Re: Why does magnesium sulfate lower the temperature of water?

Dear Larry,
Its a little difficult to explain without introducing some terms that your daughter won't be familiar with, and writing equations, so I'll apologize first.

Its all a matter of breaking bonds (endothermic) and making bonds (exothermic).

Take a look at the changes occuring when magnesium sulphate dissolves. Overall it looks like this:

MgSO₄ (+H₂O)-> Mg²⁺(aq) + SO₄^2-(aq)

Broken up into the individual changes, we see:

MgSO₄ -> Mg²⁺ + SO₄^2- (depends on lattice energy)
H₂O_n -> nH₂O (depends on liquid-liquid interaction)
Mg²⁺ + H₂O -> Mg²⁺(aq) (formation of aquo complex and solvation sphere)
SO₄^2- + H₂O ->SO₄^2- (aq) (solvation)

It requires energy (heat) to break up the magnesium sulphate lattice (solute) into individual ions, so that process is endothermic. Similarly, water (solvent) must be broken up since it is hydrogen bonded. But that requires far less energy, but is endothermic nevertheless. The solvation is exothermic, but not as much as the disruption of the solute and the solvent are endothermic. Therefore the overall reaction (in this case) is endothermic and the solution cools.

"Instant Cool Packs" are one useful, commercial example of endothermic solution formation. Such packs normally contain ammonium nitrate and water.

You might be thinking that since more energy is required to break up the magnesium sulphate (and the water) than is released by the solvation, why does it dissolve? It seems to be energetically unfavorable. It is, but enthalpy is just part of the story. The entropy change (ordered crystalline material to disordered solution) makes the whole thing go!

I hope that is not too complicated. I guess I could sum it up as follows:

the solid and the liquid want to be a solution rather than staying separate,
so heat is used up (which causes cooling) in breaking down the solid crystals.

Thanks for taking an interest in your daughter's chemistry class and turning to the Mad Scientist Network.