Henry's Law Calculator
Calculate how much gas dissolves in a liquid at a given partial pressure using Henry's law. Enter the pressure and Henry's law constant to get the dissolved concentration, or solve for the pressure instead.
Find How Much Gas Dissolves at a Given Partial Pressure
Henry's Law Formula
Where:
- \(C\) — concentration of the dissolved gas at equilibrium (mol/L)
- \(k_H\) — Henry's law constant in the solubility form (mol/(L·atm))
- \(P\) — partial pressure of the gas above the liquid (atm)
This is the solubility form of Henry's law: the amount of gas that dissolves in a liquid is directly proportional to the partial pressure of that gas in contact with the liquid. The inverse relation, \(P = C / k_H\), gives the pressure needed to reach a target concentration.
Henry's Law – Calculation Example
Example: carbon dioxide gas is in contact with water at 25 °C, where \(k_H = 0.033\) mol/(L·atm), at a partial pressure of 1 atm. How much CO₂ dissolves?
\[ C = k_H \cdot P = 0.033 \ \text{mol/(L·atm)} \times 1 \ \text{atm} = 0.033 \ \text{mol/L} \]
Result: the water holds 0.033 mol/L of dissolved CO₂ (33 mmol/L) at equilibrium.
How to use the calculator
- Open the Solve for dropdown and choose whether you want the Dissolved gas concentration (C) or the Gas partial pressure (P).
- Enter the Partial pressure (P) and pick its unit (atm, bar, kPa, mmHg or Pa) — this is the value used when solving for concentration.
- Enter the Henry's law constant (kʜ) in the solubility form \(C/P\) and select its unit. For CO₂ at 25 °C use 0.033 mol/(L·atm); for O₂ use 0.0013 mol/(L·atm).
- If you are solving for pressure, enter the Dissolved concentration (C) and its unit (mol/L, mmol/L or µmol/L) instead.
- Read the result: the main answer appears in your chosen units, with the dissolved concentration in mol/L and the partial pressure in atm shown alongside the full calculation steps.
What Henry's law describes
Henry's law states that at a constant temperature, the concentration of a gas dissolved in a liquid is directly proportional to the partial pressure of that gas above the liquid. Doubling the partial pressure doubles the amount of gas that dissolves. The relationship only holds for gases that do not react with or dissociate in the solvent — the dissolved molecules must stay in their molecular form.
The proportionality constant \(k_H\) is specific to each gas–solvent pair and changes strongly with temperature: cold water holds noticeably more dissolved gas than warm water, which is why warm soda goes flat faster and fish prefer cool, oxygen-rich water. Typical \(k_H\) values in water at 25 °C range from about 0.00037 mol/(L·atm) for helium (barely soluble) to 0.033 mol/(L·atm) for CO₂ (highly soluble) and far higher for reactive gases such as ammonia.
Interpreting the result: the calculated \(C\) is the equilibrium concentration — the level the dissolved gas eventually reaches if you wait long enough. If the liquid currently holds less gas than this, gas will keep dissolving; if it holds more, gas will bubble out until equilibrium is restored. This is the physics behind carbonated drinks (bottled under high CO₂ pressure), decompression sickness in divers, and the oxygen balance in lakes and aquariums. For temperature effects on equilibrium in general, see the Van 't Hoff equation calculator.
Reference values of Henry's law constant at 25 °C
| Gas | \(k_H\) in mol/(L·atm) |
|---|---|
| Helium (He) | 0.00037 |
| Nitrogen (N₂) | 0.00061 |
| Hydrogen (H₂) | 0.00078 |
| Oxygen (O₂) | 0.0013 |
| Argon (Ar) | 0.0014 |
| Methane (CH₄) | 0.0014 |
| Carbon dioxide (CO₂) | 0.033 |
These constants apply to pure water at 25 °C. Values in other sources may differ slightly because of different measurement methods and temperature scales.
Assumptions, limitations and practical tips
- Watch out for the two conventions. This calculator uses the solubility form \(k_H = C/P\) with units of mol/(L·atm). Many textbooks and databases instead tabulate \(k = P/C\) in L·atm/mol — the exact reciprocal. If your constant is larger than about 1, you probably have the inverse form: take its reciprocal (1 ÷ value) before entering it.
- Temperature must match. \(k_H\) increases as temperature drops, so a constant quoted at 25 °C will be wrong at 5 °C. Always check the temperature your \(k_H\) value refers to.
- Reactive gases break the rule. Gases that chemically react with water — CO₂ forming carbonic acid, NH₃, SO₂, HCl — dissolve far more than Henry's law predicts for the molecular form alone, because the reaction removes dissolved molecules and pulls more gas into solution.
- Dilute solutions and low pressures only. At very high pressures or high dissolved concentrations, deviations from proportionality appear and the simple law becomes inaccurate.
- Use partial pressure, not total pressure. For air, the oxygen concentration is governed by the partial pressure of O₂ (about 0.21 atm in dry air), not by the 1 atm total. To compare how different gases distribute between phases, see the partition coefficient calculator.
Frequently Asked Questions
How do I calculate Henry's law by hand without a calculator?
Multiply the Henry's law constant by the partial pressure: C = kH × P. For example, with oxygen at 25 °C, \(k_H = 0.0013\) mol/(L·atm) and \(P_{O_2} = 0.21\) atm in air: \(C = 0.0013 \times 0.21 = 0.00027\) mol/L. To go the other way, divide: \(P = C / k_H\). Just make sure the pressure units in \(k_H\) match the pressure you enter.
How do I convert the result from mol/L to mg/L?
Multiply the molar concentration by the gas's molar mass. For oxygen-saturated water at 25 °C: \(C = 0.00027\) mol/L × 32 g/mol ≈ 0.0087 g/L = 8.7 mg/L. Common molar masses: CO₂ 44, O₂ 32, N₂ 28, He 4 g/mol. This conversion is essential for water-quality work, since dissolved oxygen limits are usually quoted in mg/L.
My Henry's constant is in L·atm/mol or is dimensionless — how do I convert it?
If your value is in L·atm/mol, it uses the inverse convention \(k = P/C\); take its reciprocal: \(k_H = 1 / k\). A quick check: solubility-form constants for common gases are small (0.0004–0.03 mol/(L·atm)), so a value near 29 or 769 signals the inverse form. For a dimensionless constant \(K = C_{aq}/C_{gas}\), convert with \(k_H = K / (0.08206 \times T)\), where T is the temperature in kelvin.
How do I find the partial pressure of a gas in a mixture like air?
Multiply the gas's mole fraction by the total pressure: \(P_i = x_i \times P_{total}\). Dry air is about 21% oxygen and 78% nitrogen, so at 1 atm the partial pressures are 0.21 atm for O₂ and 0.78 atm for N₂. Use these partial pressures — not the total pressure — in Henry's law, because each gas dissolves according to its own pressure.
How does temperature affect Henry's law calculations?
kH depends strongly on temperature: it roughly doubles when water cools by about 20–25 °C, so cold water holds much more gas. A constant quoted at 25 °C can be off by 30% or more at 10 °C. Always look up a value measured at your temperature, or adjust it with the Van 't Hoff equation. Never pair a 25 °C constant with a different temperature and expect an accurate result.
Does Henry's law accurately describe carbonated drinks?
Only partially. Henry's law correctly predicts the physically dissolved CO₂, but soda also contains CO₂ converted to carbonic acid and bicarbonate, so the total CO₂ content is higher. In practice, bottled drinks are carbonated at roughly 2–4 atm of CO₂, and the law explains why gas escapes when the bottle is opened: the pressure drops suddenly, the equilibrium concentration falls, and the excess CO₂ fizzes out.
What is a good dissolved oxygen level, and what does Henry's law predict for natural water?
At 25 °C in contact with air, Henry's law predicts about 8.7 mg/L of dissolved O₂; at 5 °C the value rises to roughly 11–13 mg/L. Most fish need at least 5–6 mg/L to thrive, and levels below 2–3 mg/L cause stress or death. Henry's law gives the saturation value only — actual lakes and rivers can be below it when decomposition or warm weather consumes oxygen.