In order to predict how systems will behave, we need to understand their properties at different equilibrium states.
A pure substance is one that is chemically homogenous (it has the same chemical composition everywhere in the substance). Examples include
a mixture of water and steam
air (all the gaseous substances are distributed evenly, everywhere)
Examples do not include
Mixtures of oil and water (oil is not soluble in water, so the two will always be separate)
Mixtures of liquid and gaseous air (different parts of air condense at different temperatures, so the mixture is not homogeneous)
While there are three main phases a substance can be in (solid, liquid and gas), there are also a number of sub-phases within these:
As you can see, the liquid phase is broken into two bands:
subcooled (compressed) liquids are when the liquid is not about to evaporate. For example, water at 20°C
saturated liquids are when the liquid is about to evaporate. For example, water at 100°C
The gas phase is also split into two bands:
saturated vapours are vapours that are about to condense. This region overlaps with the saturated liquids phase, so at 100°C, water exists as a mixture of liquid about to vaporise and vapour about to liquify
superheated vapours are gases that are not about to condense. For example, steam at 300°C
The boiling point is also referred to as the saturation temperature and pressure. This is the given temperature and/or pressure that the liquid-vapour mixture is seen. It is important to note that the temperature remains constant during a phase change.
There are two other important point between phases: the critical and the triple point.
At the triple point, gas, solid and liquid phases can coexist.
Above the critical point, liquid and vapour can no longer be distinguished. It is the high-pressure form of the liquid-vapour mixture phase, known as the supercritical fluid phase.