The question of water’s freezing temperature seems simple, but the answer involves complexity in physics and chemistry. Most people learn that water freezes at 0° Celsius or 32° Fahrenheit. This is the basic scientific truth for pure water under a specific set of circumstances. In the real world, factors like dissolved substances and pressure constantly alter this phase transition.
The Standard Freezing Point
The standard freezing point is defined using controlled conditions: pure water and standard atmospheric pressure. Pure water is highly distilled, containing virtually no dissolved minerals or gases that could interfere with the process. Standard atmospheric pressure is defined as one atmosphere, or 101.325 kilopascals.
Under these precise conditions, the transition occurs at 0° Celsius (0°C) or 32° Fahrenheit (32°F). The Celsius scale is used in most scientific contexts globally, setting the freezing point at zero and the boiling point at 100. Scientists also reference the Kelvin scale, which measures temperature from absolute zero, where the freezing point of water is 273.15 Kelvin (K).
Understanding the Freezing Process
Freezing is a physical phase change where the kinetic energy of water molecules decreases as heat is removed from the system. As the molecules slow down, their inherent polarity allows them to form a stable, highly ordered, hexagonal crystalline structure known as ice. This structure is maintained by strong hydrogen bonds linking the individual H₂O molecules.
During the transition from liquid water at 0°C to solid ice at 0°C, the temperature of the substance does not change. Instead, the system must release a specific amount of energy, known as the latent heat of fusion, for the phase change to complete. For water, this latent heat is approximately 334 kilojoules per kilogram (kJ/kg). The removal of this heat allows the molecules to lock into the rigid ice lattice structure.
How Impurities Change the Freezing Point
In natural and engineered environments, water is seldom pure, and the presence of dissolved substances significantly changes the freezing point through a process called freezing point depression. This effect is a colligative property, meaning it depends only on the number of solute particles in the water, not on their chemical identity. The dissolved particles physically interfere with the ability of water molecules to coalesce and form the necessary ice crystal lattice.
The solute particles lower the vapor pressure of the liquid, forcing the temperature to drop lower than 0°C before the liquid and solid phases can exist in equilibrium. For example, the salt in ocean water, which has an average salinity of about 35 parts per thousand, lowers its freezing point to approximately -1.8°C (28.8°F). Road salt applied to highways exploits this principle, depressing the freezing point of water on the pavement. Antifreeze in car radiators functions the same way to prevent engine fluids from freezing in cold weather.
The Phenomenon of Supercooling
Supercooling is a distinct physical phenomenon where water remains liquid even when its temperature is well below its standard freezing point. This unstable state occurs in highly purified water that lacks nucleation sites, such as tiny dust particles or air bubbles. Without a template, the water molecules cannot initiate the formation of the first ice crystals necessary for freezing.
The liquid state can be maintained under supercooled conditions until the temperature drops to the point of homogeneous nucleation, which for pure water is around -40°C (-40°F). When supercooled water is disturbed, such as by a sudden shock or the introduction of a single ice crystal, nucleation rapidly begins. The water instantly freezes, releasing its latent heat of fusion and causing the temperature of the resulting ice to rise quickly back toward 0°C.