IB Physics SLTopic B.3 — The Behaviour of GasesPaper 1 & 2Kinetic Theory · Real vs Ideal Gases~7 min read
Kinetic Theory of Gases
The ideal gas equation works because of a simplified molecular picture sitting underneath it. Kinetic theory strips a gas down to a set of assumptions — and understanding exactly what those assumptions are is what tells you when the equation can be trusted, and when it can’t.
📘 What you need to know
Kinetic theory is a model — a simplified system of assumptions used to approximate how real gases actually behave.
It links the microscopic properties of individual molecules (mass, speed) to the macroscopic properties of the gas as a whole (pressure, volume).
A gas is treated as a huge number of identical molecules in constant, random, high-speed motion, obeying Newton’s laws and colliding elastically.
Real gases don’t perfectly obey the ideal gas equation, because some of these assumptions aren’t fully realistic — most notably, real molecules do exert forces on each other.
An ideal gas is a good approximation of a real gas at low pressure, low density, and temperatures well above the substance’s boiling point.
The Assumptions Behind an Ideal Gas
Kinetic theory rests on a specific set of simplifying assumptions about how gas molecules behave:
⚙️ The model, assumption by assumption
Enormous numbers of identical molecules — a gas consists of many molecules of the same mass, moving randomly at high speed.
Negligible molecular volume — molecules are treated as point particles, since their own size is negligible compared to the volume of the container.
Newtonian, elastic collisions — molecules obey Newton’s laws of motion and collide elastically with each other and the container walls, losing no kinetic energy.
No forces except during collisions — there are no intermolecular forces between molecules except in the brief instant they collide.
Instantaneous collisions — the time taken for a collision is negligible compared to the time between collisions.
No external forces — effects like gravity are ignored entirely.
Average behaviour matters — because the number of molecules is so large, it’s the average speed and behaviour that’s meaningful, not any individual molecule’s path.
Put together, these assumptions explain why gas pressure is uniform: every molecule contributes its own tiny, perpendicular force to the walls, and averaging over an enormous number of collisions produces one steady, consistent pressure.
Why Real Gases Aren’t Quite Ideal
No real gas perfectly satisfies every one of these assumptions. The most significant gap is the claim that molecules exert no force on each other outside of collisions — real molecules do attract and repel one another, even weakly, and they do take up a small but non-zero volume. Under everyday conditions those effects are small enough to ignore, which is why the ideal gas equation works so well for something like the air around you. Push those conditions to their limits, though, and the cracks start to show.
An ideal gas treats molecules as point particles with no interactions. A real gas at high density has molecules close enough together that their size and mutual attraction start to matter.
When the Approximation Works — and When It Doesn’t
A real gas behaves as a good approximation of an ideal gas when:
The gas pressure is low.
The gas density is low.
The temperature is sufficiently higher than the substance’s boiling point.
Under these conditions, molecules are spread far enough apart that their own volume is genuinely negligible next to the container, and any forces between them are far too weak to matter. Air at ordinary room temperature and pressure is a good real-world example — no gas is perfectly ideal, but air behaves close enough to one that the ideal gas equation gives reliable answers.
Push a gas to high pressure and density, and molecules are forced much closer together: intermolecular attractive forces become significant, and the molecules’ own volume can no longer be dismissed as negligible. Push the temperature low enough, and the gas may condense into a liquid altogether — at that point it isn’t behaving as a gas at all, let alone an ideal one.
Quick recap: kinetic theory assumes point-like molecules, elastic collisions, and no intermolecular forces except during impact. Real gases only approximate this at low pressure, low density, and temperatures well above their boiling point.
WE 1
State two assumptions of the kinetic theory of gases that are not perfectly true for real gas molecules.
Assumption 1
That molecules have negligible volume — real molecules do occupy a small but non-zero volume.
Assumption 2
That no forces act between molecules except during collisions — real molecules do exert weak attractive forces on one another even when not colliding.
Both assumptions break down most noticeably at high pressure and high density
WE 2
Explain why a real gas is more likely to obey the ideal gas equation at low pressure than at high pressure.
At low pressure
Molecules are spread far apart, so their own volume is genuinely negligible and intermolecular forces are far too weak to have any real effect.
At high pressure
Molecules are forced much closer together, so their volume becomes significant relative to the container, and intermolecular attractive forces become strong enough to affect their motion.
Low pressure keeps molecules far apart → assumptions hold → closer to ideal behaviour
💡 Top tips
If a question asks for a specific number of assumptions (e.g. “state two”), give exactly that many — extra points beyond what’s asked won’t earn extra credit.
Tie deviations from ideal behaviour to a specific condition — high pressure, high density, or low temperature — rather than a vague “the assumptions aren’t true”.
Remember that no real gas is ever perfectly ideal — some gases (like helium at room temperature) just approximate it far more closely than others.
⚠ Common mistakes
Listing every assumption from memory when the question asks for only one or two — read the command term carefully.
Saying a gas “becomes ideal” under certain conditions — it’s always an approximation, however good.
Forgetting that very low temperatures can cause a gas to condense into a liquid entirely, which is a far bigger departure from ideal behaviour than just “the equation being slightly off”.
Up next: Derivation of the Kinetic Theory of Gases Equation — where we use these exact assumptions to derive a formula connecting gas pressure directly to molecular mass and speed.
Want this to actually stick before the exam?
Book a free session and we’ll work through kinetic theory and real-vs-ideal gas questions until they’re second nature.
