As more energy is added to a liquid under constant pressure (such as in an open beaker) the vapor pressure increases in a predictable way that is related to the attractive forces operating in the liquid. As the vapor pressure reaches the external pressure--usually the air pressure--the liquid boils, i.e., it changes into a gas at a constant temperature. In the gas phase, the molecules or atoms move very far apart in comparison to the liquid state, and so any intermolecular or interatomic forces that may have existed are practically negligible. This is true at ordinary pressures.
The fact that pressure is an important factor here was mentioned when we discussed phase changes in the previous unit. For that reason simple heating or cooling curves for substances leave a lot to be desired. They give no information about the effect of pressure on the phase changes.
A phase diagram is designed to do just that.