Kinetic theory of gases
A gas's temperature is a measure of the average kinetic energy of its particles, which move randomly and collide elastically with each other and the container walls.
Practise this conceptKinetic theory explains the everyday behaviour of gases by imagining what their particles are doing: zipping around randomly, colliding elastically with each other and the walls of their container, with essentially no attraction between them except during a collision.
Gas pressure, in this picture, is just the combined effect of countless tiny collisions between particles and the container walls every second — squeeze the same number of particles into a smaller space, and they hit the walls more often, which is exactly why compressing a gas raises its pressure.
Temperature turns out to be a measure of this microscopic chaos: it's directly tied to the average kinetic energy of the particles. Heat a gas up, and you're not doing anything mysterious — you're simply making its particles move faster on average, which shows up macroscopically as a higher temperature and, if the container allows it, a higher pressure too.
Key ideas
Requires: Temperature and heat, Conservation of energy
Unlocks: Ideal gas law
Formulas
| Symbol | Name | Unit |
|---|---|---|
| speed | m/s (metre per second) | |
| temperature | K (kelvin) | |
| molar mass | kg/mol (kilogram per mole) |
Common misconceptions
- Gas particles are strongly attracted to each other, similar to particles in a liquid.
- All the particles in a gas at a given temperature move at exactly the same speed.
- Gas pressure is caused by particles pushing against each other, not against the container walls.