Dalton's Law of Partial Pressures

Total pressure of a gas mixture

John Dalton formulated this law in 1801: the total pressure exerted by a mixture of non-reacting gases equals the sum of the partial pressures of the individual gases — where the partial pressure of a gas in a mixture is the pressure that gas alone would exert if it occupied the same volume, at the same temperature, by itself.

Correcting for water vapour

Gases collected over water pick up water vapour and become moist. The pressure exerted by this saturated water vapour is called aqueous tension, and it must be subtracted to get the pressure of the dry gas:

Table 5.3 lists aqueous tension at various temperatures:

Partial pressure in terms of mole fraction

For three gases with $n_1$, $n_2$, $n_3$ moles sharing volume $V$ at temperature $T$, each contributes $p_i = \dfrac{n_iRT}{V}$, so:

Dividing $p_1$ by $p_\text{Total}$, the $RT/V$ terms cancel, leaving:

$\frac{p_{1}}{p_{\text{Total}}} = \frac{n_{1}}{n_{1}+n_{2}+n_{3}} = \frac{n_{1}}{n} = x_{1}$ …