Experiments With Kinetic Molecular Theory

Kinetic molecular theory, also known as the Kinetic Theory of Gases is a powerful model that seeks to explain the the measurable characteristics of gas in terms of the small scale movements of gas particles. Kinetic theory explains the properties of gases in terms of the motion of its particles. Kinetic theory is based on a number of assumptions and because of this it is approximate model.

Gases in the kinetic model are considered to be "perfect". Perfect gases comprise of molecules that move entirely at random and never stop moving. All gas particle collisions are completely elastic, meaning no energy is lost. (If this were not the case gas molecules would eventually run out of energy and accumulate on the floor of their container.) The next assumption is that size of the molecules is negligible meaning they are essentially have zero diameter. This is almost true for very small monoatomic gases such as helium, neon or argon. The final assumption is that gas molecules do not interact except when they collide. Kinetic theory does not consider any electrostatic forces between molecules.

A gas has three intrinsic properties, pressure, temperature, and volume. These three properties are linked to each other and can be explained using kinetic theory. Pressure is caused by particles hitting the wall of the gas container. A non-rigid container such as a balloon will expand until the gas pressure inside the balloon equals that on the outside of the balloon. When a gas is a low pressure the number of collisions is less than at high pressure. Increasing the temperature of a gas in a fixed volume also increases its pressure as the heat causes the particles to move more rapidly. Similarly expanding the volume in which a gas can move lowers both its pressure and temperature.

Robert Boyle was amongst the first to discover links between the properties of gases. Boyle's law states that a at a constant temperature the pressure of a gas is inversely proportional to its volume. Charles' law, after Jacques Charles considers temperature, finding that for a fixed pressure, the volume of a gas is directly proportional to its temperature. These equations were combined to form the perfect gas equation of state for one mole of gas, pV=RT, where p is pressure, V is volume, T is temperature and R is the universal gas constant.

The perfect gas law works well for low pressures. At high pressures or low temperatures gas molecules come into close enough proximity to interact; it is these interactions which cause gases to condense into liquids and without them all matter would be gaseous. These interactomic interactions are called Van der Waals forces. Consequently, the perfect gas equation can be modified to include a component to describe intermolecular forces. This more complicated equation is called the Van der Waals equation of state.