Partial pressure is the individual pressure that a single gas contributes within a mixture of gases. In any gas mixture, each gas pushes outward as if it were the only gas present, and that individual push is its partial pressure. The total pressure of the mixture equals all those individual pressures added together. This concept, known as Dalton’s Law, is written simply as: total pressure = the partial pressure of gas 1 + gas 2 + gas 3, and so on.
How Partial Pressure Works in Air
The atmosphere is a gas mixture, so it’s the most intuitive place to see partial pressure in action. At sea level, total atmospheric pressure is 760 mmHg. Nitrogen makes up about 78% of the atmosphere, so its partial pressure is roughly 593 mmHg. Oxygen sits at about 21%, giving it a partial pressure of 159 mmHg. Carbon dioxide, at just 0.03%, contributes only 0.2 mmHg.
The key calculation is straightforward: multiply the total pressure by the fraction of the mixture that a particular gas represents. If oxygen is 21% of the atmosphere and total pressure is 760 mmHg, oxygen’s partial pressure is 0.21 × 760, or about 159 mmHg. This relationship between a gas’s proportion (its mole fraction) and the total pressure is how partial pressures are determined in virtually every context, from weather science to hospital ventilators.
Why It Matters More Than Percentage
A common point of confusion is that the percentage of oxygen in air doesn’t change at high altitude. It’s still 21% on the summit of Everest. What changes is the total atmospheric pressure, and that drags down oxygen’s partial pressure with it. At 5,500 meters, atmospheric pressure drops to about half its sea-level value, cutting the partial pressure of inspired oxygen in half as well. At 8,900 meters (the summit of Everest), it falls to roughly 30% of sea-level values.
This is why partial pressure, not percentage, determines how gases actually behave. A lower partial pressure means less driving force to push oxygen from the air into your lungs and from your lungs into your blood. The result is a cascade of oxygen shortage that affects everything from breathing rate to blood vessel behavior to the energy-producing structures inside your cells.
Gas Exchange in Your Lungs
Your body runs on partial pressure gradients. Gases move from areas of higher partial pressure to areas of lower partial pressure, much like water flowing downhill. In the lungs, the partial pressure of oxygen in the tiny air sacs (alveoli) is about 104 mmHg. Blood arriving from the rest of the body, having already delivered most of its oxygen to tissues, carries oxygen at only about 40 mmHg. That 64 mmHg difference is what pushes oxygen out of the air and into your bloodstream.
Carbon dioxide works the same way, just in reverse. Blood returning to the lungs carries carbon dioxide at about 45 mmHg, while the alveoli sit at around 40 mmHg. That smaller gradient is enough to move carbon dioxide out of the blood and into the lungs so you can exhale it. Every breath you take exploits these pressure differences.
In healthy adults, the partial pressure of oxygen in arterial blood normally falls between 80 and 100 mmHg, and carbon dioxide between 35 and 45 mmHg. These values are measured with a blood gas test and help clinicians assess how well the lungs are working.
Oxygen and Hemoglobin
Once oxygen enters the blood, it binds to hemoglobin, the protein in red blood cells that carries it to tissues. The relationship between oxygen’s partial pressure and how much hemoglobin picks up isn’t a straight line. It follows an S-shaped curve: at higher partial pressures hemoglobin loads up quickly, but at lower levels it releases oxygen more readily. A useful reference point is the P50, the partial pressure at which hemoglobin is exactly 50% saturated with oxygen. In a healthy person, this value is about 26 mmHg. A lower P50 means hemoglobin grips oxygen more tightly; a higher one means it lets go more easily.
Gases Dissolving in Liquids
Partial pressure also governs how much gas dissolves into a liquid. Henry’s Law describes this: the amount of gas that dissolves in a liquid is directly proportional to the partial pressure of that gas above the liquid’s surface. Increase the pressure and more gas dissolves. Decrease it and gas comes back out of solution.
This principle is easy to observe with a carbonated drink. The bottle is sealed under high carbon dioxide pressure, keeping the gas dissolved. Open the cap, the pressure drops, and bubbles form as the gas escapes. The same physics plays out in far more consequential settings.
Scuba Diving and Nitrogen
Underwater, divers breathe compressed air at pressures that increase with depth. Because the total pressure rises, the partial pressure of every gas in the tank rises too. Nitrogen, which is inert at the surface, starts to cause problems. At depths of 30 meters, some divers begin to experience nitrogen narcosis, a state of impaired judgment and coordination sometimes compared to alcohol intoxication. By 60 to 70 meters, all divers breathing regular air are significantly affected. The generally accepted limit for compressed air diving is 30 to 50 meters; beyond that, divers switch to gas mixtures with reduced nitrogen content.
Henry’s Law explains another diving hazard. At depth, elevated nitrogen partial pressure forces more nitrogen to dissolve into body tissues and blood. If a diver ascends too quickly, the surrounding pressure drops faster than the body can safely release that dissolved nitrogen. The gas forms bubbles in the blood and tissues, causing decompression sickness. Both problems trace back to the same principle: higher total pressure at depth means higher partial pressures of the component gases.
Altitude and the Body’s Response
At high altitude, the low partial pressure of oxygen triggers a chain of adaptations. Breathing rate typically doesn’t increase noticeably until the inspired oxygen pressure drops to about 13.3 kPa, roughly the level found at 3,000 meters. Beyond that threshold, the body’s ventilation rises exponentially, driven by oxygen sensors in the carotid arteries near the neck.
The cardiovascular system responds differently depending on the location. In the body’s general circulation, low oxygen causes blood vessels to widen, improving flow to oxygen-starved tissues. In the lungs, however, low oxygen does the opposite: it constricts blood vessels. This pulmonary constriction can raise pressure in the lung’s blood vessels and is associated with high-altitude pulmonary edema, a dangerous buildup of fluid in the lungs. The reduced driving pressure also means blood passing through the lungs may not fully pick up oxygen before moving on, widening the gap between the oxygen levels in the air sacs and the arterial blood.
Ocean Acidification
Partial pressure doesn’t just move gases into blood. It also moves them into oceans. The ocean surface constantly equilibrates with atmospheric carbon dioxide, a process that takes roughly one year. As fossil fuel burning has raised atmospheric CO2 levels, the partial pressure of CO2 above the ocean has risen in lockstep, forcing more of the gas to dissolve into seawater.
The ocean has absorbed nearly a third of the carbon dioxide humans have added to the atmosphere, which has slowed climate change but created a separate problem. Dissolved CO2 reacts with seawater to form carbonic acid, releasing hydrogen ions that lower the water’s pH. Since preindustrial times, average ocean surface pH has dropped by about 0.1 units, from roughly 8.21 to 8.10. Monitoring stations in Hawaii, Bermuda, and the eastern Atlantic have tracked a decline of about 0.02 pH units per decade since the 1980s. If atmospheric CO2 reaches 800 parts per million by the end of the century, models project an additional drop of 0.3 to 0.4 pH units, further reducing the carbonate ions that corals and shellfish need to build their structures.
In each of these cases, from your lungs to the deep ocean, the underlying mechanism is the same. Partial pressure is the driving force that determines where gases go, how fast they get there, and how much of them end up dissolved in liquids. It connects a simple physics concept to processes as different as breathing, diving, mountain climbing, and the chemistry of the sea.