For the cell reaction 2Fe3+(aq)+2I−(aq)→2Fe2+(aq)+I2(aq)2Fe^{3+} (aq) + 2I^- (aq) \rightarrow 2Fe^{2+} (aq) + I_2 (aq)2Fe3+(aq)+2I−(aq)→2Fe2+(aq)+I2(aq), Ecell⊖=0.24 VE^{\ominus}_{cell} = 0.24 \text{ V}Ecell⊖=0.24 V at 298 K. The standard Gibbs energy (ΔrG⊖\Delta_r G^{\ominus}ΔrG⊖) of the cell reaction is : [Given that Faraday constant F=96500 C mol−1F = 96500 \text{ C mol}^{-1}F=96500 C mol−1]
-46.32 kJ mol−1^{-1}−1
-23.16 kJ mol−1^{-1}−1
46.32 kJ mol−1^{-1}−1
23.16 kJ mol−1^{-1}−1
ΔG⊖=−nFEcell⊖=−2×96500×0.24 J mol−1=−46320 J mol−1=−46.32 kJ mol−1\Delta G^{\ominus} = -nF E^{\ominus}_{cell} = -2 \times 96500 \times 0.24 \text{ J mol}^{-1} = -46320 \text{ J mol}^{-1} = -46.32 \text{ kJ mol}^{-1}ΔG⊖=−nFEcell⊖=−2×96500×0.24 J mol−1=−46320 J mol−1=−46.32 kJ mol−1.
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