Consider the following cell reaction: \[ \mathrm{Cd}_{(\mathrm{s})}+\mathrm{Hg}_{2} \mathrm{SO}_...

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Consider the following cell reaction:
\[
\mathrm{Cd}_{(\mathrm{s})}+\mathrm{Hg}_{2} \mathrm{SO}_{4(5)}+\frac{9}{5} \mathrm{H}_{2} \mathrm{O}_{(\mathrm{I})} \rightleftharpoons \mathrm{CdSO}_{4} \cdot \frac{9}{5} \mathrm{H}_{2} \mathrm{O}_{(\mathrm{s})}+2 \mathrm{Hg}_{(\mathrm{I})}
\]
The value of \( \mathrm{E}_{\text {cell }}^{0} \) is \( 4.315 \mathrm{~V} \) at \( 25^{\circ} \mathrm{C} \). If \( \Delta \mathrm{H}^{0}=-825.2 \mathrm{~kJ} \) \( \mathrm{mol}^{-1} \), the standard entropy change \( \Delta \mathrm{S}^{0} \) in \( \mathrm{J} \mathrm{K}^{-1} \) is (Nearest integer)
[Given: Faraday constant \( =96487 \mathrm{C} \mathrm{mol}^{-1} \) ]
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