Dalton's law

(physics) The law that the pressure of a gas mixture is equal to the sum of the partial pressures of the gases composing it. Also known as law of partial pressures.

The total pressure of a mixture of gases is the sum of the partial pressures of each gas in the mixture. The law was established by John Dalton (1766–1844). In his original formulation, the partial pressure of a gas is the pressure of the gas if it alone occupied the container at the same temperature. Dalton's law may be expressed as P = P_A + P_B + ···, where P_J is the partial pressure of the gas J, and P is the total pressure of the mixture; this formulation is strictly valid only for mixtures of ideal gases. For real gases, the total pressure is not the sum of the partial pressures (except in the limit of zero pressure) because of interactions between the molecules.

In modern physical chemistry the partial pressure is defined as P_J = x_JP, where x_J is the mole fraction of the gas J, the ratio of its amount in moles to the total number of moles of gas molecules present in the mixture. With this definition, the total pressure of a mixture of any kind of gases is the sum of their partial pressures. However, only for an ideal gas is the partial pressure (as defined here) the pressure that the gas would exert if it alone occupied the container. See also Avogadro number; Gas; Kinetic theory of matter.

The principle that states that the pressure of a mixture of gases equals the sum of the partial pressures of the constituent gases.

A law stating that the total pressure of a mixture of two or more gases or vapours is equal to the sum of the pressures that each gas would exert if it was present alone and occupied the same volume as the whole mixture.

The pressure exerted by a mixture of nonreacting gases is equal to the sum of the partial pressures of the separate components.

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 a gaseous mixture is equal to the sum of the partial pressures of each individual component in a gas mixture. 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

      or      

where represent the partial pressure of each component.

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

where the mole fraction of the i-th component in the total mixture of m components .

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

where: 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. Moreover, 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.

External links

• Dalton's Law

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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