Henry's Law Solubility Calculator
Introduction: why partial pressure controls dissolved gas
Henry's law describes a simple but powerful equilibrium idea: when a gas sits above a liquid, some gas molecules enter the liquid and dissolve. At the same time, some dissolved molecules leave the liquid and return to the gas phase. When those opposing processes balance, the dissolved amount depends strongly on the gas's partial pressure above the liquid. This calculator estimates that equilibrium concentration from two inputs: a Henry's law constant and a gas partial pressure. It is a compact tool, but it supports many familiar questions. How much carbon dioxide can sparkling water hold? Why does a warm soda lose fizz faster than a cold one? Why does increased pressure matter so much in diving, industrial carbonation, and environmental gas exchange?
In everyday language, the result tells you how much of a gas can remain dissolved in a liquid once the system has had time to settle at a fixed pressure and temperature. A larger pressure tends to push more gas into solution. A larger Henry's constant, in the unit system used on this page, also means the gas is more soluble in that liquid under those conditions. The law is widely taught because it gives a useful first estimate before you move on to more detailed chemistry, transport, or reaction models. It is especially helpful for dilute, nonreactive systems where the gas and liquid are near equilibrium.
The Henry's law formula C = kH × P
The version of Henry's law used here expresses dissolved concentration as proportional to gas partial pressure. In its most common form the law states that the concentration of dissolved gas is equal to a constant times the gas partial pressure :
That single line contains the entire calculation. Here, C is the dissolved concentration in mol/L, kH is Henry's constant in mol/(L·atm), and P is the gas partial pressure in atm. Because the units of atmospheres cancel correctly, the final answer naturally comes out in mol/L. This is one of the quickest ways to check your setup. If your constant has different units, the calculator's result will only be meaningful after you convert the constant into the format expected here.
The constant depends on the gas, the solvent, and the temperature. A higher constant means more gas dissolves at a given pressure in this convention. For example, carbon dioxide is much more soluble in water than helium, which is why beverages can hold noticeable carbonation while helium does not remain dissolved nearly as well. Temperature matters too. Warm liquids usually hold less dissolved gas, so the effective Henry constant changes with temperature. That is why published constants are normally reported for a specific gas-liquid pair and a specific temperature.
The constant used here is one particular formulation. In the literature you may encounter alternative definitions with reciprocal units where . Be sure you know which version you are using when comparing values. This calculator specifically expects defined as concentration divided by pressure so that . If you have a constant in the alternate form simply take its reciprocal before entering it.
How to use the Henry's law solubility calculator
Start with the Henry's constant for the exact gas, solvent, and temperature you care about. That last detail matters more than many people expect. A constant reported for oxygen in water at room temperature should not be assumed valid for colder water, salt water, or a different solvent. Once you have the correct value, enter it into the first field in units of mol/(L·atm). The calculator accepts any positive number, including decimal values and scientific notation formats that your browser allows in a number field.
Next, enter the gas partial pressure in atmospheres. Partial pressure is the pressure contributed by the gas of interest, not always the total pressure of the entire mixture. If a gas is pure above the liquid, then its partial pressure may equal the total pressure. But if the gas is only one component of air or another blend, use that component's share. For example, oxygen makes up roughly 21% of dry air, so under a total pressure of 1 atm the oxygen partial pressure is about 0.21 atm. That distinction is essential for realistic estimates.
Pick a gas from the preset list and the constant and molar mass fill themselves in from the Sander compilation; type your own constant and the selector switches to Custom automatically. Set the water temperature and the preset constant is corrected to it before anything else happens. After you press Compute solubility, the page reports the dissolved concentration in mol/L and again in mg/L, ppm and mole fraction. The answer is shown in scientific notation because many gas solubility values are quite small. That does not mean the result is wrong; it usually reflects the fact that gases such as nitrogen, oxygen, and helium are only modestly soluble in water under ordinary conditions. If you need the result in mg/L, ppm, or another unit, convert the mol/L value afterward using the gas's molar mass and any additional density assumptions required for your application.
A good habit is to do a quick reasonableness check before you trust the number. If you double the pressure while keeping the same constant, the predicted concentration should also double. If you keep the pressure fixed and use a larger Henry constant in this page's convention, the concentration should increase in the same proportion. Linear scaling is one of the reasons Henry's law is so useful for first-pass estimates, screening calculations, and classroom problem solving.
Worked example: carbonated water and dissolved oxygen
Take a bottle of carbonated water at 25 °C. Carbon dioxide has mol/(L·atm), and a typical headspace pressure is 1.8 atm of essentially pure CO2. Henry's law gives mol/L. Multiplying by the 44.01 g/mol molar mass turns that into about 2646 mg/L, or roughly 2.6 g of dissolved gas per litre — which is why opening the bottle releases so much of it.
Now warm the same bottle to 35 °C without letting any gas escape. The van 't Hoff correction with a coefficient of 2400 K multiplies the constant by , so the constant falls to about 2.57 × 10−2 mol/(L·atm) and the equilibrium concentration drops by nearly a quarter. That single number is the whole story of why a warm soft drink tastes flat and why brewers and bottlers carbonate cold.
A second example shows why partial pressure matters more than total pressure. Oxygen has a far smaller constant, 1.32 × 10−3 mol/(L·atm) at 25 °C, and in ordinary air its partial pressure is only about 0.21 atm even though the total pressure is a full atmosphere. The predicted dissolved oxygen is therefore 1.32 × 10−3 × 0.21 ≈ 2.77 × 10−4 mol/L, which is 8.85 mg/L. Anyone who has read a dissolved-oxygen meter will recognise that as the right neighbourhood, and it is the clearest demonstration that feeding this calculator 1 atm instead of 0.21 atm would overstate the answer by a factor of nearly five.
| Gas | kH (mol/L·atm) | Hcp (mol/m³·Pa) | −d ln H / d(1/T) (K) |
|---|---|---|---|
| Carbon dioxide | 3.34 × 10−2 | 3.3 × 10−4 | 2400 |
| Methane | 1.42 × 10−3 | 1.4 × 10−5 | 1600 |
| Argon | 1.42 × 10−3 | 1.4 × 10−5 | 1500 |
| Oxygen | 1.32 × 10−3 | 1.3 × 10−5 | 1500 |
| Hydrogen | 7.9 × 10−4 | 7.8 × 10−6 | 500 |
| Nitrogen | 6.5 × 10−4 | 6.4 × 10−6 | 1300 |
| Helium | 3.9 × 10−4 | 3.8 × 10−6 | 230 |
Two things in that table are worth pausing on, because both contradict a common intuition. Helium is not dramatically less soluble than nitrogen: on a per-atmosphere basis the two sit within a factor of two of each other, and helium is only the least soluble entry by a modest margin. What makes helium behave so differently in practice is the last column. Its temperature coefficient of 230 K is far smaller than any other gas here, so helium solubility barely responds to warming, while carbon dioxide at 2400 K loses solubility rapidly as the water heats. That is the real reason a warm soda goes flat and a helium-saturated liquid does not change much at all.
Reading the number: it's equilibrium, not speed
The number returned by this calculator is an equilibrium concentration. It does not tell you how quickly the gas dissolves; it tells you what dissolved amount is predicted once the system has had enough time to equilibrate at the stated conditions. In real equipment or natural systems, the approach to equilibrium depends on mixing, bubble size, surface area, turbulence, and contact time. If a measured concentration is lower than the Henry's law estimate, the liquid may simply not have had enough time or surface contact to reach equilibrium yet.
That distinction is important across many fields. Beverage production uses increased carbon dioxide pressure to drive more gas into solution before bottling. Environmental scientists use Henry's law to estimate how gases partition between the atmosphere and lakes, rivers, and oceans. Wastewater engineers look at dissolved oxygen transfer and stripping processes with the same basic idea in mind. In diving and hyperbaric medicine, elevated ambient pressure increases the amount of gas that can dissolve in blood and tissues. If the surrounding pressure drops too quickly, dissolved gas can come out of solution and form bubbles, which is one of the core physical ideas behind decompression sickness.
Scientists often measure Henry's constant by observing how the equilibrium concentration of a gas changes with controlled pressure or by using headspace analysis and related laboratory methods. Once the constant is known, the law becomes a practical design and interpretation tool. It helps answer questions such as whether an observed concentration is near saturation, how much dissolved gas is expected after a pressure change, and whether a gas is likely to remain in solution or escape back to the atmosphere.
Unit conventions and where the straight line bends
This calculator assumes the Henry constant is entered as concentration divided by pressure, specifically mol/(L·atm). If your source lists the constant in bars, pascals, mole fraction units, dimensionless forms, or a reciprocal definition, you should convert it before using the page. A very common error is copying a value from a reference table without checking the unit convention. Another common error is entering total pressure when the equation requires partial pressure for the gas of interest.
Henry's law works best when the gas does not react strongly with the solvent and when the solution is reasonably dilute. At higher concentrations, or when the dissolved gas participates in chemical reactions, the simple straight-line relation can become less accurate. Carbon dioxide in water is the classic example. Some dissolved CO2 hydrates and participates in acid-base chemistry, so rigorous models may need to track multiple species rather than just a single dissolved gas concentration. Salinity, dissolved solids, and nonideal interactions can also shift real solubility away from the basic prediction.
Temperature is another major limitation and one of the biggest reasons students and practitioners get different numbers from different tables. Henry constants can change noticeably over ordinary laboratory or environmental temperature ranges. Warm soda going flat is a familiar everyday demonstration: when temperature rises, the liquid typically holds less gas at the same pressure. So if accurate work matters, make sure the constant matches the actual temperature and solvent conditions rather than using a generic room-temperature value from memory.
If you want a quick reality check, ask whether the trend matches intuition. Increasing partial pressure should increase dissolved concentration, and a larger Henry's constant should also increase dissolved concentration in this convention. If either input doubles while the other stays fixed, the predicted concentration doubles too. That clean proportionality is exactly why Henry's law remains one of the most useful introductory models in physical chemistry, environmental science, and process engineering.
Reading the extra outputs and the two charts
The result table restates one answer in the units different fields actually use. Chemists want mol/L, water-quality work is done in mg/L, food and beverage specifications are often written in ppm by mass, and thermodynamic tables want a mole fraction. Because a litre of dilute water weighs close to a kilogram, mg/L and ppm by mass are numerically the same here; that convenience fails for concentrated brines or non-aqueous solvents. The mole fraction is computed against 55.345 mol of water per litre, the value for pure water at 25 °C.
The left chart is deliberately a straight line through the origin, because that is the entire content of Henry's law: double the partial pressure and you double the dissolved concentration. Your operating point is marked on it. The right chart is the part that is easy to forget — the constant itself is not constant, and the curve shows it falling as the water warms. Comparing the steepness of that curve between carbon dioxide and helium is the fastest way to see why the temperature coefficient matters as much as the constant.
Limitations of this Henry's law model
- Dilute, non-reacting solutions only. Henry's law is the low-concentration limit. Once dissolved gas is abundant enough to interact with itself or with the solvent, the straight line bends.
- Chemical speciation is ignored. Carbon dioxide, ammonia and sulfur dioxide all react with water. This page reports the physically dissolved gas, not the total of every species it becomes.
- Pure water is assumed. Salinity lowers gas solubility appreciably; seawater holds roughly 20 % less oxygen than freshwater at the same temperature and pressure.
- The temperature fit is local. The van 't Hoff coefficients are fitted near 298 K. Extrapolating to near-boiling water is an extrapolation, and the page will still return a number.
- Water vapour is not subtracted. Published saturation tables reduce the total pressure by the vapour pressure of water before splitting it among the gases; this page takes the partial pressure you supply at face value.
- It is an equilibrium, not a rate. Nothing here tells you how long the system takes to get there, which depends on mixing, surface area, bubble size and contact time.
- Custom constants carry no temperature model. If you type your own value the temperature field is inert, because no coefficient is known for it.
Sources. Every preset constant and temperature coefficient on this page comes from the same peer-reviewed compilation, converted to mol/(L·atm) by multiplying the tabulated Hcp by 101.325.
- Henry's law constants and their temperature dependence for water as solvent: R. Sander, Compilation of Henry's law constants (version 4.0) for water as solvent, Atmospheric Chemistry and Physics 15, 4399–4981 (2015).
- Definition of the standard atmosphere as exactly 101325 Pa, used for every unit conversion here: BIPM, The International System of Units (SI Brochure).
- The van 't Hoff form of the temperature dependence, and the concentration versus inverse constant conventions: IUPAC Compendium of Chemical Terminology, Henry's law.
Common questions about Henry's law solubility
Which Henry's law convention does this calculator use?
The concentration form, written C = k_H P, with k_H in mol per litre per atmosphere. Multiply your dissolved concentration by nothing and read it straight off. Many references instead publish the inverse constant K_H = P / C in litre-atmospheres per mole, or the SI form H^cp in mol per cubic metre per pascal. The results table shows your constant converted into both of those so you can check it against whatever source you are holding.
Why did the helium constant on this page change?
The earlier value of 1.0 x 10^-4 mol per litre per atmosphere was too low by roughly a factor of four. The Sander compilation gives H^cp for helium as 3.8 x 10^-6 mol per cubic metre per pascal, which converts to 3.9 x 10^-4 mol per litre per atmosphere. Helium is the least soluble gas in the preset list, but only modestly less soluble than nitrogen rather than several times less.
How is the temperature correction calculated?
With the van 't Hoff relation H(T) = H(298.15) exp(C (1/T - 1/298.15)), where C is the tabulated quantity minus d ln H over d one-over-T in kelvin. Each preset carries its own coefficient, so carbon dioxide at 2400 K responds strongly to warming while helium at 230 K barely moves. The correction is a fit around 298 K and grows less reliable far from it, especially approaching the boiling point.
Should I enter total pressure or partial pressure?
Partial pressure, always. If the gas above the liquid is pure then the two coincide, but for a mixture you need that component's share. Dry air is about 21 percent oxygen, so at a total pressure of one atmosphere the oxygen partial pressure is roughly 0.21 atmospheres. Entering the total pressure instead is the single most common way to get an answer that is several times too large.
Why does the calculator disagree with a dissolved oxygen saturation table?
Saturation tables usually build in two corrections this page leaves out. They subtract the vapour pressure of water from the total pressure before taking the oxygen share, which matters more as the water warms, and many are fitted to measurements rather than derived from a single Henry constant. Expect agreement within several percent near room temperature and a wider gap in hot water.
Does Henry's law work for carbon dioxide in water?
For the physically dissolved gas, yes, and that is what this page reports. But a fraction of dissolved carbon dioxide hydrates to carbonic acid and then dissociates to bicarbonate and carbonate, and how much depends on pH. In alkaline water the total inorganic carbon can be far larger than the Henry's law figure. Treat the result as the dissolved CO2 gas only, not as total carbon.
| Constant used at this temperature (mol/L·atm) | — |
|---|---|
| Dissolved concentration (mol/L) | — |
| Dissolved concentration (mg/L) | — |
| Mass fraction (ppm by mass) | — |
| Mole fraction in the liquid | — |
| Same constant as Hcp (mol/m³·Pa) | — |
| Inverse convention KH = P/C (L·atm/mol) | — |
Mini-Game: Henry's Law Pressure Lock
This optional mini-game turns the same calculation into a fast pressure-tuning challenge. Each sample gives you a gas, a Henry's constant, and a target dissolved concentration. Your job is to tune the chamber pressure so the live concentration meter lands inside the glowing band and stays there long enough to lock. The farther you go, the tighter the target becomes, and later phases introduce temperature swings that change the effective Henry constant on the fly.