Electrode Potential and Standard Electrode Potential

Two half-cell reactions of an electrochemical cell are given below :

 ${\mathrm{MnO}}_4^{-}\left(\mathrm{aq}\right)+8{\mathrm{H}}^{+}\left(\mathrm{aq}\right)+5{\mathrm{e}}^{-}\to {\mathrm{Mn}}^{2+}\left(\mathrm{aq}\right)+4\text{\hspace{0.17em}}{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right),\mathrm{E}°=1.51\mathrm{V}$

 ${\mathrm{Sn}}^{2+}\left(\mathrm{aq}\right)\to {\mathrm{Sn}}^{4+}\left(\mathrm{aq}\right)+2{\mathrm{e}}^{-}\mathrm{E}°=+0.15\mathrm{V}$

What would be the final cell potential?

A copper-silver cell is set up. The copper ion concentration in it is 0.10 M. The concentration of silver ion is not known. The cell potential measured is 0.422 V. Determine the concentration of silver ion in the cell. Given :  ${\mathrm{E}}_{{\mathrm{Ag}}^{+}/\mathrm{Ag}}^{\mathrm{o}}=+0.80\hspace{0.17em}\mathrm{V}, \hspace{0.17em}{\mathrm{E}}_{{\mathrm{Cu}}^{2+}/\mathrm{Cu}}^{\mathrm{o}}=+0.34\hspace{0.17em}\mathrm{V}$.

A solution containing one mole per litre of each  $Cu{\left(N{O}_3\right)}_2;AgN{O}_3;H{g}_2{\left(N{O}_3\right)}_2;$  is being electrolysed by using inert electrodes. The values of standard electrode potentials in volts are :

 $\begin{array}{l}\mathrm{Ag}/{\mathrm{Ag}}^{+}=+0.80,{\mathrm{Hg}}_2/{\mathrm{Hg}}_2^{2+}=+0.79\\ \mathrm{Cu}/{\mathrm{Cu}}^{2+}=+0.34,\mathrm{Mg}/{\mathrm{Mg}}^{2+}=-2.37\end{array}$

With increasing voltage, the sequence of deposition of metals on the cathode will be :

The equilibrium constant for the reaction,  $2{\mathrm{Fe}}^{3+}+3{\mathrm{I}}^{-}\underset{}{\rightleftharpoons }2{\mathrm{Fe}}^{2+}+{{\mathrm{I}}_3}^{-}$ would be, if the standard reduction potentials in acidic conditions are 0.77 V and 0.54 V respectively for  ${\mathrm{Fe}}^{3+}/{\mathrm{Fe}}^{2+}\hspace{0.17em}\mathrm{and}\hspace{0.17em}\mathrm{\hspace{0.17em}}{{\mathrm{I}}_3}^{-}/{\mathrm{I}}^{-}$couple.

The standard potential of the following cell is 0.23 V at  $\mathrm{}\mathrm{}\mathrm{}\mathrm{}\mathrm{}\mathrm{}\mathrm{}\mathrm{}\mathrm{}\mathrm{}\mathrm{}\mathrm{}15°C$  and 0.21 V at  $\mathrm{}\mathrm{}\mathrm{}\mathrm{}\mathrm{}\mathrm{}\mathrm{}\mathrm{}\mathrm{}\mathrm{}\mathrm{}\mathrm{}35°C$ .

 $\mathrm{Pt}\left(\mathrm{s}\right)\left|{\mathrm{H}}_2\right(\mathrm{g}\left)\right|\mathrm{HCl}\left(\mathrm{aq}\right)\left|\mathrm{AgCl}\right(\mathrm{s}\left)\right|\mathrm{Ag}\left(\mathrm{s}\right)$

 (ii) Calculate  $\Delta H°and\Delta S°$ for the cell reaction by assuming that these quantities remain unchanged in the range $15°\mathrm{C} \mathrm{to} 35°\mathrm{C}$.

(iii) Calculate the solubility of AgCl in water at  $\mathrm{}\mathrm{}\mathrm{}\mathrm{}\mathrm{}\mathrm{}\mathrm{}\mathrm{}\mathrm{}\mathrm{}\mathrm{}\mathrm{}25°C$ .

Given: The standard reduction potential of the  ${\mathrm{Ag}}^{+}\left(\mathrm{aq}\right)/\mathrm{Ag}\left(\mathrm{s}\right)$  couple is 0.80 V at $25°\mathrm{C}$.

Standard electrode potential data are useful for understanding the suitability of an oxidant in a redox titration. Some half cell reactions and their standard potentials are given below :

 $\begin{array}{l}{\mathrm{MnO}}_4^{-}(\mathrm{aq}.)+8{\mathrm{H}}^{+}(\mathrm{aq},)+5{\mathrm{e}}^{-}\to {\mathrm{Mn}}^{2+}(\mathrm{aq}.)+4{\mathrm{H}}_2\mathrm{O}\left(1\right)\\ \mathrm{E}°=1.51\mathrm{V}\\ {\mathrm{Cr}}_2{\mathrm{O}}_7^{2-}(\mathrm{aq}.)+14{\mathrm{H}}^{+}(\mathrm{aq}.)+6{\mathrm{e}}^{-}\to 2{\mathrm{Cr}}^{3+}(\mathrm{aq}.)+7{\mathrm{H}}_2\mathrm{O}\left(1\right)\\ \mathrm{E}°=1.38\mathrm{V}\\ {\mathrm{Fe}}^{3+}(\mathrm{aq}.)+{\mathrm{e}}^{-}\to {\mathrm{Fe}}^{2+}(\mathrm{aq}.)\mathrm{E}°=0.77\mathrm{V}\\ {\mathrm{Cl}}_2\left(\mathrm{g}\right)+2{\mathrm{e}}^{-}\to 2{\mathrm{Cl}}^{-}(\mathrm{aq}.)\mathrm{E}°1.40\mathrm{V}\end{array}$

Identify the only incorrect statement regarding the quantitative estimation of aqueous  $Fe{\left(N{O}_3\right)}_2$