Summary

Samples of matter are either solid, liquid, or gaseous. Each state has charactreictics of shape, compressibility, anddensity that are dependent on the bonds between the component particles of the sample. Each particle also has a kinetic energy (KE = (1/2) mv², where m = mass and v = velocity). A collection of particles will have a characteristic distribution of kinetic energy, the average of which is dependent on the temperature of the sample.

The kinetic molecular theory for gases describes the behavior of an ideal gas as a collection of very tiny molecules moving rapidly and freely in a very large volume. There are no attractive forces between these molecles. Collisions occur with no loss of kinetic energy whether the collision is with other molecules or with the walls of the container. The pressure a gas exerts is caused by collisions with the walls of the container. The temperature of a gas is proportional to the average kinetic energy of its molecules.

An ideal gas obeys certain laws that are predictable from the kinetic molecular theory. They are:

Boyle's Law: P₁V₁ = P₂V₂
Charles' Law: V₁/T₁ = V₂/T₂
Combined Gass Law: P₁V₁/T₁ = P₂V₂/T₂
Avogadro's Hypothesis and molar volume: 22.4 l at STP
Dalton's Law of Partial Pressures: P_Total = P₁ + P₂ + P₃ + ...
Ideal Gas Equation: PV = nRT

Application of these laws makes possible a wide variety of calculations.

Real Gases differ from the ideal gas described by the kinetic molecular theory in that their molecules have real volume and there are attractive forces between the molecules due to permanent or temporary dipoles.