According to Kinetic Theory, the pressure exerted by the gas in a container is caused by the collisions of its molecules with the walls of the container. The pressure changes directly with the number of molecules colliding with the wall per unit of time. When pressure on a given mass of gas is increased at constant temperature, the distance between molecules decreases and the volume decreases. If volume is reduced to half, the number of molecules per unit volume doubles, thus collisions per unit time double, and pressure doubles → Boyle’s Law.
Robert Boyle verified this using a J-shaped tube and mercury. By adding mercury until the pressure reached ~2 atm (extra 760 mm Hg), the trapped air volume became half of the original, confirming the inverse relationship. The experiment maintained constant temperature and fixed gas mass.
According to kinetic theory, increasing gas temperature increases average speed & kinetic energy, causing more frequent and harder collisions → pressure increases if volume fixed. To keep pressure constant, volume must expand. Visualize a piston moving upward when heated. Jacques Alexander Charles stated: “Volume of a given mass of gas varies directly with absolute temperature at constant pressure.” Mathematically: V ∝ T (Kelvin scale).
Kinetic theory: pressure depends on number of particles and their speed. At constant T & P, increasing gas amount requires larger volume to keep pressure unchanged. Equal volumes of different gases at same T & P contain equal number of molecules. V ∝ n (moles) at constant T & P.
Test your understanding of Kinetic Theory, Boyle’s Law, Charles’ Law & Avogadro’s Law. Click on an option — correct turns green, wrong options turn red, then submit for detailed result & answer key.