Kinetic Molecular Theory

A gas looks empty, but it's full of particles moving in straight lines at high speed until they hit something, usually the container wall. Each collision delivers a tiny push. One hit is nothing. Trillions of hits per second, spread evenly over the wall, blur together into a steady force per area that a pressure gauge reads as constant. Pressure is really the statistical average of countless individual impacts, not one dramatic shove.

Heat the gas at fixed volume and the particles move faster on average. Faster particles hit the walls harder and more often in a given second, so the force per collision goes up and the collision rate goes up too. Both effects push pressure higher, which is exactly what the gas laws describe when they link temperature and pressure at constant volume.

This is also where a common confusion lives: temperature is not speed. Temperature tracks the average kinetic energy of the whole population of particles. Kinetic energy depends on mass and on speed squared, and in any real sample particles are flying at a huge range of speeds, some slow, some fast, most in between. That spread is described by a distribution of speeds, not a single value. Average kinetic energy is directly proportional to Kelvin temperature, which is why doubling the Kelvin temperature exactly doubles the average kinetic energy, even though the average speed itself does not simply double. This proportionality is the bridge between kinetic molecular theory and the gas laws covered right alongside it.