Henry's Law Calculator

Calculate gas solubility in liquids at different pressures using Henry's law.

Gas

Partial pressure P (atm) (atm)

Henry's law constant kH (L.atm/mol, custom) (L.atm/mol)

Temperature correction factor (optional)

1=25C reference; higher T reduces solubility

Please check your inputs and try again.

Why Gas Solubility Behaves Backwards

Henry's law states that dissolved gas concentration is proportional to partial pressure above the liquid. Two consequences run counter to intuition about solids.

C = P / k_H

First, what matters is partial pressure, not total pressure. Oxygen dissolves according to its own 0.21 atm share of the atmosphere, unaffected by the nitrogen alongside it.

Second, gas solubility falls as temperature rises — the opposite of most salts. Gas dissolution is exothermic, since no energy is needed to separate gas molecules from each other, so heating drives them back out.

Gas	k_H (L·atm/mol)	Relative solubility
CO₂	29.4	High — reacts with water
O₂	769	Low
N₂	1,600	Very low
He	2,700	Lowest — relevant to deep diving

Three Everyday Consequences

Carbonated drinks. Bottled under 2–4 atm of CO₂, opening drops the partial pressure to 0.0004 atm. The dissolved gas is now vastly supersaturated and escapes. Warm drinks go flat faster because k_H rises with temperature.

Decompression sickness. At 30 m a diver breathes at 4 atm, so nitrogen dissolves in tissue at four times the surface amount. Ascending too fast drops the partial pressure faster than the nitrogen can be exhaled, and bubbles form in tissue. Helium is used in deep mixes partly because it is less soluble and diffuses out faster.

Thermal pollution. Warm water holds less oxygen — about 14.6 mg/L at 0°C against 7.6 at 30°C. Discharging cooling water into a river reduces dissolved oxygen precisely when fish metabolism demands more.

Worked Examples

Example 1: O2 in water at atmospheric P=0.21 atm

C=0.21/769

Result: C=2.73e-4 mol/L = 8.7 mg/L

Dissolved O2 in surface water

Example 2: CO2 at P=0.0004 atm (atmospheric)

C=0.0004/29.4

Result: C=1.36e-5 mol/L = 0.6 mg/L

CO2 in equilibrium with air

Example 3: Why divers get the bends

Nitrogen at 4 atm (30 m) versus 1 atm at surface

Result: Four times as much dissolved in tissue

Rapid ascent drops partial pressure faster than nitrogen can diffuse out, so it comes out of solution as bubbles within the tissue itself.

Example 4: Oxygen and water temperature

Dissolved O₂ at 0°C versus 30°C

Result: 14.6 versus 7.6 mg/L

Solubility roughly halves. Thermal discharge into rivers therefore stresses aquatic life, reducing oxygen just as metabolic demand rises.

Common Mistakes

⚠️

Using total pressure instead of partial pressure

Each gas dissolves according to its own partial pressure. Oxygen at sea level dissolves according to 0.21 atm, not 1 atm.

⚠️

Expecting gases to dissolve better when heated

Gas dissolution is exothermic, so solubility falls with temperature — opposite to most solids. This is why warm drinks go flat quickly.

⚠️

Applying Henry's law at high pressure

The linear relationship holds for dilute solutions at moderate pressure. At high pressure, deviations become significant.

⚠️

Ignoring chemical reaction

CO₂ reacts with water to form carbonic acid, so total dissolved carbon exceeds what physical solubility alone predicts. Henry's law covers only the physically dissolved fraction.

Frequently Asked Questions

Why carbonated drinks go flat?

CO2 dissolved under high pressure (2-5 atm). When opened, P drops to 0.0004 atm atmosphere. Henry's law: much less CO2 can dissolve at this P. Excess CO2 escapes as bubbles.

Scuba diving decompression?

N2 dissolves at higher pressure underwater. Rapid ascent releases N2 as bubbles in blood/tissues (decompression sickness). Must ascend slowly to allow safe off-gassing via lungs.

Why does gas solubility decrease with temperature?

Gas dissolution releases heat, since no energy is needed to separate gas molecules from one another. Warming shifts the equilibrium back toward the gas phase.

Why do carbonated drinks go flat?

Bottling uses 2–4 atm of CO₂. Opening drops the partial pressure to atmospheric levels, so the dissolved gas is supersaturated and escapes.

What causes decompression sickness?

Nitrogen dissolves in tissue at elevated partial pressure during a dive. Ascending too quickly drops that pressure faster than the gas can be exhaled, forming bubbles.

Does total pressure or partial pressure matter?

Partial pressure of the specific gas. Other gases present do not affect how much of a given gas dissolves.

Formula Explorer connections

Interpretation: This formula describes how reactants, products, ions or phases distribute when opposing processes reach equilibrium. Assumption: Use equilibrium rather than initial concentrations, correct stoichiometric exponents, and the specified temperature; activities may replace concentrations in nonideal systems.

ICE Table Calculator →Ionic Strength Calculator →Kp to Kc Calculator →Chemistry Formula Explorer →