Question

$d\left(\Delta G\right)=-\left(\Delta S_{\text{reaction}}\right).dT$ $\Delta G=\Delta H-T\Delta S$ $\Delta G=\Delta H+T{\left(\frac{d\left(\Delta G\right)}{dT}\right)}_P$
Calculate $\Delta S$ for the given cell reaction in the previous question.

• A. $-73.53{\text{JK}}^{-1}{\text{mol}}^{-1}$
• B. $83.53{\text{JK}}^{-1}{\text{mol}}^{-1}$
• C. $100{\text{JK}}^{-1}{\text{mol}}^{-1}$
• D. None of these

Answer 1

The correct option is A $-73.53{\text{JK}}^{-1}{\text{mol}}^{-1}$

$d\left(\Delta G\right)=\Delta H-T\Delta S$

$\Delta G=-nFE$

$-2\times 96500\times 1.015=-21800-300\Delta S$

$-195895+21800=-300\Delta S$

$\frac{-22105}{300}=\Delta S$

$\Delta S=-73.53J{K}^{-1}mo{l}^{-1}$