Summary

This chapter built up chemical thermodynamics as one continuous, logical chain. It began by carefully defining a system and its surroundings, and classifying systems as open, closed or isolated, together with the distinction between extensive properties (which scale with the amount of matter) and intensive properties (which do not), and between state functions (path-independent) and path functions (path-dependent). The Zeroth Law was shown to be the logical basis for defining temperature itself.

The First Law of Thermodynamics, $\Delta U = q + w$, established internal energy as a state function that conserves total energy across any process, using the IUPAC sign convention for heat and work. Because most reactions occur at constant pressure rather than constant volume, enthalpy, $H = U + PV$, was introduced as the state function whose change equals the heat exchanged at constant pressure, $\Delta H = q_p$, related to $\Delta U$ through $\Delta H = \Delta U + \Delta n_g RT$. A family of named, standardized enthalpy changes — formation, combustion, atomization, bond dissociation, ionization, solution, and sublimation — was defined, each describing one precisely specified process.

Because enthalpy is a state function, Hess's law allows the enthalpy of a reaction to be calculated by combining known enthalpies of other reactions algebraically, regardless of the actual reaction pathway. The Born-Haber cycle applied this same principle to ionic solids, allowing the otherwise unmeasurable lattice enthalpy to be calculated from sublimation, ionization, bond dissociation, electron gain, and formation enthalpies. …