ABSTRACT
Gases are an ensemble of molecular entities (atoms or molecules) without definite shape or volume, and where each particle is widely separated from the others. As a dispersed phase, the particles are treated as point masses with no particle–particle interaction. The temperature of a gas is proportional to the mean velocity of its particles, the pressure of a gas is a measure of the frequency and energy of collisions with the container, and the volume of a gas (in a flexible container) depends on the amount of gas and the mean kinetic energy of the gas. Because all of these properties (pressure, temperature, and volume) relate to the kinetic energy of a gas, they can be and are all related together in the ideal gas law (pV = nRT). In this chapter, the origins of the ideal gas law from the energy of a gas is derived and this understanding is used to explain the different gas laws (Boyle’s, Charles’, Amonton's [Gay-Lussac’s], and Avogadro’s).
