At any temperature, equilibrium exists when the escaping tendency (or free energy per mole) of molecules in the liquid and vapor is the same. If the temperature is raised: the escaping tendency of the liquid increases.

More liquid will evaporate until the vapor pressure rises to the point at which the escaping tendency of vapor molecules back into the liquid matches the tendency of the liquid molecules to evaporate. This equilibrium partial pressure of vapor above a liquid is known as the equilibrium vapor pressure of the substance.

The vapor pressure of water at room temperature (25° C) is 0.0313 atm, or 23.8 mm of mercury (760 mm Hg = 1 atm). This means that if a still body of air over a lake is saturated with moisture at 25° C, there will be 0.0313 atm of water vapor in the air, and 0.969 atm of O₂, N₂, and other gases. The way in which equilibrium vapor pressure changes with temperature is shown in the graph on the right. At 0° C the molecules of liquid water move slowly, their escaping tendency is small, and the equilibrium vapor pressure above the liquid is only 4.6 mm Hg. At 50° C it increases to 92.5 mm Hg, and at 100° C it equals 760 mm Hg or 1 atm pressure. This is the definition of the boiling point of a liquid -the temperature at which its vapor pressure equals the external pressure.