Kinetic Theory of Gases

What is kinetic theory of gases

Kinetic theory of gases is a physically-significant explanation for why gases are usually denser than air. The speed of an arrow fired from the muzzle of an air-powered gun is important, and can be represented as an arrow in a curve of temperature vs. density. Although other factors may change the trajectory, most air resistance causes the arrow to deviate to one side from the arrow's original trajectory and towards the force of gravity.

The total energy of the projectile and the air is generally not appreciably different, so the difference is due to density and entropy changes in the two media, respectively. It is also important to note that arrows were not historically used in artillery because they were believed to have too short a flight time for accurate volleys.

The first and second laws of thermodynamics

The second law of thermodynamics—energy conservation—is the basis for understanding that there is no net energy gain, that is, no energy is lost in a gas. This law states that for a given state of matter, the amount of matter and energy in an isolated system remains constant.

This law is an approximation, which was not accepted for many years by physicists. A clear understanding of the second law came only with the development of quantum mechanics, which led to an understanding of the uncertainty principle.

In physics, the law of thermodynamics is a fundamental law of thermodynamics, where the specific entropy of a gas is defined as the logarithm of the pressure and temperature of the gas. The name thermodynamic state of matter (SOM) is also defined in terms of pressure and temperature.

Applications of kinetic theory to gas behavior

The modern understanding of thermodynamics was set by Jakob Lelius of Copenhagen in the 1770s. At the time, he showed that a reversible process of exchanging heat between gas molecules was impossible and, in a sense, irrelevant.

Today, in the form of the Second Law of Thermodynamics, we have reason to believe that irreversible processes have physical implications.

For example, if an object of one volume is placed in a volume of a different volume at a given temperature, the volume can be emptied or warmed up, but the temperature cannot. Kinetic theory explains how one can study gas behavior. With it, we can determine that the basic behavior of gases is also determined by temperature and pressure, rather than by volume.

Kinetic Theory in Everyday Life

Like the familiar passage about how steam creates white noise and muffle the sound of a steam whistle, many everyday phenomena are "mechanical," in the sense that their mechanisms are mechanical in origin. The reaction of a chemical reaction can be read off a gas chromatograph or a mass spectrometer, the energy required to lift a water molecule off the surface of an atom can be measured, and the motion of a spinning barberpole is kinetic, as it applies rotating momentum to the axis of rotation.

This kinetic theory of the common phenomena of thermodynamics underpins the study of low temperature physics, as well as of several other physical processes. And that's all because a liquid of a gas starts out like an extremely viscous liquid and becomes less viscous as it is cooled.

Conclusion

In this article, I am going to discuss the importance of the kinetic theory of gases in a much more limited context. This small step helps to reinforce the valuable point that the kinetic theory of gases is indeed a fundamental topic in classical thermodynamics.

Before I proceed, let me say the following. Yes, I know, there are various other classical thermodynamic principles that are of no particular importance. For example, there is the second law of thermodynamics, which suggests that the entropy of a system will grow with time.

The law is a simplification, or a more obvious statement than the kinetic theory of gases. It also suggests that a system never reaches maximum entropy and remains in a state of maximum entropy at all times. However, it doesn't imply anything about equilibrium.