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When you look at the sky, the air seems like a single substance. But the Earth's atmosphere is actually a mixture, composed of 78.08% nitrogen and 20.95% oxygen by volume (with the rest made of argon, carbon dioxide and trace gases). Because these gases have different molecular weights, you might wonder why the heavy gases do not settle at the bottom and the light ones float to the top.
The answer to that question was provided in October 1801 when English chemist and meteorologist John Dalton presented a series of essays to the Manchester Literary and Philosophical Society. His findings laid the groundwork for modern chemistry.
What is Dalton's law of partial pressures?
Dalton's law of partial pressures states that in a mixture of non-reacting gases, the total pressure exerted is equal to the sum of the partial pressures of the individual gases.
This means that in a sealed container, oxygen pushes against the walls exactly as hard as it would if the nitrogen were completely removed. Mathematically, it is written as:
P_total = P₁ + P₂ + P₃ + ...
Every component gas behaves like a vacuum to the others. If you want to explore related gas properties, see our guide on finding the average molar mass of a gas mixture.
Why doesn't the atmosphere separate into layers?
Dalton was not originally looking for a fundamental law of chemistry; he considered himself a meteorologist and kept a daily weather diary for 57 years.
He was fascinated by water vapour. Before 1801, many scientists thought water dissolved into the air like salt dissolving into a cup of water. Dalton proved that evaporation depends only on temperature, not on the other gases present. Water vapour pushes its way into the air independently.
Because every gas acts entirely on its own, they mix uniformly rather than layering by density. This led him to ask what the air was made of. While studying why water does not absorb all gases equally, he concluded that the difference depended on the weight and number of the ultimate particles. This was the birth of his atomic theory in 1803: different gases must be made of distinct, independent atoms.
How does the law apply in scuba diving?
Dalton's law is not just for laboratories; it is a matter of survival for scuba divers. As a diver goes deeper underwater, ambient pressure increases. Because a diver breathes air delivered at the surrounding water pressure, the partial pressure of every gas in the tank increases with depth.
For example, at a depth of 30 metres (where the total pressure is about 4 atm), the partial pressure of nitrogen becomes four times higher than it is at the surface. High partial pressures of nitrogen act as an anaesthetic on the brain, causing a condition known as nitrogen narcosis. To avoid these dangers, technical divers use Dalton's law to custom-blend their tanks with helium, effectively lowering the partial pressures of both nitrogen and oxygen.
How do you calculate partial pressure?
If you know the total pressure of a gas mixture and the mole fraction of the gas, you can find its partial pressure using Dalton's law.
For example, if you have a mixture of gases at 10 atm total pressure, and it is made of 20% gas A and 80% gas B, what are the partial pressures?
Find the mole fraction of gas A (0.20).
Multiply it by the total pressure: P_A = 0.20 × 10 = 2 atm.
Multiply the mole fraction of gas B (0.80) by the total pressure: P_B = 0.80 × 10 = 8 atm.
Check that they add up to the total: 2 atm + 8 atm = 10 atm.
If you are practising these calculations for an upcoming test, the Fermi Problem Library has worked problems with a step-by-step solver that explains the reasoning rather than printing an answer.
Understanding how to apply this law is essential for chemistry exams. You can also explore how it connects to calculating the γ of a gas mixture. If you need to review the ideal gas law and mole fractions first, you can find them in the Concept Library, which covers every concept in physics, chemistry, biology, mathematics and computing, each with its own lesson and its place in the map.
