In chemistry and physics, Dalton's law (also called Dalton's law of partial pressures) states that in a mixture of non-reacting gases, the total pressure exerted is equal to the sum of the partial pressures of the individual gases. This empirical law was observed by John Dalton in 1801 and published in 1802. and is related to the ideal gas laws.
Mathematically, the pressure of a mixture of non-reactive gases can be defined as the summation:
where p1, p2, ..., pn represent the partial pressures of each component.
where xi is the mole fraction of the ith 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
where ci is the concentration of the ith component.
Dalton's law is not strictly followed by real gases, with the deviation increasing with pressure. Under such conditions the volume occupied by the molecules becomes significant compared to the free space between them. In particular, the short average distances between molecules increases intermolecular forces between gas molecules enough to substantially change the pressure exerted by them, an effect not included in the ideal gas model.
- Raoult's law
- Henry's law
- Amagat's law
- Boyle's law
- Combined gas law
- Gay-Lussac's law
- Mole (unit)
- Partial pressure
- Vapour pressure
- Silberberg, Martin S. (2009). Chemistry: the molecular nature of matter and change (5th ed.). Boston: McGraw-Hill. p. 206. ISBN 9780073048598.
- J. Dalton (1802), "Essay IV. On the expansion of elastic fluids by heat," Memoirs of the Literary and Philosophical Society of Manchester, vol. 5, pt. 2, pages 595–602; see page 600.