Understanding Phase Transitions: A Thermodynamic Perspective
A phase transition is the transformation of a substance from one state of matter (or phase) to another—such as from solid to liquid, or liquid to gas—under specific conditions of temperature and pressure. These transitions are fundamental processes governed by thermodynamics and often visualized through phase diagrams, which map the regions where different phases are stable.
Each phase transition occurs at a characteristic transition temperature for a given pressure, where the Gibbs free energy of the two phases becomes equal. At this point, the system is at equilibrium, and both phases can coexist. For example, ice and liquid water coexist at 0 °C under 1 atm pressure. The transition can be endothermic (absorbing heat, like melting) or exothermic (releasing heat, like freezing).
Some transitions are easily observable, like boiling, while others, such as solid–solid transitions, may require sensitive techniques like differential scanning calorimetry or X-ray diffraction for detection. Importantly, not all thermodynamically favorable transitions happen quickly—metastable phases like diamond can persist for long periods because the transition to a more stable phase (graphite) is kinetically hindered.
The phase rule $F = C − P + 2$ describes how many variables (like temperature and pressure) can be changed independently while maintaining equilibrium between phases, providing a powerful tool to predict phase behavior in pure substances and mixtures alike.