The $emf$ of the following cell $Mg|Mg^{2+}(0.01 \ M)||Sn^{2+}(0.1 \ M)|Sn$ at $298 \ K$ in $V$ is: (Given: $E^{\circ}_{Mg^{2+}|Mg} = -2.34 \ V, E^{\circ}_{Sn^{2+}|Sn} = -0.14 \ V$)

At what $pH$,given half cell $MnO_4^{-} (0.1 \ M) \mid Mn^{2+} (0.001 \ M)$ will have electrode potential of $1.282 \ V$? (Nearest Integer) Given $E_{MnO_4^{-} / Mn^{2+}}^{o} = 1.54 \ V, \frac{2.303 RT}{F} = 0.059 \ V$

Calculate the cell potential for the cell $Zn_{(s)} | Zn^{2+} (0.6 \ M) || Cd^{2+} (0.85 \ M) | Cd_{(s)}$ at $298 \ K$. (Given: $E^{\circ}_{Zn^{2+}/Zn} = -0.76 \ V$ and $E^{\circ}_{Cd^{2+}/Cd} = -0.40 \ V$) (in $V$)

The cell potential for the following cell notation is approximately
$M_{(s)} | M^{3+}(aq, 0.01 \ M) || N^{2+}(aq, 0.1 \ M) | N_{(s)}$
$E_{M^{3+} / M}^0 = 0.6 \ V$ and $E_{N^{2+} / N}^0 = 0.1 \ V$ (in $V$)

$2Ag^{+}_{(aq)} + Cu_{(s)} \longrightarrow Cu^{2+}_{(aq)} + 2Ag_{(s)}$
The standard potential for this reaction is $0.46 \ V$. Which change will increase the potential the most?

Write the Nernst equation for the anode and cathode of a Daniell cell.