E deg(Mg^2+/Mg) = -2.37 V and E deg(Cu^2+/Cu) = +0.34 V. For the cell Mg | Mg^2+ (1 M) || Cu^2+ (1 M) | Cu at 298 K (F = 96500 C mol^-1), which statements are correct?

Delta_r G deg = -523.03 kJ per mole of the reaction Mg + Cu^2+ -> Mg^2+ + Cu

Delta_r G deg per mole of electrons transferred is -261.5 kJ

For the reversed reaction Mg^2+ + Cu -> Mg + Cu^2+, E deg = -2.71 V and Delta_r G deg = +523.03 kJ

Delta_r G deg = -523.03 kJ per mole of the reaction Mg + Cu^2+ -> Mg^2+ + Cu

Delta_r G deg per mole of electrons transferred is -261.5 kJ

For the reversed reaction Mg^2+ + Cu -> Mg + Cu^2+, E deg = -2.71 V and Delta_r G deg = +523.03 kJ

When the cell reaction reaches equilibrium, E deg_cell becomes zero

JEE Advanced levelOne or more correct optionsDifficulty 3 of 3ec054

$E^{\circ} (Mg X^{2 +} / Mg) = - 2.37 V$ and $E^{\circ} (Cu X^{2 +} / Cu) = + 0.34 V$ . For the cell $Mg ∣ Mg X^{2 +} (1 M) ∣∣ Cu X^{2 +} (1 M) ∣ Cu$ at 298 K ( $F = 96500 C mo l^{- 1}$ ), which statements are correct?

A $Δ_{r} G^{\circ} = - 523.03 kJ$ per mole of the reaction $Mg + Cu X^{2 +} Mg X^{2 +} + Cu$
B $Δ_{r} G^{\circ}$ per mole of electrons transferred is $- 261.5 kJ$
CFor the reversed reaction $Mg X^{2 +} + Cu Mg + Cu X^{2 +}$ , $E^{\circ} = - 2.71 V$ and $Δ_{r} G^{\circ} = + 523.03 kJ$
DWhen the cell reaction reaches equilibrium, $E_{cell}^{\circ}$ becomes zero

E_{cell}^{\circ} = 0.34 - (- 2.37) = 2.71 V

n = 2

Δ_{r} G^{\circ} = - 2 \times 96500 \times 2.71 = - 523030 J = - 523.03 kJ

(A), i.e.

- 261.5 kJ

per mole of electrons (B). Reversing the reaction reverses both signs (C). D confuses

E

with

E^{\circ}

: at equilibrium the actual EMF

E

is zero, but

E^{\circ}

is a fixed constant (

2.71 V

) referring to standard conditions.