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The Kinetic Molecular Theory

The kinetic molecular theory describes the properties of molecules in terms of motion (kinetic energy) and of temperature. The theory is most often applied to gases but is helpful in explaining molecular behavior in all states of matter. As applied to gases, the kinetic molecular theory has the following postulates:

  1. Gases are composed of very tiny particles (molecules). The actual volume of these molecules is so small as to be negligible compared with the total volume of the gas sample. A gas sample is, then, mostly empty space. This fact explains the compressibility of gases.

  2. There are no attractive forces between the molecules of a gas. This postulate explains why, over a period of time, the molecules of a gas do not cluster together at the bottom of its container.

  3. The molecules of a gas are in constant, rapid, random, straight-line motion. This postulate explains why a gas spreads so rapidly through the available space - for example, why the smell of hot coffee can spread quickly from the kitchen throughout the house.

  4. During their motion, the gas molecules constantly collide with one another and with the walls of the container. (The collision with the walls provides the pressure exerted by a gas.) None of these collisions is accompanied by any loss of energy; instead, they are what is known as elastic collisions. A "new" tennis ball collides more elastically than a "dead" tennis ball.

  5. The average kinetic energy of the molecules in a gas sample is proportional to its temperature (Kelvin) and is independent of the composition of the gas. In other words, at the same temperature, all gases have the same average kinetic energy. It also follows from this postulate that at zero Kelvin all molecular motion has ceased.

These postulates and the experimental evidence for them are summarized in Table 9.2.

TABLE 9.2 The kinetic molecular theory
Postulate Evidence
1. Gases are tiny molecules in mostly empty space. The compressibility of gases.
2. There are no attractive forces between molecules. Gases do not clump.
3. The molecules move in constant, rapid, random, straight-line motion. Gases mix rapidly.
4. The molecules collide elastically with container walls and one another. Gases exert pressure that does not diminish over time.
5. The average kinetic energy of the molecules is proportional to the Kelvin temperature of the sample. Charles' Law (Section 9.5B)


Clearly, the actual properties of individual gases vary somewhat from these postulates, for their molecules do have a real volume and there is some attraction between the molecules. However, our discussion will ignore these variations and concentrate on an ideal gas, one that behaves according to this model.

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