Partial pressure is the pressure a single gas would generate if it alone filled the whole volume of a mixture. It lets us calculate how much of each gas is actually present and how they affect breathing and other processes.
How gas pressure works
Inside a container, gas molecules move fast and collide with the walls. Every collision pushes a tiny bit of force onto the wall. The sum of all those tiny pushes is what we feel as pressure. The amount of pressure depends on how many molecules there are, how fast they move (temperature), and how much space they have (volume).
Air is a mixture, not a single gas
Dry air at sea level is about 78 % nitrogen, 21 % oxygen, 0.9 % argon, and 0.1 % trace gases. Each type of molecule behaves on its own; they do not chemically interact in the air we breathe.
Defining partial pressure
Because every gas in a mixture behaves independently, each one exerts its own pressure on the container walls. If we could isolate one gas and let it occupy the entire volume, the pressure it would produce is its partial pressure.
Dalton’s Law of partial pressures
Dalton discovered that the total pressure of a mixture of non‑reacting gases equals the sum of their partial pressures:
| Gas | Fraction of air | Partial pressure (mmHg) |
|---|---|---|
| Oxygen | 0.21 | ≈160 |
| Nitrogen | 0.78 | ≈593 |
| Argon + traces | 0.01 | ≈7 |
| Total | 1.00 | 760 |
Why partial pressure matters in real life
Our lungs rely on the partial pressure of oxygen to move it from the air into blood. At high altitudes the total atmospheric pressure falls, but the oxygen fraction stays about the same. The reduced total pressure lowers oxygen’s partial pressure, making it harder for the body to take in enough oxygen.
Scuba divers also use partial pressures to design breathing mixes. Too much nitrogen at depth can cause narcosis, while too little oxygen can lead to hypoxia. By calculating partial pressures, divers keep each component gas within safe limits.