Physical Chemistry: Thermodynamics, Kinetics and Electrochemistry
Core concepts
Thermodynamics: distinguish state functions (U, H, S, G) from path functions (q, w). Apply the first law, ΔU = q + w, and know the sign convention used. Use ΔG = ΔH − TΔS for spontaneity, remembering that a reaction can be spontaneous at high temperature (positive ΔH, positive ΔS) or low temperature (negative ΔH, negative ΔS). Link ΔG° to the equilibrium constant through ΔG° = −RT ln K and to cell potential through ΔG° = −nFE°.
Kinetics: rate laws are determined experimentally, not from stoichiometry. Know order, molecularity, half-life (first order t½ = 0.693/k, independent of concentration), and the Arrhenius equation k = A e^(−Ea/RT). Understand that a catalyst lowers Ea without changing ΔH or the equilibrium position.
Electrochemistry: assign oxidation numbers, balance redox half-reactions, and use the Nernst equation E = E° − (0.059/n) log Q at 25 °C. A positive E°cell means a spontaneous galvanic reaction; electrolysis is driven by an external supply.
Traps to watch
- Assuming doubling concentration always doubles rate (true only for first order in that species).
- Forgetting that ΔG° = 0 at equilibrium, not ΔH = 0.
- Mixing up anode and cathode: oxidation at the anode in both galvanic and electrolytic cells.
- Using degrees Celsius in the Nernst equation.
Worked example
A first-order reaction has k = 0.0693 min⁻¹. Its half-life is 0.693/0.0693 = 10 min. After 30 min (three half-lives) 12.5% of reactant remains. Rule of thumb: for first order, fraction remaining = (½)^n where n is the number of half-lives elapsed.