Sodium acetate (also known as sodium ethanoate) is currently my favourite chemical. That's only partly because it gets bloody cold in my lab sometimes, and it's nice to have a hand warmer at the ready.
One thing I like about it is how mundane it is. When you react vinegar with sodium bicarb (baking soda) in the classic kitchen chemistry 'volcano' demonstration, it's the chemical left behind once all the carbon dioxide has bubbled away. As you might expect from something easily produced from common food ingredients, it's not particularly harmful. It is a mild irritant though, and we want it much more concentrated than you'd normally get in the kitchen, so wash your hands if you get it on your skin.
The really fun thing about sodium acetate is that you can dissolve far more of it in hot water than in cold - but once it cools down it stays dissolved, until it has something to crystallise around. That's because any change of state requires a nucleation point, a place to get started. In this case, the beginnings of a crystal structure provide a scaffold for more particles to crystallise around. Without that, the solution stays supersaturated, containing more of the chemical than it could normally dissolve. Once a seed crystal is introduced, crystals quickly spread out in all directions, like a cloud of smoke made out of needles of ice. You may have seen this in action, in those re-usable hand warmers - sodium acetate is the main chemical used in them. It only takes a few seconds for the crystals to silently spread through a whole hand warmer or flask, radiating a pleasing warmth as they go.
If you have seen reusable hand warmers in action, you may have wondered where that heat comes from. The answer is that when crystals form, particles are attracted together - electrostatic forces pull them into place, and that gives them energy, which becomes heat. A shoe will drop to the the floor with a thump for essentially the same reason: forces cause things to move, and that movement energy has to go somewhere. So crystallisation is almost always exothermic. The energy it releases is called the latent heat of fusion. By the same token it takes energy to break a crystal's bonds, whether by melting a solid or dissolving it, because that means tearing apart mutually attracted particles. With sodium acetate, this isn't too hard - you can re-dissolve the crystals by just heating the container in boiling water or on a Bunsen burner for a couple of minutes. Anything that heats it up past 58°C for long enough should do the trick.
There are at least two very satisfying ways of demonstrating sodium acetate's power of rapid crystallisation. One way is to introduce a seed crystal to the container, which is what commercial hand warmers do when you pop the little metal thing inside. In the lab I've found that shaking the flask usually triggers the process, thanks to imperfections in the bung, but it's probably more fun to start it off by sprinkling magic dust on top. The other way to do it is to pour the supersaturated solution out onto a seed crystal, creating an instant stalagmite which just keeps growing as long as you keep pouring.
The chemical has various other uses besides making hand warmers and great chemistry demonstrations. It has antifungal and antibacterial properties. It's used as a sealant for concrete. Together with acetic acid, it acts as a buffer, helping keep a solution's pH fairly steady when acids or bases are added to it. Perhaps most excitingly, compounding it with acetic acid allows the formation of sodium diacetate crystals, suitable for flavouring salt and vinegar crisps.
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