When 0.1 mol MnO₄^2– is oxidized, the quantity of electricity required to completely oxidise MnO₄^2– to MnO₄^– is:

1.	96500 C	2.	2 × 96500 C
3.	9650 C	4.	96.50 C

The cell potential will be:

The molar conductance of $\frac{\mathrm{M}}{32}$ solution of a weak monobasic acid is 8.0 ohm^-1 cm² and at infinite dilution is 400 ohm^-1 cm². The dissociation constant of this acid is: 

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

If ${\mathrm{E}}_{{\mathrm{Fe}}^{2+}/\mathrm{Fe}}^{\mathrm{o}}$ = -0.441 V and  ${\mathrm{E}}_{{\mathrm{Fe}}^{3+}/{\mathrm{Fe}}^{2+}}^{\mathrm{o}}$ = 0.771 V, the standard emf of the reaction: 

Fe + 2Fe³⁺→ 3Fe²⁺ will be:

Given the following cell reaction:
 \(\mathrm{2Fe^{3+}(aq) \ + \ 2I^{-}(aq)\rightarrow 2Fe^{2+}(aq) \ + \ I_{2}(aq)}\)

[Given: \(F  = 96500\) \(C\) \(mol^{- 1}\)]

1. \(23 . 16\) \(kJ\) \(mol^{- 1}\)

2. \(- 46 . 32\) \(kJ\) \(mol^{- 1}\)

3. \(- 23 . 16\) \(kJ\) \(mol^{- 1}\)

4. \(46 . 32\) \(kJ\) \(mol^{- 1}\)

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

The pressure of H₂ required to make the potential of H₂ - electrode zero in pure water at 298 K is: