Kinetic theory and gas behaviourIB MYP Physics: Revision notes

The kinetic theory says that gas particles are in constant, random motion. As they move they collide with the walls of their container. Each collision exerts a tiny force on the wall. Millions of collisions every second together produce a steady gas pressure.

Pressure is measured in pascals (Pa) or kilopascals (kPa).

The temperature of a gas is a measure of the average kinetic energy of its particles. When a gas is heated:

• the particles gain kinetic energy
• they move faster on average

The particles do not get bigger. They just move faster.

As a gas is cooled the particles slow down. The lowest temperature possible is absolute zero, 0 K or −273 °C, where particles have the least possible kinetic energy.

The Kelvin scale starts at absolute zero. To convert:

K = °C + 273

Worked example: 27 °C = 27 + 273 = 300 K. And 250 K = 250 − 273 = −23 °C.

For a fixed mass of gas in a rigid container (constant volume), raising the temperature increases the pressure:

1. Particles gain kinetic energy and move faster.
2. They hit the walls more often.
3. They hit the walls with more force.
4. So the pressure increases.

Cooling the gas has the opposite effect. This is why a sealed can can burst if it is heated.

For a fixed mass of gas at constant temperature, reducing the volume increases the pressure. In a smaller space the particles travel shorter distances, so they hit the walls more often. The speed does not change.

Boyle's law: pressure × volume stays constant, so

p₁V₁ = p₂V₂

Worked example: 120 cm³ of air at 100 kPa is squashed to 40 cm³ at constant temperature. 100 × 120 = p₂ × 40, so p₂ = 12 000 ÷ 40 = 300 kPa.