Given $E^{\circ} (Cl X_{2} / Cl X^{-}) = 1.36 V$ and $E^{\circ} (Fe X^{3 +} / Fe X^{2 +}) = 0.77 V$ , and $F = 96500 C mo l^{- 1}$ , the standard Gibbs energy change for $Cl X_{2} + 2 Fe X^{2 +} 2 Cl X^{-} + 2 Fe X^{3 +}$ (per mole of $Cl X_{2}$ ) is

Answer: B

E_{ce ll}^{\circ} = 1.36 - 0.77 = 0.59 V

. The balanced equation transfers

n = 2

electrons per

Cl X_{2}

Δ G^{\circ} = - n F E^{\circ} = - 2 \times 96500 \times 0.59 = - 113.9 kJ

(negative, so the reaction is spontaneous).