Chemical Thermodynamics
Chemical Thermodynamics tests the first and second laws, enthalpy cycles, entropy, Gibbs energy and spontaneity. Expect sign-convention questions and calculations linking ΔG, ΔH, ΔS and K.
Core rules
- First law: ΔU = q + w, with w = -pΔV for work done by the system (IUPAC convention).
- ΔH = ΔU + Δn_gRT for reactions involving gases.
- Hess's law: enthalpy change is path-independent.
- ΔS = q_rev/T; entropy increases with temperature, phase change to gas, and more particles.
- ΔG = ΔH - TΔS; spontaneous at constant T and p when ΔG < 0; equilibrium when ΔG = 0; ΔG° = -RT ln K.
Worked example
For a reaction ΔH = +30 kJ and ΔS = +100 J K^-1, find whether it is spontaneous at 400 K. ΔG = 30,000 - 400(100) = 30,000 - 40,000 = -10,000 J = -10 kJ, so spontaneous. Reasoning: convert kJ to J because ΔS is in J K^-1; the sign of ΔG decides direction, not the sign of ΔH alone.
Traps
Mixing the old and new work conventions, forgetting unit conversion, and assuming ΔH < 0 always means spontaneous. At the temperature where ΔG = 0, the system is at equilibrium, not necessarily spontaneous.