Summary
Highlights
Under most conditions, gas particles occupy negligible volume compared to the total volume, meaning gases are mostly empty space and are highly compressible, unlike solids or liquids.
Pressure arises from gas particles colliding with and imparting kinetic energy to the container walls. The sheer number of particles, even if tiny, generates significant force, with more and faster particles leading to higher pressure.
Gas particles are assumed to have no gravitational or electromagnetic influence on each other. Collisions are perfectly elastic, meaning no kinetic energy is lost, much like billiard balls bouncing.
The average kinetic energy of gas particles is directly proportional to the gas's temperature in Kelvin. Increasing temperature means increased kinetic energy and faster-moving particles; temperature is indicative of average molecular velocity.
The kinetic molecular theory explains ideal gas laws. Boyle's law (inverse pressure-volume) is explained by particles needing to travel further in larger volumes, reducing collisions. Charles's law (direct volume-temperature) is explained by faster particles needing more space to maintain constant collision frequency, leading to volume expansion. Amontons's law (direct pressure-temperature) is explained by faster particles hitting container walls more frequently and forcefully.
Gas laws describe observed relationships but don't explain why. Kinetic molecular theory provides a powerful explanation for the behavior of gases, deriving all ideal gas laws from its five postulates.
Gas consists of particles (atoms or molecules) in constant, straight-line motion. They collide with each other and container walls, changing direction. This implies particles follow laws of motion like macroscopic objects.