Chemistry · Ch 2 — Electrochemistry

Summary

Electrochemical cells convert chemical energy to electrical (galvanic) or vice versa (electrolytic).
Galvanic cell: spontaneous redox reaction ( $\Delta G < 0$ ); anode (oxidation, negative), cathode (reduction, positive).
Cell representation: $Zn(s) | Zn^{2+}(aq) || Cu^{2+}(aq) | Cu(s)$ — single vertical line for phase boundary, double for salt bridge.
Cell potential ( $E_{\text{cell}}$ ): $E_{\text{cell}} = E_{\text{cathode}} - E_{\text{anode}}$ (in volts).
Standard electrode potential ( $E^\ominus$ ): measured under 1 M, 1 bar, 298 K relative to SHE ( $E^\ominus = 0$ V).
Nernst equation: $E = E^\ominus - \frac{0.0591}{n} \log Q$ at 298 K; relates potential to concentration.
Equilibrium constant ( $K$ ): $\log K = \frac{n E^\ominus_{\text{cell}}}{0.0591}$ at 298 K.
Gibbs free energy: $\Delta G = -n F E_{\text{cell}}$ ; spontaneous if $E_{\text{cell}} > 0$ .
Conductance: $\kappa = \frac{1}{\rho}$ (specific conductivity); molar conductivity $\Lambda_m = \frac{\kappa}{c}$ .
Kohlrausch’s law: $\Lambda_m^\infty = \lambda_+^\infty + \lambda_-^\infty$ ; used to find $\Lambda_m^\infty$ for weak electrolytes.
Faraday’s laws: mass deposited $m = \frac{Q M}{n F}$ ; $Q = I t$ . …