Chemical Kinetics and Electrochemistry
Scope
This domain joins reaction rate laws and mechanisms with electrochemical cells: cell potentials, spontaneity, the Nernst equation, and electrolysis stoichiometry.
Core rules
- Rate = k[A]^m[B]^n; orders come from initial-rate data, not from balanced coefficients.
- Integrated forms: zero order [A] = [A]₀ − kt; first order ln[A] = ln[A]₀ − kt with t½ = 0.693/k; second order 1/[A] = 1/[A]₀ + kt.
- Arrhenius: k = Ae^(−Ea/RT); higher Ea means stronger temperature sensitivity.
- The rate-determining step sets the overall rate; intermediates may not appear in the final rate law.
- E°cell = E°cathode − E°anode; ΔG° = −nFE°; at 25 °C, E = E° − (0.0592/n)logQ.
- Electrolysis: moles of electrons = It/F, with F = 96,485 C/mol.
Worked example
A first-order reaction has t½ = 20.0 min. What fraction remains after 60.0 min? First, k = 0.693/20.0 = 0.0347 min⁻¹. Then 60.0 min is exactly three half-lives, so remaining fraction = (1/2)³ = 0.125, or 12.5%. Equivalently, ln([A]/[A]₀) = −0.0347 × 60.0 = −2.08, giving [A]/[A]₀ = e^(−2.08) = 0.125.
Traps
- Multiplying E° values by stoichiometric coefficients; electrode potentials are intensive and are never scaled.
- Mixing anode and cathode conventions: oxidation is always at the anode, but its sign differs between galvanic and electrolytic cells.
- Forgetting that first-order half-life is independent of starting concentration.
- Using log Q without excluding pure solids and liquids.