This diagram is a highly simplified representation of the phase diagram of water. At high pressures (greater than 2000 atmospheres) various allotropes of solid ice have been observed – these are omitted for clarity.

The liquid-vapour phase boundary in the diagram summarises the variation in the vapour pressure of liquid water with temperature. Conversely, we can look at it as representing the variation in the boiling temperature of water with pressure.

The solid-liquid boundary represents the variation of the melting point with pressure. Its very steep slope indicates that the pressure changes required to noticeably affect the melting point are enormous. The line also has a negative gradient up to about 2000 atm, which is highly unusual, indicating as it does that an increase in pressure lowers the melting point. The reason behind this behaviour can be traced to the fact that ice has a larger molar volume than liquid water close to its freezing point. (Due to the hydrogen bonding between water molecules in the solid which enforces a fairly open cage-like structure.) Raising the pressure thus makes it more favourable for the solid to transform into the liquid, as it can reduce its volume (and the pressure acting upon it) by doing so.

This is an example of Le Chatelier’s principle, which states that a system in equilibrium (such as one at the point of a phase transition) will alter its position of equilibrium in such a way as to minimise the effect of any applied constraint. i.e. the increase in pressure upon the ice may be somewhat offset if it can reduce its volume, as this will increase the volume (and hence decrease the pressure) of its surroundings. The volume reduction is achieved by a shift in the position of equilibrium to favour liquid water, the denser phase, with the overall result that a pressure increase converts ice to liquid water.