Hot and cold packs are used by athletes to minimise swelling of injuries such as muscle and joint sprains. But have you ever wondered why squeezing or pounding on the hot or cold pack causes the temperature of the pack to change? Now, we’ll tell you the reason!

Hot Packs

One of the chemicals commonly used in hot packs is:

CaCl₂(s) + H₂O → CaCl₂(aq) + heat

In the case of hot packs, there is an exothermic chemical reaction taking place when the contents are mixed together. In general, within the hot pack, there will be small capsules containing dry chemicals. When these capsules are crushed, the dry chemicals mix with the water inside the pack and undergo an exothermic reaction. Thus, heat is released and causes the temperature of the pack to increase.

Cold Packs

On the other hand, in the case of cold packs, an endothermic chemical reaction takes place when the contents are mixed together. Similar to the hot packs, cold packs also contain small capsules contained dry chemicals and water. When the capsules are crushed, the dry chemicals mix with the water inside the pack, and undergo an endothermic reaction, absorbing heat from the environment. This causes the temperature of the pack to drop.

Common chemicals used in cold packs are:

NH₄NO₃(s) + H₂O + heat → NH₄NO₃(aq)

NH₄Cl(s) + H₂O + heat → NH₄Cl(aq)

These two cases describe one-time use, or disposable, packs. Once the capsules are crushed and the contents are mixed, there is no way to “unmix” the aqueous solution and reform the dry chemical without opening the packs. Hence these packs are not reusable.

Reusable Heat Packs

Nowadays you can buy reusable heat packs. In these packs, there are no dry chemicals to dissolve in water. Instead, there is a tiny metal disc embedded in the pack induces crystallisation in the fluid upon flexing, giving rise to a heating effect of the pack.

The fluid commonly used in these reusable heat packs is a super-cooled, supersaturated solution of sodium acetate trihydrate. The freezing point of sodium acetate is 54^o C, which is also the temperature at which it crystallises.

The paragraph above seems a little far-fetched. What is a super-cooled solution? If the freezing point of sodium acetate is 54^o C, then why is it still not a solid at room temperature? A super-cooled fluid is a fluid whose temperature is below its freezing point, but the fluid remains a fluid – it has not become a solid. This can be achieved if there are no nucleation points, which are points where solid crystal structures form, in the solution. The solution cannot crystallise into the solid state, in other words.

The general equation for the resuable heat pack is as follows:

Na⁺(aq) + CH₃COO⁻(aq) + 3H₂O(l) → NaCH₃COO·3H₂O(s) + heat

When the tiny iron disk embedded in the liquid is flexed, rough surfaces on the disk are exposed and very tiny adhered crystals of sodium acetate are released into the solution. These tiny crystals act as nucleation sites for the crystals of sodium acetate trihydrate to form and grow. Because the liquid is supersaturated, this makes the solution crystallise (or freeze) very rapidly, thereby forming the crystal lattice and releasing energy in the form of heat (this is the heat of formation of the crystal lattice). The temperature of the pack thus rises.

The pad can be reused by boiling it in water until all the crystals of sodium acetate trihydrate have melted. The supersaturated solution is recreated, and once the pad has cooled down, the fluid returns to its super-cooled state and can be used again.