Ask OxSciBlog: under pressure

As part of our 'Any questions?' campaign a question sent in by Silvan Griffith is answered by Claire Vallance from Oxford's Department of Chemistry.

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Q: I have always understood that pressure, temperature and appearance of a substance are directly related: the higher the pressure, the lower the temperature, the more solid it becomes due to the atoms moving less. Water reaches its highest density at +4°C. If one would decrease the pressure, would it become colder or warmer?

Claire Vallance: You are correct that both the pressure and the temperature affect the ‘appearance’ or ‘solidness’ of a substance. However, temperature and pressure are not locked together in quite the way you describe.

An equation known as the phase rule (described in detail in any thermodynamics textbook) predicts that for pure water we can vary both pressure and temperature independently. For example, liquid water may be turned into solid water (ice) either by reducing the temperature or by increasing the pressure sufficiently.

When we reduce the temperature we reduce the kinetic energy of the water molecules, so that they move around more slowly. Within a liquid, molecules are constantly colliding with each other, and at high temperatures the collisions are energetic enough that the relatively weak attractive forces between individual molecules have little effect and the molecules simply bounce off each other.  However, once the temperature approaches the freezing point, collisions occur so slowly and with so little energy that the intermolecular forces take over and the molecules start to stick together.

As the temperature falls even further we eventually end up with the molecules locked into the lattice structure of solid ice. A similar result may be obtained by reducing the pressure, but in this case the mechanism for crystallisation into the ice structure is not that the collisions become less energetic (assuming that we keep the temperature the same as we increase the pressure), but that the molecules are forced closer together on each collision.

Intermolecular forces are strongly dependent on distance, and are much stronger at smaller separations. At high enough pressures, water can be made to form ice even at room temperature and beyond.

This behaviour can be summarised on a phase diagram. The phase diagram for water [part of which is shown below] reveals which phase (solid, liquid or gas) is most stable at a given temperature and pressure.

 At low temperatures and high pressures (top left of the diagram), the solid phase is most stable, while at low pressures and high temperatures (bottom right of the diagram) the gas phase is most stable.

At intermediate temperatures and pressures we have the liquid phase. The lines, or ‘phase boundaries’, on the diagram show conditions under which two phases can exist together in equilibrium. For example, the line separating solid and liquid allows us to determine the freezing point of a substance at any pressure, and the line separating liquid and gas does the same for the boiling point. We can use the diagram to explore what would happen in your example of water at 4 °C.

Starting at a pressure of 1 atmosphere and a temperature of 4 °C (filled circle), reducing the pressure corresponds to following the vertical line in the direction of the arrow – note that the temperature doesn’t change. As we reduce the pressure, the liquid will become less dense, until when we reach a low enough pressure (just less than 0.01 atm on the diagram, the pressure you would find at an altitude of 32,000 m!) we cross the phase boundary between liquid and gas, and the water evaporates.

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Dr Claire Vallance is based at Oxford's Physical & Theoretical Chemistry Laboratory.