The partial pressure of a gas is a measure of the thermodynamic activity of the gas's molecules. The pressure that a gas would create if it occupied the total volume available is called the gas's partial pressure. If two or more gases are mixed together in a container, the molecules move randomly and collide with each other, causing them to reach thermal equilibrium. When the gases have the same temperature, their molecules have the same average kinetic energy. Thus, each gas obeys the ideal gas law separately and exerts the same pressure on the walls of a container individually. Therefore, in a mixture of gases, the total pressure in the container is the sum of the partial pressures of the component gases, assuming ideal gas behavior and no chemical reactions between the components. This law is known as Dalton's law of partial pressures. The theory was established by the English scientist John Dalton (1766–1844). Dalton's law is consistent with the fact that pressures add up, according to Pascal's principle. In a mixture of ideal gases in thermal equilibrium, the number of molecules of each gas is proportional to its partial pressure.

Another important application of partial pressure is vapor pressure, which is the partial pressure of a vapor at which it is in equilibrium with the liquid phase of the same substance. At any temperature, the partial pressure of the water in the air cannot exceed the vapor pressure of the water at that temperature, as whenever the partial pressure reaches the vapor pressure, water condenses out of the air. Dew is an example of this condensation. The temperature at which condensation occurs for a sample of air is called the dew point. It is easily measured by slowly cooling a metal ball; the dew point is the temperature at which condensation first appears on the ball.