Assume a cell with the following reaction:Cu_ | vedclass

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(C) The Nernst equation for the cell reaction is:$E_{cell} = E_{cell}^{\ominus} - \frac{0.0591}{n} \log \frac{[Cu^{2+}]}{[Ag^{+}]^2}$Here,$n = 2$,$[Cu

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$3$

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(C) The Nernst equation for the cell reaction is:$E_{cell} = E_{cell}^{\ominus} - \frac{0.0591}{n} \log \frac{[Cu^{2+}]}{[Ag^{+}]^2}$Here,$n = 2$,$[Cu^{2+}] = 0.250 \, M$,and $[Ag^{+}] = 1 	imes 10^{-3} \, M$.$E_{cell} = 2.97 - \frac{0.0591}{2} \log \frac{0.250}{(1 	imes 10^{-3})^2}$$E_{cell} = 2.97 - 0.02955 \log \frac{0.250}{10^{-6}}$$E_{cell} = 2.97 - 0.02955 \log (2.5 	imes 10^5)$$E_{cell} = 2.97 - 0.02955 (\log 2.5 + \log 10^5)$$E_{cell} = 2.97 - 0.02955 (0.3979 + 5)$$E_{cell} = 2.97 - 0.02955 (5.3979)$$E_{cell} = 2.97 - 0.1595 \approx 2.81 \, V$The nearest integer value is $3 \, V$.

Assume a cell with the following reaction:
$Cu_{(s)} + 2 Ag^{+} (1 \times 10^{-3} \, M) \rightarrow Cu^{2+} (0.250 \, M) + 2 Ag_{(s)}$
$E_{Cell}^{\ominus} = 2.97 \, V$
$E_{cell}$ for the above reaction is $.... \, V.$ (Nearest integer)
[Given: $\log 2.5 = 0.3979, T = 298 \, K]$

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Assume a cell with the following reaction:$Cu_{(s)} + 2 Ag^{+} (1 \times 10^{-3} \, M) \rightarrow Cu^{2+} (0.250 \, M) + 2 Ag_{(s)}$$E_{Cell}^{\ominus} = 2.97 \, V$$E_{cell}$ for the above reaction is $.... \, V.$ (Nearest integer)[Given: $\log 2.5 = 0.3979, T = 298 \, K]$