For the reaction taking place in the cell:
$Pt_{(s)} | H_{2(g)} | H^{+}_{(aq)} || Ag^{+}_{(aq)} | Ag_{(s)}$
$E_{Cell}^0 = +0.5332 \ V$.
The value of $\Delta_f G^0$ in $kJ \ mol^{-1}$ (nearest integer) is:

Which of the following metals does not liberate hydrogen from dilute hydrochloric acid?

The standard electrode potential for a $Daniell$ cell is $1.1 \ V$. Calculate the standard Gibbs energy for the reaction: $[Zn_{(s)} + Cu^{2+}_{(aq)} \rightarrow Zn^{2+}_{(aq)} + Cu_{(s)}]$

For the half-reactions $Zn \rightarrow Zn^{2+} + 2e^{-}$ and $Fe \rightarrow Fe^{2+} + 2e^{-}$,the standard oxidation potentials are given as $E^{0}_{Oxi} = +0.76 \ V$ and $E^{0}_{Oxi} = +0.41 \ V$ respectively. Calculate the cell potential for the reaction $Fe^{2+} + Zn \rightarrow Zn^{2+} + Fe$ in $V$.

Review the $SRP$ (at $25\,\text{°C}$) data in acidic medium:
$Ti^{4+} + e^- \to Ti^{3+}, \, E^o = -x \text{ V}$
$Fe^{3+} + e^- \to Fe^{2+}, \, E^o = -y \text{ V}$
where $x < y$,point out the wrong statement.

Standard electrode potential for the cell with cell reaction $Zn_{(s)} + Cu^{2+}_{(aq)} \rightarrow Zn^{2+}_{(aq)} + Cu_{(s)}$ is $1.1 \ V$. Calculate the standard Gibbs energy change for the cell reaction. (Given $F = 96487 \ C \ mol^{-1}$)