Kinetic Theory of Gases

About this topic

Kinetic theory models a gas as many tiny particles in constant random motion colliding elastically with each other and the container walls. Pressure arises from the combined force of countless collisions, and the average kinetic energy is proportional to absolute temperature.

The particles are very far apart compared to their size and exert almost no forces except during collisions. Average kinetic energy per particle is 3/2 kT, where k is Boltzmann's constant 1.38 × 10⁻²³ J/K. Squeezing a gas into a smaller volume makes collisions with the walls more frequent, raising pressure, and heating it makes them harder and faster, also raising pressure.

Worked example: at room temperature, air molecules move at roughly 500 m/s on average, faster than a rifle bullet, but collide billions of times a second and so travel randomly rather than in straight lines. Lighter molecules move faster at the same temperature, which is why hydrogen escapes more easily than nitrogen.

The theory explains diffusion, such as perfume spreading across a room, and Brownian motion, the jiggling of pollen grains seen by Robert Brown, which was Einstein's evidence that atoms exist. A common misconception is that gas particles stop at absolute zero in a literal sense; they reach minimum energy. Another is that temperature relates to the speed of the whole gas, when it measures random motion.