Henry's Law Calculator
Calculate dissolved gas concentration from partial pressure using C = kH × P.
Presets for CO₂, O₂, N₂, H₂ at 25 °C.
Explains carbonation, scuba bends, and blood gases.
Henry’s law says the amount of a gas that dissolves in a liquid is proportional to the partial pressure of that gas above the liquid:
C = k_H × P
C is the concentration of dissolved gas (mol/L), P is the partial pressure of the gas above the liquid (atm), and k_H is Henry’s constant for that specific gas-solvent pair at a given temperature. Different gases dissolve very differently in the same solvent, and the same gas dissolves differently in different solvents.
Henry’s constants for common gases in water at 25 °C
| Gas | k_H (mol/L/atm) | k_H (mg/L/atm) |
|---|---|---|
| CO₂ (carbon dioxide) | 0.034 | 1,500 |
| H₂S (hydrogen sulfide) | 0.10 | 3,400 |
| SO₂ (sulfur dioxide) | 1.2 | 77,000 |
| NH₃ (ammonia) | 57 | 970,000 |
| O₂ (oxygen) | 0.0013 | 42 |
| N₂ (nitrogen) | 0.00065 | 18 |
| H₂ (hydrogen) | 0.00078 | 1.6 |
| He (helium) | 0.00037 | 1.5 |
| CH₄ (methane) | 0.0014 | 22 |
| CO (carbon monoxide) | 0.00095 | 27 |
The 4-orders-of-magnitude spread (ammonia 57 vs helium 0.00037) is why gas-handling equipment varies so dramatically between species. Ammonia is essentially absorbed on contact with water; helium barely dissolves at all.
Three things you should always notice in the formula
- Linear in P — doubling the pressure doubles the dissolved amount. Triple it, you get triple. No exponential weirdness.
- Strongly temperature-dependent — k_H values above are at 25 °C only. Gas solubility decreases with temperature (opposite of most solids), which is why warm beverages go flat fast and why warm rivers in summer struggle to hold enough dissolved oxygen for fish.
- Partial pressure, not total pressure — for air (21% O₂), P_O₂ = 0.21 × P_total. At 1 atm total pressure, dissolved O₂ comes from a partial pressure of only 0.21 atm, not 1 atm.
Why warm soda goes flat
CO₂ at 25 °C has k_H ≈ 0.034 mol/L/atm. At 5 °C (cold from the fridge), k_H rises to roughly 0.063 mol/L/atm, about 85% more. The same bottle pressure holds nearly twice as much CO₂ in solution when cold. Open the cold bottle and the pressure drops; CO₂ comes out of solution slowly because it has further to travel before becoming supersaturated. Warm the bottle, and that same dissolved CO₂ exceeds the new (lower) equilibrium concentration immediately. It boils out quickly. Cold beer keeps its fizz; warm beer goes flat in minutes.
Worked example: scuba diver at 30 meters
At 30 m depth, total pressure is about 4 atm absolute (1 atm atmospheric + 3 atm from water column). Breathing air through a regulator:
P_N₂ at depth = 0.79 × 4 = 3.16 atm Dissolved N₂ in blood = 0.00065 × 3.16 ≈ 0.00205 mol/L (≈ 57 mg/L)
Compare to surface dissolved N₂: 0.00065 × 0.79 ≈ 0.00051 mol/L (≈ 14 mg/L).
A diver at 30 m has roughly 4× more dissolved nitrogen than at the surface. Surfacing too quickly drops the pressure faster than the body can offgas; dissolved N₂ comes out of solution as bubbles in the bloodstream and tissues, causing decompression sickness (the bends).
This is the entire reason for dive tables and slow ascent rates: blood and tissues need time to equilibrate to lower partial pressures as the diver ascends.
Worked example: O₂ in blood
Arterial blood has about 100 mmHg partial pressure of O₂ (0.13 atm). Per Henry’s law:
Dissolved O₂ = 0.0013 × 0.13 ≈ 0.00017 mol/L ≈ 5.4 mg/L of plasma
That’s tiny. Almost all the O₂ carried by blood is bound to hemoglobin, not dissolved freely (the binding raises the effective capacity about 70-fold). Henry’s law sets the lower bound; hemoglobin makes oxygen delivery actually work.
Where this calculator helps you
- Carbonation design: what bottle pressure to use for a target dissolved-CO₂ level
- Aquarium / pond management: minimum O₂ partial pressure needed to keep fish alive
- Wastewater stripping: how much air pressure or temperature change to drive off dissolved gases
- Hyperbaric oxygen therapy: dissolved O₂ in plasma at 2 to 3 atm of pure O₂
- Coffee and tea: dissolved CO₂ (cold brew) vs the lack of it (hot pour-over)
Limitations
- Accurate only for dilute solutions and low to moderate pressures. Above about 10 atm, real gases deviate from ideal-solution behavior.
- The constant k_H here is for pure water at 25 °C. Seawater (high salt), biological fluids, and organic solvents have different constants.
- The law assumes no chemical reaction. CO₂ in water actually equilibrates with H₂CO₃ (carbonic acid), so a small fraction of “dissolved CO₂” is chemically reacted; the calculator’s number is the total of both. SO₂ and NH₃ have even larger chemical-reaction components.
- Below the freezing point of the solvent or above its boiling point, the formula does not apply.