State the main points in the kinetic-molecular theory of gases.

Answer 1

The kinetic-molecular theory rests on five postulates. Learn them as a set, because each one is doing specific work — and knowing which assumption fails is how you explain real-gas behaviour later.

1. A gas consists of a very large number of tiny particles in constant, random motion. They travel in straight lines until they collide with something.

2. The particles' own volume is negligible compared with the volume of the container. A gas is mostly empty space — which is why gases are compressible and liquids are not.

3. There are no attractive or repulsive forces between particles except during the instant of a collision. Each molecule moves independently of the others.

4. Collisions are perfectly elastic. Total kinetic energy is conserved. If it were not, molecules would gradually slow down and the gas would settle to the bottom of its container — which does not happen.

5. The average kinetic energy is directly proportional to absolute temperature. KE_avg = (3/2)kT per molecule. Note average — individual molecules span a wide distribution of speeds.

What each postulate buys you

• Pressure is molecules striking the walls. More collisions, or harder ones, means higher pressure.
• Boyle's law: shrink the volume and the walls are hit more often, so P rises.
• Charles's law: raise T and postulate 5 says molecules move faster, so the gas must expand to hold pressure constant.
• Avogadro's law: since particle volume and identity do not matter (postulates 2 and 3), equal numbers of any gas occupy equal volumes.

A consequence worth knowing

Postulate 5 says temperature fixes the average kinetic energy — not the average speed. Since KE = ½mv², at the same temperature a lighter molecule must move faster:

u_rms = √(3RT/M)

That is why helium escapes a balloon faster than nitrogen, and it is the basis of Graham's law of effusion.

Where the theory breaks — and this is the exam favourite

Postulates 2 and 3 are idealisations, and they fail under two conditions:

• High pressure — molecules are crowded, so their own volume is no longer negligible (postulate 2 fails).
• Low temperature — molecules move slowly enough that intermolecular attractions matter (postulate 3 fails).

Real gases behave most ideally at high temperature and low pressure. The van der Waals equation patches exactly these two failures, with a b term correcting for molecular volume and an a term for attractions — so if you can name which postulate breaks, you already understand where van der Waals comes from.