The phenomenon of salt lowering the temperature at which water freezes is commonly observed during winter de-icing efforts. This effect, which allows liquid water to persist below its normal freezing point of 0°C (32°F), is not a chemical reaction but a physical change. Understanding this process requires looking closely at how pure water molecules organize themselves to form ice and how dissolved particles interfere with that structure.
The Physics of Pure Water Freezing
Freezing is the process where liquid water transitions into a solid state, requiring water molecules to slow their movement. As the temperature of pure water decreases, the molecules lose kinetic energy, allowing them to establish stable, structured connections.
At the standard freezing point, water molecules bond together in a precise arrangement, forming a highly ordered, three-dimensional hexagonal lattice. This crystalline structure (ice) is energetically favorable once the temperature drops to 0°C.
The structure of ice is less dense than liquid water, which is why ice floats. This lower density results from the specific, open arrangement required by the hydrogen bonds when molecules lock into fixed positions. For the phase change to occur, water molecules must align perfectly to build this stable crystalline framework.
How Dissolved Salt Disrupts Freezing
When sodium chloride (NaCl) is introduced into water, it immediately dissociates into its component ions: positively charged sodium (\(Na^+\)) and negatively charged chloride (\(Cl^-\)). These dissolved particles disperse throughout the liquid water, physically interrupting ice formation.
For the water to freeze, its molecules must push these foreign ions out of the way to join the growing ice lattice. This interference raises the energy barrier for the water molecules to achieve the necessary ordered arrangement. The disruptive presence of the ions means the water must reach a lower temperature before the crystal structure can form.
This mechanism is classified as a colligative property, meaning the reduction in the freezing point depends only on the number of solute particles dissolved, not their chemical identity. The more ions present, the greater the disruption and the lower the required freezing temperature.
The Limits of Salt’s Effectiveness
Adding more salt does not lower the freezing point indefinitely, as there is a practical limit. As salt concentration increases, the freezing point drops to a specific minimum temperature known as the eutectic point. For sodium chloride, this point is approximately -21.2°C (-6.2°F), achieved at a concentration of about 23.3% salt by mass.
Beyond this concentration, the solution becomes saturated and can no longer dissolve more salt. If additional salt is added, it remains as solid crystals, and the freezing point of the mixture begins to rise.
The ability to lower the freezing point is also impacted by the specific salt used, due to the colligative property. A salt that dissociates into more particles per molecule is more effective at disrupting the water’s structure. For instance, calcium chloride (\(CaCl_2\)) breaks down into three ions, making it more potent than sodium chloride, which yields only two. This allows calcium chloride to achieve a significantly lower eutectic point, around -51°C (-60°F), useful in much colder conditions.
Where We See This Effect
This principle of freezing point depression has several applications. Most commonly, it is exploited in winter maintenance when salt is spread on roads. The salt dissolves into the thin layer of liquid water already present on the ice, creating a brine with a lower freezing point that melts the surrounding ice.
The same mechanism explains why seawater remains liquid at temperatures below the freezing point of pure freshwater. The dissolved solids in the ocean maintain the water in a liquid state down to a lower temperature.
Another practical example is making homemade ice cream, where rock salt is mixed with ice surrounding the cream mixture. The salt causes the ice to melt and form a super-cold brine that efficiently draws heat away from the cream, causing it to freeze.