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. $1.0\times {10}^{10}$

If the E^o_cell for a given reaction has a negative value, which of the following gives correct relationships for the values of  ∆G^o and K_eq?

1. $∆{\mathrm{G}}^{\mathrm{o}} > 0; {\mathrm{K}}_{\mathrm{eq}}< 1$

2. $∆{\mathrm{G}}^{\mathrm{o}} > 0; {\mathrm{K}}_{\mathrm{eq}}> 1$

3. $∆{\mathrm{G}}^{\mathrm{o}} < 0; {\mathrm{K}}_{\mathrm{eq}}> 1$

4. $∆{\mathrm{G}}^{\mathrm{o}} < 0; {\mathrm{K}}_{\mathrm{eq}}< 1$

A hydrogen gas electrode is made by dipping platinum wire in a solution of HCl of pH = 10 and by passing hydrogen gas around the platinum wire at one atm pressure. The oxidation potential of the electrode would be: 

The Gibb's energy for the decomposition of $A{l}_2{O}_3$ at $500°C$ is as follows: 

2/3Al₂O₃ → 4/3Al + O₂ ; ∆_rG = + 960 k J mol^-1

The potential difference needed for the electrolytic reduction of aluminium oxide (Al₂O₃) at $500°C$ is at least,

1. 3.0 V 

2. 2.5 V 

3. 5.0 V 

4. 4.5 V 

If the  $E_{cell}^{}$ $0$for a given reaction has a negative value, then which of the following gives the correct relationship for the values of $∆G^0$ $$ $and$ $$ $$ $K_{eq}$?

1.  $∆{\mathrm{G}}^0<0;$ ${\mathrm{K}}_{\mathrm{eq}}>1$

2.  $∆{\mathrm{G}}^0<0;$ ${\mathrm{K}}_{\mathrm{eq}}<1$

3.  $∆{\mathrm{G}}^0>0;$ ${\mathrm{K}}_{\mathrm{eq}}<1$

4.  $∆{\mathrm{G}}^0>0;$ ${\mathrm{K}}_{\mathrm{eq}}>1$

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