What will happen to the melting point temperature of ice if some common salt is added to it? How to justfy the answer?

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    $\begingroup$ What is your initial impression? Please show some effort on your part to answer the question. Hint, see your textbook under "molality". $\endgroup$ – jonsca Jun 20 '13 at 9:37

From http://antoine.frostburg.edu/chem/senese/101/solutions/faq/why-salt-melts-ice.shtml:

Consider that ice, a crystalline solid, forms via weak partial-polar forces between water molecules, known as hydrogen bonding. As liquid water cools, it loses the energy that keeps these forces from attracting and capturing neighboring water molecules, and so these bonds form in a crystalline structure. Exactly what that structure is depends on the temperature and pressure at which the ice forms; your everyday freezer ice as well as most of the stuff outside in January is Ice Ih, which is a very spacious hexagonal arrangement that makes ice less dense than liquid water.

Also consider that even within what looks like solid ice, individual molecules "melt" and "freeze" all the time as they transiently gain enough energy to break free of the lattice, and then lose it and are recaptured. When the rate at which molecules melt and freeze are equal, the ice is stable. This happens at 0.01*C; below this temperature, more water will freeze, while above this temperature, more water will melt. Right at 0.01*C, water remains water and ice remains ice (the other half of this coexistence, besides temperature, is that water like most substances requires more energy gain or loss to change its phase, known as the Latent Heat of Vaporization, and for water that's 334J/mL, 80 times what's needed to cool liquid water from 2*C to 1*C).

Introduce salt, and these dissolved ions interfere with this hydrogen bonding within the ice, preventing it from happening at the normal freezing point of water. Now the melting molecules don't refreeze, because their partial polar charges are more strongly attracted to the full charges of the sodium and chloride ions than to other water molecules. That upsets the balance; more water is melting than freezing at the normal freezing temperature, and so the ice melts. It takes more energy loss within this system before the water is too unenergetic to stay liquid and settles into crystal, releasing the salt, which then finds other liquid water at the boundary layer between liquid and ice.

This creates a vicious cycle; as the salt concentrates within the remaining liquid water, its freezing point is lowered even further as the sodium and chlorine, which won't resolidify as salt crystals until there's just too little water to remain dissolved, will become even more attracted to the remaining water, even further lowering the remaining liquid's freezing point as the water must free itself from even more ions before it can be captured by the ice crystal.

So, an aqueous salt solution, in nearly any concentration, will lower the boiling point of the solution below 0*C, but the minimum freezing point is at -21.1*C, which is the freezing point of a fully-saturated salt solution. Below this temperature, salt that loses its water to freezing can't stay dissolved, and precipitates out as a separate crystal solid suspended within the ice.


protected by Loong Aug 18 '16 at 12:13

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