12.2.1 Kinetic Theory · 12. Particle Model | Edexcel GCSE Combined Science 1SC0: Physics

Section 1 of 9The kinetic model of a gas

The kinetic theory says that all matter is made of tiny particles (atoms or molecules) that are always moving. In a gas the particles are far apart compared with their own size, so there are almost no forces between them except during collisions. They move in straight lines at high speed until they hit something, then change direction.

Typical air molecules at room temperature travel at several hundred metres per second.

Random

Their motion is random: any direction, and a wide spread of speeds.

The volume of the particles themselves is tiny compared with the space between them, which is why gases are easy to compress and have low density.

Section 2 of 9Evidence: Brownian motion

Smoke particles viewed through a microscope in a glass cell jiggle in an irregular way.

The smoke grains are large enough to see, but air molecules are far too small.

The grains are knocked unevenly from different sides by the invisible, fast-moving air molecules.

Because the hits from one side are sometimes more frequent than from another, the grain moves along a zig-zag path.

Brownian motion

This random jerking is called Brownian motion and is direct evidence that gas molecules exist and move randomly.

Section 3 of 9How a gas produces pressure

Pressure is the force acting normally (at right angles) on each unit of area:

p = F / A

What the symbols mean: p = pressure (Pa = N m⁻²) · F = force acting at right angles to the surface (N) · A = area (m²)

where p is pressure in pascals (Pa, or N m⁻²), F is force in newtons (N) and A is area in square metres (m²).

When a molecule bounces off a wall its momentum changes direction, so the wall exerts a force on the molecule and, by Newton's third law, the molecule exerts an equal and opposite force on the wall.

How a gas produces pressure

A single collision is tiny, but there are billions of collisions every second.

Together they give a steady, almost constant force on the wall and therefore a steady pressure.

The force is always at right angles to the surface, whatever the shape of the container.

Section 4 of 9Changing the pressure

The pressure on the walls depends on two things: how often molecules hit the wall and how hard each hit is.

Changing the pressure

More molecules in the same volume

(higher density): more collisions each second, so higher pressure. This is why pumping air into a tyre raises its pressure.

Smaller volume

Smaller volume for the same number of molecules: molecules reach the walls more often, so pressure rises (see Boyle's law in 5.3.6).

Higher temperature

Higher temperature molecules move faster, hit the walls more often and with greater force, so pressure rises if the volume is fixed (see 5.3.5).

Section 5 of 9Gases compared with solids and liquids

Property	Solid	Liquid	Gas
Particle spacing	Close, regular	Close, irregular	Very far apart
Movement	Vibrate about fixed positions	Slide past each other	Fast, random, free
Compressible?	No	Almost not	Yes, easily

A gas fills any container because its particles are not held together. A sealed gas exerts pressure on every part of the container, not just the bottom, because the particles move in all directions.

Section 6 of 9Why gases are easy to compress and spread out

Because the particles in a gas are so far apart, there is mostly empty space between them.

1

When you push on a gas, you are not squashing the molecules themselves; you are reducing the empty space between them.

2

That is why a gas can be compressed to a fraction of its volume, whereas a liquid, whose molecules are already touching, can hardly be compressed at all.

3

It is also why a gas, unlike a liquid, will always expand to fill whatever container you put it in, however large.

4

The molecules keep moving until they have spread through the whole space.

The same model explains diffusion.

Why gases are easy to compress and spread out

If you open a bottle of perfume in a room, molecules of scent travel outwards and are slowly carried across the room by their random motion and many collisions with air molecules.

The process is slow compared with the speed of the molecules because each molecule suffers an enormous number of collisions, changing direction each time.

Link to atmospheric pressure

The air around us is a gas, and the pressure it exerts, about 1.0 × 10⁵ Pa at sea level, is due to the same collisions.

Why gases are easy to compress and spread out

Every square metre of your skin is hit by air molecules with a total force of roughly 100 000 N.

You do not feel crushed because the fluids inside your body push outwards with the same pressure.

At high altitude the air is thinner, so fewer molecules strike each square metre per second and the pressure is lower.

Section 7 of 9Worked examples

A rectangular wall of a gas tank has area 0.40 m². The gas exerts a force of 40 000 N on it. Calculate the pressure.
A sealed syringe of air is pushed so that its volume halves while the temperature stays the same. Explain, using particles, why the pressure rises.

Section 8 of 9Key terms

Term	Meaning
Kinetic theory	Model in which matter consists of particles in constant motion.
Brownian motion	Random jiggling of visible particles caused by collisions with fast-moving invisible molecules.
Pressure	Force acting normally per unit area, p = F / A.
Pascal (Pa)	Unit of pressure, equal to 1 N m⁻².
Random motion	Movement in no fixed direction with a range of speeds.

Section 9 of 9Exam technique

1

Saying molecules in a hot gas 'expand'. The molecules do not get bigger; they move faster and spread further apart.

2

Thinking gas pressure is caused by molecules pushing each other. It is caused by collisions with the container walls.

3

Saying pressure acts only downwards in a gas. Gas pressure acts equally in all directions on every surface.

4

Forgetting that a single collision is not the pressure: pressure is the combined effect of enormous numbers of collisions.