The standard reduction electrode potential of ${\mathrm{Zn}}^{2+}/\mathrm{Zn}$ is $-0.76 \mathrm{V}$. This means:

The half cell reactions for rusting of iron are:

$2{\mathrm{H}}^{+}+\frac{1}{2}{\mathrm{O}}_2+2{\mathrm{e}}^{-}⟶{\mathrm{H}}_2\mathrm{O}; {\mathrm{E}}^2=+1.23 \mathrm{V}$, and ${\mathrm{Fe}}^{2+}+2{\mathrm{e}}^{-}⟶\mathrm{Fe}; \mathrm{E}°=-0.44 \mathrm{V}$.

$\mathrm{\Delta G}°$ (in $\mathrm{kJ} {\mathrm{mol}}^{-1}$) for the overall reaction is: 

All the energy released from the reaction $\mathrm{X}\to \mathrm{Y}$,${\mathrm{\Delta }}_{\mathrm{r}}{\mathrm{G}}^{\mathrm{\Theta }}=-193 \mathrm{kJ} {\mathrm{mol}}^{-1}$ is used for oxidizing ${\mathrm{M}}^{+}$ as ${\mathrm{M}}^{+}\to {\mathrm{M}}^{3+}+2\stackrel{-}{\mathrm{e}} {\mathrm{E}}^{\mathrm{\Theta }}=-0.25 \mathrm{V}$ under standard conditions. The number of moles of ${\mathrm{M}}^{+}$ ions oxidized when one mole of X is converted to Y is...... 

$\left(\mathrm{F}=96500{\mathrm{Cmol}}^{-1}\right)$

For the following electrochemical cell at $298 \mathrm{K}$, $\mathrm{Pt}\left(\mathrm{s}\right)\overline{){\mathrm{H}}_2\left(\mathrm{g}\right)(1 \mathrm{bar})\overline{){{\mathrm{H}}^{+}}_{\left(\mathrm{aq}\right)}(1 \mathrm{M})\parallel {{\mathrm{M}}^{4+}}_{\left(\mathrm{aq}\right)},{{\mathrm{M}}^{2+}}_{\left(\mathrm{aq}\right)}\overline{)\mathrm{Pt}\left(\mathrm{s}\right)}}}$, ${\mathrm{E}}_{\mathrm{cell}}=0.092 \mathrm{V}$ when $\frac{\left[{{\mathrm{M}}^{2+}}_{\left(\mathrm{aq}\right)}\right]}{\left[{{\mathrm{M}}^{4+}}_{\left(\mathrm{aq}\right)}\right]}={10}^{\mathrm{x}}$. Given : ${{\mathrm{E}}^{°}}_{({\mathrm{M}}^{4+}/{\mathrm{M}}^{2+})}=0.151 \mathrm{V}; 2.303\frac{\mathrm{RT}}{\mathrm{F}}=0.059 \mathrm{V}$. The value of $\mathrm{x}$ is

Consider an electrochemical cell: $\mathrm{A}\left(\mathrm{s}\right)\overline{){{\mathrm{A}}^{\mathrm{n}}}^{+}}(\mathrm{aq}, 2\mathrm{M})\parallel {\mathrm{B}}^{2\mathrm{n}+}(\mathrm{aq}, 1 \mathrm{M})\overline{)\mathrm{B}\left(\mathrm{s}\right)}$, the value of $∆\mathrm{H}°$ for the cell reaction is twice that of $∆\mathrm{G}°$ at $300 \mathrm{K}$. If the emf of the cell is zero, $∆\mathrm{S}°$(in ${\mathrm{JK}}^{-1}{\mathrm{mol}}^{-1}$) of the cell reaction per mole of  $\mathrm{B}$ formed at $300 \mathrm{K}$ is (Given : $\ln \left(2\right)=0.7, \mathrm{R}=8.31 {\mathrm{JK}}^{-1}{\mathrm{mol}}^{-1}.\mathrm{H},\mathrm{S}$and $\mathrm{G}$ are enthalpy, entropy and Gibbs energy, respectively)

For the following cell with hydrogen electrodes at two different pressure

$\mathrm{Pt}\underset{{\mathrm{p}}_1}{\left|{\mathrm{H}}_2\left(\mathrm{g}\right)\right|}\underset{1\mathrm{M}}{{\mathrm{H}}^{+}\left(\mathrm{aq}\right)}\underset{{\mathrm{p}}_2}{\left|{\mathrm{H}}_2\left(\mathrm{g}\right)\right|}\mathrm{Pt}$

$\mathrm{emf}$ is given by: