In a dilute binary solution, the solvent (major component) approximately obeys Raoult’s law, while the solute (minor component) obeys Henry’s law. Henry’s law states the solute’s partial pressure is proportional to its mole fraction ([latex]P_{solute} = K_H x_{solute}[/latex]), where [latex]K_H[/latex] is the Henry’s law constant. Raoult’s law is a limiting case where [latex]K_H = P_{solvent}^*[/latex].
This relationship provides a more complete thermodynamic description of real, dilute solutions. Raoult’s law works well for the solvent because its molecules are predominantly surrounded by other solvent molecules, an environment similar to the pure liquid. The mole fraction [latex]x_{solvent}[/latex] is close to 1, and its behavior is nearly ideal. Its chemical environment is essentially unchanged from its pure state.
Conversely, the solute molecules are scarce and are entirely surrounded by solvent molecules. This environment is very different from that of the pure solute. Therefore, its tendency to escape into the vapor phase is not proportional to its pure vapor pressure but to an empirical constant, [latex]K_H[/latex], which reflects the specific solute-solvent interactions. Henry’s law captures this behavior. The Gibbs-Duhem equation mathematically proves that if one component in a binary mixture obeys Raoult’s law over a certain concentration range, the other component must obey Henry’s law in the same range. The two laws thus describe the limiting behaviors at the two extremes of the concentration range for any binary mixture.
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