The chemistry of cold packs - John Pollard

So you just strained a muscle and the inflammation is unbearable.

You wish you had something ice-cold to dull the pain,

but to use an ice pack, you would have had to put it in the freezer hours ago.

Fortunately, there's another option.

A cold pack can be left at room temperature until the moment you need it,

then just snap it as instructed and within seconds you'll feel the chill.

But how can something go from room temperature to near freezing

in such a short time?

The answer lies in chemistry.

Your cold pack contains water and a solid compound,

usually ammonium nitrate, in different compartments separated by a barrier.

When the barrier is broken, the solid dissolves

causing what's known as an endothermic reaction,

one that absorbs heat from its surroundings.

To understand how this works,

we need to look at the two driving forces behind chemical processes:

energetics and entropy.

These determine whether a change occurs in a system and how energy flows if it does.

In chemistry, energetics deals with the attractive and repulsive forces

between particles at the molecular level.

This scale is so small that there are more water molecules in a single glass

than there are known stars in the universe.

And all of these trillions of molecules are

constantly moving, vibrating and rotating at different rates.

We can think of temperature as a measurement of the average motion,

or kinetic energy, of all these particles,

with an increase in movement meaning an increase in temperature,

and vice versa.

The flow of heat in any chemical transformation

depends on the relative strength of particle interactions

in each of a substance's chemical states.

When particles have a strong mutual attractive force,

they move rapidly towards one another, until they get so close,

that repulsive forces push them away.

If the initial attraction was strong enough,

the particles will keep vibrating back and forth in this way.

The stronger the attraction, the faster their movement,

and since heat is essentially motion,

when a substance changes to a state in which these interactions are stronger,

the system heats up.

But our cold packs do the opposite,

which means that when the solid dissolves in the water,

the new interactions of solid particles and water molecules with each other

are weaker than the separate interactions that existed before.

This makes both types of particles slow down on average,

cooling the whole solution.

But why would a substance change to a state where the interactions were weaker?

Wouldn't the stronger preexisting interactions keep the solid from dissolving?

This is where entropy comes in.

Entropy basically describes how objects and energy

are distributed based on random motion.

If you think of the air in a room, there are many different possible arrangements

for the trillions of particles that compose it.

Some of these will have all the oxygen molecules in one area,

and all the nitrogen molecules in another.

But far more will have them mixed together,

which is why air is always found in this state.

Now, if there are strong attractive forces between particles,

the probability of some configurations can change

even to the point where the odds don't favor certain substances mixing.

Oil and water not mixing is an example.

But in the case of the ammonium nitrate, or other substance in your cold pack,

the attractive forces are not strong enough to change the odds,

and random motion makes the particles composing the solid separate

by dissolving into the water and never returning to their solid state.

To put it simply, your cold pack gets cold because random motion

creates more configurations where the solid and water mix together

and all of these have even weaker particle interaction,

less overall particle movement,

and less heat than there was inside the unused pack.

So while the disorder that can result from entropy

may have caused your injury in the first place,

its also responsible for that comforting cold that soothes your pain.