Mike Leahy and Jonathan Hare Copyrighted image Credit: Production team Challenge
To make a safe, portable form of heating to take to the rest of the team working up on the glacier.


Exothermic and endothermic chemical reactions
When a chemical reaction takes place, chemical bonds in the reactants are broken and new bonds are formed. Depending on the nature of the bonds involved, a reaction will either release energy to, or absorb energy from, the surroundings, as heat. When a process releases energy to the surroundings it’s called exothermic (from the Greek thermo meaning heat, and exo meaning outside). The most obvious effect of an exothermic chemical reaction is generally a rise in the temperature of the reaction mixture. However, some processes absorb heat energy from the surroundings in which case they’re called endothermic reactions (endo in Greek, means within). The effect of an endothermic chemical reaction is generally a fall in the temperature of the reaction mixture.

All combustion reactions (e.g burning of coal) are exothermic. Incredibly, the reaction between iron and moist air to produce rust is a very exothermic process that generates lots of heat. Unfortunately, this particular reaction takes place so slowly that the liberation of heat is undetectable. Pyrotechnics, explosives and fuels, on the other hand, all involve very fast and hugely exothermic chemical reactions.

For our hand warmer, do we need to use an exothermic process or an endothermic one?

To generate heat and release it to the surroundings, we need an exothermic process and one that’s suitable for use in a hot pack for warming the hands.

What properties do you think are necessary for a suitable chemical reaction for use in a hand warmer?

Well, the reaction must be portable and easily reproducible. It must generate and be able to maintain a temperature that’s neither too hot nor too cold. It must also be safe, and not involve the use of hazardous chemicals.

We decided to try three different exothermic processes to see if any met the above criteria:

Heat from cold wood ashes
Slaking lime
The thermit(e) reaction.

Heat from cold wood ashes
The first ‘reaction’ that we tried for our hand warmer involved dissolving potassium carbonate (K2CO3) in water. Potassium carbonate is one of the chemicals left behind in the ashes of burnt wood. You can extract it from the ashes by boiling them in water and filtering off any undissolved solids. The potassium carbonate dissolves in the water, along with some other salts. Reducing the aqueous solution down by evaporating off much of the water, followed by filtering, will leave a clear solution containing predominantly potassium carbonate. If you evaporate off nearly all the water, solid potassium carbonate will eventually fall out of solution as it becomes increasingly concentrated. It can then be filtered off and dried.

When we dissolved our dry, solid potassium carbonate in water, a temperature rise of only a few degrees Celsius occured . The process wasn’t therefore suitable for our purposes. We considered alternatives involving a chemical reaction of some sort.

Although dissolving potassium carbonate (K2CO3) in water is an exothermic process, it isn’t a chemical reaction, because no chemical bonds are broken or formed as it takes place. But it generates heat, so can be described as exothermic.

Slaking lime
The second exothermic process we tried for our hand warmer involved the ‘slaking’ of quicklime, which involves a chemical reaction or two.. We took ordinary chalk (calcium carbonate, CaCO3) and heated it to between 1 200 and 1 400 ºC. At these temperatures, the calcium carbonate releases carbon dioxide gas (CO2) and is converted into quicklime (calcium oxide, CaO). When you add water to the quicklime, a process called ‘slaking’, a vigorous reaction takes place, and lots of heat is generated. The product of the reaction is a compound called ‘slaked’ lime (calcium hydroxide, Ca(OH)2).

Although aqueous solutions of 'slaked' lime are alkaline, its solubility in water is so low that the solution would not represent a caustic hazard in a hand warmer. Given the right amount of lime, the amount of heat generated would be suitable for our purposes. The reaction’s also portable and easily reproducible. But let’s try one more exothermic reaction.

The thermit(e) reaction
The third process we tried was a chemical reaction called the thermit reaction (sometimes spelt thermite). This is a vigorous and highly exothermic reaction, which was used in the past to weld railway lines together. We managed to generate enough heat in our thermit reaction to weld two pieces of steel plate. The temperature generated by the reaction must therefore have been higher than 1 600 ºC.

The thermit reaction involves the reduction (Originally, reduction involved the chemical removal of oxygen from, or the addition of hydrogen to, a compound. Reduction is more accurately defined as a process in which atoms, molecules or ions gain electrons) of certain metal oxides by aluminium (Al) powder. We got our aluminium powder by filing down old drinks cans. We chose to use iron(III) oxide, or Fe2O3, known more commonly as rust. We got ours off an old corrugated iron shed at the sawmill. When we mixed our aluminium powder and rust in the right proportions, and kick-started the reaction, a glowing mass of molten iron was formed, along with an amazing amount of heat and light.

This was a little too vigorous a reaction to be used in making hand warmers so instead we decided to stick to the slaking of lime.

Making your own hand warmer
The best reaction for use in our hand warmers was the ‘slaking’ of lime. It generated enough heat for our purposes, and involved chemicals (water, calcium oxide and calcium hydroxide) that are comparatively safe, and readily available., The reaction’s also portable and easily reproducible.


2 sealable polythene bags (approx. 10cm × 10cm)
14g powdered lime (calcium oxide)
10cm3 of tap water
plastic pipette or eye dropper
bucket of cold water

Before you try the following experiment, make sure that you are wearing protective glasses nd latex or rubber gloves. Do NOT handle these chemicals with your bare hands. Don’t use a Nylon garment.

NOTE: Be careful: when mixing the contents you may find they become so hot, they may cause serious burns.

Take a small, self-sealing polythene bag and place it inside another bag of the same type and size. Using the teaspoon, carefully add 14g of powdered lime to the inner bag, and use the pipette or eye-dropper to add about 10cm3 of tap water to it. Seal up both bags securely, and mix the contents together by carefully manipulating the powder and the water with your fingers. After a few seconds, you will notice that the temperature of the hand warmer starts to rise dramatically.

When we took our ‘slaked’-lime hand warmers to the rest of the Team on the Franz Josef Glacier, we found that so much heat was generated by them that some of the plastic bags actually melted. If this happens when you try the experiment, drop the bags into a bucket of cold water and dispose of the resulting solution by flushing it down the toilet. Be sure to wash your hands immediately and thoroughly in plenty of cold running water.