \log K_{eq}$ $\log K_{eq} = \frac{0.591}{0.0591} = 10$ $K_{eq} = 10^{10} = 1.0 	imes 10^{10}$ Therefore,the correct option is $C$.

The standard $e.m.f.$ of a cell,involving one electron change is found to be $0.591 \ V$ at $25^{\circ} C$. The equilibrium constant of the reaction is :
$(F=96500 \ C \ mol^{-1} ; R=8.314 \ JK^{-1} \ mol^{-1})$

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The standard $e.m.f.$ of a cell,involving one electron change is found to be $0.591 \ V$ at $25^{\circ} C$. The equilibrium constant of the reaction is :$(F=96500 \ C \ mol^{-1} ; R=8.314 \ JK^{-1} \ mol^{-1})$