Solubility of Gases in Liquids: Henry's Law

When a gas is brought into contact with a liquid, some of it dissolves until the rate at which gas molecules enter the liquid equals the rate at which dissolved molecules escape back into the gas phase — a dynamic equilibrium whose extent is described quantitatively by Henry's law. Henry's law states that, at a constant temperature, the partial pressure of a gas in the vapour phase above a solution is directly proportional to the mole fraction of the gas dissolved in the solution:

Here $p$ is the partial pressure of the gas above the solution, $x$ is the mole fraction of the gas actually dissolved in the liquid, and $K_H$ is the Henry's law constant, a proportionality constant that is characteristic of the particular gas-solvent pair and that changes with temperature. A graph of $p$ against $x$ for a gas obeying Henry's law is a straight line through the origin, whose slope is $K_H$; different gases dissolving in the same solvent give straight lines of different slopes, so a gas with a larger $K_H$ requires a higher pressure to achieve the same mole fraction dissolved — in other words, a larger $K_H$ corresponds to lower solubility at a given pressure.

Gas solubility, and hence $K_H$, depends on the identity of both the gas and the liquid, and — critically — on temperature: since dissolving a gas in a liquid is normally an exothermic process (the gas loses translational freedom as it becomes solvated), Le Chatelier's principle predicts, and experiment confirms, that gas solubility decreases as temperature rises, which appears in Henry's law as $K_H$ increasing with temperature for a fixed mole fraction. …