IB Physics HLThe Behaviour of GasesPaper 1 & 2Kinetic Theory~8 min read
Kinetic Theory of Gases
So far we’ve used PV = nRT as if a gas just knows how to behave. Kinetic theory is the model that explains why — picture a gas as a huge swarm of tiny balls, forever bouncing around, obeying a short list of simple rules. From that picture, all the gas behaviour we’ve seen falls out naturally. Here we set up the model and check when real gases actually follow it.
📘 What you need to know
Kinetic theory models a gas as many tiny particles in constant random motion
It links microscopic properties (mass, speed) to macroscopic ones (pressure, volume, temperature)
It rests on a set of simplifying assumptions (point particles, elastic collisions, no forces except in collisions, and so on)
Pressure arises because countless particles collide with the container walls
A real gas behaves most like an ideal gas at low pressure, low density and high temperature
It deviates at high pressure/density and low temperature, where molecules crowd together and forces matter
What kinetic theory is
Kinetic theory is a model: a deliberately simplified picture we use to approximate how real gases behave. The idea is that a gas is nothing more than a crowd of atoms or molecules flying around randomly at high speed. By treating those particles with basic mechanics, the model bridges two very different worlds — the microscopic (individual particles, their mass and speed) and the macroscopic (the pressure, volume and temperature you actually measure).
Identical particles fly in random directions at random speeds, colliding elastically with the walls. Each wall collision gives a tiny push — and billions per second add up to the steady gas pressure.
The assumptions
The model only works because we agree on some simplifying rules. These are the assumptions of kinetic theory — worth knowing by heart, as “state the assumptions” is a classic exam ask.
A gas is made of many identical molecules, all with the same mass.
The molecules are so small that their own volume is negligible compared with the container — treat them as points.
They are in continuous, random motion at high speed.
They obey Newton’s laws of motion.
All collisions — with each other and with the walls — are perfectly elastic (no kinetic energy is lost).
There are no forces between molecules except during a collision.
The time of a collision is negligible compared with the time between collisions.
External forces such as gravity are ignored.
There are so many molecules that we work with averages, and their many wall collisions produce a steady, uniform pressure.
A quick word on “elastic”: it doesn’t mean the particles don’t collide — they collide constantly. It means that when they do, no kinetic energy is lost. And “point particles” doesn’t mean massless: each molecule still has mass, we just treat its size as negligible.
Real gases vs ideal gases
Here’s the catch: those assumptions aren’t perfectly true. Real molecules do take up space, and they do attract one another. So a real gas only obeys PV = nRTapproximately — and how good that approximation is depends on the conditions.
An ideal gas obeys PV/T = constant with point-like, non-interacting molecules. A real gas has molecules of real size that attract one another, so it only approximates that behaviour.
A real gas is a good approximation to an ideal gas when:
the pressure is low — molecules are far apart,
the density is low — again, plenty of space between them, and
the temperature is well above the boiling point — so it stays firmly a gas.
It breaks down at high pressure and density (molecules are forced close together, so their volume is no longer negligible and their attractions start to matter) and at low temperature (where the gas may condense into a liquid and stop behaving like a gas at all). No gas is perfectly ideal — but air at everyday room temperature and pressure comes very close.
EXAM Q
State two assumptions of the kinetic theory of gases, and explain why a real gas deviates from ideal behaviour at high pressure.
Two assumptions (any two)
• the molecules’ own volume is negligible
• there are no forces between molecules (except in collisions)
Why it deviates at high pressure
At high pressure the molecules are squeezed close together. Now their own volume is no longer negligible compared with the container, and the attractive forces between them become significant.
Both assumptions break down, so PV = nRT no longer holds accurately.
💡 Top tips
Learn the assumptions — “state the assumptions of kinetic theory” is a frequent, easy-marks question.
“Elastic” = no KE lost, not “no collisions”. Particles collide all the time.
Point particles still have mass — only their volume is treated as negligible.
Ideal conditions: low pressure, low density, high temperature.
Link deviations to assumptions — explain which assumption fails (volume, or forces).
⚠ Common mistakes
Saying molecules “don’t collide” — they collide constantly; the collisions are elastic
Thinking point particles are massless — they have mass, just negligible size
Claiming a real gas is exactly ideal — it’s only an approximation
Listing deviation conditions without saying which assumption breaks
Forgetting that a very cold gas may condense and stop behaving as a gas
Quick recap: Kinetic theory pictures a gas as many identical particles in random motion, colliding elastically and exerting pressure on the walls — the bridge between the microscopic and the macroscopic. Its assumptions (negligible volume, no forces except in collisions, elastic collisions…) make a gas “ideal”. Real gases follow this best at low pressure, low density and high temperature, and deviate when crowded or cold.
Here’s the exciting part: this simple bouncing-particle picture isn’t just hand-waving — you can turn it into real equations. Next we’ll actually derive the kinetic theory equation, following a single molecule as it bounces around a box and building up to a formula for pressure. It’s one of the most satisfying derivations in the whole course.
Want this to actually click before the exam?
Book a free meeting and let’s work through the tricky bits together.
