This is a good demonstration for the start of a discussion of thermochemistry. Dissolving lithium chloride in water is exothermic. Dissolving ammonium nitrate in water is endothermic.

two 500 mL erlenmeyer flasks
two scoops for chemicals
thermometer
500 g of lithium chloride
500 g of ammonium nitrate
one liter of distilled water

This should go along with a discussion of exo- and endothermic reactions and processes. The simplest way to introduce these ideas is in terms of heat being either a product or a reactant. When heat is a product (exothermic), the temperature of the system increases. When a reactant (endothermic), the temperature decreases because heat must be absorbed by the system itself.

Show the class a 500 mL flask, the water, and the LiCl. Have a couple of students touch these objects to confirm they are at room temperature. Pour some water in the flask and add a few scoops of LiCl. Swirl to dissolve and have the same students re-examine the flask to confirm the increase in temperature. If you feel like it, measure the water temperature with the thermometer before and after making the solution.

Go through the same process with the ammonium nitrate. The solution becomes quite cold.

Review the ideas connecting exo- and endothermic with the observation of temperature increase or decrease in the system. In the real world such heat of solution phenomena are used in commercial products like the heat packs and cold packs available at most drug stores. These products simply consist of separate bags of water and appropriate salt in a larger plastic lined bag. The end user simply manipulates the outer bag to break open the two interior bags so that the water and salt can mix.

Hard to think of any. Don’t burn yourself or get frostbite.