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How does salt melt ice?

  1. Jan 19, 2005 #1


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    Of course I mean my table salt, specificly, but how does salt melt ice?
  2. jcsd
  3. Jan 19, 2005 #2


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    Unfortunately, I do not know what physical or chemical process occurs. All I know is that salt lowers the melting point of water so that it is able to melt at lower temperatures. I've seen many icy sidewalks to which salt had been recently applied. The ice looks as though it is "dissolving" ...becoming more of a "slush". This is presumably because it is melting partially.
  4. Jan 19, 2005 #3
    From what I know about inter-molecular forces My best guess is that the sodium and chloride ions "gets in the middle" of the strong hydrogen bonds. You might think of it like the small beads in some rotating device (have no idea what it's called) designed to reduce friction.
  5. Jan 19, 2005 #4


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    But what does reducing friction have to do with my question? The sodium and chloride bond with the hydrogen? Could you be more specific please.
  6. Jan 19, 2005 #5
  7. Jan 19, 2005 #6


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    And why does when you put a cup of ice with salt, frost forms on the outside of the cup from condensated water? And a cup without salt doesn't?
  8. Jan 19, 2005 #7
    It is a colligative property-lowering of melting point.
    it is proportional to no:moles of salt per 1000g of ice.
  9. Jan 19, 2005 #8
    Decrease of free energy of salt due to decrease in temperature is less than the corresponding decrease for water.

    So, when salt is mixed in water, the free energy of the system decreases less with decrease in temperature than it would have decreased if water were present alone. As a result, the free energy for the corresponding solid state (to be formed if the mixture freezes) becomes lower than the free energy of pure water only at a relatively lower temperature. That is why the melting (or freezing) point decreases.

  10. Jan 19, 2005 #9
    Adding to the discussion that salt also causes osmosis. Whether melting ice or causing osmosis, salt "pulls water molecules towards itself".
  11. Jan 19, 2005 #10


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    Since salt lowers the freezing/melting point of water, the outside of the cup is below freezing - thus, frost.
  12. Jan 20, 2005 #11
    its actually an exothermic process
    when we put salt the bonds break and also results in liberation of heat
  13. Jan 20, 2005 #12
    First, the melting point of saltwater is much lower than fresh water. This is because the van der Waals forces between water molecules that hold the water in its solid state are weakened by the presence of the sodium+ and chloride ions that salt dissolves into in water. As a result, salt added to ice melts the ice.

    Second, there is latent heat involved in the transition from the solid to the liquid state; thus, when salt is added to ice, as the ice melts, it not only loses heat to provide the latent heat of liquifaction, but it also absorbs heat from its surroundings. Thus, not only is the saltwater cooler than the ice was, even though it is now a liquid because it is still above the liquifaction temperature of salt water, it has also cooled the air around it and any solid material it might be in contact with.

    So if you want to draw heat out of your ice cream, add rock salt to the ice in the outside container, and it will melt the ice and draw out the the heat!
  14. Jan 20, 2005 #13
    does anyone know at what temperature salt stops working and the area would remain icy?
  15. Jan 20, 2005 #14
    Depends on the concentration of salt. I couldn't find a table that will give the answer to your question; this seems to go over how to approximate it. You should be able to derive the freezing point for salt water by concentration of salt from this.
    Last edited: Jan 20, 2005
  16. Jan 20, 2005 #15


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    Freezing point depression for water is 1.86 K/m, or thereabouts; NaCl is a dissociating solute, 2m ions/m; solubility of NaCl decreases as T decreases, something like 5.4 m at 25 C, decreasing to 4.5, if memory serves at the eutectic point (the lower limit jjjsarnis is asking for), 4.5 might be at 0 C; 4.5 X 1.86 x 2 gives us -16 or 17 C. This ain't what's observed since ion association and other non-ideal behavior in concentrations exceeding 0 m reduce the 1.86 K/m freezing point depression constant which is defined only in the limit of zero concentration.
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