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.
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Wow talk about a unique way to create your own outdoor hand warmer. I'm pretty sure this is past my abilities though. I took chemistry not long ago and loved it, but I can't remember anything from it.
Is this dangerous then? Should I get a portable water filter, or is it safe to drink? THanks for the post. It was sweet!
So cool! I am going to have to try this! I love being outdoors but living where I do sometimes it can get chilly. This would definitely come in handy!
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.
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.
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
Hmm, is there anything you should use your finger for to start crystallization?
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.
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
in the directions i read it says you need to buy sodium acetate and then further on it says use sodium acetate trihydrate crystals
which do you use
are they the same
help jaz
Hey where can i find sodium acetate trihydrate crystals for a science fair project? I dont need a VERY large amount but i do need it to make a few batched
How interesting! I was always so curious as to how hand-warmers worked, and now it makes sense. Science is such a mind boggling subject to me, but the way you worded things helped me understand rather easily. So thank you!