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

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

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

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Fortunately, there's another option.

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

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

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But how can something go from 
room temperature to near freezing

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in such a short time?

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The answer lies in chemistry.

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Your cold pack contains water
and a solid compound,

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

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When the barrier is broken,
the solid dissolves

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causing what's known as an 
endothermic reaction,

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one that absorbs heat from its surroundings.

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To understand how this works,

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

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energetics and entropy.

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

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In chemistry, energetics deals with
the attractive and repulsive forces

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between particles at the molecular level.

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

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than there are known stars in the universe.

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And all of these trillions 
of molecules are

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constantly moving, vibrating 
and rotating at different rates.

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

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or kinetic energy, of all these particles,

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with an increase in movement
meaning an increase in temperature,

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and vice versa.

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The flow of heat in any 
chemical transformation

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depends on the relative strength
of particle interactions

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in each of a substance's chemical states.

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When particles have a strong mutual
attractive force,

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

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that repulsive forces push them away.

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If the initial attraction was 
strong enough,

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

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The stronger the attraction,
the faster their movement,

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and since heat is essentially motion,

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

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the system heats up.

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But our cold packs do the opposite,

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which means that when 
the solid dissolves in the water,

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

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are weaker than the separate interactions
that existed before.

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This makes both types of particles 
slow down on average,

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cooling the whole solution.

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

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

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This is where entropy comes in.

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Entropy basically describes 
how objects and energy

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are distributed based on random motion.

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

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for the trillions of particles 
that compose it.

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

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and all the nitrogen molecules in another.

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But far more will have them 
mixed together,

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which is why air is always
found in this state.

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Now, if there are strong
attractive forces between particles,

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the probability of some configurations
can change

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

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Oil and water not mixing is an example.

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

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

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and random motion makes the particles
composing the solid separate

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

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

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creates more configurations where 
the solid and water mix together

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and all of these have even weaker
particle interaction,

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less overall particle movement,

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and less heat than there was 
inside the unused pack.

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So while the disorder that can result
from entropy

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may have caused your injury
in the first place,

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