Kinetic theory of gases

The kinetic theory of gases is a simple classical model of the thermodynamic behavior of gases. It helps us understand how gases work by treating them as made up of tiny particles—like atoms or molecules—that are always moving around randomly. These particles are so small that we can’t see them, even with a microscope!

This theory explains how things like the volume, pressure, and temperature of a gas are related to each other. It does this by looking at how these tiny particles bump into each other and the walls of the container they’re in. Even though we can’t see these particles, their movement and collisions help us understand a lot about gases.

The basic version of this theory describes an ideal gas. In this model, the particles don’t lose any energy when they collide—they bounce off each other perfectly, like rubber balls. The particles are also assumed to be much smaller than the space between them, so they don’t get in each other’s way very much.

Because of these ideas, the kinetic theory of gases became important in the development of statistical mechanics. It was one of the first clear examples of using statistics to understand physics, and it helps explain the behavior of gases that are not very crowded. This theory gives us a strong foundation for studying how gases move and spread out.

History

See also: Heat § History, Atomism, and History of thermodynamics

Long ago, around 50 BCE, the Roman philosopher Lucretius thought that everything solid is made of tiny, fast-moving particles. This idea was not popular for many years.

Later, in the 1600s, scientists began to connect heat with movement. Francis Bacon said that heat is the motion of tiny parts of matter. Galileo Galilei agreed, saying that heat comes from particles moving inside objects.

In 1738, Daniel Bernoulli published a book called Hydrodynamica. He suggested that gases are made of many tiny molecules moving in all directions. He said that when these molecules hit a surface, they create gas pressure.

Later scientists like James Clerk Maxwell and Ludwig Boltzmann developed these ideas further. Their work helped explain many things about gases and heat.

Assumptions

The kinetic theory of gases uses a few simple ideas to explain how gases act. It says that gases are made of very tiny particles that are far apart, so their size doesn’t matter much. These particles move fast and bounce off each other and the walls of their container without losing energy — these bounces are called elastic collisions.

We also assume there are so many particles that we can use statistics to guess how they will behave. Between bounces, the particles don’t push or pull on each other. This makes it easier to use basic physics to describe their movement. Sometimes we also assume all particles have the same mass, but the theory can work for particles with different masses too.

Equilibrium properties

The kinetic theory of gases explains how gases act by thinking about tiny particles that are always moving.

In this theory, gas pressure happens when tiny particles hit the walls of the container they are in. These particles move around randomly and bounce off each other and the walls. When they hit the walls, they push on them. This push is what we feel as pressure.

The theory also connects temperature to how fast these particles move. When the temperature is higher, the particles move faster. This affects the pressure and energy of the gas. This helps us understand how gases behave in different situations.

Transport properties

See also: Transport phenomena

The kinetic theory of gases looks at gases not just when they are still, but also when they are moving. This helps us learn about "transport properties" like viscosity, thermal conductivity, and mass diffusivity.

Basic kinetic theory works best for very thin gases. For thicker gases, a special method called Revised Enskog Theory was made in the 1980s by E. G. D. Cohen, J. M. Kincaid, and M. Lòpez de Haro. They built on older ideas from H. van Beijeren and M. H. Ernst.