Partial Pressure Converter
Calculate partial pressures of gas components using Dalton Law. Enter total pressure and mole fractions to get partial pressures in atm, kPa, mmHg, psi,
Partial Pressure Converter
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Formula: P_i = x_i x P_total (Dalton Law)
Worked example โ N2: 0.7808 atm | O2: 0.2095 atm | Ar: 0.0093 atm
Formula
P_i = x_i x P_total (Dalton Law)
The partial pressure of each gas component equals its mole fraction (x_i) multiplied by the total pressure of the mixture. Mole fractions must sum to 1.0 for a complete mixture description. Convert between pressure units using standard conversion factors.
Worked Examples
Example 1: Air Composition at Sea Level
Problem:Calculate partial pressures of N2 and O2 in air at 1 atm total pressure.
Solution:P(N2) = 0.7808 x 1 atm = 0.7808 atm = 593.4 mmHg P(O2) = 0.2095 x 1 atm = 0.2095 atm = 159.2 mmHg P(Ar) = 0.0093 x 1 atm = 0.0093 atm = 7.07 mmHg
Result:N2: 0.7808 atm | O2: 0.2095 atm | Ar: 0.0093 atm
Example 2: Diving Gas Mix at Depth
Problem:A diver breathes Nitrox (32% O2, 68% N2) at 2.5 atm (15m depth). Find partial pressures.
Solution:P(O2) = 0.32 x 2.5 = 0.80 atm P(N2) = 0.68 x 2.5 = 1.70 atm Note: O2 toxicity threshold is ~1.6 atm
Result:O2: 0.80 atm | N2: 1.70 atm
Frequently Asked Questions
What is Dalton Law of Partial Pressures?
Dalton Law states that the total pressure of a gas mixture equals the sum of the partial pressures of each individual gas component. The partial pressure of each gas is the pressure it would exert if it alone occupied the entire volume at the same temperature. Mathematically, P_total = P_1 + P_2 + P_3 + ... for all gas components. This law holds rigorously for ideal gases and is a good approximation for real gases at moderate pressures and temperatures.
How do you calculate the partial pressure of a gas in a mixture?
The partial pressure of a gas equals its mole fraction multiplied by the total pressure of the mixture: P_i = x_i times P_total. The mole fraction is the number of moles of that gas divided by the total moles of all gases. For example, in dry air at 1 atm, nitrogen has a mole fraction of 0.7808, so its partial pressure is 0.7808 times 1 = 0.7808 atm, or about 593.4 mmHg. This relationship follows directly from the ideal gas law applied to each component.
Why are partial pressures important in chemistry and biology?
Partial pressures are critical in many areas of science. In respiratory physiology, the partial pressure of oxygen (pO2) determines how effectively oxygen binds to hemoglobin in the blood. In chemical engineering, partial pressures are used to calculate chemical equilibrium constants (Kp) for gas-phase reactions. In diving medicine, nitrogen partial pressure determines decompression requirements. In environmental science, the partial pressure of CO2 in the atmosphere drives climate change models and ocean acidification calculations.
Do mole fractions always need to sum to exactly 1?
Yes, for a complete description of a gas mixture, the mole fractions of all components must sum to exactly 1.0 (or 100%). If the sum is less than 1, it means there are additional gas components not accounted for. If the sum exceeds 1, there is an error in the input values. In practice, minor components like noble gases or trace pollutants are sometimes omitted for simplicity, but the partial pressures of the listed gases will still be correctly calculated as long as their individual mole fractions are accurate.
How do I convert a percentage gas composition into a partial pressure?
Divide the percentage by 100 to obtain the mole fraction, then multiply by the total pressure. Dry air is about 20.95 percent oxygen, so at a total pressure of 101.325 kPa the oxygen partial pressure is 0.2095 x 101.325 = 21.23 kPa. This works only when the percentage is stated by volume or by mole, since Dalton's law is written in mole fractions. A percentage by mass must first be converted using the molar masses of each component.
What is the difference between partial pressure and vapour pressure?
Partial pressure is the share of total pressure contributed by one component of a gas mixture, set by its mole fraction. Vapour pressure is the equilibrium pressure a liquid's own vapour exerts above it at a given temperature, and it is a property of the substance rather than of the mixture. A gas can hold a water partial pressure anywhere from zero up to the saturation vapour pressure, and the ratio of the two is the relative humidity.
How does altitude change the partial pressure of oxygen?
The mole fraction of oxygen stays near 0.209 throughout the lower atmosphere, but total barometric pressure falls with altitude, so the oxygen partial pressure falls in direct proportion. At half the sea-level barometric pressure the oxygen partial pressure is also halved even though the air composition is unchanged. This is why physiological effects at altitude track partial pressure rather than percentage, and why supplemental oxygen works by raising the mole fraction.
Why do divers work in partial pressures rather than gas percentages?
Underwater the ambient absolute pressure rises by roughly one atmosphere for every ten metres of seawater, so a fixed gas fraction produces a rising partial pressure with depth. A 32 percent nitrox mix at four atmospheres absolute delivers an oxygen partial pressure of 0.32 x 4 = 1.28 atm. Both oxygen and inert gas exposure limits are framed in partial pressure terms, so a mix percentage means nothing without the depth. Follow the limits taught by your training agency.
How does Henry's law link partial pressure to dissolved gas concentration?
Henry's law states that at equilibrium the concentration of a dissolved gas is proportional to its partial pressure above the liquid, through a constant that depends on the gas, the solvent and the temperature. Doubling the partial pressure roughly doubles the dissolved amount. This underpins carbonation, blood gas transport and aeration calculations in water treatment, and it explains why gas comes out of solution when the pressure above it is released.
Should I use mole fraction or mass fraction in Dalton's law?
Dalton's law requires mole fractions. If a composition is given by mass, divide each component's mass by its molar mass to obtain moles, then divide by the total moles. The two fractions diverge whenever component molar masses differ: in a mixture of oxygen at 32 g/mol and helium at 4 g/mol, an equal-mass blend is far from an equal-mole blend, so using the mass fraction directly would badly misstate each partial pressure.
References
Reviewed for accuracy by Manoj Kumar, Mathematics Educator ยท Editorial policy
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