Freezing is the physical phase transition in which a liquid changes into a solid when the thermodynamic conditions favour the solid phase. For a pure crystalline substance at a specified pressure, freezing normally occurs at the same equilibrium temperature as melting, although the observed temperature can differ because of kinetic effects such as supercooling. During freezing, the chemical identity of the substance is ordinarily preserved.
Liquid water becomes solid ice, liquid aluminium becomes solid aluminium and molten wax becomes solid wax without requiring the substance to undergo a chemical reaction. The transition involves changes in particle arrangement, mobility, energy and entropy. In a liquid, molecules or atoms can move relatively freely and continuously change neighbours, while in a crystalline solid they become constrained by an ordered lattice or other stable structure. Freezing therefore reduces the configurational freedom of the material and normally produces a more organised phase. The temperature at which equilibrium freezing occurs depends on pressure and composition, and dissolved substances can alter freezing behaviour substantially.
At the microscopic level, freezing occurs when particles in a liquid form stable arrangements corresponding to the solid phase and those arrangements grow into larger regions of solid material. Thermal motion constantly creates temporary clusters and local structures in a liquid, but most small clusters disappear because surface energy makes them unstable. Once a sufficiently large and appropriately structured nucleus forms, continued crystallisation becomes favourable and particles can join the growing solid. This process is called nucleation.
Heterogeneous nucleation occurs when crystallisation begins at surfaces, impurities, container walls or suspended particles, while homogeneous nucleation occurs within the bulk liquid without an external nucleation site and generally requires greater supercooling. As crystallisation proceeds, latent heat is released to the surroundings. This release occurs because the solid phase has lower enthalpy than the liquid under the relevant equilibrium conditions. The temperature can therefore remain approximately constant during freezing if heat is removed at a rate that permits the system to remain near phase equilibrium.
Freezing is strongly affected by pressure and by the composition of the liquid. For water at ordinary atmospheric pressure, the equilibrium freezing point is approximately 0 degrees Celsius, or 273.15 kelvin, for pure water under standard conditions. Dissolved substances generally lower the freezing point because the liquid solution has a different chemical potential from pure liquid solvent. This phenomenon is called freezing point depression and is an example of a colligative property when considering dilute solutions.
Salt spread on roads can lower the freezing point of water and help prevent or remove ice when temperatures are within the effective range of the salt. Antifreeze solutions containing substances such as ethylene glycol or propylene glycol are used in cooling systems because they reduce the freezing point and can also raise the boiling point of the coolant. In phase equilibrium terms, freezing occurs when the chemical potentials of the liquid and solid phases become equal at the transition condition.
Freezing releases latent heat of fusion to the surroundings. The same quantity that must be supplied to melt a specified mass of a substance at its equilibrium melting temperature must be removed, under corresponding conditions, to freeze that mass. For water, the latent heat associated with freezing is substantial, which is why melting and freezing can act as important thermal buffers.
During the freezing of a pure substance at constant pressure, the temperature may remain near the freezing point while the phase transition proceeds, even though energy is continuously being removed. The extracted energy includes the latent heat associated with the change in phase. After all liquid has frozen, continued removal of heat lowers the temperature of the solid. In a mixture, however, the temperature often changes throughout the freezing process because the composition of the remaining liquid changes as the solid phase preferentially forms with a particular composition. This distinction is important in food science, metallurgy and industrial crystallisation.
Water provides several unusual examples of freezing behaviour. Unlike most common substances, ordinary ice is less dense than liquid water, so ice floats. This arises from the open hydrogen bonded structure of common ice, which occupies more volume than the more compact arrangements present in liquid water. Floating ice has major consequences for aquatic ecosystems because surface ice can form an insulating layer that reduces heat loss from the water beneath it.
Freezing also plays a major role in Earth's climate system, glaciers, sea ice and seasonal snow cover. The formation and melting of ice influence albedo, ocean circulation, freshwater availability and habitat conditions. Sea ice formation additionally rejects much of its dissolved salt into surrounding seawater, increasing the salinity and density of the remaining water and contributing to ocean mixing. Thus freezing is not merely a laboratory phase transition but a major process in planetary environmental systems.
Freezing is distinct from solidification in some contexts, although the words are often used interchangeably. Freezing usually refers to a liquid becoming a solid, especially in everyday and physical contexts, whereas solidification is a broader term that can include a molten material becoming solid. Some substances do not form ordinary crystalline solids readily and can become glasses or amorphous solids when cooled rapidly.
In such cases there may be no single sharp freezing point because the material undergoes a glass transition over a temperature range rather than ordinary crystallisation. Rapid cooling can also produce supercooled liquids in which the temperature falls below the equilibrium freezing point without immediate solidification. If nucleation eventually occurs, crystallisation can happen rapidly and release latent heat, sometimes causing the temperature to rise towards the equilibrium freezing temperature. These effects demonstrate that thermodynamic equilibrium and the actual observed rate of a phase transition are related but not identical concepts.
Freezing has extensive practical significance in engineering, biology, medicine, food preservation and materials processing. Freezing food slows microbial growth and chemical reaction rates, extending shelf life, although ice crystal formation can damage cellular structures and alter texture. Controlled freezing techniques seek to manage crystal size and distribution, which is particularly important in frozen foods and cryopreservation. In cryogenic engineering, materials must be selected according to their freezing and phase behaviour because unwanted solidification can obstruct pipes or damage equipment.
In construction and geotechnical engineering, water freezing in pores and cracks can generate expansion and mechanical stresses that contribute to frost weathering and pavement damage. In biology, ice formation inside cells can be severely damaging, which is why cryopreservation methods use controlled cooling and cryoprotective agents. Industrial crystallisation similarly uses controlled freezing or cooling to separate and purify substances. The practical outcome of freezing therefore depends not only on temperature but also on cooling rate, composition, pressure, nucleation conditions and material structure.
Thermodynamically, freezing illustrates the relationship between enthalpy, entropy and Gibbs free energy. The solid phase generally has lower entropy than the liquid because its particles have fewer accessible configurations, but it can have a favourable enthalpy because of stronger or more organised interactions. At the equilibrium freezing temperature and pressure, the Gibbs free energies of the two phases are equal.
Below that temperature, the solid has the lower Gibbs free energy and is thermodynamically stable, while above it the liquid is favoured. The entropy decrease associated with freezing is balanced by entropy transferred to the surroundings through released heat, so the total entropy of an isolated system does not decrease. This is an important correction to the misconception that every part of a freezing process must involve entropy simply disappearing. Freezing instead represents an ordering transition within the material accompanied by compensating thermodynamic changes in the surroundings.
