Pure water cannot freeze at \(34^\circ\)F (\(1.1^\circ\)C), as its freezing point is \(32^\circ\)F (\(0^\circ\)C). The freezing process is not always a simple switch at this temperature. Various physical and chemical factors cause water to behave non-standardly near its freezing threshold. Impurities or the absence of a trigger can either lower the temperature required for freezing or allow the water to remain liquid below \(32^\circ\)F.
Establishing the Baseline: The Standard Freezing Point
The standard freezing point for pure water (\(\text{H}_2\text{O}\)) is precisely \(32^\circ\)F, or \(0^\circ\)C. This value is established under standard atmospheric pressure, the average air pressure at sea level. At this temperature and pressure, the liquid and solid phases of water exist in equilibrium.
The transition from liquid water to solid ice occurs when the water molecules lose enough kinetic energy. As the temperature drops, the molecules slow down, allowing the attractive forces between them to dominate. These forces, primarily hydrogen bonds, pull the molecules into a highly ordered, three-dimensional crystalline structure known as ice.
This crystalline lattice structure is less dense than liquid water, which is why ice floats. The standard freezing point represents the temperature at which this orderly arrangement becomes thermodynamically favorable. The process is typically straightforward for pure water, but deviations occur when conditions change.
The Supercooling Phenomenon: Why Water Stays Liquid Below 32 Degrees
Supercooling occurs when liquid water remains fluid even when its temperature is below the standard freezing point of \(32^\circ\)F. This phenomenon is possible because freezing requires a starting point for ice crystals to grow, known as a nucleation site.
In the absence of these sites, water molecules cannot properly align themselves to begin the formation of the solid crystalline structure. The liquid remains in a metastable state, which can persist for a duration. Highly pure water, free of dissolved gases and suspended particles, can be supercooled significantly, with recorded examples down to approximately \(-55^\circ\)F (about \(-48.3^\circ\)C) before spontaneously freezing through homogeneous nucleation.
Nucleation sites are typically microscopic impurities such as dust particles, air bubbles, or tiny imperfections on the container surface. These sites act as a template, providing a stable surface upon which the initial ice crystal, or nucleus, can form. Once a nucleus forms, the remaining supercooled water can freeze almost instantly in a rapid transformation.
Any sudden disturbance, such as shaking the container or introducing a small seed crystal of ice, can trigger solidification. This sudden crystallization releases latent heat, which briefly raises the temperature of the newly formed ice back up to \(32^\circ\)F.
Freezing Point Depression: How Impurities Change the Threshold
The presence of dissolved substances, or solutes, in water chemically interferes with the freezing process, lowering the temperature at which the water can solidify. This effect is known as freezing point depression, meaning the water must be colder than \(32^\circ\)F to freeze. The most relatable example is the addition of salt, such as sodium chloride, to water.
When salt dissolves, it dissociates into ions—sodium (\(\text{Na}^+\)) and chloride (\(\text{Cl}^-\))—which disperse throughout the water. These ions physically disrupt the water molecules’ ability to form the highly ordered, stable crystalline lattice required for ice. The dissolved particles essentially get in the way, making it energetically more difficult for the water molecules to lock into the solid structure.
Freezing point depression is a colligative property, meaning the extent of the temperature drop depends only on the number of dissolved particles, not the type of particle. For example, a 10% solution of sodium chloride by weight can lower the freezing point to approximately \(21^\circ\)F (\(-6^\circ\)C). This principle is widely used for de-icing roads in winter, where the salt lowers the freezing point of the ice, allowing it to melt even when the air temperature is below \(32^\circ\)F.
The Effect of Pressure
Pressure is a secondary factor that also slightly influences the freezing point of water. Unlike most substances, water expands when it freezes, so increasing the external pressure slightly opposes this expansion. Increasing the pressure slightly lowers the freezing point, requiring a colder temperature for the water to freeze. This effect is minimal at typical atmospheric pressures but becomes significant in high-pressure environments.