The electrochemical cell shown below is a concentration cell. M mid M 2+ (saturated solution of a sparingly soluble salt, . MX 2)| M 2+(0.001 mol dm -3)| M The emf of the cell depends on the difference in concentrations of M 2+ ions at the two electrodes. The emf of the cell at 298 K is 0.059 V The value of Δ G ( kJ mol -1) for the given cell is (take 1 F =96500 C mol -1 )

Question

The electrochemical cell shown below is a concentration cell. M mid M 2+ (saturated solution of a sparingly soluble salt, . MX 2)| M 2+(0.001 mol dm -3)| M The emf of the cell depends on the difference in concentrations of M 2+ ions at the two electrodes. The emf of the cell at 298 K is 0.059 V The value of Δ G ( kJ mol -1) for the given cell is (take 1 F =96500 C mol -1 ) (A) -5.7 (B) 5.7 (C) 11.4 (D) -11.4

Tardigrade Admin · Accepted Answer

At anode: M ( s )+2 X -( aq ) arrow MX 2( aq )+2 e - At cathode: M +2( aq )+2 e - arrow M ( s ) n-factor of the cell reaction is 2 . Δ G =- nFE text cell =-2 × 96500 × 0.059 =-113873 / mole =-11.387 KJ / mole -11.4 KJ / mole

-5.7

5.7

11.4

-11.4

At anode: M(s)+2X−(aq)→MX2​(aq)+2e− At cathode: M+2(aq)+2e−→M(s) n-factor of the cell reaction is 2 . ΔG=−nFEcell ​=−2×96500×0.059 =−113873/mole=−11.387KJ/mole−11.4KJ/mole

At anode: $M (s) + 2 X^{-} (a q) \to M X_{2} (a q) + 2 e^{-}$
At cathode: $M^{+ 2} (a q) + 2 e^{-} \to M (s)$
n-factor of the cell reaction is 2 .
$Δ G = - n F E_{cell} = - 2 \times 96500 \times 0.059$
$= - 113873/ m o l e = - 11.387 K J / m o l e - 11.4 K J / m o l e$