Question

The standard emf of the cell,

$\text{Cd}(s)\text{ CdC}\text{l}_{2}(\text{aq})\ (\text{0.1 M})\ ||\text{ AgCl}(s)\text{ Ag}(s)$ in which the cell reaction is,

Cd(s) + 2AgCl(s) $\longrightarrow$ 2Ag(s) + Cd+2 (aq) + 2Cl–(aq) is 0.6915 V at 0⁰C and 0.6753 V at 25⁰C. The $\Delta H^{0}$ of the reaction at 25⁰C is

• A. – 176 Kj
• B. – 234.7 kJ
• C.
  • 123.5 kJ
• D. – 167.26 kJ

Answer 1

Answer: (D)

– 167.26 kJ

Answer 2

$\frac{dE}{dT} = \frac{(0.6753 - 0.6915)}{(25 - 0)} = - 6.48 \times 10^{- 4}V$

$\Delta$H298 = – neF + nFT $\frac{dE}{dT}$= – 2 × 0.6753 × 96500 + 2

× 96500 × 298 × (–6.48×10–4)

= 2 × 96500 (– 0.6753 – 0.1931)= – 167.6 KJ