The transformation of water from its liquid state to solid ice is a fundamental physical change known as freezing. This phase transition occurs at a specific, measurable temperature. While this temperature is a precise physical constant for pure water, it can be affected by external factors.
The Standard Freezing Point of Water
The standard freezing temperature for pure water is a universally accepted reference point in science, defined at standard atmospheric pressure (average pressure at sea level). In the Celsius scale, the freezing point is precisely $0^\circ\text{C}$.
The Fahrenheit scale sets this change at $32^\circ\text{F}$. For scientific applications, the absolute temperature scale, Kelvin, is often used, where the freezing point corresponds to $273.15\text{ K}$. These values serve as a baseline for scientific and engineering applications.
Understanding the Phase Change
Freezing is the result of water molecules losing energy as heat is removed from the liquid. When water is cooled, the molecules slow down, and their kinetic energy decreases. This reduction in motion allows the attractive forces between the molecules to organize them into a structured arrangement.
The organized structure that water molecules form when they solidify is a crystal lattice, or ice. This process does not happen instantly upon reaching $0^\circ\text{C}$, but begins with nucleation, the initial formation of tiny, stable crystal structures. As more energy is drawn away, the crystals grow, releasing a specific amount of energy known as the latent heat of fusion.
Latent heat is the energy released or absorbed during a phase change without a change in the substance’s temperature. For water, this heat must be continuously removed for the liquid to fully convert to ice at $0^\circ\text{C}$. The freezing point of a substance is equivalent to its melting point, representing the temperature at which the solid and liquid phases coexist in equilibrium.
How Impurities Change the Freezing Point
The presence of foreign substances, or solutes, dissolved in water alters the temperature at which it freezes, a phenomenon known as freezing point depression. When a solute is added, like salt or alcohol, its particles interfere with the water molecules’ ability to join together and form the ordered crystal lattice structure required for ice. The water must then be cooled to an even lower temperature to overcome this disruption and begin the crystallization process.
This principle has many practical applications, such as using sodium chloride (salt) to de-ice roads in winter. The salt mixes with water on the ice surface, lowering the freezing point so the ice can melt even when the ambient temperature is below $0^\circ\text{C}$. Ethylene glycol is also added to water in car radiators as antifreeze, preventing the coolant from freezing in cold weather. Adding enough common salt can depress the freezing point of water to approximately $-21^\circ\text{C}$ ($-6^\circ\text{F}$).