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Challenge: Make Ice

Updated Monday, 28th January 2008

The science behind using cooling liquids to make ice, part of the BBC/OU's programme website for Rough Science 2

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WARNING: The methods used by the Rough Scientists in their attempts to make ice involved the use of methanol, ethanol, acetone and ether. These are highly flammable liquids and their vapours form explosive mixtures with air.

They should only be used by experienced chemists who are taking the appropriate safety precautions.

The following notes are provided as background to the Rough Science programme, not as a set of instructions for carrying out the experiments.


There's nothing more refreshing on a tropical island than an ice-cold drink. But would we be able to overcome the heat and high humidity to succeed in the ice challenge?

water molecule Firstly, what are we trying to make?

Water, like most matter, comes in three distinct forms: solid (ice), liquid (water) and gas (vapour). In all three forms, water is made up of molecules consisting of two hydrogen atoms and one oxygen atom (H2O).

The difference between ice, water and vapour is to do with the bonds between these molecules.

How do the molecules behave in the three states?

In a solid there are strong bonds linking the molecules into a lattice structure. The molecules can vibrate but cannot move around, so the solid has a fixed shape. However in a liquid some of these bonds have been broken, allowing the molecules to move freely and take up the shape of the container. In water vapour even more of these bonds have been broken so that not only does the gas take up the shape of the container but it is able to expand and fill the entire volume of the container, no matter how big it is.

ice - water - water vapour

A crucial point is that in liquids molecules are moving at different speeds all the time. How fast a molecule moves depends on how much energy it has. This is called kinetic energy.

It's a liquid's average energy that we measure when we find the temperature. More energy - and so faster-moving molecules - means the liquid will be hotter, less energy means the liquid will be colder.

Heat Transfer

So, if we are going to freeze our water to make ice, we will need to slow the water molecules down by removing energy to the point where the hydrogen bonds can hold the molecules closely together in a lattice.

The First Law of Thermodynamics says that energy cannot be created or destroyed, only converted from one form to another. We can remove energy from the water by placing it in contact with something cold. We'll be using a cooling liquid.

heat transfer diagram Conduction transfers heat between the water and the cooling liquid. We don't want to contaminate the water with the cooling liquid so we keep them separate by putting the cooling liquid inside a container. The container then sits in the water.

Energy from the water is transferred through the walls of the container to the cooling liquid.

This is all in accordance with The Second Law of Thermodynamics.

We need to make sure that as we 'suck' heat out of the water and into our cooling liquid, we don't get more heat creeping into the water from outside the container. If we use a cup made of polystyrene as the container, it will insulate the water and prevent any more heat getting in.


As you might expect, we want our cooling liquid to be colder than 0°C to turn our water into ice. As the cooling liquid absorbs energy from the water there's a danger it will heat up. We can stave off this rise in temperature by using evaporation which takes heat energy away by forming vapour.

Ever wondered why you feel cold as you get out of the swimming pool, even on a hot day? That's evaporation: the water on your skin evaporates, taking heat energy from your skin, and using it to break the hydrogen bonds between water molecules as the liquid turns to vapour.

magnified molecules escaping from liquid Evaporation occurs when there is sufficient energy to enable the fastest moving molecules to break the inter-molecular bonds in a liquid at its surface.

Since the temperature we measure depends on the average speed of the molecules, if the faster molecules escape from the liquid the average molecular speed is reduced and the temperature falls. The hotter the liquid, the more molecules there are moving rapidly so the faster it evaporates and the greater the cooling effect.

Obviously there is a top limit involved here - the boiling point. At this temperature we have the maximum possible rate of evaporation and hence the fastest cooling.

Can the evaporation process occur at a lower temperature?

It is possible to make the evaporation process happen at an even lower temperature - by reducing the chance of the escaping molecules bouncing back into the liquid. To do this we need to make the molecules in the air above the liquid spread out even further, in other words reduce the air pressure.

This is the reason why water boils at a lower temperature up a high mountain where air pressure is lower than at sea level. But we don't have to climb Mount Everest to reduce the air pressure. Instead, we can use a siphon pump.

A siphon pump increases the space a fixed mass of air has to fill, thereby reducing its pressure.

What liquids could we use as a coolant?

There are a number of possibilities but it depends on two key properties: boiling point and latent heat of vaporisation.

The latent heat of vapourisation of a liquid is the amount of energy in joules required to turn one gram of the liquid into a vapour. When you get out of the shower or bath and feel cold it is because for every gram of water that evaporates from your skin 2264 joules (539 calories) of energy are being used to evaporate that gram of water. The energy is taken in the form of heat from you and therefore you feel cold.

Boiling temperature and latent heat of vapourisation of some liquids.

Liquid at atmospheric pressure Boiling Temperature (oC) Approximate Latent Heat of Vapourisation, (joules/gram)
Ether (C2H5)2O 34.5 378 ( 90 calories)
Acetone (CH3COCH3) 56.2 987 (235 calories)
Methanol (CH3OH) 64.5 1008 (240 calories)
Ethanol (C2H5OH) 78.3 848 (202 calories)
Water (H2O) 100 2264 (539 calories)


Ideally, for the best removal of heat by an evaporating liquid we need one with a high latent heat of evaporation (so that each gram evaporated removes a lot of heat) and a low boiling temperature (so that it boils easily when we apply our vacuum and therefore removes heat rapidly).

However, maybe methanol looks promising. It's got a reasonably low boiling temperature and quite a high latent heat. We knew that you can make methanol and acetone from the destructive distillation of wood. So that's what we did.

Once we had the mixture of methanol/acetone we tried to evaporate it, under vacuum, in the presence of water, to try and cool the water to ice.

The methanol/acetone mixture evaporates taking heat from the water, which is enclosed inside an insulating container with lid.

The siphon pump turned out to be too weak to evaporate the methanol fast enough to achieve a temperature low enough to turn our water to ice. So we moved on to try making ice using ether as a cooling liquid.

With a lower boiling point of 34.5°C it might just tip the balance between water and ice, even though the latent heat values show that more ether than methanol needs to be evaporated to remove the same amount of heat.

Ether can be made by heating ethanol with sulfuric acid at 140°C.

How does this reaction happen?

At 140°C in the presence of sulfuric acid two molecules of ethanol are dehydrated (lose a molecule of water) to form ether. The sulphuric acid is regenerated and reused. Above about 160°C ethene is produced instead, by loss of water from just one molecule of ethanol. Some sort of temperature control is needed.

water molecule

As the ethanol and sulfuric acid reacted together ether vapour was given off. It needed to be cooled to condense it into liquid ether, so the apparatus was joined to a long tube that went through cold water to a receiver in a cold water bath.

Because ether boils at 34.5°C it needs to be cooled to condense it from a gas to a liquid and kept cool so it remains liquid. Below around 120°C the reaction does not take place, above around 160°C ethene gas instead of ether is formed, so 140°C was chosen as a compromise between no reaction and the wrong reaction.

Any experiment with ether must be carried out in the absence of any naked flames.

After redistilling the ether (to purify it from unreacted ethanol and sulphuric acid) it was ready for making ice. The space above the ether in a small glass bottle was joined to a pipe leading to the siphon pump. As the vacuum was generated, the ether evaporated and removed heat from the water in the insulated container.

Although in the end ice eluded us, we came incredibly close and proved that a cooling effect, even in such high temperatures, could be achieved.

Web Links

The BBC is not responsible for the content of external websites.

Physics lecture notes on Pressure - from the University of New South Wales, Australia website

How do Snowflakes Form - from the site (gives lots of links to other sites about ice)

Wood distillation
Distillation of Wood Lab - on the Interactive Lab Works section of the ThinkQuest Library of Entries site - explains what happens when you heat wood splints in a test tube without burning them.

Pine Tar - History and Uses - by Theodore P. Kaye on the San Francisco Maritime National Park Association site.

A Comprehensive Look at Alcohol Distillation - an excerpt from ‘The Household Cyclopedia of General Information’ published in 1881, from the Public Bookshelf site

Mother's Mash Recipes for Alcohol Production - from Mother’s Alcohol Fuel Seminar on the Journey to Forever site

Alcohol Chemistry - from the Dawson College website

Ether and Alcohol - from The Silver Sunbeam by John Towler on the Albumen website

The Composition and Structure of Ether - from the Classes pages on the Yale University site


Advanced Physics by Tom Duncan, John Murray

Physics of Ice by Victor F. Petrenko and Robert W. Whitworth, Oxford University Press

Handbook of Chemistry and Physics by David R. Lide PhD, CRC Press


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