Question

Write Nernst equation and its one application.

Hint:

At $298 \, K$, always remember the simplified form: $E = E^\circ - \frac{0.0591}{n} \log Q$.

Answer 1

Step 1: General form of the Nernst equation.
The Nernst equation relates the electrode potential of a half-cell to the standard electrode potential, temperature, and the activities (or concentrations) of the chemical species involved.
\[ E = E^\circ - \frac{2.303RT}{nF} \log \frac{[Products]}{[Reactants]} \] At $298 \, K$, the equation simplifies to: \[ E = E^\circ - \frac{0.0591}{n} \log \frac{[Products]}{[Reactants]} \] where,
$E$ = electrode potential
$E^\circ$ = standard electrode potential
$R$ = gas constant (8.314 J K$^{-1}$ mol$^{-1}$)
$T$ = temperature (in Kelvin)
$n$ = number of electrons involved in reaction
$F$ = Faraday constant (96500 C mol$^{-1}$)
Step 2: One application.
The Nernst equation is used to calculate the electrode potential of a cell under non-standard conditions. For example, the potential of a Daniell cell: \[ Zn | Zn^{2+} (aq) || Cu^{2+} (aq) | Cu \] can be calculated under different concentrations of Zn$^{2+}$ and Cu$^{2+}$. Conclusion:
The Nernst equation provides a direct relation between electrode potential and ionic concentrations, helping in predicting the feasibility and direction of redox reactions.