kinetic-molecular theory of gases, physical theory that explains the behavior of gases on the basis of the following assumptions: (1) Any gas is composed of a very large number of very tiny particles called molecules; (2) The molecules are very far apart compared to their sizes, so that they can be considered as points; (3) The molecules exert no forces on one another except during rare collisions, and these collisions are perfectly elastic, i.e., they take place within a negligible span of time and in accordance with the laws of mechanics. A gas corresponding to these assumptions is called an ideal gas; as the temperature of a real gas is lowered, or its pressure is raised, its behavior no longer resembles that of an ideal gas because one or more of the assumptions of the theory is no longer valid. The analysis of the behavior of an ideal gas according to the laws of mechanics leads to the general gas law, or ideal gas law: The product of the pressure and volume of an ideal gas is directly proportional to its absolute temperature, or PV = kT (see gas laws). Boyle's law, Charles's law, and Gay-Lussac's law, which are special cases of the general gas law, may also be easily derived. The theory further shows that the absolute temperature is directly proportional to the average kinetic energy of the molecules, thus providing an interpretation of the nature of temperature in general in terms of the detailed structure of matter (see temperature; Kelvin temperature scale). Pressure is seen to be the result of large numbers of collisions between the molecules and the walls of the container in which the gas is held. See thermodynamics.
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