Partial Pressure Calculator
Find a gas's partial pressure with Dalton's law. Enter its mole fraction (0 to 1) and the total pressure in atm, kPa, mmHg, psi or bar to get the partial pressure and its value in atmospheres.
Example: with Mole fraction of the gas (0–1) 0.21 · Total pressure 1 · Pressure unit atm → Partial pressure: 0.21 atm.
- In atmospheres0.21 atm
- Share of total21% of the 1 atm total
Computed by the calculator below using its default values. Change any input to see your own numbers.
Dalton's law: the partial pressure of a gas equals its mole fraction times the total pressure, Pᵢ = xᵢ × P_total. The partial pressures of all components add up to the total.
How Dalton's law works
In a mixture of gases that do not react, each gas behaves as if it alone occupied the container. The pressure it contributes — its partial pressure — is simply its share of the molecules times the total pressure. That share is the mole fraction: moles of the gas divided by the total moles. So Pᵢ = xᵢ × P_total.
Dry air near sea level is a clean example. Oxygen is about 21% of the molecules, so at 1 atm its partial pressure is 0.21 atm, or roughly 159 mmHg. Nitrogen, at 78%, contributes 0.78 atm. Add every component's partial pressure and you recover the total, which is Dalton's law stated in reverse.
How it’s calculated
Partial pressure Pᵢ = xᵢ × P_total, with the mole fraction xᵢ between 0 and 1. The result is reported in the entered unit and converted to atmospheres using 1 atm = 101.325 kPa = 760 mmHg = 14.6959 psi = 1.01325 bar. A value above 1 is read as a percent and divided by 100.
The gases are treated as ideal and non-reacting, so partial pressures simply add. Mole fraction, not volume or mass fraction, drives the result.
Partial pressures in dry air at 1 atm
| Gas | Mole fraction | Partial pressure |
|---|---|---|
| Nitrogen | 0.781 | 0.781 atm (594 mmHg) |
| Oxygen | 0.209 | 0.209 atm (159 mmHg) |
| Argon | 0.0093 | 0.0093 atm (7.1 mmHg) |
| Carbon dioxide | 0.0004 | 0.0004 atm (0.3 mmHg) |
Computed as Pᵢ = xᵢ × 1 atm from standard dry-air composition; rounded.
Common mistakes
- Entering a percent (21) instead of a fraction (0.21); this tool reads values above 1 as a percent, but the formula needs the fraction.
- Using a volume or mass fraction where the mole fraction is required — for ideal gases volume fraction equals mole fraction, but mass fraction does not.
- Forgetting that the partial pressures must add up to the total pressure.
- Mixing pressure units between the gas and the total.
Frequently asked questions
What is the partial pressure formula?
Dalton's law gives Pᵢ = xᵢ × P_total: the gas's mole fraction times the total pressure. The partial pressures of all gases in the mix sum to the total.
How do I find the mole fraction?
Divide the moles of that gas by the total moles of all gases. For ideal gases this also equals its volume fraction, so 21% oxygen by volume means a mole fraction of 0.21.
Do partial pressures add up to the total?
Yes. That is the heart of Dalton's law: the total pressure is the sum of every component's partial pressure, because each gas pushes independently.
What is the partial pressure of oxygen in air?
With a mole fraction near 0.21, oxygen's partial pressure at 1 atm is about 0.21 atm, or roughly 159 mmHg. It falls with altitude as total pressure drops.