Henry’s law describes a simple relationship: the amount of gas that dissolves in a liquid is directly proportional to the pressure of that gas above the liquid. Push more gas against a liquid’s surface, and more of it dissolves. Reduce the pressure, and gas escapes. First described by English chemist William Henry in 1803, this principle explains everything from the fizz in a soda bottle to how oxygen enters your bloodstream.
The Basic Formula
Henry’s law is expressed as C = kP, where C is the concentration of dissolved gas, P is the partial pressure of the gas above the liquid, and k is a proportionality constant specific to each gas-liquid pair. The constant k (often written as kH) is determined experimentally by measuring how much of a particular gas dissolves in a particular liquid at a known pressure and temperature.
A useful rearrangement compares two different conditions: S1/P1 = S2/P2, where S represents solubility and P represents pressure. This form lets you predict what happens when pressure changes. If you double the pressure of a gas over a liquid, the amount of gas that dissolves also doubles.
For oxygen dissolving in water at 25°C, the Henry’s law constant is approximately 0.0013 mol/(kg·bar). That number is small because oxygen is not very soluble in water, which is why aquatic organisms need specialized structures like gills to extract enough of it.
Why Temperature Matters
The “constant” in Henry’s law is not truly constant. It changes significantly with temperature, and this is one of the most common mistakes people make when applying the law. The constant typically increases with temperature at lower temperatures, reaches a maximum, and then decreases at higher temperatures. The exact peak depends on the specific gas and liquid involved.
The temperature sensitivity is dramatic. A shift of just 10°C can cause the Henry’s law constant to change by a factor of two. Near room temperature, most of this variation comes from changes in the gas’s vapor pressure. At higher temperatures, interactions between the gas and the liquid itself start to matter more, making the relationship increasingly nonlinear. For any serious calculation, you need a constant measured at or near your actual temperature.
The Soda Bottle Example
Carbonated drinks are the most intuitive demonstration of Henry’s law. During bottling, the beverage is exposed to carbon dioxide at high pressure, forcing a large amount of CO2 into solution. The container is then sealed, trapping the high-pressure gas above the liquid and keeping the CO2 dissolved.
When you crack the cap, that familiar hiss is the pressurized gas escaping. The pressure above the liquid drops suddenly, and Henry’s law dictates that the liquid can no longer hold as much dissolved CO2. Gas molecules come out of solution, forming the bubbles you see rising through the drink. Leave the bottle open long enough and the soda goes flat, because the pressure of CO2 above the liquid has equalized with the tiny amount of CO2 in the surrounding air.
How It Works in Your Lungs
Every breath you take depends on Henry’s law. When air reaches the tiny sacs in your lungs (alveoli), oxygen has a higher partial pressure in the air than in the blood flowing past. That pressure difference drives oxygen to dissolve into the blood. Carbon dioxide moves in the opposite direction, from the blood (where its partial pressure is higher) into the air you exhale.
An important detail: gases move between your lungs and blood based on differences in partial pressure, not concentration. This distinction matters because only freely dissolved gas molecules contribute to partial pressure. Oxygen bound to hemoglobin and CO2 converted into bicarbonate don’t count. Your body exploits this: hemoglobin grabs dissolved oxygen, lowering the blood’s oxygen partial pressure, which pulls even more oxygen across from the lungs. It is an elegant system built on the physics Henry described over two centuries ago.
Decompression Sickness in Divers
Scuba divers breathe air at pressures that increase with depth. At greater depth, the higher pressure forces more nitrogen (which makes up about 78% of air) to dissolve into the blood and tissues, exactly as Henry’s law predicts. The nitrogen doesn’t do anything useful in your body. It just accumulates.
The danger comes during ascent. If a diver rises too quickly, the surrounding pressure drops faster than the dissolved nitrogen can be carried back to the lungs and exhaled. The nitrogen supersaturates in the blood and tissues, then forms bubbles, much like opening a shaken soda. These bubbles can block blood vessels and damage tissues, causing decompression sickness (often called “the bends”). The standard prevention is ascending slowly and making decompression stops, giving the body time to off-gas nitrogen gradually through normal breathing.
Oceans and Climate
Henry’s law also governs how much CO2 the world’s oceans absorb from the atmosphere. The ocean surface is in constant contact with atmospheric gases, and the amount of CO2 that dissolves depends on both the gas’s partial pressure in the air and the water temperature. Cold water dissolves more gas than warm water, which is why polar oceans are more effective carbon sinks than tropical ones.
This creates a concerning feedback loop. As global temperatures rise, ocean water warms and its capacity to absorb CO2 decreases. Some dissolved CO2 may even be released back into the atmosphere, raising concentrations further and trapping more heat. At the same time, rising atmospheric CO2 levels increase the partial pressure of the gas, which pushes more CO2 into the ocean. These two forces work against each other, but the net effect of warming is a reduction in the ocean’s ability to buffer our emissions.
When the Law Breaks Down
Henry’s law works well under a specific set of conditions: relatively low gas concentrations, moderate pressures, and gases that don’t react chemically with the liquid they’re dissolving into. When any of these conditions are violated, the simple proportional relationship starts to fall apart.
At very high pressures, gas molecules in solution begin to interact with each other and with the solvent in ways the law doesn’t account for. Gases that react with their solvent also deviate. CO2 in water, for instance, partially reacts to form carbonic acid, meaning the total amount of carbon dioxide species in solution is higher than what Henry’s law alone would predict from the dissolved gas. For gases like oxygen and nitrogen that don’t react significantly with water, the law holds well across a wide range of everyday conditions.