# Assume a cell with the following reaction

Question:

Assume a cell with the following reaction

$\mathrm{Cu}_{(\mathrm{s})}+2 \mathrm{Ag}^{+}\left(1 \times 10^{-3} \mathrm{M}\right) \rightarrow \mathrm{Cu}^{2+}(0.250 \mathrm{M})+2 \mathrm{Ag}_{(\mathrm{s})}$ $\mathrm{E}_{\text {cell }}^{\ominus}=2.97 \mathrm{~V}$

$\mathrm{E}_{\text {cell for the above reaction is }}$_____________ V. (Nearest integer)

[Given : $\log 2.5=0.3979, \mathrm{~T}=298 \mathrm{~K}$ ]

Solution:

$\mathrm{E}=\mathrm{E}^{\circ}-\frac{0.059}{2} \log \frac{\left[\mathrm{Cu}^{+2}\right]}{\left[\mathrm{Ag}^{+}\right]^{2}}$

$=2.97-\frac{0.059}{2} \log \frac{0.25}{\left(10^{-3}\right)^{2}}=2.81 \mathrm{~V}$