0.21
0.21
atm
21.278
kPa
21
%
0.21
0.21
atm
21.278
kPa
21
%
The Partial Pressure Calculator applies Dalton's Law to determine the pressure contribution of an individual gas in a mixture. Enter the total pressure, moles of the component gas, and total moles of the mixture to find the mole fraction, partial pressure, and percentage composition. This tool is essential for respiratory physiology, environmental science, industrial gas handling, and atmospheric chemistry.
Dalton's Law of Partial Pressures states that in a mixture of non-reacting gases, the total pressure equals the sum of the partial pressures of each component. Each gas behaves independently as if it alone occupied the entire volume, and its partial pressure is proportional to its mole fraction.
Dalton's Law defines partial pressure as:
$$P_i = \chi_i \times P_{total}$$
where the mole fraction is:
$$\chi_i = \frac{n_i}{n_{total}}$$
The total pressure is the sum of all partial pressures:
$$P_{total} = P_1 + P_2 + P_3 + \ldots = \sum_i P_i$$
And mole fractions always sum to unity:
$$\sum_i \chi_i = 1$$
This law follows from the ideal gas assumption: in an ideal gas mixture, each component exerts pressure independently of the others because there are no intermolecular interactions. Each gas contributes to the total pressure in proportion to the number of its molecules present. The law is most accurate at low to moderate total pressures where ideal behavior holds.
The mole fraction represents the proportion of the component gas in the mixture (0 to 1). The partial pressure is the pressure that component would exert if it alone occupied the entire volume. The percentage shows the composition in an easily understood format. For atmospheric calculations, the partial pressure of oxygen in dry air at sea level is about 0.21 atm (21% of 1 atm).
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Oxygen makes up 21% of air. At 1 atm total pressure, the partial pressure of O₂ is 0.21 atm (21.3 kPa, or 160 mmHg). This value is critical in respiratory and aviation medicine.
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At 5,500 m altitude, the oxygen partial pressure drops to 0.105 atm (10.6 kPa, 80 mmHg) — half of sea level. This reduced O₂ pressure causes altitude sickness and explains why supplemental oxygen is needed at high elevations.
Dalton's Law states that the total pressure of a gas mixture equals the sum of the partial pressures of each individual gas: P_total = P₁ + P₂ + P₃ + .... Each gas contributes pressure independently, proportional to its mole fraction. Published by John Dalton in 1801.
Partial pressure is the pressure that one component of a gas mixture would exert if it alone occupied the entire volume at the same temperature. It represents that gas's contribution to the total pressure. Mathematically, P_i = χ_i × P_total.
Mole fraction (χ) is the ratio of moles of one component to the total moles in the mixture: χ_i = n_i/n_total. It is dimensionless and ranges from 0 to 1. All mole fractions in a mixture sum to exactly 1. It can also be expressed as a percentage (volume %).
The partial pressure of oxygen (PO₂) in inspired air determines how much oxygen enters the blood. At sea level, PO₂ ≈ 160 mmHg in dry air, but decreases to ~100 mmHg in the alveoli (after humidification and mixing with CO₂). Blood gas analysis measures PO₂ and PCO₂ to assess respiratory function.
The mole fraction of O₂ remains constant at 21% at all altitudes, but the total atmospheric pressure decreases. Since PO₂ = 0.21 × P_total, the oxygen partial pressure drops proportionally. At 5,500 m, PO₂ is half of sea level, causing hypoxia and altitude sickness.
When collecting a gas over water, the total pressure equals the sum of the collected gas pressure and water vapor pressure: P_total = P_gas + P_water. To find the dry gas pressure: P_gas = P_total - P_water. The water vapor pressure depends on temperature and must be looked up.
Dalton's Law is exact for ideal gases and a good approximation for real gas mixtures at moderate pressures. Deviations occur at high pressures or when components have strong intermolecular interactions (e.g., ammonia and water vapor). Lewis-Randall and fugacity corrections handle real gas mixtures.
Henry's Law states that the dissolved gas concentration in a liquid is proportional to the partial pressure of that gas above the liquid: C = k_H × P_gas. This explains why carbonated drinks fizz when opened (CO₂ partial pressure drops) and why deep-sea divers get the bends (N₂ comes out of solution as pressure decreases).
Anesthetic gases are administered at specific partial pressures to achieve desired blood concentrations. The partial pressure gradient between the inspired mixture and the blood drives gas exchange. MAC (Minimum Alveolar Concentration) is expressed as a percentage (essentially a mole fraction) at 1 atm.
Yes, if you know the total pressure and partial pressures of components, you can determine the composition: χ_i = P_i/P_total. Gas chromatography and mass spectrometry measure partial pressures to determine the composition of unknown gas mixtures.
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