The description of the behavior of a gas in terms of the macroscopic state variables such as pressure, volume and temperature can be related to simple averages of microscopic quantities, such as the mass and speed of the molecules in the gas. The resulting theory, called the kinetic molecular theory of gases, provides a detailed model of dilute gases. From the point of view of kinetic theory, a confined gas consists of a large number of rapidly moving particles. In a monatomic gas, like helium and neon, these particles are single atoms, but in polyatomic gases, like oxygen and carbon dioxide, the particles are molecules. In kinetic theory, it is common practice to refer to the constituent particles of a gas as “molecules”, even if they are actually atoms.
In a gas at room temperature, a very large fraction of molecules are moving at speeds of a few hundred meters per second. These molecules are in constant motion and make elastic collisions, both with each other and with the walls of the container. In the context of the kinetic theory, any effects due to gravity are neglected, so there are no preferred velocity vectors (i.e., directions) either. The molecules are separated, on average, by distances that are large compared with their diameters. They also exert no forces on each other except when they undergo elastic collisions. This assumption is equivalent to assuming a very low gas density, which is the same as assuming that the gas is an ideal gas. Because momentum is conserved, the collisions that the molecules make with each other have no effect on the total momentum in any direction. Thus, such collisions may be neglected.
The kinetic molecular theory of gases begins with five postulates that describe the behavior of molecules in a gas. These postulates are based upon some basic scientific notions, but they also involve some simplifying assumptions.
They are:
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A gas consists of a collection of small particles traveling in straight-line motion and obeying Newton’s Laws.
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The diameter of a molecule in a gas can be assumed to be negligible compared to the distance between collisions.
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Gas molecules exert no attractive or repulsive force on each other except for when colliding and collisions between molecules are perfectly elastic, which is to say that no energy is gained or lost during the collision.
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The particles are considered to be in constant random motion. The pressure exerted by a gas in a container is a result of collisions between the gas molecules and the container walls.
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The average kinetic energy of a molecule is proportional to the absolute temperature of the gas. More specifically, KE= 3kT/2, where T is the absolute temperature in Kelvin and k is the Boltzmann constant.