"Fake" latent heat and supersaturation

Set-up
We begin with a flask with a small amount of a solution of sodium acetate trihydrate (Na2C2H3O2 linked to 3-H20). The sodium acetate is broken down into sodium ions and acetate ions which are dissolved in the water.

Supercooled
The melting/freezing point for this substance is above room temperature (something on the order of 54°C/130°F). Since we start with it still in liquid phase at room temperature which is below the freezing temperature, it is referred to as "supercooled", just like liquid cloud droplets in the atmosphere where the temperature is below 0°C.

How did it get supercooled?
The sodium acetate was in its frozen crystalline form, but the flask was heated to above melting point, and the crystals melted into liquid. Well, actually, with some heating (to around 54°C) the water molecules (the 3-H20) were lost from the crystals, and after further heating (to around 79°C) the sodium acetate salt broke down into its constituent ions and dissolved in that water.

When allowed to cool back down below its freezing point to room temperature the substance did not freeze. What happens in this case is the sodium and acetate ions can stay dissolved in the water in the solution, as the water is still above its freezing point. The solution does not freeze into the sodium acetate trihydrate crystals because the arrangement and structure of the water and the sodium ions and acetate ions must be just right to form the crystals.

Supersaturated
The liquid solution is saturated with respect to sodium acetate (in the form of dissolved sodium ions and acetate ions). However, it is supersaturated with respect to sodium acetate trihydrate (the crystalline form). But unless there is something to activate the crystallization process the solution can remain as it is. Adding a seed crystal of sodium acetate trihydrate jump-starts the rapid freezing process. Dropping a tiny crystal into the solution causes it to rapidly freeze, just like ice crystals can cause supercooled liquid droplets to freeze.

Latent heat release
When molecules go from the liquid phase to the less energetic solid phase, that energy is released from the molecules to the environment. That energy of the phase change is called latent heat. The flask becomes quite warm when the crystallization (freezing) begins.

Why this is technically "fake" latent heat
This is really an exothermic ("heat-releasing") chemical reaction and not just a simple substance phase change. This demonstration depends on ions forming crystals with water molecules. A true phase change like water freezing is simply water itself crystallizing. Of course there is likely a surface or ice nucleus involved with water freezing, but that is not chemically active in the process. Still it is a demonstration illustrating latent heat at work as the substance goes from liquid to solid and gives off a great deal of heat.

More about this reaction
You may have seen hand warmers or baby bottle warmers or something of that sort that work using this reaction. Something jarring the container or a flexible metal piece that bends back and forth may be enough to begin crystallization by getting a few molecules situated appropriately. The freezing occurs very rapidly, but it may take much time for the heat to be completely given off. Then the process can be repeated by heating the substance back up above freezing, typically by boiling the container in water.

Also see these links:
http://www.howstuffworks.com/question290.htm
http://jchemed.chem.wisc.edu/JCESoft/CCA/CCA3/MAIN/ACETATE/PAGE1.HTM


Another demonstration of supersaturation

Take a glass (something which you can clearly see through the side) and put some sort of clear or light-colored carbonated beverage. Sprinkle some ordinary table salt into the liquid in the glass. Carbonated beverages are supersaturated with respect to dissolved carbon dioxide (just as the atmosphere can be supersaturated with respect to water vapor), then the salt provides a starting point for the carbon dioxide to stick to (just like CCN serve as starting points for liquid water to condense).