The potential of the cell for the reaction $M_{(s)} + 2H^{+}(1 \ M) \to H_2(1 \ atm) + M^{2+}(0.1 \ M)$ is $1.5 \ V$. The standard reduction potential for $M^{2+}/M_{(s)}$ couple is ............. $V$.

For a cell,$Cu_{(s)} \mid Cu^{2+}(0.001\,M) \mid\mid Ag^{+}(0.01\,M) \mid Ag_{(s)}$,the cell potential is found to be $0.43\,V$ at $298\,K$. The magnitude of standard electrode potential for $Cu^{2+}/Cu$ is $......... \times 10^{-2}\,V$. $[\text{Given}: E^{\Theta}_{Ag^{+}/Ag} = 0.80\,V \text{ and } \frac{2.303RT}{F} = 0.06\,V]$

For a cell reaction involving a two-electron change,the standard emf of the cell is found to be $0.295 \ V$ at $25 \ ^oC$. The equilibrium constant of the reaction at $25 \ ^oC$ will be

Calculate the equilibrium constant $(K_c)$ of the reaction: $Ni_{(s)} + 2Ag_{(aq)}^{+} \rightarrow Ni_{(aq)}^{2+} + 2Ag_{(s)}$; $E_{cell}^{\circ} = 1.05 \ V$. (Given: $\frac{2.303 \ RT}{F} = 0.06$)

Consider the cell at $25^{\circ} C$:
$Zn | Zn^{2+}_{(aq)} (1 \ M) || Fe^{3+}_{(aq)}, Fe^{2+}_{(aq)} | Pt_{(s)}$
The fraction of total iron present as $Fe^{3+}$ ion at the cell potential of $1.500 \ V$ is $X \times 10^{-2}$. The value of $X$ is $.....$ (Nearest integer).
(Given $E^{0}_{Fe^{3+} / Fe^{2+}} = 0.77 \ V, E^{0}_{Zn^{2+} / Zn} = -0.76 \ V$)

For the galvanic cell,
$Zn_{(s)} + Cu^{2+}(0.02 \ M) \rightarrow Zn^{2+}(0.04 \ M) + Cu_{(s)}$
$E_{cell} = ...... \times 10^{-2} \ V \text{ (Nearest integer) }$
$[\text{Use}: E_{Cu^{2+}/Cu}^{0} = 0.34 \ V, E_{Zn^{2+}/Zn}^{0} = -0.76 \ V]$
$[\frac{2.303 \ RT}{F} = 0.059 \ V]$