Calculate how much gas dissolves in water using Henry's law. Change the gas, its partial pressure, the water temperature and the liquid volume to see the dissolved concentration and total dissolved amount update in real time, and build intuition for fizzy drinks and dissolved oxygen.
Parameters
Gas
Sets the Henry's constant k_H and molar mass M
Partial pressure P
kPa
Partial pressure of the gas above the liquid. Atmospheric pressure is about 101 kPa
Water temperature T
°C
The higher the temperature, the less gas dissolves
Liquid volume (water) V
L
Used to compute the total dissolved amount of gas
Results
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Henry's constant k_H(T) (mol/(L·kPa))
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Dissolved concentration (mol/L)
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Mass concentration (mg/L)
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Total dissolved gas (mg)
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Solubility vs 25 °C (×)
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Solubility rating
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Gas-liquid equilibrium — dissolution animation
Gas molecules at partial pressure P fill the headspace above; dissolved gas molecules fill the liquid below. Molecules crossing the interface represent equilibrium, and the density tracks the partial pressure and dissolved concentration.
The dissolved concentration C is proportional to the partial pressure P of the gas (Henry's law). The constant k_H differs from gas to gas and decreases as the temperature T rises (solubility falls).
T: absolute temperature [K], M: molar mass of the gas [g/mol]. The mass concentration C_mass is in mg/L; the total dissolved amount is the mass concentration multiplied by the liquid volume V.
What is the Henry's Law Gas Solubility Simulator?
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I have heard the name "Henry's law", but in a nutshell, what does it actually describe?
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In a nutshell, it tells you how much of a gas will dissolve in water. There is really just one idea: at equilibrium, the concentration of a gas dissolved in the liquid is directly proportional to the pressure that gas pushes on the surface — its partial pressure. As a formula, C = k_H·P. So double the partial pressure and you double the amount dissolved. Beautifully simple, isn't it?
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So it just scales with partial pressure. Does the type of gas not matter, then?
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Oh, it matters — that is the interesting part. The proportionality constant k_H is completely different from gas to gas. Switch the gas on the left. At the same 100 kPa, CO₂ dissolves to over a thousand mg per litre, while oxygen and nitrogen reach only a few to a dozen mg. CO₂ reacts weakly with water to form carbonic acid, so it dissolves well; a gas like nitrogen that does not mix with water barely dissolves at all.
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When I raise the water-temperature slider, the solubility keeps dropping. Does warming the water make gas dissolve less?
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Yes — and that is the opposite of solids. Sugar dissolves more easily when you heat the water, but gases dissolve less. Dissolving a gas in water is exothermic, so raising the temperature shifts the equilibrium toward "not dissolved". That is why warm rivers and lakes hold less dissolved oxygen in summer and fish can suffocate. The small bubbles that appear first when you heat a pot of water are warmed water expelling the air it had dissolved.
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I see! Can this also explain why opening a fizzy drink makes bubbles?
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That is the textbook example of Henry's law. A carbonated drink is bottled by forcing CO₂ into the water under a high CO₂ partial pressure of 2-4 atmospheres. The instant you open the cap, the CO₂ partial pressure above the liquid drops to atmospheric level. By C = k_H·P the solubility drops too, and the CO₂ that can no longer stay dissolved escapes as bubbles — the fizz. And if the drink is warm, k_H is even smaller, so a lukewarm soda goes flat in no time.
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It is striking how everyday it is. Is this law used in engineering practice too?
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Constantly. Designing the carbonation equipment for soft drinks, the absorption columns (scrubbers) that dissolve harmful gases from flue gas into water, the deaeration of boiler feedwater, and even decompression sickness in divers — the nitrogen that bubbles out of the blood on a too-rapid ascent — all come back to Henry's law. Whenever you ask "how much gas dissolves, or comes out", this law is the first tool you reach for.
Frequently Asked Questions
Henry's law states that, at a constant temperature, the concentration of a gas dissolved in a liquid is proportional to the partial pressure of that gas above the liquid. Written out, C = k_H·P, where C is the dissolved concentration, P is the partial pressure of the gas above the liquid and k_H is the Henry's-law (solubility) constant. Double the partial pressure and the dissolved concentration doubles. The constant k_H differs greatly between gases and falls as the temperature rises.
Dissolving a gas in a liquid is an exothermic process, so solubility falls as the temperature rises. This tool applies a van't Hoff form, k_H(T)=k_H(298)·exp[C_vH(1/T−1/298)]. For CO₂, the k_H at 0 °C is about twice the value at 25 °C and about a quarter of it at 80 °C. The same temperature dependence is why warm rivers and lakes hold less dissolved oxygen for fish in summer, and why a warm, opened soda goes flat so quickly.
A carbonated drink is bottled with CO₂ dissolved under a high CO₂ partial pressure, typically 2-4 atmospheres. When you open it, the CO₂ partial pressure above the liquid drops abruptly to atmospheric level, so by Henry's law the solubility drops too, and the CO₂ that can no longer stay dissolved escapes as bubbles. That is the fizz. In a warm drink, k_H is smaller still, so it goes flat faster.
In water at 25 °C the Henry's constant of CO₂ is about 26 times that of O₂ and about 52 times that of N₂, so CO₂ is far more soluble. At the same partial pressure, CO₂ dissolves to hundreds or over a thousand mg per litre, while O₂ and N₂ reach only a few to a dozen mg per litre. CO₂ reacts weakly with water to form carbonic acid, whereas the low-polarity O₂, N₂ and CH₄ do not mix well with water and dissolve poorly.
Real-World Applications
Beverages and carbonation: In making soft drinks, beer and sparkling water, a high CO₂ partial pressure is applied to chilled water or syrup to raise the solubility. From Henry's law, manufacturers back-calculate the pressure and temperature needed to reach a target carbonation level (gas volumes). Because k_H is larger at lower temperatures, carbonation is always done while the liquid is cold.
Environment and water treatment: The dissolved-oxygen content (DO) of rivers, lakes and aquaculture ponds is critical to the survival of aquatic life. As water temperature rises, O₂ solubility falls, so oxygen depletion (hypoxic zones) is more likely in summer or where warm effluent is discharged. Aeration equipment is sized with Henry's law to estimate how far oxygen can dissolve at a given water temperature and pressure.
Gas absorption columns and scrubbers: Chemical plants and power stations use absorption columns to remove CO₂, SO₂ and ammonia from flue gas by dissolving them into water or an absorbent. A gas with a larger Henry's constant (more soluble) can be removed with a smaller liquid flow, which feeds directly into the column height and circulation rate. Poorly soluble gases need chemical absorption, where a chemical reaction is added.
Diving and decompression sickness: As a diver descends, the surrounding pressure rises and, by Henry's law, more nitrogen dissolves into the blood and tissues. On a rapid ascent the pressure drops sharply, and the nitrogen that can no longer stay dissolved forms bubbles in the blood, causing decompression sickness — "the bends". Ascending slowly with decompression stops releases the gas gradually and prevents bubble formation.
Common Misconceptions and Pitfalls
The most common mistake is confusing total pressure with partial pressure. What drives Henry's law is the partial pressure of that gas alone, not the total pressure of a gas mixture. Air is about 78% nitrogen and 21% oxygen, so even under 101 kPa of atmospheric pressure the partial pressure of oxygen is only about 21 kPa and nitrogen about 79 kPa. To calculate dissolved oxygen you must use the 21 kPa partial pressure, not the 101 kPa total — using the total gives a value roughly five times too high. This tool uses the partial pressure you enter directly, so enter the partial pressure of the gas of interest correctly.
Next, assuming Henry's law holds at any pressure or concentration. The law is an approximation that works well in the dilute regime where a gas simply dissolves physically into the liquid. At very high pressures, when a chemical equilibrium is involved (CO₂ reacting with water to form carbonic acid and bicarbonate ions), or in the high-concentration regime where a large amount of gas dissolves, the proportional relationship breaks down. The CO₂ result in this tool uses a Henry's constant for physical dissolution and is an estimate; it does not track the exact speciation under conditions where the pH shifts strongly.
Finally, do not assume the Henry's constant is a single number. The Henry's constant comes in a solubility form (k_H = C/P) and a volatility form (k_H = P/C, with units of pressure over concentration), and references not uncommonly use opposite definitions and inverted units. When you cite a value, always check the definition, the units and the reference temperature (usually 298.15 K, 25 °C). This tool adopts the k_H = C/P form, in mol/(L·kPa), with the van't Hoff temperature correction.
How to Use
Enter the partial pressure of gas (kPa) in the partialPNum field; range 0.1–500 kPa covers atmospheric to pressurized vessel conditions.
Set temperature (°C) in tempCNum; the simulator recalculates Henry's constant k_H(T) for CO₂, O₂, N₂, or custom gases using temperature-dependent coefficients.
Specify liquid volume (L) in liquidVolNum to compute total dissolved mass in milligrams; typical lab scales use 0.5–10 L.
Review output metrics: dissolved concentration (mol/L), mass concentration (mg/L), and solubility ratio versus 25 °C reference.
Worked Example
CO₂ sparging into water at 20 °C: partial pressure 85 kPa, liquid volume 2.5 L. Henry's constant k_H(20 °C) ≈ 0.034 mol/(L·kPa) for CO₂. Dissolved concentration = 0.034 × 85 = 2.89 mol/L. Mass concentration = 2.89 × 44 g/mol = 127.2 mg/L. Total dissolved CO₂ = 127.2 × 2.5 = 318 mg. At 25 °C, k_H ≈ 0.033, so solubility ratio ≈ 1.03× reference.
Practical Notes
Temperature sensitivity dominates for volatile gases: O₂ solubility decreases ~2% per °C rise in aqueous systems between 10–40 °C; use precise thermometer readings in bioreactors.
Ionic strength effects—adding 0.5 M NaCl reduces CO₂ solubility by ~8%; account for salinity in marine or fermentation broth contexts.
Pressure cycling in dissolved oxygen measurement: pressurized headspace (>1 atm) increases k_H linearly; depressurization strips gas rapidly, critical for anaerobic culture transfers.