Kinetic Theory of Gases

Having built a picture of ordered solids, the chapter now turns to the opposite extreme of molecular organisation

— the gaseous state, where particles are in constant, chaotic motion with almost no fixed structure at all. The

kinetic theory of gases is the model that explains the bulk, measurable properties of a gas (pressure,

temperature, volume, diffusion) in terms of the ceaseless motion of the tiny particles that make it up.

The theory rests on a small set of postulates. First, a gas consists of a very large number of extremely small

particles (atoms or molecules) that are in continuous, random, straight-line motion in all directions, colliding

with one another and with the walls of their container. Second, the actual volume of the gas particles

themselves is negligible compared with the total volume occupied by the gas — the particles are treated as

point masses. Third, there is no significant force of attraction or repulsion between gas particles, or between

the particles and the walls of the container — an ideal gas particle moves independently of every other. Fourth,

collisions between gas particles, and between particles and the container walls, are perfectly elastic — no

kinetic energy is lost in a collision, only redistributed among the colliding particles. Fifth, between

collisions, particles travel in straight lines at constant speed (Newton's first law applies between collisions,

since there are no intermolecular forces to curve their path). Sixth, at any given instant, different particles

in the gas possess a wide range of speeds — the distribution of these speeds is described statistically (the

Maxwell-Boltzmann distribution) rather than by a single value, though the average kinetic energy of the

particles is directly proportional to the absolute (Kelvin) temperature of the gas, and is the same for all gases

at a given temperature regardless of their molar mass.

From this picture, pressure is explained as the cumulative effect of an enormous number of gas particles

continually striking the container walls and exerting a tiny force at each collision — summed over the huge number

of particles and collisions per second, this produces the steady, measurable pressure of the gas. Temperature

is explained microscopically as a direct measure of the average kinetic energy of the gas particles: a hotter gas

has, on average, faster-moving particles.

A gas that obeys every one of these postulates exactly is called an ideal gas — a useful theoretical …