John Dalton, an English chemist, experimented on the behaviour of gas mixtures in 1801 and concluded that:

**The total pressure exerted by a mixture of non-reacting gases is equal to the sum of the partial pressures of each gas.**

This became the **Dalton’s law of partial pressures**. The law can be rationalised for an ideal gas, which is described by the ideal gas law, *pV* = *nRT*.

Consider a mixture of ideal gases in a rigid container at constant temperature. Since ideal gas particles are assumed to be point masses with elastic collisions and no intermolecular forces of interaction, they are expected to collide onto the walls of the container independently from one another in the mixture. The total pressure exerted by the gas mixture on the walls of the container is therefore proportional to the total number of particles of ideal gases in the container. The ideal gas law for a mixture of ideal gases then becomes:

where *p _{T}* is the total pressure of the gas mixture and

*n*is the number of moles of gas

_{i}*i*in the mixture. Expanding eq16, we have:

where *p _{i}* is known as the

**partial pressure**of gas

*i*:

Dividing eq18 with eq16 and rearranging, we have

where , i.e. the **mole fraction** of gas *i*.

Even though eq17 and eq19 are derived from the ideal gas law, they are applicable to real gases in many situations.

###### Question

What is the partial pressure of *He* in a balloon that is filled with *He* and 0.410 atm of *N _{2}*, assuming atmospheric pressure is 760.0 mmHg?

###### Answer

Since 760.0 mmHg = 1.000 atm, *p _{He}* = 1.000 – 0.410 = 0.59 atm