$\begin{aligned} & \text { Calculate } \mathrm{E}_{\text {cell }}^{\circ} \text { for }…

$\begin{aligned} & \text { Calculate } \mathrm{E}_{\text {cell }}^{\circ} \text { for } \mathrm{Cd}_{(\mathrm{s})}\left|\mathrm{Cd}_{(\mathrm{IM})}^{++}\right|\left|\mathrm{Ag}_{(\mathrm{MM})}^{+}\right| \mathrm{Ag}_{(\mathrm{s})} . \\ & {\left[\mathrm{E}_{\mathrm{Cd}}^{\circ}=-0.403 \mathrm{~V} ; \mathrm{E}_{\mathrm{Ag}}^{\circ}=0.799 \mathrm{~V}\right]}\end{aligned}$
  1. $1.202 V$
  2. $-1.202 V$
  3. $0.396 V$
  4. $-0.396 V$

Solution

For the given cell, anode is $\mathrm{Cd}$ and cathode is Ag. $\begin{aligned} \mathrm{E}_{\text {cell }}^{\circ} & =\mathrm{E}_{\text {cathode }}^{\circ}-\mathrm{E}_{\text {anode }}^{\circ} \\ & =0.799-(-0.403) \\ & =1.202 \mathrm{~V} \end{aligned}$

Asked in: MHT CET 2023 (14 May Shift 2)

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