Molecules: You'd better learn to live with them.
Certain compounds occur as "hydrates"; that is, with one or more molecules of water. Sometimes there is just water trapped in the crystal structure at a specific stoichiometry (i.e., creatine), sometimes the water is actually covalently incorporated into the molecule, as in formaldehyde hydrate.
Sodium acetate trihydrate is the first case - just crystals of sodium acetate + 3 molecules of water. If you heat this up, it "melts," but it's arguably not "molten," you're just dissolving the NaOAc in the water and making a ~10M solution - well above the saturation limit, but it'll stay in solution until you give it a reason not to. You can even cool this down to room temperature and it'll stay liquid indefinitely.
If you induce crystallization, though, by, say, dropping a speck of dust in, the solution, you get the solid back, and it gets pretty hot.
Alternatively, you can put it in a plastic packet with a metal disc with a tiny crack in the middle. Flex it, and it acts as a nucleation site for starting crystallization. Do this and you have a handy, reusable handwarmer.
Comments
I was told that it was not a supersaturated sol. but a supercooled solution of molten NaOAc and when you induce crystallization it jump to its mp. just like supercooled water, but unlike water it has a mp. of 56°C-ish.
Posted by: Oskar | March 28, 2008 5:46 AM
One thing about sodium acetate - don't try to induce crystallization using your finger. Because sodium acetate will happily start crystallizing on _your_ _finger_ giving you a painful burn.
Posted by: Alex Besogonov | March 28, 2008 7:17 AM
The crystallization happens remarkably quickly once it's induced. There are some pretty cool videos of the effect on YouTube, such as this one:
http://www.youtube.com/watch?v=7Z5Yx9ULfD0&feature=related
Posted by: Russ | March 28, 2008 9:06 AM
Hmm, is there anything you should use your finger for to start crystallization?
Posted by: katie | March 28, 2008 1:33 PM
If you want to realize how little the general public knows about chemistry, read the comments on the youtube video posted in the comment above.
Most people (including the person who posted the video) believe this is "hot ice" or water freezing at room temperature. The concept of a super-saturated solution means nothing.
Posted by: Vince Noir | March 31, 2008 1:13 AM
Try the same trick with Na2S2O5 pentahydrate. The melting point is about 49C so you will not burn yourself and the gorgeous producerd rhombs are gem-quality. If you dip in a crystal on a string into supercooled melt you can get 3-inch near-monocrystal chunk in seconds
Posted by: milkshake | April 2, 2008 5:04 AM