Phase Transitions Explained

Types of Phase Transitions

First-order phase transitions involve a discontinuous change in entropy and volume at the transition point. Melting, boiling, and sublimation are all first-order transitions. During a first-order transition, the two phases coexist at the transition temperature, and a definite amount of latent heat must be added or removed to complete the transformation. Water at 100 degrees Celsius and 1 atmosphere can exist as both liquid and vapor simultaneously, with the proportion of each phase determined by how much heat has been added.

Second-order (continuous) phase transitions have no latent heat and no discontinuity in entropy or volume, but the heat capacity, compressibility, or other second derivatives of the free energy show discontinuities or divergences. The transition from ferromagnetic to paramagnetic behavior at the Curie temperature, the onset of superconductivity, and the lambda transition in helium-4 are examples of second-order transitions. These transitions are characterized by critical exponents and universality classes, concepts from the field of critical phenomena.

The liquid-gas transition has a critical point (for water, at 374 degrees Celsius and 218 atmospheres) above which the distinction between liquid and gas disappears. Beyond the critical point, the substance exists as a supercritical fluid with properties intermediate between liquid and gas. Supercritical fluids are used industrially for extraction, cleaning, and as reaction media because their density, viscosity, and solvent properties can be tuned by adjusting temperature and pressure.

Latent Heat and Energy During Phase Changes

During a first-order phase transition, heat flows into or out of a substance without changing its temperature. This energy, called latent heat, goes into changing the arrangement and interactions of molecules rather than increasing their kinetic energy. The latent heat of fusion for water is 334 J/g (the energy needed to melt ice at 0 degrees Celsius), and the latent heat of vaporization is 2260 J/g (the energy needed to boil water at 100 degrees Celsius).

The large difference between these values reflects the different physical processes involved. Melting loosens the rigid crystal structure but maintains significant intermolecular contact. Vaporization requires completely separating molecules from their neighbors, overcoming all attractive forces. This is why sweating is such an effective cooling mechanism: evaporating one gram of sweat absorbs 2260 joules of heat from the skin, nearly seven times more than melting the same mass of ice.

Heating curves (plots of temperature versus heat added) show flat plateaus at phase transition temperatures, where all added heat goes into the phase change rather than raising the temperature. The length of the plateau is proportional to the latent heat and the mass of the substance. Only after the transition is complete does the temperature begin rising again, with a slope determined by the heat capacity of the new phase.

Phase Diagrams

A phase diagram maps the stable phase of a substance as a function of temperature and pressure. The boundaries between phase regions (called phase boundaries or coexistence curves) show the conditions under which two phases can coexist in equilibrium. The triple point is where all three phase boundaries meet, marking the unique temperature and pressure at which solid, liquid, and gas coexist simultaneously. For water, the triple point is at 273.16 K and 611.73 Pa.

The Clausius-Clapeyron equation describes the slope of phase boundaries: dP/dT = delta S / delta V = delta H / (T delta V), where delta S, delta H, and delta V are the entropy, enthalpy, and volume changes of the transition. For most substances, melting increases volume, so the solid-liquid boundary slopes to the right (higher pressure raises the melting point). Water is an exception: ice is less dense than liquid water, so the slope is slightly to the left, meaning pressure lowers the melting point.

Phase diagrams become more complex for mixtures and for substances with multiple solid phases. Iron has several solid phases (ferrite, austenite, delta-ferrite) with different crystal structures, making its phase diagram essential for steel production. Binary phase diagrams for alloys show regions of complete mixing, partial mixing, and eutectic compositions where mixtures melt at lower temperatures than either pure component.

Phase Transitions in Nature and Technology

The water cycle is the largest phase transition process on Earth. Solar energy evaporates about 500,000 cubic kilometers of water per year from oceans, lakes, and land surfaces. This water vapor carries latent heat into the atmosphere, where condensation releases that heat, driving atmospheric circulation and weather systems. The global water cycle transfers about 40 percent of solar energy absorbed by Earth surface into the atmosphere via latent heat.

Phase change materials (PCMs) exploit latent heat for thermal energy storage. By melting and freezing at convenient temperatures, PCMs can absorb and release large amounts of heat while maintaining a nearly constant temperature. Applications include building temperature regulation (paraffin wax panels that melt during the day and freeze at night), spacecraft thermal management, and cold chain logistics for pharmaceuticals and food transport.

In metallurgy, controlled phase transitions are the basis of heat treatment processes such as annealing, quenching, and tempering. The properties of steel depend critically on its microstructure, which is determined by the sequence of phase transitions during heating and cooling. Rapid quenching traps the metal in a metastable phase (martensite) that is extremely hard but brittle, while slow cooling produces softer, more ductile phases (pearlite, ferrite).