Science, Tech, Math › Science Why Does Salt Melt Ice? Science of How It Works Understand the Chemistry of Why Salt Melts Ice Share Flipboard Email Print A man salts the sidewalk in Paris to prevent ice. Herv?? de Gueltzl / Getty Images Science Chemistry Basics Chemical Laws Molecules Periodic Table Projects & Experiments Scientific Method Biochemistry Physical Chemistry Medical Chemistry Chemistry In Everyday Life Famous Chemists Activities for Kids Abbreviations & Acronyms Biology Physics Geology Astronomy Weather & Climate By Anne Marie Helmenstine, Ph.D. Chemistry Expert Ph.D., Biomedical Sciences, University of Tennessee at Knoxville B.A., Physics and Mathematics, Hastings College Dr. Helmenstine holds a Ph.D. in biomedical sciences and is a science writer, educator, and consultant. She has taught science courses at the high school, college, and graduate levels. our editorial process Facebook Facebook Twitter Twitter Anne Marie Helmenstine, Ph.D. Updated December 04, 2019 You know that you can sprinkle salt on an icy road or sidewalk to help keep it from becoming icy, but do you know how salt melts ice? Take a look at freezing point depression to understand how it works. Key Takeaways: Why Salt Melts Ice Salt melts ice and help prevent re-freezing by lowering the freezing point of water. This phenomenon is called freezing point depression.The working temperature range isn't the same for all types of salt. For example, calcium chloride lowers the freezing point more than sodium chloride.In addition to melting ice, freezing point depression can be used to make ice cream without a freezer. Salt, Ice, and Freezing Point Depression Salt melts ice essentially because adding salt lowers the freezing point of the water. How does this melt ice? Well, it doesn't, unless there is a little water available with the ice. The good news is you don't need a pool of water to achieve the effect. Ice typically is coated with a thin film of liquid water, which is all it takes. Pure water freezes at 32°F (0°C). Water with salt (or any other substance in it) will freeze at some lower temperature. Just how low this temperature will be depends on the de-icing agent. If you put salt on ice in a situation where the temperature will never get up to the new freezing point of the salt-water solution, you won't see any benefit. For example, tossing table salt (sodium chloride) onto ice when it's 0°F won't do anything more than coat the ice with a layer of salt. On the other hand, if you put the same salt on ice at 15°F, the salt will be able to prevent melting ice from re-freezing. Magnesium chloride works down to 5°F while calcium chloride works down to -20°F. If the temperature gets down to where the salt water can freeze, energy will be released when bonds form as the liquid becomes a solid. This energy may be enough to melt a small amount of the pure ice, keeping the process going. Use Salt to Melt Ice (Activity) You can demonstrate the effect of freezing point depression yourself, even if you don't have an icy sidewalk handy. One way is to make your own ice cream in a baggie, where adding salt to water produces a mixture so cold it can freeze your treat. If you just want to see an example of how cold ice plus salt can get, mix 33 ounces of ordinary table salt with 100 ounces of crushed ice or snow. Be careful! The mixture will be about -6°F (-21°C), which is cold enough to give you frostbite if you hold it too long. Table salt dissolves into sodium and chloride ions in water. Sugar dissolves in water, but does not dissociate into any ions. What effect do you think adding sugar to water would have on its freezing point? Can you design an experiment to test your hypothesis? Beyond Salt and Water Putting salt on water isn't the only time freezing point depression occurs. Any time you add particles to a liquid, you lower its freezing point and raise its boiling point. Another good example of freezing point depression is vodka. Vodka contains both ethanol and water. Ordinarily, vodka does not freeze in a home freezer. The alcohol in the water lowers the freezing point of the water. Sources Atkins, Peter (2006). Atkins' Physical Chemistry. Oxford University Press. pp. 150–153. ISBN 0198700725.Petrucci, Ralph H.; Harwood, William S.; Herring, F. Geoffrey (2002). General Chemistry (8th ed.). Prentice-Hall. p. 557-558. ISBN 0-13-014329-4.