Question

For the cell reaction
$2{\mathrm{Fe}}^{3+}\left(\mathrm{aq}\right)+2I^{-}\left(\mathrm{aq}\right)\to 2{\mathrm{Fe}}^{2+}\left(\mathrm{aq}\right)+I_2\left(\mathrm{aq}\right)E{°}_{\mathrm{cell}}=0.24V\mathrm{at}298K$ The standard Gibbs energy $△G^o$ of the cell reaction is:
[Given that Faraday constant F = 96500 C mol–1]

• A. – 46.32 kJ $mo{l}^{-1}$
• B. – 23.16 kJ $mo{l}^{-1}$
  
• C. 46.32 kJ$mo{l}^{-1}$
• D. 23.16 kJ$mo{l}^{-1}$

Answer 1

The correct option is A – 46.32 kJ $mo{l}^{-1}$

$△G^o$ = -nF$E_{cell}^o$
= -2 $\times$96500$\times$0.24 J $mo{l}^{-1}$
= -46320 J $mo{l}^{-1}$
= -46.32 kJ$mo{l}^{-1}$