Partial Pressure Calculator
Choose a calculation mode: Dalton's law (mole fraction times total pressure), the ideal gas law (moles, temperature, volume), or Henry's law (dissolved gas concentration). Switch between atmospheres, kPa, bar, mmHg, and psi at any time. The show-your-work panel traces every step.
Formula
Worked example
Oxygen in dry air: mole fraction 0.21, total pressure 101.325 kPa. P(O₂) = 0.21 × 101.325 = 21.28 kPa (= 0.21 atm = 159.6 mmHg). The other gases account for the remaining 80.05 kPa.
Four ways to find partial pressure
This calculator covers the three most common approaches. Dalton's law (Pᵢ = Xᵢ × P_total) is the fastest when you know the mole fraction and total pressure: every component's partial pressure is simply its share of the moles times the total pressure, and they all add back to that total. The ideal gas law (Pᵢ = nᵢRT / V) is useful when you know the actual moles, temperature, and container volume but not the total pressure. Henry's law (p = KH × c or p = KH × x) connects the partial pressure of a gas above a liquid to how much of it is dissolved, making it essential for solubility, carbonation, and blood-gas problems. All three modes share the same pressure-unit selector so you can read the answer in atm, kPa, bar, mmHg, or psi without manual conversion.
Unit conversions and multi-component breakdowns
Pressure can be reported in many units: 1 atm = 101.325 kPa = 1.01325 bar = 760 mmHg = 14.696 psi. This calculator converts internally through atmospheres and reports the result in whichever unit you select. The cross-unit summary in the output always shows kPa and mmHg alongside your chosen unit, which is handy for comparing with clinical or meteorological references. In Dalton mode you can also switch on the multi-component breakdown, enter the mole fractions of up to three additional gases, and see all four partial pressures side by side. Their mole fractions should sum to 1 (any residual is lumped into "all other components") and their partial pressures should sum to the total.
Henry's law and dissolved gases
Henry's law applies when a gas is in equilibrium with a liquid. In the concentration form, p = KH × c, the Henry constant KH has units of L·atm/mol; a larger KH means the gas is less soluble (more partial pressure is needed to keep a given concentration dissolved). In the mole-fraction form, p = KH × x, KH is in pure atmospheres. The constants pre-loaded here are for 25 °C; they fall as temperature rises because gases become less soluble in warmer water. CO₂ (KH about 29 L·atm/mol) dissolves roughly 26 times more readily than O₂ (KH about 769 L·atm/mol) at the same partial pressure, which is why CO₂ can be carbonated into drinks at modest pressures while oxygen stays mostly in the gas phase.
Where partial pressures matter
In respiratory physiology the alveolar partial pressure of oxygen (PaO₂) drives uptake into blood; it falls at high altitude because total pressure drops while the 21% mole fraction of oxygen in air stays constant. Divers breathe gas mixtures where the partial pressure of oxygen must stay between roughly 16 kPa and 160 kPa: too low causes hypoxia, too high causes oxygen toxicity. In chemical engineering, equilibrium constants for gas-phase reactions are written in terms of partial pressures (the Kp formulation), and reactor design requires tracking each component's partial pressure separately. Dalton's law assumes ideal behaviour, which holds well at moderate pressures and temperatures but breaks down for real gases at very high pressures where intermolecular forces become significant.
Henry's law constants for common gases at 25 °C
| Gas | Formula | KH1 (L·atm/mol) | Solubility note |
|---|---|---|---|
| Oxygen | O₂ | 769.2 | Sparingly soluble |
| Nitrogen | N₂ | 1639.3 | Very sparingly soluble |
| Hydrogen | H₂ | 1282.1 | Sparingly soluble |
| Helium | He | 2702.7 | Very sparingly soluble |
| Argon | Ar | 714.3 | Slightly more soluble than O₂ |
| Neon | Ne | 2222.2 | Very sparingly soluble |
| Carbon monoxide | CO | 1052.6 | Sparingly soluble |
| Carbon dioxide | CO₂ | 29.41 | Highly soluble (reacts with water) |
KH1 in L·atm/mol (concentration form). Larger KH1 = lower solubility. Source: Sander (2015), NIST WebBook.
Frequently asked questions
What is Dalton's law of partial pressures?
Dalton's law states that in a mixture of non-reacting ideal gases the total pressure equals the sum of the partial pressures of the individual gases. Each component exerts pressure as if it alone occupied the container, and those independent contributions simply add together. From this it follows that a component's partial pressure equals its mole fraction multiplied by the total pressure: Pᵢ = Xᵢ × P_total, which is the relationship used in the Dalton mode of this calculator.
How do I find a mole fraction?
A mole fraction is the moles of one component divided by the total moles of all components in the mixture. For example, if a flask holds 2 mol of oxygen and 8 mol of nitrogen the total is 10 mol, so the mole fraction of oxygen is 2 / 10 = 0.20 and nitrogen is 8 / 10 = 0.80. All the mole fractions in a mixture must add to exactly 1.
How do I convert between pressure units?
1 atm = 101.325 kPa = 1.01325 bar = 760 mmHg = 14.696 psi. Use the pressure-unit selector at the top of the calculator to switch all inputs and the result to your preferred unit. The output panel always also shows the answer in kPa and mmHg for easy cross-reference with clinical or meteorological data.
When should I use the ideal gas law mode instead of Dalton's law?
Use the ideal gas law mode (Pᵢ = nᵢRT / V) when you know the moles, temperature, and volume of the component but not the total pressure or the mole fraction. This is common in laboratory problems where you add a known mass of gas to a container of known volume at a known temperature. The partial pressure it produces is that component's contribution; add all components' partial pressures to get the total.
What is Henry's law and when does it apply?
Henry's law relates the partial pressure of a gas above a liquid to the amount dissolved in that liquid, at equilibrium and at dilute concentration. In the concentration form, p = KH × c, where KH is the Henry constant in L·atm/mol. Gases with large KH (like N₂) dissolve little; gases with small KH (like CO₂) dissolve a lot. Henry's law is the basis for blood-gas analysis, carbonated beverage production, and environmental modelling of atmospheric gas absorption by oceans.
Does the pressure unit affect the calculation?
In Dalton's law mode the mole fraction is dimensionless, so the partial pressure comes out in whatever unit you entered for the total pressure. In ideal gas law mode the formula uses the gas constant R = 0.08206 L·atm/(mol·K), so the internal calculation is always in atm and then converted to your chosen unit. In Henry's law mode the pre-loaded constants are in atm-based units and are converted after the calculation. The displayed result always matches the unit you selected.