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Two Dalton's laws of partial pressure characterize the behavior of nonreacting, ideal gas mixtures. One of them (1801) states that the pressure of gas mixture (p) under given conditions is equal to the sum of the partial pressures of components (∑i pi), i.e., p = ∑i pi. It follows that pi = pxi, where xi is the molar concentration of the given component in a mixture. According to the other law (1803), the solubility in a liquid of each component of the ideal gas mixture in equilibrium with this liquid is proportional to the partial pressure of the given component. Both laws reflect the fact that each of the gas mixture components is a separate independent gas in the nonreacting ideal gas approximation. In this case, the properties of a mixture are additive with respect to the components' properties (with the exception of configuration functions, which include entropy).

The partial pressure of a mixture component is the pressure of a corresponding individual gas at the mixture temperature, if this gas alone occupies the given volume of a mixture. In the absence of concentration gradients in a mixture, i.e., when the mixture is in equilibrium (the diffusion processes are absent), the spatial distribution of a partial pressure of each component is uniform.

When Dalton's law applies, the relationship between the partial pressure of a component (pi) of gas mixture and its equilibrium solubility (the molar concentration xi) in a liquid obeys quantitatively the Henry Law (1807). For each individual component, pi = Kixi where Ki is the Henry constant of the given component. The applicability of Dalton's law, as well as that of Henry's law, is at low pressures. As pressure grows, the forces of intermolecular interaction and the size of the molecules become influential; in this case, the pressure of a mixture loses its additivity with respect to partial pressures of components and the solubility of components in a liquid ceases to be linear with respect to these pressures. The latter circumstance may become apparent even in the region of low-gas solubilities, i.e., in that region of concentrations from which the Henry law has been just derived. The main reason for this is the dependence of the Henry constant on the total pressure of a gas mixture.

Just like Dalton's law for pressures, the additivity of volumes of nonreacting ideal gas mixture gives rise to Amagat's Law, which states that at a given pressure the total volume of a gas mixture is equal to the sum of the volumes of individual gases. In other words, the partial molar volume of a component is equal to the molar volume of a corresponding individual gas.

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