The kinetic theory of the gases is a simple, historically significant classical model of the thermodynamic behaviour of gases that established many fundamental thermodynamic concepts. A gas is described by the model as a large number of identical submicroscopic particles (atoms or molecules) that are all in constant, rapid, random motion. Their size is assumed to be much smaller than the average particle distance. Random elastic collisions occur between the particles and with the enclosing walls of the container.
The basic version of the model describes the ideal gas and takes into account no other particle interactions. The kinetic theory of gases explains macroscopic properties of gases like volume, pressure, and temperature, as well as transport properties like viscosity, thermal conductivity, and mass diffusivity. The model also takes into account related phenomena like Brownian motion. Historically, the kinetic theory of gases was the first explicit application of statistical mechanics ideas.
The behaviour of gases is studied by taking into account either the small scale action of individual molecules or the large scale action of the gas as a whole. The large-scale action of the gas can be easily studied and measured; however, a theoretical model is required to study the action of the gas molecules. The kinetic theory of gases is the name given to this theoretical model. Let us look at the kinetic theory of gases and the assumptions that go into it in this article.
Kinetic theory of gases is a theory that is based on a simplified molecular or particle description of a gas and can derive many gross properties of the gas from this. It is a theory derived from the fact that particles in a gas move freely and rapidly along straight lines but frequently collide, resulting in changes in velocity and direction. The impacts of these particles on the walls of a container are also interpreted as causing pressure. The greater the density of a gas, the more collisions between molecules and the surface there will be, and the greater the pressure exerted.
The following assumptions are made when applying kinetic theory to ideal gases:
Thus, particle motion dynamics can be treated classically, and the equations of motion are time-reversible. As a simplifying assumption, the particles are usually assumed to have the same mass; however, the theory can be generalised to mass distribution, with each mass type contributing to the gas properties independently of one another, in accordance with Dalton’s Law of partial pressures.
Many of the model’s predictions are the same whether or not particle collisions are included, so they are frequently ignored as a simplifying assumption in derivations. These assumptions are relaxed in more recent developments, which are based on the Boltzmann equation. These can accurately explain the features of dense gases because they incorporate the volume of the particles as well as contributions from intermolecular and intramolecular interactions, quantized molecular rotations, quantum rotational-vibrational symmetry effects, and electronic excitation.
The following are the postulates of gas kinetic theory:
The three main components of the gas kinetic theory are as follows:
Kinetic theory explains the behaviour of gases by assuming that the gas is made up of rapidly moving atoms or molecules. This is possible because in gas, interatomic forces between molecules are ignored.
The macroscopic properties of gases, such as volume, pressure, and temperature, are explained by the kinetic theory of gases.
The three main components of gas kinetic theory are as follows:
1) There is no gain or loss of energy when molecules collide.
2) The amount of space occupied by gas molecules in a container is negligible.
3) These molecules always move in a straight line.
1) All gases are composed of molecules that move in random directions all the time.
2) All collisions between molecules, as well as collisions between molecules and walls, are considered elastic.
3) All molecules in a given gas sample follow Newton’s laws of motion.