Remember that the pressure of a gas is due to the elastic shocks of the molecules with the walls of the container.

As an example we look at one mol of oxygen (O₂) in a container of 22.4 liter at 0 °C. Then the oxygen pressure is one bar. If we replace half of the oxygen by nitrogen the pressure in the container will not change, because the kinetic energy of the molecules depends on the temperature only (E_cin = 3/2 ? k ? T) and not on their mass. The collisions of the oxygen molecules with the wall and of the nitrogen molecules with the wall are responsible for exactly the same pressure. Half of the total pressure of 1 bar will be caused by the oxygen and the other half by the nitrogen. In this example the partial pressure of O₂ is 0.5 bar as is the partial pressure of N₂. One may generalize: In a gas mixture, the partial pressure exerted by one component is proportional to its concentration.

The total pressure is the sum of the partial pressures.

This is called Dalton?s law (John Dalton, 1766 ? 1844). It may also be expressed in another way: In a mixture the partial pressure of a gas corresponds to the pressure the gas would exert if it were alone in the container.

     Example: Air consists of 78 % N₂, 21 % d?O₂, 1 % Ar and 0.03 % CO2 (percent of volume = percent of moles). When the pressure of the air is 1 bar, the partial pressures are: P(N₂) = 0.78 bar, P(O₂) = 0.21 bar, P(Ar) = 0.01 bar and P(CO₂) = 0.0003 bar.