Ideal gases: (b) model of kinetic theory of gases

The idea

The kinetic theory model treats a gas as a large number of identical point molecules in continuous random motion. Their volume is negligible compared with the container, and collisions between molecules and with the walls are perfectly elastic. Between collisions no forces act, except during contact. Pressure arises from the rate of change of momentum when molecules collide with the walls. For example, doubling the absolute temperature of a fixed volume of gas increases the mean square speed of its molecules, so they strike the walls more often and harder, raising the pressure. This model explains the macroscopic gas laws in terms of molecular behaviour.

What earns the marks

• Gas consists of a very large number of identical molecules moving in random directions with a range of speeds.
• The volume of the molecules is negligible compared with the volume of the container.
• Collisions between molecules, and between molecules and the container walls, are perfectly elastic.
• Intermolecular forces are negligible except during collisions.
• Pressure is caused by the rate of change of momentum of molecules colliding with the walls.

Common mistakes and how to put them right

• Thinking molecules all move at the same speed.

  Put it right: The model uses a distribution of speeds and refers to mean square speed.

• Believing intermolecular forces act continuously.

  Put it right: The model assumes they are negligible except during collisions.

• Confusing the model with a real gas.

  Put it right: The ideal gas model ignores molecular volume and attractive forces, which real gases have.

Exam tips

• State assumptions precisely: 'volume of molecules negligible' not just 'molecules are small'.
• Link pressure to rate of change of momentum, not just to collisions.
• Use the phrase 'random motion' when describing molecular movement.