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ChemMatters February 1987 Page 7

© Copyright 1987, American Chemical Society

Hot & Cold Packs

by Gail Marsella
It is late in the fourth quarter of the football championship, and the
score is tied. Thirty seconds remain. With second down and goal to go,
the home team calls its last time out. Robert, the quarterback who led
the team all season, has sprained his thumb, and it’s too painful to
throw the ball. The trainer quickly takes a white plastic bag from his
pocket, gives it a sharp punch with his fist, and applies it to Robert’s
injured hand. Within seconds the bag is ice cold.
The coach briefs the second-string quarterback and sends him into
the game. The replacement calls a pass play, but is sacked even before
he can raise his arm to throw. Third and goal to go…fifteen seconds left.
The coach looks at Robert, who nods and drops the plastic bag. His
thumb is numb enough for one more play.
Back in the game, Robert takes the snap, fakes right, steps out of the
pocket, then runs to his left. Three defenders rush him. He throws the
ball, missing the defenders’ out-stretched hands by inches. The receiver
catches it in the end zone. Touchdown! Seconds later the clock runs
out...it is the winning score.

Cold in a bag
The plastic bag that the trainer used to cool the quarterback’s thumb is
an “instant cold pack.” It does not need refrigeration and can be stored
for months in a first aid kit, yet it produces cold the moment it is
needed. How does the instant cold pack work? As shown in Figure 1,
the pack has two sealed bags, one inside the other. The outer bag is
made of thick plastic and is relatively strong. It contains two things: a
white powder, and a second plastic bag. The inner bag is made of
weaker plastic and contains water. When the trainer punches the pack
the inner bag breaks, and the water mixes with the powder. As the
water dissolves the powder—a substance called ammonium nitrate—
the solution becomes very cold. The reaction couldn’t be simpler: a
powder dissolving in water. This particular dissolving reaction absorbs
heat, which is a technical way of saying it gets cold. (For why it gets
cold, see “The cold facts”). Other compounds that get cold on
dissolving in water are potassium nitrate, potassium chloride, and, to a
very slight extent, plain old table salt (sodium chloride). Some other
salts, such as sodium hydroxide, give off heat on dissolving. A reaction
that absorbs heat is called an endothermic reaction, whereas heat-
producing reactions are called exothermic.

Hot stuff
Most familiar chemical reactions give off heat. Light a match. Hot, isn’t
it? Where does the heat energy come from? It wasn’t in the match
before you struck it, was it?
Yes it was. It was stored in the match—in the various chemicals that
make up the match. When the match burned, a series of vigorous
reactions took place. Combustion occurs in many steps. To break the
original bonds, heat must be absorbed. As new bonds are formed, heat
is released. In this case, the amount absorbed is less than the amount
released. Overall, heat is given off—the surroundings get hot.
Not all exothermic reactions are as vigorous as a burning match.
Instant hot packs use slower reactions that take place at lower
temperatures. The “Heat Factory” is a brand of hot pack that is sold at
many camping stores. It has an outer plastic envelope and an inner
paper bag perforated by minute holds (see Figure 3). The paper bag
contains a mixture of powdered iron, sodium chloride, activated
charcoal, and sawdust, all dampened with water. Remove the envelope
from the outer plastic bag and shake it vigorously. It gets hot! What’s
going on here?
Everyone knows what happens when an iron shovel is left out in the
rain for a couple of days—it rusts. The chemical reaction of iron and
oxygen (oxidation), produces iron (III) oxide, or rust.

4Fe + 3O2 ⇒ 2Fe2O3 + Heat

In this reaction, ionic bonds are formed between iron and oxygen and
heat is released (197 kilocalories per mole of iron (III) oxide). The
rusting goes faster if the iron is wet, and faster still if the iron is wet
with a salt solution. The shovel left out in the rain rusts too slowly for
the heat to be noticeable. In the Heat Factory, though, the ingredients
are mixed in precise proportions and ground up finely to make the
oxidation go much faster. The Heat Factory is activated by shaking the
envelope to get the oxygen in the air circulating through the small
holes. The heat is the result of fast rusting.
Another brand of heat pack, called the “Heat Solution,” works on a
different principle. The Heat Solution looks like a small air mattress and
is filled with a liquid the consistency of honey (see Figure 3). To activate
it, you squeeze a special compartment in one corner of the pack, which
releases a triggering crystal. The liquid then gradually solidifies and
gives off heat for several hours.
This heat generator uses a phase change instead of a chemical
reaction. We are all familiar with the three common phases of matter:
solid, liquid, and vapor. The most common example is water, which can
exist as ice, liquid water, or steam—same chemical, different phases.
Like other substances, water has specific temperatures at which it
changes phase—boiling point (100 °C) and freezing point (0 °C).
Under certain conditions, a phase can exist outside of the normal
temperature limits. Water, for example, can be cooled below 0 °C.
Meteorologists have learned that, high in the atmosphere, the tiny
droplets of water in clouds may be as cold as -30 °C and still be liquid.
When a chemical remains liquid at temperatures below its normal
freezing point, it is called a supercooled liquid.
What does this have to do with hot packs? The liquid in the Heat
Solution contains supercooled sodium thiosulfate. To make the
supercooled liquid solidify, a seed crystal is needed. This is simply a
small piece of solid sodium thiosulfate around which more solid can
crystallize. When a seed crystal is added, it triggers the change from
supercooled liquid to solid. As the sodium thiosulfate solidifies around
the seed crystal, the pack heats up. The heat is the result of bonds being
formed as the substance crystallizes. The temperature rises to the
freezing point of the sodium thiosulfate, a pleasingly warm 48 °C (118
°F). The valuable feature of phase change systems is that they can’t
overheat. When a supercooled substance crystallizes, the temperature
rises to the freezing point and stays there. It goes no higher or lower
until all of the material has solidified. Notice that the inventors picked
their chemical carefully. Unlike many compounds that cannot be
supercooled at all, sodium thiosulfate supercools easily, and its freezing
point is comfortably warm, but not hot enough to burn the person who
has a pulled muscle.
Unlike the other hot or cold packs, the Heat Solution is reusable.
Simply heat the pack in boiling water for a while to return the sodium
thiosulfate to its supercooled state; let it cool, and it’s ready for the next
emergency. The pack can be recycled until the supply of seed crystals is
used up. (A home-made heat pack is described in “Experimenter’s
Notebook,” page 12.)

Bond energy
We have examined three thermal first-aid packs. The instant cold pack
uses a dissolving reaction (ammonium nitrate in water) that is
endothermic. The Heat Factory uses an exothermic chemical reaction
(iron and oxygen rust to iron (III) oxide). The Heat Solution uses an
exothermic phase change (crystallization of supercooled sodium
thiosulfate). Only in the Heat Factory “fast-rust” system was a new
chemical compound formed. Nevertheless, the underlying theory is the
same. Chemical processes always involve breaking and making bonds,
which cause heat to be absorbed and released. It is the relative balance
of heat change that determines whether the overall process feels hot or
cold to the touch.
It is a week after the football game. Having won the championship,
Robert and some friends are on a weekend camping trip. They had
planned a lot of fishing but, at the lake, they mostly talk about the
game, eat, tell jokes, and relax. What if it gets cold during the night, and
the fire won’t start? What if one of them strains a muscle chopping
wood? Robert is not worried. In addition to the food, tents, and
sleeping bags, he brought some chemical hot and cold packs.

SIDE BARS

The cold facts

Ammonium nitrate (NH4NO3), is classified as a salt. Chemically
speaking, there are thousands of salts in addition to sodium chloride,
common table salt. Salts contain ions, particles with electrical charges.
Because the ions with positive charge are strongly attracted to those
with negative charge, they form a solid crystal. Below is a diagram of a
salt crystal dissolving in water. To the eye, it looks like a simple process.
In fact, there are two distinct steps, and each involves energy changes.
In the first step, the solid crystal separates into ions. Breaking the ionic
bonds requires a lot of energy, which means that heat must be absorbed
from the surroundings. In the second step the water molecules, which
are attracted to the charged ions, attach themselves to the ions. This
step releases energy, which means that heat flows to the surroundings.
The steps can be written like this:
 Step 1:
NH4NO3 + Heat ⇒ NH4+ + NO3-

Step 2:
NH4+ + NO3- + H2O ⇒ NH4+ (H2O)x + NO3- (H2O)x + Heat

Several water molecules may bond to each ion, as indicated by (H2O)x.

In the first step, heat is absorbed; in the second step, heat is released.
Overall, because more heat is involved in the first step than in the
second step, heat is absorbed from the surroundings (6 kilocalories per
mole of ammonium nitrate). This leaves the surroundings with less
thermal energy—colder.

The tendency to mess up

Portable hot and cold packs depend on reactions that are spontaneous.
Because the packs must be quick and easy to use, they require reactions
that begin as soon as the reactants are placed together and that will
continue on their own. Most spontaneous chemical reactions are
exothermic—they give off heat. This is because chemical bonds have a
tendency to shed their stored energy and release it as heat. The people
who designed hot packs found that this natural flow of energy suited
their needs perfectly. They selected the appropriate reactants, put them
in the same package, but kept them separated. When warmth is needed,
the reactants are simply mixed, and heat is produced automatically.
The tendency of stored bond energy to emerge as heat would seem
to rule out cold packs. Because endothermic reactions absorb heat, the
bonds end up with more stored energy than they started with—which
is against the natural flow of things. Yet, this occasionally happens.
When ammonium nitrate dissolves in water it gets very cold—
spontaneously. Why does this occur? Scientists explain it with a concept
called entropy.
Entropy is the degree of disorder in a system. Chemical changes
tend to go from orderly arrangements of molecules and ions to
disorderly arrangements. Nature tends to increase the amount of
messiness, or disorder, or entropy.
The natural tendency to increase entropy sometimes opposes the
tendency to release heat. When the increase in entropy is great enough,
it can drive the heat flow “backward.” The drive for high entropy
overpowers the drive to release heat. Endothermic reactions happen
spontaneously only when the reactions permits a large increase in
entropy.
 In the case of the instant cold pack, the starting material were highly
ordered: The water was pure and sealed in its own container, and the
ammonium ions and nitrate ions were arranged in an orderly pattern in
solid crystals. The substances were sorted and organized—everything
in its place. When the inner plastic bag was broken and the water
dissolved the ammonium nitrate, the orderly arrangement of the ions
was disrupted. The ions were dispersed randomly throughout the
water, and the once-pure water became “contaminated.” Disorder
reigned. The system went from very ordered to very disordered, and
the reaction was partly driven by this increase in entropy.

CAPTIONS
Figure 1. Jack Frost brand instant cold pack. The ammonium nitrate crystals and a plastic bag
of water are contained inside a heavier plastic bag. Punching the pack bursts open the inner
bag, allowing the water and ammonium nitrate to mix.
Figure 2. An instant heat pack. Remove the inner bag from the plastic envelope and shake it to
start the heat-releasing reaction. Because the reaction needs oxygen from the air, you can stop
it by returning it to the airtight envelope, then restart it later.
Figure 3. A constant-temperature heat pack. A triangular pocket contains seed crystals that can
be released through a valve by squeezing the corner. This triggers the crystallization that
warms the pack and maintains a comfortable 48 °C (118 °F).

BIOGRAPHY

Gail Marsella teaches chemistry at Muhlenberg College, Allentown, Pa.

REFERENCES
Mortimer, Charles. Chemistry, 6th ed.; Wadsworth Publishing: Belmont, Calif.
“Exothermic Composition and Warming Bag Containing Same,” U.S. Patent 4,268,272, May
19, 1981.
“Constant Temperature Device,” U.S. Patent 3,951,127, April 20, 1976.
“Refrigerating Package,” U.S. Patent 2,925,719, February 23, 1960.
Shakhashiri, Bassam Z. Chemical Demonstrations: A Handbook for Teachers of Chemistry;
University of Wisconsin Press: Madison, Wis., 1983, Vol. 1.