Tiny Particles, Big Pressure: KMT and the Gas Laws

Goggles on first, always. Now picture the world as countless tiny particles in constant motion. That single idea is kinetic molecular theory, and it explains almost everything about states of matter and gases. In a solid, particles are packed close and only vibrate in place, so solids hold their shape. In a liquid, particles are close together but not rigidly held — they are free to move past one another, which is why liquids flow and take their container's shape. Note that liquid particles are not in continuous contact the way a solid is; they are simply much closer than gas particles and are not locked in fixed positions. In a gas, particles are far apart and zoom freely, so gases fill any container and are easily squeezed. Temperature is the key. The Kelvin temperature is proportional to the average kinetic energy of the particles. Heat them up and they move faster; cool them and they slow down. Pressure comes from particles colliding with the walls of their container. More collisions, or harder collisions, means more pressure. Squeeze a gas into half the volume and the particles hit the walls twice as often, so pressure doubles. That is Boyle's law: pressure times volume stays constant when temperature and amount are fixed. Heat a sealed, rigid container and faster particles strike the walls harder and more frequently, raising the pressure. This is why you must never heat a sealed can. If you ever feel stuck on a problem, write down what stayed constant first, then track which way the particles' motion or spacing changes.