Partial Pressure Calculator

About the Partial Pressure Calculator

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.

When This Page Helps

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.

How to Use the Inputs

1. Enter the total pressure of the gas mixture (in atm, kPa, mmHg, or psi).
2. Add gas components with either their mole amounts or volume/mole percentages.
3. Use presets for common mixtures: dry air, exhaled breath, natural gas, etc.
4. Review partial pressures for each component in your chosen pressure unit.
5. Optionally enter altitude or depth to see adjusted pressures.
6. Check the dissolved concentration using Henry's law constants.
7. Compare results against physiological or safety thresholds.

Example Calculation

Tips & Best Practices

• For humid gas mixtures, subtract water vapor pressure from total pressure before calculating dry gas partial pressures.
• At body temperature (37°C), water vapor pressure is 47 mmHg — always account for this in respiratory calculations.
• Mole fraction equals volume fraction ONLY for ideal gases at the same T and P.
• For breathing gas calculations, multiply PO₂ by the respiratory quotient to estimate alveolar partial pressures.
• When collecting gases over water, subtract the vapor pressure of water to get the dry gas pressure.
• Henry's law constants are highly temperature-dependent — always use values at your actual temperature.

Dalton's Law in Atmospheric Chemistry

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.

Breathing and Physiology

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.

Industrial Gas Processing

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.

Frequently Asked Questions

• What is partial pressure?

  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.

• When does Dalton's law fail?

  At very high pressures or with gases that interact strongly (like NH₃ + HCl), Dalton's law becomes less reliable because it assumes no intermolecular interactions between different gas species.

• Why are partial pressures important in diving?

  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.

• What is the partial pressure of O₂ in the lungs?

  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.

• How does altitude affect partial pressures?

  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.

• What is Henry's law and how does it relate?

  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.