Henry's law states that the solubility of a gas in a liquid is:

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Directly proportional to the partial pressure of the gas above the liquid. Henry's law states that the solubility of a gas in a liquid is directly proportional to the partial pressure of that gas above the liquid surface, at constant temperature. Mathematically, C = kH × P, where C is the concentration of dissolved gas, kH is Henry's law constant (specific to each gas-solvent pair at a given temperature), and P is the partial pressure of the gas. Increasing the gas pressure forces more gas molecules into solution, increasing solubility proportionally.

Gas solubility is not independent of pressure; pressure has a significant and direct effect on how much gas dissolves, as Henry's law quantifies. While temperature does affect gas solubility (gases are generally less soluble at higher temperatures), Henry's law specifically describes the pressure dependence at constant temperature, not the temperature dependence. The molecular weight of the gas alone does not determine its solubility; Henry's law constants vary widely among gases of similar molecular weights because other factors (like gas-solvent interactions and molecular polarity) also play important roles.

Henry's law has many practical applications: carbonated beverages are produced by dissolving CO₂ at high pressure, and when the bottle is opened (reducing pressure), the gas comes out of solution as bubbles. Scuba divers must ascend slowly to allow dissolved nitrogen (which enters the blood at high pressure underwater) to exit gradually; rapid ascent can cause decompression sickness ("the bends") as nitrogen bubbles form in the blood. Understanding Henry's law and its applications is useful for the CLEP Chemistry exam.

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