Henrys Law Gas Solubility Interactive Calculator

If you need to know how much gas is going to stay dissolved in a liquid under pressure, you’ll hit this in a lot of practical settings—carbonating drinks, keeping fish alive, or running a gas scrubber at a factory. This Henry's Law calculator lets you run the numbers for dissolved gas, needed pressure, or the Henry's constant using data like partial pressure, temperature, volume, and molar mass. These numbers are important when you're making beverages, managing aquaculture, treating wastewater, or dealing with gas balance in diving situations. The rest of the page covers the basic formula, a typical carbonation example, how temperature correction works, and a FAQ.

What is Henry's Law Gas Solubility?

Henry's Law is pretty direct: the amount of gas dissolved in a liquid is proportional to the partial pressure of that gas above the liquid. If you double the pressure, you double the dissolved concentration. That's it.

Simple Explanation

With a sealed soda bottle, the CO₂ above the liquid pushes down and keeps gas in the solution. Raise the pressure and you'll keep more gas dissolved; lower the pressure, and you'll see bubbles. Henry's Law gives you an equation that tells you exactly how much gas is hanging around in your liquid at a given pressure.

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Visual Diagram

How to Use This Calculator

1. Select your Calculation Mode from the dropdown — choose from dissolved concentration, required partial pressure, Henry's constant, temperature-corrected concentration, total dissolved gas mass, or solubility ratio.
2. Enter the values for the input fields that appear — partial pressure (atm), Henry's constant (mol/L·atm), concentration (mol/L), temperatures, enthalpy, volume, molar mass, or second pressure depending on your selected mode.
3. Check your units — partial pressure in atm, temperatures in °C, enthalpy in kJ/mol, volume in litres, and molar mass in g/mol.
4. Click Calculate to see your result.

Henry's Law Gas Solubility Calculator

Equations & Variables

Run dissolved gas calculations using this formula for concentration, partial pressure, and Henry’s constant.

If you need to factor in temperature, use the Van't Hoff equation for correcting Henry’s constant.

For total dissolved gas mass, it’s a straight multiplication of concentration, volume, and molar mass.

Simple Example

Oxygen dissolving in water at 25°C, 1 atm partial pressure, k_H = 0.00129 mol/L·atm:

• Partial pressure (P) = 1.0 atm
• Henry's constant (k_H) = 0.00129 mol/L·atm
• C = 0.00129 × 1.0 = 0.00129 mol/L (1.29 mmol/L)

Theory & Engineering Applications

Henry's Law dates back to 1803 and lays out—in pretty basic terms—that gas dissolved in a liquid tracks directly with the partial pressure above it as long as the temperature stays put. It’s fundamental for phase equilibrium and turns up wherever gases and liquids meet in industry, the environment, or biology.

Fundamental Principles and Molecular Basis

At the surface, gas molecules are always entering and leaving the liquid. When the rate going in matches the rate coming out, you have equilibrium. Henry's constant gives you the proportionality—unique for every gas and solvent pair, and very sensitive to temperature.

For example, at 25°C, oxygen in water has k_H ≈ 1.29 × 10^-3 mol/L·atm, while CO₂ is much more soluble at roughly 3.41 × 10^-2 mol/L·atm—about 26 times higher. CO₂ can actually react with water, so the numbers are not just about physical dissolution. Henry's constant units vary. Sometimes you'll see atm·L/mol (the inverse), other times mol/L·atm, or even as a unitless mole fraction.

Temperature Dependence and Thermodynamic Considerations

As a rule, gas solubility drops as temperature rises. This has to do with most gas dissolutions being exothermic—dissolving the gas gives off heat. So if you heat up the system, you actually drive gas out of solution. For oxygen in water, dissolution enthalpy is about -13 kJ/mol. Practically, if water heats up, less oxygen stays dissolved—bad for fish, and why boiling water removes dissolved gases so quickly.

This temperature effect isn’t just academic. If you’re carbonating at 4°C, you can hold nearly double the CO₂ as you could at 25°C under the same pressure. On an industrial scale, gas scrubbers can lose 15–30% absorption efficiency if temperatures rise and you don’t manage the heat—something you notice when the scrubber needs extra cooling.

Non-Ideal Behavior and Limitations

Don’t expect Henry's Law to work everywhere. The equation assumes the gas and liquid behave ideally. Go much above 5–10 atm and molecular effects start mattering, so activity coefficients become necessary. Gases that react significantly (like ammonia or HCl in water) need a chemical equilibrium approach instead. CO₂ falls somewhere in-between, since it reacts with water but is often treated with Henry’s Law for a good-enough engineering estimate.

Also, adding salt—raising the water’s salinity—pushes gas out of solution. Seawater at 35 ppt salinity holds about 20% less oxygen than pure water under the same conditions. The ions tie up the water, giving gas fewer places to dissolve. If you’re building aeration for salty waste streams, you’ll need to size your pumps and diffusers for the lower solubility.

Industrial Gas Absorption Processes

In real-world chemical engineering—from absorption towers to stripping units—Henry’s Law helps size equipment, but you always compare actual concentrations to equilibrium values given by Henry’s Law to get your transfer driving force. For example, stripping 5% CO₂ from a gas stream at 10 atm puts the CO₂ partial pressure at 0.5 atm, which decides your maximum dissolved CO₂ in the liquid.

Modern CO₂ capture isn’t just about physical solubility. Chemical solvents like amines react with the gas, driving capture 50-100× higher than physical Henry's Law would allow. Regeneration cycles use heat to reverse this, so even though Henry's Law still governs the physical part, the chemistry dominates the design.

Comprehensive Worked Example: Beverage Carbonation System Design

If you need to carbonate 1000 liters of a beverage to 4.0 volumes of CO₂ (4 L of CO₂ gas per L beverage at STP), and you’re working at 4°C with k_H = 6.15 × 10^-2 mol/L·atm, here’s how you’d do it:

Step 1: Convert volume specification to molar concentration

At STP, 1 mol gas is 22.4 L. So for 4.0 L per L beverage:

Moles CO₂/L = 4.0 ÷ 22.4 = 0.1786 mol/L

Step 2: Apply Henry's Law

P = C ÷ k_H = 0.1786 ÷ 0.0615 = 2.904 atm

Step 3: Account for vapor pressure

Water vapor at 4°C is about 0.008 atm. So total system pressure = 2.904 + 0.008 = 2.912 atm (~42.4 psig).

Step 4: CO₂ needed in total

For 1000 L: 0.1786 × 1000 = 178.6 mol; 178.6 × 44.01 g/mol = 7,860 g or 7.86 kg.

Step 5: What if temperature rises?

If the liquid warms to 20°C (k_H drops to about 0.0341), then at 2.904 atm pressure: C = 0.0341 × 2.904 = 0.0990 mol/L

Excess CO₂ in solution: (0.1786 - 0.0990) × 1000 = 79.6 mol comes out as headspace gas. In small bottle headspaces, this can send internal pressure very high—hence the need for sturdy packaging. Actual pressures are lower than this max, but you’ll see 5–6 atm in real bottles if the beverage is warm.

Environmental and Biological Significance

Dissolved oxygen levels in water limit what fish and other organisms can survive. For reference, at 15°C with normal pressure (0.21 atm O₂), you get about 10.2 mg/L saturation. Trout want at least 6–7 mg/L, so there’s not much margin, especially as temperatures rise—each +10°C can cut capacity by ~20%.

In diving or physiology, Henry’s Law is behind the gas exchange story. Blood at sea level dissolves about 0.003 mL O₂/mL plasma, way less than the oxygen carried on hemoglobin. Scuba divers at 30 m face about 4 atm total pressure—this rapidly increases dissolved nitrogen in tissues, and coming up too fast leads to decompression sickness because dissolved nitrogen becomes bubbles.

Advanced Applications in Process Engineering

For semiconductor work, sub-ppb dissolved oxygen is often needed, which requires degassing at near-vacuum. Pulling the pressure down to 0.007–0.013 atm drives most O₂ out per Henry’s Law. In pharmaceuticals, hollow fiber membrane contactors use a high area-to-volume ratio to drive out dissolved gases reliably, all using these same principles.

Sometimes you’ll need more calculations across process engineering. For more on fluid dynamics and mass transfer, check out the engineering calculator hub.

Practical Applications

Frequently Asked Questions

📹 Video Walkthrough — Henrys Law Gas Solubility Interactive Calculator

📹 Video Walkthrough — Henrys Law Gas Solubility Interactive Calculator