Last updated: 2026-07-01

Understanding Henry's law

Henry's law describes how much of a gas dissolves in a liquid. At a fixed temperature, the concentration of a dissolved gas is directly proportional to the partial pressure of that gas in contact with the liquid:

where $C$ is the dissolved concentration, $P$ is the partial pressure of the gas above the liquid, and $k_H$ is the Henry constant for that gas–solvent pair at the given temperature. Double the partial pressure and twice as much gas dissolves; release the pressure and the excess gas comes back out of solution.

The calculator solves the same relation in three directions. Given any two of the three quantities, it returns the third: the dissolved concentration from $C = k_H \cdot P$, the partial pressure from $P = C / k_H$, or the Henry constant from $k_H = C / P$.

The Henry constant and its units

The constant $k_H$ captures how soluble a gas is. A large $k_H$ means a lot of gas dissolves per unit of pressure. This tool uses the solubility form, in which

so concentration is measured in mol/L and partial pressure in atm. Pressure inputs given in kPa, bar, or mmHg are converted to atm before the constant is applied.

These values fall as temperature rises — warm water holds less dissolved gas, which is why a warm carbonated drink goes flat faster than a cold one.

Conventions for the constant

Several conventions for the Henry constant coexist, and mixing them is a common source of error. They are reciprocals or rescalings of one another:

• Solubility form ($k_H = C / P$, mol·L⁻¹·atm⁻¹) — used here; increases with solubility.
• Volatility form ($k_H = P / C$, L·atm·mol⁻¹) — the reciprocal; decreases with solubility.
• Mole-fraction and molality forms — replace molarity with mole fraction or molality.
• Dimensionless form — a gas–liquid partition ratio.

Before substituting a tabulated value, confirm which convention and which units it uses, and check that it applies at the temperature of interest.

Worked example

A sealed bottle of carbonated water is pressurized with CO₂ at a partial pressure of 2 atm at 25 °C, where $k_H = 0.034\ \text{mol·L}^{-1}\text{·atm}^{-1}$ for CO₂ in water. Find the concentration of dissolved CO₂.

When the cap is removed, the CO₂ partial pressure above the liquid drops toward its value in ordinary air (about 0.0004 atm). The equilibrium concentration falls with it:

The dissolved concentration drops by more than three orders of magnitude, so the excess CO₂ escapes as bubbles until the new, much lower equilibrium is reached.

Where Henry's law applies

The proportionality between dissolved concentration and partial pressure shows up across chemistry, biology, and engineering:

• Carbonation. Soft drinks are bottled under elevated CO₂ pressure; opening the container lowers the partial pressure and the gas fizzes out.
• Aquatic life. Dissolved oxygen in lakes and aquariums tracks the oxygen partial pressure in the air above, so warming or reduced surface exchange lowers it.
• Respiration. Gas exchange in the lungs moves oxygen and carbon dioxide across membranes according to their partial pressures.
• Diving. Nitrogen dissolves into tissues at the high pressures encountered at depth and must be released gradually during ascent to avoid decompression sickness.

Limits of the law

Henry's law is a limiting law that works best for sparingly soluble gases at low partial pressures. It becomes less accurate when the gas is very soluble, when the partial pressure is high, or when the gas reacts with the solvent rather than simply dissolving. Carbon dioxide, for instance, partly reacts with water to form carbonic acid, so its apparent solubility is slightly higher than the simple physical dissolution the law describes. Because $k_H$ varies strongly with temperature, a constant measured at one temperature should not be used at another without correction.