Enthalpy Changes during Phase Transformations

Enthalpy Changes during Phase Transformations

Phase transformations involve energy changes. When ice melts, heat is absorbed. This melting occurs at constant pressure (atmospheric pressure) and during the phase change, temperature remains constant at 273 K.

Here $\Delta_{\text{fus}}H^\circ$ is the standard enthalpy of fusion. If water freezes, the process reverses and an equal amount of heat is released to the surroundings.

Water also requires heat for evaporation. At the constant temperature of its boiling point $T_b$ and at constant pressure:

$\Delta_{\text{vap}}H^\circ$ is the standard enthalpy of vaporisation — the amount of heat required to vaporise one mole of a liquid at constant temperature and under standard pressure (1 bar).

Sublimation is the direct conversion of a solid into its vapour. Solid CO$_2$ (dry ice) sublimes at 195 K with $\Delta_{\text{sub}}H^\circ = 25.2\ \text{kJ mol}^{-1}$. Naphthalene sublimes slowly with $\Delta_{\text{sub}}H^\circ = 73.0\ \text{kJ mol}^{-1}$.

The magnitude of the enthalpy change depends on the strength of intermolecular interactions in the substance undergoing the phase transformation. Strong hydrogen bonds between water molecules hold them tightly in the liquid phase. For an organic liquid like acetone, the intermolecular dipole-dipole interactions are significantly weaker. Thus, less heat is required to vaporise 1 mol of acetone than to vaporise 1 mol of water. …