If we consider a mixture of gases in a container, then they each exert a pressure on the walls of that container. ie: one gas provides a component of the total pressure, and another gas provides another component. The component that any one gas provides is called the partial pressure and is measured in Pascals (or other units of pressure).

The definition of partial pressure (for **any** gas ‘j’) is **p _{j} = x_{j}.P** where x

_{j}is the mole fraction, and P is the total pressure exerted by the sample. That is, p

_{j}is the product of the mole fraction and the total pressure.

So, if we have a container with 3 different gases in it, and we add together the partial pressure of each, we will have the total pressure. This makes sense from the definition of mole fraction, and gives rise to Dalton’s Law:

In general, the sum of the partial pressures of all the gases in a sample is always equal to the total pressure. This is true for both perfect and real gases.

Note that for a perfect gas, its partial pressure in a mixture of gases is identical to the pressure it would exert if it were alone in the same container (ie if it alone were occupying the same volume as the mixture). This is logical, as one of the assumptions for perfect gases is that there are no interactions between their molecules, so the presence or absence of other perfect gases has no effect upon the pressure that a perfect gas exerts. This is not true for real gases, due to the mutual interactions that exist between molecules.