Question

The standard electrode potentials are:
\(E^\circ(\text{Zn}^{2+}/\text{Zn}) = -0.76 \, \text{V}, \quad E^\circ(\text{Cu}^{2+}/\text{Cu}) = +0.34 \, \text{V}\).
The equilibrium constant for the reaction
\(\text{Zn} + \text{Cu}^{2+} \rightleftharpoons \text{Cu} + \text{Zn}^{2+}\)
at \(25^\circ \text{C}\) is of the order of

• A. \(10^{-37}\)
• B. \(10^{37}\)
• C. \(10^{-17}\)
• D. 10¹7

Hint:

Large positive \(E^\circ\) implies a very large equilibrium constant.

Answer 1

The Correct Option is B

Step 1: Cell emf:
\[ E^\circ = 0.34 - (-0.76) = 1.10 \, \text{V} \]
Step 2:
\[ \log K = \frac{n E^\circ}{0.0591} = \frac{2 \times 1.10}{0.0591} \approx 37 \]
Step 3:
\[ K \approx 10^{37} \]