Placing salt directly on top of ice effectively initiates melting, even when the air temperature is below the freezing point of pure water. This widely observed phenomenon is not due to a chemical reaction generating heat, but rather to a fundamental change in the properties of water itself. This alteration modifies the conditions under which water can exist as a solid, turning ice into a liquid brine solution.
Why Salt Melts Ice
The mechanism behind this effect is known as freezing point depression. Ice surfaces are typically coated with a very thin layer of liquid water, even when the surrounding air temperature is below 32°F (0°C). When a compound like rock salt, or sodium chloride (NaCl), is sprinkled onto the ice, it first dissolves in this pre-existing liquid water film.
Once dissolved, the salt breaks apart into its constituent ions (sodium and chloride). These ions physically interfere with the ability of water molecules to bond together. Water molecules naturally form a highly organized, lattice-like structure to solidify into ice. The presence of the dispersed salt ions blocks the water molecules from aligning correctly to maintain this rigid crystalline structure.
Consequently, the temperature must drop significantly lower than 32°F for the water to refreeze into ice. This lowered freezing point causes the surrounding ice to melt into the newly formed saline solution. The process of melting ice is endothermic, meaning it absorbs energy, or heat, from the environment.
Optimizing the Melting Process
While effective, the process of using salt to melt ice is constrained by temperature. Standard rock salt, which is primarily sodium chloride, has a practical working temperature limit. As the temperature drops, the salt’s effectiveness slows dramatically because its capacity to melt ice decreases.
For most de-icing applications, sodium chloride is efficient only down to temperatures of about 15°F (-9.4°C) to 20°F (-6.7°C). Below this threshold, the melting process becomes too slow to be practical for road safety. The lowest theoretical temperature at which a sodium chloride solution can remain liquid is its eutectic point, which is around -6°F (-21°C).
When dealing with more extreme cold, other salt compounds are used because they have lower eutectic points. For example, magnesium chloride (MgCl₂) can remain effective down to approximately -10°F (-23°C), and calcium chloride (CaCl₂) is utilized for its ability to work at temperatures as low as -20°F (-29°C).
Using too much salt is wasteful once the liquid layer reaches its saturation point, where no more salt can dissolve. The most efficient application involves using just enough salt to form a brine solution that prevents refreezing.
Practical Uses of Salt and Ice
The principle of freezing point depression is widely applied in two seemingly different but scientifically related ways. The most common use is in de-icing public roads and sidewalks during winter weather. Salt is spread across these surfaces to lower the freezing point of any moisture present.
This action prevents a strong bond from forming between the ice and the pavement, allowing traffic and plows to more easily clear the surface. By dissolving in the thin water film, the salt creates a liquid layer that remains unfrozen at typical winter temperatures, significantly improving safety.
A counterintuitive use is in the creation of very cold mixtures, such as those used in traditional hand-cranked ice cream makers. Salt is packed around a container holding the ice cream base and ice. The salt rapidly dissolves the ice, and the energy required for this phase change is drawn from the surroundings, including the ice cream mixture. This energy absorption causes the temperature of the ice and salt mixture to drop well below the normal freezing point of pure water, sometimes as low as -5°F (-21°C).