Kinetic Molecular Theory of Gases

From experimental laws to a molecular picture

Laws like Boyle's and Charles' are concise summaries of what scientists observe in the lab. But observing that pressure rises on compression only tells us what happens — it doesn't explain why, at the level of individual molecules. To answer that "why", scientists build a theory: a mental model that makes sense of the observations. The theory that explains gas behaviour at the molecular level is the kinetic molecular theory of gases.

Because it describes what atoms and molecules — things we cannot directly see — are doing, this theory gives us a microscopic model of gases. Its postulates:

The postulates

• Gases consist of a very large number of tiny, identical particles, so small and so widely spaced that their actual volume is negligible compared to the empty space between them; they can be treated as point masses. This explains gases' great compressibility.
• There is no force of attraction between gas particles at ordinary temperature and pressure — supported by the fact that a gas expands to fill whatever space is available to it.
• Gas particles are always in constant, random motion. If they were fixed in place, a gas would have a definite shape, which is never observed.
• Particles move in straight lines in all directions, colliding with each other and with the container walls; pressure arises from these wall collisions.
• Collisions between gas molecules are perfectly elastic — total kinetic energy before and after a collision is unchanged (individual molecules' energies may change, but the sum is conserved). If energy were lost in collisions, molecules would gradually slow to a stop and the gas would settle — which is not observed.
• At any moment, different particles have different speeds and hence different kinetic energies — a reasonable consequence of ongoing collisions constantly redistributing energy, even if all particles started out identical. …