Standard electrode potentials for a few half cells are mentioned below :
$$\begin{aligned} & \mathrm{E}_{\mathrm{Cu}^{2+} / \mathrm{Cu}}^{\circ}=0.34 \mathrm{~V}, \mathrm{E}_{\mathrm{Zn}^{2+} / \mathrm{Zn}}^{\circ}=-0.76 \mathrm{~V} \\ & \mathrm{E}_{\mathrm{Ag}^{+} / \mathrm{Ag}}^{\circ}=0.80 \mathrm{~V}, \mathrm{E}_{\mathrm{Mg}^{2+} / \mathrm{Mg}}^{\circ}=-2.37 \mathrm{~V} \end{aligned}$$
Which one of the following cells gives the most negative value of $\Delta \mathrm{G}^{\circ}$ ?
Solution
<p>$$\begin{gathered}
\because \Delta \mathrm{G}^{\circ}=-\mathrm{nFE}^{\circ} \\
\text { Option }(1) \mathrm{E}^{\circ}=0.8+0.76 \\
=1.56 \mathrm{~V} \\
\therefore \Delta \mathrm{G}^{\circ}=-2 \times \mathrm{F} \times 1.56 \\
=-3.12 \mathrm{~V} \\
\text { Option (2) } \mathrm{E}^{\circ}=-2.37+0.76 \\
=-1.61 \mathrm{~V} \\
\therefore \Delta \mathrm{G}^{\circ}=-2 \times \mathrm{F} \times(-1.61) \\
=+3.22 \mathrm{~V} \\
\text { Option (3) } \mathrm{E}^{\circ}=-2.37-0.8 \\
=-3.17 \mathrm{~V} \\
\begin{array}{c}
\therefore \Delta \mathrm{G}^{\circ}=-2
\end{array} \\
=\mathrm{F} \times(-3.17) \\
=+6.34 \\
\text { Option }(4) \mathrm{E}^{\circ}=0.8-0.34 \\
=0.46 \mathrm{~V} \\
\Delta \mathrm{G}^{\circ}=-2 \times \mathrm{F} \times 0.46 \\
=-0.92 \mathrm{~V}
\end{gathered}$$</p>
About this question
Subject: Chemistry · Chapter: Electrochemistry · Topic: Electrochemical Cells
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