The Kinetic Molecular Theory: Assuming Things about Gases

A theory is useful to scientists if it describes the physical system they’re examining and allows them to predict what will happen if they change some variable. The Kinetic Molecular Theory of Gases does just that. It has limitations (all theories do), but it’s one of the most useful theories in chemistry. Here are the theory’s basic postulates — assumptions, hypotheses, axioms (pick your favorite word) you can accept as being consistent with what you observe in nature:

✓ Postulate 1: Gases are composed of tiny particles, either atoms or molecules. Unless you’re discussing matter at really high temperatures, the particles referred to as gases tend to be relatively small. The more-massive particles clump together to form liquids or even solids, so gas particles are normally small with relatively low atomic and molecular weights.

✓ Postulate 2: The gas particles are so small when compared to the distances between them that the volume the gas particles themselves take up is negligible and is assumed to be zero. Individual gas particles do take up some volume — that’s one of the properties of matter.

But if the gas particles are small (which they are), and there aren’t many of them in a container, you say that their volume is negligible when compared to the volume of the container or the space between the gas particles.

Sure, they have a volume, but it’s so small that it’s insignificant (just like a dollar on the street doesn’t represent much at all to a multimillionaire; it may as well be a piece of scrap paper).

This explains why gases are compressible. There’s a lot of space between the gas particles, so you can squeeze them together. This isn’t true in solids and liquids, where the particles are much closer together.

✓ Postulate 3: The gas particles are in constant random motion, moving in straight lines and colliding with the inside walls of the container. The gas particles are always moving in a straight-line motion. They continue to move in these straight lines until they collide with something, either each other or the inside walls of the container.

The particles also all move in different directions, so the collisions with the inside walls of the container tend to be uniform over the entire inside surface. A balloon, for instance, is relatively spherical because the gas particles are hitting all points of the inside walls the same. The collision of the gas particles with the inside walls of the container is called pressure.

This postulate explains why gases uniformly mix if you put them in the same container. It also explains why, when you drop a bottle of cheap perfume at one end of the room, the people at the other end of the room are able to smell it right away.

✓ Postulate 4: The gas particles are assumed to have negligible attractive or repulsive forces between each other. In other words, you assume that the gas particles are totally independent, neither attracting nor repelling each other. That said, if this assumption were correct, chemists would never be able to liquefy a gas, which they can. However, the attractive and repulsive forces are generally so small that you can safely ignore them. The assumption is most valid for nonpolar gases, such as hydrogen and nitrogen, because the attractive forces involved are London forces, weak forces that have to do with the ebb and flow of the electron orbitals. However, if the gas molecules are polar, as in water and HCl, this assumption can become a problem, because the forces are stronger. (Turn to Chapter 6 for the scoop on London forces and polar things — all related to the attraction between molecules.)

✓ Postulate 5: The gas particles may collide with each other. These collisions are assumed to be elastic, with the total amount of kinetic energy of the two gas particles remaining the same. When gas particles hit each other, no kinetic energy — energy of motion — is lost.

That is, the type of energy doesn’t change; the particles still use all that energy for movement. However, kinetic energy may be transferred from one gas particle to the other. For example, imagine two gas particles — one moving fast and the other moving slow — colliding. The one that’s moving slow bounces off the faster particle and moves away at a greater speed than before, and the one that’s moving fast bounces off the slower particle and moves away at a slower speed. But the sum of their kinetic energy remains the same.

✓ Postulate 6: The Kelvin temperature is directly proportional to the average kinetic energy of the gas particles.

The gas particles aren’t all moving with the same amount of kinetic energy. A few are moving relatively slow, and a few are moving very fast, but most are somewhere in between these two extremes. Temperature, particularly as measured using the Kelvin temperature scale, is directly related to the average kinetic energy of the gas. If you heat the gas so that the Kelvin temperature (K) increases, the average kinetic energy of the gas also increases.

A gas that obeys all the postulates of the Kinetic Molecular Theory is called an ideal gas. Obviously, no real gas obeys the assumptions made in the second and fourth postulates exactly. But a nonpolar gas at high temperatures and low pressure (concentration) approaches ideal gas behavior.