The e.m.f. of a cell whose half cells are given below is .............. $V$
$Mg^{2+} + 2e^- \to Mg_{(s)}$; $E^o = - 2.37 \ V$
$Cu^{2+} + 2e^- \to Cu_{(s)}$; $E^o = + 0.34 \ V$

In which metal container can an aqueous solution of $CuSO_4$ be stored?
$E^0_{Cu^{2+}/Cu} = 0.34 \ V$,$E^0_{Fe/Fe^{2+}} = 0.44 \ V$,$E^0_{Al/Al^{3+}} = 1.66 \ V$,$E^0_{Ni/Ni^{2+}} = 0.25 \ V$,$E^0_{Ag^{+}/Ag} = 0.80 \ V$

$A$ standard hydrogen electrode has zero electrode potential because

The standard reduction potentials at $25^o C$ are given below. Which is the strongest reducing agent?
$Zn^{2+}_{(aq)} + 2e^{-} \rightleftharpoons Zn_{(s)}$,$E^o_{RP} = -0.762 \, V$
$Cr^{3+}_{(aq)} + 3e^{-} \rightleftharpoons Cr_{(s)}$,$E^o_{RP} = -0.740 \, V$
$2H^{+}_{(aq)} + 2e^{-} \rightleftharpoons H_{2(g)}$,$E^o_{RP} = 0.00 \, V$
$Fe^{3+}_{(aq)} + e^{-} \rightleftharpoons Fe^{2+}_{(aq)}$,$E^o_{RP} = 0.77 \, V$

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 $F = 96487 \ C \ mol^{-1}$)

Four alkali metals $A$,$B$,$C$ and $D$ have standard electrode potentials of $-3.05 \ V$,$-1.66 \ V$,$-0.40 \ V$ and $0.80 \ V$ respectively. Which one will be the most reactive?