\(Fe^{2+}(aq) \ + \ Ag^{+}(aq) \ \rightarrow \ Fe^{3+}(aq) \ + \ Ag(s)\)

\(E_{Fe^{3+}/Fe^{2+}}^{o} \ = \ 0.77 \ V; \ E_{Ag^{+}/Ag}^{o} \ = \ 0.80 \ V\)

The value of \(\Delta G_{r}^{o} \)  in the above reaction will be:

1. +$2.89$ $\mathrm{kJ}$

2. $-2.89$ $\mathrm{kJ}$

3. $-2.89$ $\mathrm{J}$

4. -186.83 J

Given:
(i) ${\mathrm{Cu}}^{2+}+2{\mathrm{e}}^{-}\to \mathrm{Cu}$    E^o = 0.337 V 
(ii) ${\mathrm{Cu}}^{2+}+{\mathrm{e}}^{-}\to {\mathrm{Cu}}^{+}$  E^o = 0.153 V 
Electrode potential, E^o for the reaction, 
${\mathrm{Cu}}^{+}+{\mathrm{e}}^{-}\to \mathrm{Cu}$, will be: 

1. 0.52 V

2. 0.90 V

3. 0.30 V

4. 0.38 V

For the cell reaction \(2Fe^{3+}(aq) \ + \ 2I^{-}(aq)\rightarrow 2Fe^{2+}(aq) \ + \ I_{2}(aq)\)

\(E_{cell}^{o} \ = \ 0.24 \ V\) at $298$ $\mathrm{K}$. The standard Gibbs energy ∆_rG^⊝ of the cell reaction is:

[Given: $\mathrm{F }= 96500$ $\mathrm{C}$ ${\mathrm{mol}}^{-1}$]

1. $23.16$ $\mathrm{kJ}$ ${\mathrm{mol}}^{-1}$

2. $-46.32$ $\mathrm{kJ}$ ${\mathrm{mol}}^{-1}$

3. $-23.16$ $\mathrm{kJ}$ ${\mathrm{mol}}^{-1}$

4. $46.32$ $\mathrm{kJ}$ ${\mathrm{mol}}^{-1}$

For a cell involving one electron ${\mathrm{E}}_{\mathrm{cell}}^{⊝}=0.\mathrm{59 }\mathrm{V}$ at 298 K.
The equilibrium constant for the cell reaction is :
( $\left[\mathrm{Given}\right]$ $\mathrm{that}$ $\frac{2.303}{}$ $\mathrm{RT}\mathrm{F}=0.059$ $\mathrm{V}$ $\mathrm{at}$ $\mathrm{T}=298$ $\mathrm{K)}$

1. $1.0\times {10}^{30}$

2. $1.0\times {10}^2$

3. $1.0\times {10}^5$

4. $1.0\times {10}^{10}$