Question

The electrochemical cell shown below is a concentration cell. $M \mid M ^{2+}$ (saturated solution of a sparingly soluble salt, $\left. MX _{2}\right)\left| M ^{2+}\left(0.001 mol dm ^{-3}\right)\right| M$ 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 $\Delta G \left( kJ mol ^{-1}\right)$ for the given cell is (take $1 F =96500 C mol ^{-1}$ )

• A. -5.7
• B. 5.7
• C. 11.4
• D. -11.4

Answer 1

Answer: (D)

At anode: $M ( s )+2 X ^{-}( aq ) \rightarrow MX _{2}( aq )+2 e ^{-}$
At cathode: $M ^{+2}( aq )+2 e ^{-} \rightarrow M ( s )$
n-factor of the cell reaction is 2 .
$\Delta G =- nFE _{\text {cell }}=-2 \times 96500 \times 0.059$
$=-113873 / mole =-11.387 KJ / mole -11.4 KJ / mole$