Thermodynamics and Kinetics
What the exam expects
This domain pairs energy changes with reaction rates. Thermodynamics concepts: enthalpy (ΔH), exothermic vs. endothermic, bond energy, Hess's law, entropy (ΔS), Gibbs free energy (ΔG = ΔH − TΔS), and spontaneity. Kinetics concepts: collision theory, activation energy, catalysts, concentration, temperature, surface area, and rate laws.
Traps to watch for
- Thermodynamics vs. kinetics confusion. A spontaneous reaction (negative ΔG) can be very slow; a catalyst speeds kinetics without changing ΔH, ΔS, or equilibrium position.
- Sign conventions. Exothermic: ΔH negative (energy released). Endothermic: ΔH positive.
- Spontaneity conditions. ΔG < 0 is spontaneous. Temperature can flip spontaneity when ΔH and ΔS have the same sign.
- Catalysts lower activation energy for both forward and reverse reactions equally; they do not shift equilibrium.
- Bond breaking absorbs energy; bond forming releases energy.
- Rate-determining step controls overall rate in a mechanism.
Rule of thumb
Use ΔG = ΔH − TΔS. If ΔH is negative and ΔS is positive, spontaneous at all temperatures. If both positive, spontaneous only at high T. If both negative, spontaneous only at low T.
Worked example
A reaction has ΔH = +50 kJ and ΔS = +150 J/K. At what temperature does it become spontaneous? Set ΔG = 0: T = ΔH/ΔS = 50,000 J / 150 J/K ≈ 333 K. Above ~333 K the reaction is spontaneous; below, it is not. Watch the unit conversion (kJ to J).