For the following cell, $\mathrm{Zn}\left(\mathrm{s}\right)\overline{){{\mathrm{ZnSO}}_4}_{\left(\mathrm{aq}\right)}\parallel {{\mathrm{CuSO}}_4}_{\left(\mathrm{aq}\right)}\overline{)\mathrm{Cu}\left(\mathrm{s}\right)}}$ when the concentrations of ${\mathrm{Zn}}^{2+}$ is $10$ times the concentration of ${\mathrm{Cu}}^{2+}$, the expression for $∆\mathrm{G}$ (in $\mathrm{J} {\mathrm{mol}}^{-1}$) is $[$$\mathrm{F}$ is Faraday constant; $\mathrm{R}$ is gas constant; $\mathrm{T}$is temperature; ${{\mathrm{E}}^{°}}_{\mathrm{cell}}=1.1 \mathrm{V} ]$

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: