Henry's Law Solubility Calculator
Introduction
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.
Formula
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
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.
After you click Compute Solubility, the script multiplies the two inputs and reports the dissolved concentration in mol/L. 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.
Example
Suppose you want an equilibrium estimate for carbon dioxide dissolved in water near room temperature. A typical Henry's constant is about 3.4 × 10−2 mol/(L·atm). If the carbon dioxide partial pressure above the liquid is 1.8 atm, then the concentration predicted by Henry's law is 0.034 × 1.8 = 0.0612 mol/L. Interpreted physically, that means each liter of liquid would contain about 0.0612 moles of dissolved CO2 once equilibrium is reached, assuming Henry's law remains a good approximation over that range. Raise the pressure and the dissolved concentration rises in direct proportion.
A second example shows why partial pressure matters. Oxygen has a much smaller constant in water, roughly 1.3 × 10−3 mol/(L·atm) near room temperature. In ordinary air, oxygen's partial pressure is only about 0.21 atm. The estimated dissolved oxygen concentration is therefore 1.3 × 10−3 × 0.21 ≈ 2.73 × 10−4 mol/L. That small number is normal, not a mistake. Many environmentally important gases are present at low partial pressure and are only sparingly soluble, yet even those small dissolved amounts can matter a great deal in natural waters, fermentation tanks, and biological systems.
| 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 |
Applications and Interpreting the Result
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.
Units, Assumptions, and Limitations
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.
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.