Phase Transitions

A phase transition is the process in which a substance changes from one state of matter to another, like from a solid to a liquid, liquid to gas, or vice versa, at a specific temperature and under given pressure conditions. This change is spontaneous and is affected by alterations in temperature and pressure. These parameters impact the strength of the forces between molecules (intermolecular forces) in the substance.

During a phase transition, both the initial and final phases of the substance coexist. An interesting characteristic of phase transitions is that while they occur, the temperature of the substance remains constant, making them isothermal processes. The Gibbs energy, a thermodynamic potential that measures the maximum reversible work that a system can perform at constant temperature and pressure, becomes zero for an isothermal phase transition. This is because, at equilibrium, the chemical potentials of different phases of the same component are equal. By applying this condition to the equation that connects the Gibbs free energy, enthalpy (heat content), and entropy (degree of randomness), we can find that the entropy for phase change is the ratio of the enthalpy of phase change to the transition temperature.

The heat absorbed or released during a phase transition corresponds to the product of the mass of the substance and its enthalpy of phase change. By replacing the mass with the number of moles, we can express the heat exchange in terms of moles.