Question

Write Nernst equation and its one application in chemical cells.

Hint:

Always remember: At \(25^\circ C\), the simplified form is \(E = E^\circ - \tfrac{0.0591}{n} \log Q\).

Answer 1

Step 1: General form of Nernst equation.
The Nernst equation relates the cell potential under non-standard conditions to the standard electrode potential and the reaction quotient: \[ E = E^\circ - \frac{RT}{nF} \ln Q \] where: - \( E \) = electrode potential under given conditions
- \( E^\circ \) = standard electrode potential
- \( R \) = gas constant (\(8.314 \, J \, mol^{-1}K^{-1}\))
- \( T \) = absolute temperature (in K)
- \( n \) = number of electrons transferred
- \( F \) = Faraday constant (\(96500 \, C \, mol^{-1}\))
- \( Q \) = reaction quotient At \( 298 \, K \), the equation becomes: \[ E = E^\circ - \frac{0.0591}{n} \log Q \] Step 2: Application.
One application of the Nernst equation is in calculating the electrode potential of the hydrogen electrode: \[ E = 0 - 0.0591 \, \log \frac{1}{[H^+]} \] Thus, the Nernst equation helps determine the pH of a solution using hydrogen electrode. Conclusion:
The Nernst equation is essential in electrochemistry for predicting cell potentials under non-standard conditions and is widely used in calculating pH values and electrode potentials.