The $E^{0}_{Red}$ values of $A, B, C,$ and $D$ are $0.8 \, V, 0.79 \, V, 0.34 \, V,$ and $-2.37 \, V$ respectively. Which element can displace the other three from their salt solutions?

Which metal among the following is most sensitive towards oxidation?

For the cell reaction:
$2 Fe^{3+}_{(aq)} + 2 I^{-}_{(aq)} \rightarrow 2 Fe^{2+}_{(aq)} + I_{2(aq)}$
$E^{\ominus}_{cell} = 0.24 \ V$ at $298 \ K$. The standard Gibbs energy $(\Delta_r G^{\ominus})$ of the cell reaction in $kJ \ mol^{-1}$ is:
[Faraday constant $F = 96500 \ C \ mol^{-1}$]

If $Cu^{+} + e^- \to Cu$ ; $E^o = X_1$ and $Cu^{2+} + 2e^- \to Cu$ ; $E^o = X_2$,then the value of $E^o$ for $Cu^{2+} + e^- \to Cu^{+}$ will be:

If the emf of the cell $Cu(s) | Cu^{2+}(1M) || Ag^+(1M) | Ag(s)$ is $0.463 \text{ V}$ at $25^{\circ} \text{C}$ and the standard electrode potential of the $Cu$ electrode is $0.337 \text{ V}$, find the standard electrode potential of the $Ag$ electrode. (in $\text{ V}$)

The standard electrode potentials of two half-cells are given as: $Ni^{2+}_{(aq)} + 2e^{-} \rightarrow Ni_{(s)}$; $E^o = -0.25 \ V$ and $Zn^{2+}_{(aq)} + 2e^{-} \rightarrow Zn_{(s)}$; $E^o = -0.76 \ V$. What is the standard $emf$ of the cell formed by connecting these two half-cells in $volt$?