Partial Pressure Calculator

In a mixture of gases each component pushes on the walls of the container as if it were alone. That contribution is its partial pressure, and John Dalton noticed in 1801 that the partial pressures add up to the total pressure. Because each gas behaves independently, its partial pressure equals its mole fraction times the total pressure – or, from the ideal gas law, its own amount of substance times R·T divided by the volume.

Partial pressure, not percentage, is what drives physiology and chemistry. Air is 20.95% oxygen everywhere in the lower atmosphere, but at the top of Everest the total pressure is only about a third of sea level, so the oxygen partial pressure – and the oxygen reaching your blood – drops by two thirds. Divers face the opposite problem: at 40 m depth (5 atm absolute) the oxygen partial pressure in air passes 1.0 atm, and above about 1.4–1.6 atm oxygen becomes toxic.

Henry’s law links the partial pressure of a gas above a liquid to the amount dissolved in it. A carbonated drink is bottled under several atmospheres of CO₂; open it and the CO₂ partial pressure falls, so the dissolved gas escapes as bubbles.

How to use the Partial Pressure Calculator

1. Method
2. Total pressure of the mixture
3. Mole fraction of the gas (0–1)
4. Amount of the gas
5. Temperature
6. Volume of the container
7. Gas (dissolved in water, 25 °C)
8. Henry’s law constant k_H,pc
9. Henry’s law constant k_H,px
10. Concentration of dissolved gas
11. Mole fraction of dissolved gas

Results update as you type, every field has its own unit menu, and everything is calculated in your browser.

How it works

Pick the method that fits your data. Dalton’s law needs the total pressure and the mole fraction; the ideal gas mode needs the moles of the gas, the temperature and the container volume; the two Henry’s law modes need the dissolved concentration (or mole fraction) and the Henry constant, which is filled in automatically for common gases in water at 25 °C.

Worked example

Dalton’s law: dry air at 1 atm contains 20.95% oxygen, so p(O₂) = 0.2095 × 1 atm = 0.2095 atm = 21.23 kPa. The other gases together exert 0.7905 atm.

Ideal gas law: 0.5 mol of a gas in a 10 L container at 25 °C (298.15 K) exerts p = 0.5 × 0.082057 × 298.15 / 10 = 1.223 atm, whatever else is in the container.

Henry’s law: water at 25 °C holding 2.6 × 10⁻⁴ mol/L of dissolved O₂ is in equilibrium with an oxygen partial pressure of 769.23 × 2.6 × 10⁻⁴ = 0.200 atm – about the oxygen pressure in air, so this is roughly air-saturated water (8.3 mg/L).

Henry’s law constants in water at 25 °C

The values come from the standard compilation by Sander; k_H,px ≈ k_H,pc × 55.3 mol/L (the molar concentration of water). Different sources differ by a few percent.

Tips

• Keep temperature in kelvin and pressure in consistent units.
• For gases, ideal-gas behaviour is a good approximation near room conditions.
• Half-life and rate constant are linked by t½ = ln 2 / k for first-order processes.
• Everything is calculated in your browser.

Common mistakes to avoid

• Using the mass number instead of the molar mass in Graham’s law.
• Mixing partial pressures and mole fractions.
• Forgetting that osmotic pressure uses the total particle concentration (van ‘t Hoff factor).

Inputs explained

What each result means

• Partial pressure in kPa = [Partial pressure]×101.325
• Partial pressure in mmHg = [Partial pressure]×760
• Partial pressure of all other gases = [Total pressure of the mixture] – [Partial pressure]
• Share of the total pressure = ([Mole fraction of the gas (0–1)])×100
• Molar concentration of the gas = [Amount of the gas]÷[Volume of the container]