Calculate ΔG, ΔH, and ΔS for chemical reactions.
Thermodynamics governs the energy flow of all chemical processes and determines whether reactions occur spontaneously. For AP Chemistry Unit 9 and the USNCO and IChO examinations, thermodynamics requires mastery of enthalpy, entropy, Gibbs free energy, and how these quantities interrelate.
The first law of thermodynamics states that energy is conserved: ΔU = q + w, where q is heat transferred and w is work done on the system. For constant-pressure processes, the heat transferred equals the enthalpy change ΔH. Reactions with ΔH < 0 are exothermic (release heat to surroundings); those with ΔH > 0 are endothermic (absorb heat).
Entropy S measures the dispersal of energy and matter. The second law states that the entropy of the universe always increases in any spontaneous process. Phase changes from solid to liquid to gas increase entropy dramatically; dissolving a solid ionic compound into ions also increases entropy. Reactions that decrease entropy can still be spontaneous if they are sufficiently exothermic.
The unifying quantity is Gibbs free energy: ΔG = ΔH − TΔS. A negative ΔG indicates a spontaneous process at constant temperature and pressure. When ΔH and ΔS have opposite signs, temperature determines spontaneity. At ΔG = 0, the system is at equilibrium, and the temperature at which this occurs is T = ΔH/ΔS — a commonly tested calculation on AP Chemistry free response.
The relationship between ΔG and the equilibrium constant K is: ΔG° = −RT ln K. This connects thermodynamics directly to equilibrium. A large negative ΔG° gives a large K (products heavily favored); a positive ΔG° gives K << 1 (reactants dominate). The Nernst equation extends this to electrochemical cells.
Hess's law allows ΔH values for complex reactions to be calculated by combining known ΔH values for simpler reactions. Bond energies provide an alternative estimate: ΔH ≈ Σ(bonds broken) − Σ(bonds formed). This simulator covers all these topics with interactive visualizations.