Calculate partial pressures of gases in mixtures using Dalton's law, mole fractions, vapor pressure corrections, and real gas adjustments.
Dalton's law of partial pressures states that the total pressure of a gas mixture equals the sum of the partial pressures of each individual gas component. Each gas's partial pressure equals its mole fraction multiplied by the total pressure: Pᵢ = xᵢ × P_total. This fundamental principle is essential in chemistry, engineering, biology, and atmospheric science.
Understanding partial pressures is critical for: respiratory physiology (oxygen delivery to tissues depends on PO₂, not total pressure), diving safety (partial pressures of O₂ and N₂ determine toxicity risks at depth), industrial gas processing (distillation, absorption, membrane separations), and atmospheric chemistry (pollutant concentrations, greenhouse gas levels).
This calculator handles gas mixtures with up to 8 components, automatically calculates mole fractions and partial pressures from either moles or volume percentages, and provides context for common applications like breathing gases, atmospheric composition, and industrial processes. It includes altitude-adjusted atmospheric calculations and dissolved gas concentrations via Henry's law.
This calculator quickly determines partial pressures in complex gas mixtures, converts between pressure units, and applies physiological/safety context — invaluable for lab work, diving planning, medical gas calculations, and atmospheric chemistry. This partial pressure calculator helps you compare outcomes quickly and reduce avoidable mistakes when making day-to-day care decisions. Use the estimate as a planning baseline and confirm final decisions with a qualified professional when risk is high.
Dalton's Law: P_total = Σ Pᵢ = Σ (xᵢ × P_total), where xᵢ = nᵢ/n_total is the mole fraction. For ideal gases, the volume fraction equals the mole fraction. Henry's Law: C = k_H × P, where C = dissolved concentration, k_H = Henry's constant.
Result: PO₂ = 0.2095 atm = 159.2 mmHg
In dry air at 1 atm: PO₂ = 0.2095 × 1.0 = 0.2095 atm = 159.2 mmHg. This is the inspired oxygen partial pressure at sea level.
Earth's atmosphere is a natural example of Dalton's law. The total atmospheric pressure at sea level (~101.325 kPa) is the sum of partial pressures from nitrogen (79.1 kPa), oxygen (21.2 kPa), argon (0.94 kPa), carbon dioxide (0.04 kPa), and trace gases. Weather reporting of barometric pressure reflects this sum. Changes in water vapor content alter the total pressure and shift the partial pressures of other components.
In medicine and diving, partial pressures determine gas exchange rates. Oxygen diffuses from alveoli into blood because alveolar PO₂ (~100 mmHg) exceeds venous PO₂ (~40 mmHg). At altitude, reduced PO₂ triggers hyperventilation and eventual acclimatization through increased red blood cell production. Hyperbaric oxygen therapy exploits elevated PO₂ to enhance wound healing and treat decompression sickness.
Chemical engineers use partial pressures extensively in vapor-liquid equilibrium calculations (Raoult's law, modified Raoult's law), absorption column design, and membrane separation modeling. The driving force for gas absorption into a liquid is the difference between the gas-phase partial pressure and the equilibrium partial pressure above the liquid — directly connecting Dalton's law with Henry's law and mass transfer operations.
The pressure that each gas in a mixture would exert if it alone occupied the entire volume at the same temperature. It's proportional to the gas's mole fraction.
At very high pressures or with gases that interact strongly (like NH₃ + HCl), Dalton's law gives inaccurate results because it assumes no intermolecular interactions between different gas species. This keeps planning practical and lowers the chance of preventable errors.
At depth, total pressure increases (~1 atm per 10 m). PO₂ above 1.6 atm causes oxygen toxicity; PN₂ above ~3.2 atm causes nitrogen narcosis. Divers adjust breathing gas composition accordingly.
In alveoli at sea level: ~100 mmHg (not 159 mmHg) because inspired air is humidified (PH₂O = 47 mmHg) and mixed with CO₂ (PCO₂ = 40 mmHg). This keeps planning practical and lowers the chance of preventable errors.
Total atmospheric pressure decreases exponentially with altitude. At 5500 m (18,000 ft), pressure is ~half sea level, so PO₂ is only ~80 mmHg, making supplemental O₂ necessary.
Henry's law states that the dissolved concentration of a gas is proportional to its partial pressure above the solution: C = k_H × P. Higher partial pressure means more dissolved gas.