Kinetic Molecular Theory and Gas Laws Explain Properties of Gas Molecules

The observations of different gas properties, as expressed by the various gas laws derived by Boyle, Charles, Gay-Lussac, and Avogadro, follow conceptually from kinetic molecular theory.

The pressure exerted by a gas results from the impact of constantly moving particles on the walls of its container.

Decreasing the volume of the container, while keeping the number of moles and temperature constant, brings the gas particles closer together, reducing their interparticle spacing.

In this smaller volume, the density of the gas and collision frequency — the frequency of molecule-wall collisions — increases. Therefore, the pressure exerted by the gas increases. The inverse relationship between pressure and volume is given by Boyle’s law.  

Adding more moles of gas to the container at constant temperature increases the gas density, and hence, the collision frequency. 

To maintain the initial pressure, the volume must expand. This direct relationship between volume and moles is given by Avogadro’s law. 

Now, consider the number of moles is kept constant and the temperature is raised. Because the average kinetic energy of gas particles increases proportionally with temperature, the particles collide more frequently and forcefully.

If the volume is held constant while the temperature is increased, the density of the gas and the collision frequency increases, and hence the pressure will also increase. The direct relationship between pressure and temperature is given by Gay-Lussac’s law. 

If, on the other hand, the pressure must remain constant along with a constant number of moles, then a rise in temperature must be accompanied by an increase in volume to spread the collisions out over a greater surface area. This direct relationship between volume and temperature is given by Charles’s law. 

Finally, according to kinetic molecular theory, gas particles do not attract or repel one another. In a mixture of different gases, the components act independently and their individual pressures remain unaffected by the presence of another gas. 

The total pressure of the mixture is, therefore, the sum of individual partial pressures. This is Dalton’s law.