When a gas contacts a liquid it can dissolve in that liquid. The amount of dissolved gas is often directly proportional to the partial pressure of the gas above the solution. This simple proportionality is known as Henry's law and it is a cornerstone of physical chemistry and environmental science. In its most common form the law states that the concentration of dissolved gas \(C\) is equal to a constant \(k_H\) times the gas partial pressure \(P\):
The constant \(k_H\) depends on the gas, the solvent, and the temperature. A higher constant means more gas dissolves at a given pressure. For example, carbon dioxide is quite soluble in water with a constant near 3.4×10−2 mol/(L·atm) at room temperature, which is why beverages can hold so much fizz. In contrast, noble gases like helium are far less soluble with constants on the order of 1×10−4 mol/(L·atm).
Temperature has a pronounced effect: warm soda goes flat because gas solubility decreases as temperature rises. Henry's constant can vary by 20% or more over a small temperature range. Pressure also matters. In scuba diving the elevated pressures at depth cause nitrogen to dissolve in a diver's blood. If a diver ascends too quickly, that dissolved nitrogen comes out of solution forming dangerous bubbles—decompression sickness.
Scientists often measure Henry's constant by observing how the equilibrium concentration of a gas changes with controlled pressure or by using techniques such as headspace analysis. The constant can be tabulated for many common gases in water or other solvents. Some typical values are shown below.
| Gas | kH (mol/L·atm) |
|---|---|
| Oxygen | 1.3 × 10−3 |
| Carbon Dioxide | 3.4 × 10−2 |
| Nitrogen | 6.0 × 10−4 |
| Helium | 1.0 × 10−4 |
Henry's law is critical for predicting how much oxygen dissolves in a fish tank, how carbon dioxide behaves in a bottling plant, or how pollutants like ammonia move between the atmosphere and water bodies. Environmental engineers use the law to model gas exchange across oceans and lakes, helping to estimate greenhouse gas fluxes. Beverage companies rely on Henry's law to control carbonation levels for soft drinks and sparkling water. The oil and gas industry uses it to quantify the behavior of dissolved gases in crude oil or drilling fluids.
Another area where Henry's law is indispensable is hyperbaric medicine. At high pressure more gas dissolves in blood and tissues. For divers breathing compressed air, nitrogen narcosis and the bends are direct results of Henry's law working on the human body. Hyperbaric chambers treat these conditions by carefully manipulating pressure to force gases back into solution and then removing them slowly. Understanding solubility helps physicians prevent or manage decompression sickness.
This tool accepts any Henry constant expressed in mol per liter per atmosphere along with the gas partial pressure in atmospheres. After entering the values, click the Compute Solubility button. The script multiplies the two numbers together to give the dissolved concentration in moles per liter. You can then copy the result for later use. The calculator performs no unit conversion; if you prefer other units such as mg/L or ppm you must convert the final value yourself based on the molar mass.
For example, consider carbon dioxide with a constant of 3.4×10−2 mol/(L·atm) at 25 °C. If the partial pressure is 1 atm, the solubility is simply 0.034 mol/L. At a colder temperature where the constant might rise to 4.0×10−2, the solubility at the same pressure would be 0.040 mol/L. This difference illustrates why warm soft drinks lose their carbonation so quickly. Similarly, if a diver breathes air at 3 atm pressure, the dissolved nitrogen concentration will triple compared to the surface, which underpins safe diving decompression tables.
Henry's law assumes the gas does not react chemically with the solvent and that the solution is dilute. At high concentrations or when gases dissociate or react, the linear relationship breaks down. For instance, carbon dioxide forms carbonic acid in water, so at higher pressures a more complex model is required. Real solutions can also deviate from ideal behavior due to salinity or the presence of other solutes. Still, the law provides a remarkably good approximation across many everyday scenarios.
The constant used here is one particular formulation. In the literature you may encounter alternative definitions with reciprocal units where \(k_H = P/C\). Be sure you know which version you are using when comparing values. This calculator specifically expects \(k_H\) defined as concentration divided by pressure so that \(C = k_H P\). If you have a constant in the alternate form simply take its reciprocal before entering it.
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