MadSci Network: Chemistry |
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) + SO42-(aq)
Broken up into the individual changes, we see:
MgSO4 -> Mg2+ +
SO42- (depends on lattice energy)
H2On -> nH2O (depends on liquid-liquid
interaction)
Mg2+ + H2O -> Mg2+(aq) (formation of
aquo complex and solvation sphere)
SO42- + H2O ->SO42-
(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.
Try the links in the MadSci Library for more information on Chemistry.