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    When a solute dissolves in a solvent the new intermolecular bonds in the solution are stronger that those in the solvent alone. Right? Otherwise the solute would be insoluble.

    So, the reaction is exothermic and the new strong bonds are going to result in a higher melting point which, considering water and salt (sodium chloride) the aqueous salt solution could be expected to have a melting/freezing point around 3 or 4 degrees C. But it doesn't - the melting point is, in fact, less than that of pure water :eek:

    Is water a special case or am I getting something quite fundamentally wrong?

    Please set me straight as this has been bothering me for the better part of a decade!
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    http://antoine.frostburg.edu/chem/se...elts-ice.shtml
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    Consider adding salt to ice. Ice has a nice ordered structure thanks to the way the polar water molecules hydrogen bond to one another.

    If you add charged ions to this lattice, the water molecules become more attracted to these ions than to themselves. The water molecules then align themselves around the ions instead of sitting in their own lattice - effectively breaking up the lattice and so the ice melts. So the effect is a lower melting point.

    In order to re-freeze the water these solvated ions need to be slowed down sufficiently so they can pack, which is done by lowering the temperature further.
 
 
 
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