0.4
0.4
atm
40.53
kPa
304
mmHg
4.8932
atm
40.00%
0.4
0.4
atm
40.53
kPa
304
mmHg
4.8932
atm
40.00%
The Partial Pressure Calculator applies Dalton's Law of Partial Pressures to determine the pressure contribution of an individual gas within a mixture. John Dalton established in 1801 that the total pressure of a gas mixture equals the sum of the partial pressures of each component:
$$P_{total} = P_1 + P_2 + P_3 + \cdots = \sum_i P_i$$
Each gas in a mixture behaves independently and exerts pressure as if it alone occupied the entire container. The partial pressure of gas i is calculated using the mole fraction:
$$P_i = \chi_i \cdot P_{total}$$
where the mole fraction $$\chi_i = \frac{n_i}{n_{total}}$$ represents the fraction of total moles contributed by gas i. This principle is fundamental in chemistry, respiratory physiology, scuba diving, and atmospheric science. For example, at sea level (1 atm), oxygen's partial pressure is $$P_{O_2} = 0.2095 \times 1 = 0.2095$$ atm because oxygen constitutes 20.95% of air by mole fraction.
Enter the moles of your target gas, the total moles of all gases in the mixture, and the total pressure. The calculator computes:
Mole fraction: $$\chi = \frac{n_{gas}}{n_{total}}$$
Partial pressure: $$P_{partial} = \chi \times P_{total}$$
The result is displayed in three common pressure units: atm, kPa, and mmHg (torr). The calculator also provides an independent verification using the ideal gas law directly: $$P = \frac{n_{gas}RT}{V}$$, which should match the partial pressure from Dalton's law when total pressure is consistent with the container conditions.
Dalton's law assumes ideal gas behavior — no intermolecular interactions between different gas species. This holds well at moderate pressures and temperatures. At very high pressures or with polar gas mixtures, deviations occur and more sophisticated mixing rules are needed.
The mole fraction ranges from 0 to 1 and represents the proportion of the gas in the mixture. A mole fraction of 0.40 means 40% of the gas molecules are the target species. The partial pressure directly determines reaction rates in gas-phase chemistry (via equilibrium constants expressed in partial pressures), gas solubility in liquids (Henry's law), and the effective breathing gas in physiology (hypoxia occurs when $$P_{O_2}$$ drops below ~0.16 atm).
Inputs
Results
Oxygen makes up 20.95% of air, giving a partial pressure of 0.2095 atm (159 mmHg) at sea level.
Inputs
Results
3 mol H₂ out of 8 mol total: mole fraction 0.375, partial pressure 0.75 atm out of 2 atm total.
Dalton's Law states that in a mixture of non-reacting gases, the total pressure equals the sum of the partial pressures of each individual gas: $$P_{total} = \sum P_i$$. Each gas contributes pressure proportional to its mole fraction.
Mole fraction ($$\chi$$) is the ratio of moles of one component to the total moles in the mixture: $$\chi_i = n_i / n_{total}$$. It is dimensionless and always between 0 and 1. All mole fractions in a mixture sum to exactly 1.
Henry's Law states that gas solubility in a liquid is proportional to its partial pressure above the liquid: $$C = k_H \cdot P_{partial}$$. This is why carbonated drinks fizz when opened — reducing CO₂ partial pressure lowers its solubility.
Blood gas analysis measures $$P_{O_2}$$ and $$P_{CO_2}$$ to assess respiratory function. Normal arterial $$P_{O_2}$$ is 75–100 mmHg. Hypoxia occurs when tissue $$P_{O_2}$$ drops below critical thresholds. Anesthesia dosing is based on partial pressures of anesthetic gases.
At high pressures, real gas interactions cause deviations. The Lewis-Randall rule and fugacity coefficients are used for accurate calculations in high-pressure systems such as industrial chemical reactors.
Simply add all partial pressures: $$P_{total} = P_1 + P_2 + P_3 + \cdots$$. If you know the moles and conditions, use $$P_{total} = \frac{n_{total}RT}{V}$$ and then find each partial pressure via mole fractions.
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