Why do we put salt on icy sidewalks in the winter?

[Editor's note: In his answer to this question, the late John Margrave argued that salt dissolves in water as ions of sodium and chlorine, and these ions hydrate, or join to, the water molecules. This process gives off heat, which thaws ice. A number of readers alerted us to problems with this explanation. Chemical engineering professor Arthur Pelton of the University of Montreal provided a representative correction. His explanation follows, and Margrave's original answer appears below that.]

Although the hydration process gives off heat, this is more than compensated for by the heat absorbed during the initial decomposition of the salt into ions. In other words, the total process of dissolution--decomposition into ions plus hydration--absorbs heat. This can easily be demonstrated: pour some water into a glass and test its temperature with your finger. Add some salt, stir, and test it again. The temperature will have decreased.

The actual reason that the application of salt causes ice to melt is that a solution of water and dissolved salt has a lower freezing point than pure water. When added to ice, salt first dissolves in the film of liquid water that is always present on the surface, thereby lowering its freezing point below the ice¿s temperature. Ice in contact with salty water therefore melts, creating more liquid water, which dissolves more salt, thereby causing more ice to melt, and so on. The higher the concentration of dissolved salt, the lower its overall freezing point. There is a limit, however, to the amount of salt that can be dissolved in water. Water containing a maximum amount of dissolved salt has a freezing point of about zero degrees Fahrenheit. Therefore, the application of salt will not melt the ice on a sidewalk if the temperature is below zero degrees F.

To understand why water containing dissolved salt has a lower freezing point than pure water, consider that when ice and water are in contact there is a dynamic exchange at the interface of the two phase states. Because of thermal vibrations in the ice, a large number of molecules per second become detached from its surface and enter into the water. During the same period of time, a large number of water molecules attach themselves to the surface of the ice and become part of the solid phase. At higher temperatures, the former rate is faster than the latter and the ice melts. At lower temperatures the reverse is true. At the freezing point the two rates are equal. If salt is dissolved in the water, the rate of detachment of the ice molecules is unaffected but the rate at which water molecules attach to the ice surface is decreased, mainly because the concentration of water molecules in the liquid (molecules per cubic centimeter) is lower. Hence, the melting point is lower.

John Margrave, a chemistry professor at Rice University, explains.

All icy surfaces in fact contain small puddles of water. Because salt is soluble in water, salt applied to such surfaces dissolves. Liquid water has what is known as a high dielectric constant, which allows the ions in the salt (positively charged sodium and negatively charged chlorine) to separate. These ions, in turn, react with water molecules and hydrate¿that is, form hydrated ions (charged ions joined to water molecules). This process gives off heat, because hydrates are more stable than the individual ions. That energy then melts microscopic parts of the ice surface. Thus a substantial amount of salt spread over a large surface can actually thaw the ice. In addition, if you drive over the ice in your automobile, the pressure helps force the salt into the ice and more of this hydration occurs.

The rock salt applied to icy roads in the winter is the same substance that comes out of your salt shaker. The only difference is the size. Rock salt is the material that has crystalized in larger pieces, whereas table salt has been ground up and pulverized to a more or less uniform size distribution. Calcium chloride is just as commonly used to melt ice on the streets as sodium chloride is. In fact, it's cheaper than sodium chloride. Companies manufacture large amounts of calcium chloride from brines and other natural materials that can be used for the same purpose.

Originally published on December 8, 2003.

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