Henry's Law and Gas Solubility

Core Idea

Henry's law states that the solubility of a gas is directly proportional to the partial pressure of that gas above the solution (at constant temperature). Increased pressure forces more gas molecules into solution; increased temperature decreases gas solubility by increasing molecular motion. This law applies to ideal solutions of gases.

Explainer

From Dalton's law and the concept of partial pressure, you know that each gas in a mixture exerts its own pressure independently, proportional to its mole fraction. Henry's law extends this idea to the liquid phase: the amount of gas that dissolves in a liquid is directly proportional to the partial pressure of that gas above the liquid. The mathematical expression is simple: C = kH × P, where C is the concentration of dissolved gas, P is the partial pressure of the gas above the solution, and kH is the Henry's law constant — a value specific to each gas-solvent pair at a given temperature.

The physical picture is straightforward. Gas molecules are constantly striking the liquid surface. Some bounce off; some penetrate and dissolve. At the same time, dissolved gas molecules near the surface escape back into the gas phase. At equilibrium, the rate of dissolution equals the rate of escape. If you increase the partial pressure of the gas — by pumping more gas into the space above the liquid — more molecules strike the surface per second, and more dissolve until a new, higher equilibrium concentration is reached. Double the pressure, double the dissolved concentration. This linear relationship is Henry's law.

The most familiar example is carbonated beverages. Carbon dioxide is dissolved in soda under high pressure (typically 2–4 atmospheres of CO₂). When you open the bottle, the partial pressure of CO₂ above the liquid drops to atmospheric levels (about 0.0004 atm), and Henry's law predicts a dramatic decrease in solubility. The excess dissolved CO₂ comes out of solution as bubbles — that is the fizz. If you leave the bottle open, it eventually goes flat as dissolved CO₂ equilibrates with the tiny partial pressure of CO₂ in the atmosphere. Divers encounter the same principle: at depth, the elevated pressure causes more nitrogen to dissolve in blood. Rising too quickly drops the pressure faster than nitrogen can leave the blood gradually, forming bubbles that cause decompression sickness (the bends).

Temperature works against gas solubility — unlike most solids, gases become *less* soluble as temperature rises. Higher temperature gives dissolved gas molecules more kinetic energy, making them more likely to escape the liquid. This is why a warm soda goes flat faster than a cold one, and why aquatic organisms can be stressed by warm water that holds less dissolved oxygen. Henry's law applies best to dilute solutions of gases that do not react chemically with the solvent. Gases like HCl or NH₃ that react extensively with water deviate from Henry's law because the chemical reaction removes dissolved gas, pulling more into solution than pressure alone would predict.