Thermodynamics

Enthalpy and Hess's Law

Enthalpy change $\Delta H$ is negative for exothermic and positive for endothermic processes. Because enthalpy is a state function, Hess's law lets you add reaction steps, use standard enthalpies of formation, or approximate with bond enthalpies (bonds broken minus bonds formed).

Entropy and Spontaneity

Entropy $S$ measures energy dispersal: it rises with gas formation, more particles, and dissolution. Spontaneity depends on both drives through Gibbs free energy: $\Delta G = \Delta H - T\Delta S$ A reaction with $\Delta H < 0$ and $\Delta S > 0$ is spontaneous at all temperatures; mixed cases flip at $T = \Delta H/\Delta S$.

Free Energy, Equilibrium, and Phases

Standard free energy connects to the equilibrium constant and to electrochemistry: $\Delta G^\circ = -RT\ln K_{eq} = -nFE^\circ_{cell}$ Coupled reactions (e.g., ATP hydrolysis driving biosynthesis) sum their $\Delta G$ values. Heating curves alternate $q = mc\Delta T$ segments with flat phase-change plateaus; phase diagrams locate the triple and critical points, and colligative properties depend only on particle count.

Key Takeaways

• $\Delta G = \Delta H - T\Delta S$ decides spontaneity; $\Delta G^\circ < 0$ means $K_{eq} > 1$.
• State functions permit Hess's law bookkeeping.
• Temperature is constant during a phase change; all heat goes into breaking intermolecular forces.
• Spontaneous does not mean fast — thermodynamics ignores kinetics.