enthalpy

Enthalpy is a thermodynamic property that represents the heat content of a system at constant pressure. It is defined by the equation H = U + pV, where U is internal energy, p is pressure, and V is volume. Enthalpy is a state function, meaning its value depends only on the current state of the system, not on the path taken to reach that state.

For a simple compressible system, the differential form is dH = dU + p dV + V d p. Using the first law, dU = δQ − δW, and assuming only PV work (δW = p dV), this reduces to dH = δQ + V d p. Therefore, at constant pressure (d p = 0), the heat exchanged with the surroundings equals the enthalpy change: δQ_p = ΔH. This makes enthalpy particularly useful in chemistry, where it provides a direct measure of heat under constant-pressure conditions.

Enthalpy changes are central to thermochemistry. The standard molar enthalpy change, ΔH°, describes heat absorbed or released during a process at 1 bar and a specified temperature. Standard enthalpy of formation, ΔH_f°, is the enthalpy change when forming 1 mole of a substance from its elements in their standard states. Enthalpies of fusion, vaporization, and combustion are common examples used to characterize phase transitions and chemical reactions.

Measurements and units: enthalpy is an extensive property, with common units of joules (J) or kilojoules (kJ); molar enthalpies are expressed in J/mol or kJ/mol. In calorimetry, constant-pressure experiments yield ΔH directly, while constant-volume experiments measure ΔU, with the relationship ΔH = ΔU + Δ(pV) used to relate the two when needed.