Dalton's law

For the law of stoichiometry, see Law of multiple proportions.

In chemistry and physics, Dalton's law (also called Dalton's law of partial pressures) states that the total pressure exerted by the mixture of non-reactive gases is equal to the sum of the partial pressures of individual gases. This empirical law was observed by John Dalton in 1801 and is related to the ideal gas laws.

Mathematically, the pressure of a mixture of gases can be defined as the summation

\[P_{\text{total}} = \sum_{i=1} ^ n {p_i}\]       or      \(P_{\text{total}} = p_1 +p_2 + \cdots + p_n\)

where \(p_{1},\ p_{2},\dots,\ p_{n}\) represent the partial pressure of each component.

It is assumed that the gases do not react with each other.

\[\ p_{i} =P_{\text{total}}y_i \]

where \(y_i\) is the mole fraction of the i-th component in the total mixture of n components .

The relationship below provides a way to determine the volume based concentration of any individual gaseous component.

\[P_i =\frac{P_{\text{total}}C_i}{1,000,000}\]

where \(C_i\) is the concentration of the ith component expressed in ppm.

Dalton's law is not exactly followed by real gases. Those deviations are considerably large at high pressures. In such conditions, the volume occupied by the molecules can become significant compared to the free space between them. In particular, the short average distances between molecules raises the intensity of intermolecular forces between gas molecules enough to substantially change the pressure exerted by them. Neither of those effects are considered by the ideal gas model.

See also

• Henry's law
• Amagat's law
• Boyle's law
• Combined gas law
• Gay-Lussac's law
• Mole (unit)
• Partial pressure
• Raoult's law

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