Kinetic theory facts for kids
The kinetic theory of gases is a model of the thermodynamic behavior of gases. It attempts to explain overall properties of gases, such as pressure, temperature, or volume, by considering their molecular composition and motion. The theory basically states that pressure is not caused by molecules pushing each other away, like earlier scientists thought. Instead, pressure is caused by the molecules colliding with each other and their container. Kinetic theory is also known as kinetic-molecular theory or collision theory. The kinetic theory of gases helped establish many principal concepts of thermodynamics.
There are three main components to kinetic theory:
- No energy is gained or lost when molecules collide
- The molecules in a gas take up a negligible (able to be ignored) amount of space in relation to the container they occupy
- The molecules are in constant, linear motion.
History
In about 50 BCE, the Roman philosopher Lucretius proposed that apparently static macroscopic bodies were composed on a small scale of rapidly moving atoms all bouncing off each other. This point of view was rarely considered in the subsequent centuries, when Aristotlean ideas were dominant.
In 1738 Daniel Bernoulli published Hydrodynamica, which laid the basis for the kinetic theory of gases. In this work, Bernoulli posited the argument, that gases consist of great numbers of molecules moving in all directions, that their impact on a surface causes the pressure of the gas, and that their average kinetic energy determines the temperature of the gas. The theory was not immediately accepted.
Other pioneers of the kinetic theory, whose work was also largely neglected by their contemporaries, were Mikhail Lomonosov (1747), Georges-Louis Le Sage (ca. 1780, published 1818), John Herapath (1816) and John James Waterston (1843). In 1856 August Krönig created a simple gas-kinetic model. In 1857 Rudolf Clausius developed a similar, but more sophisticated version of the theory. In this same work he introduced the concept of mean free path of a particle.
In 1859, after reading a paper about the diffusion of molecules by Clausius, Scottish physicist James Clerk Maxwell formulated the Maxwell distribution of molecular velocities. This was the first-ever statistical law in physics.
In 1871, Ludwig Boltzmann generalized Maxwell's achievement and formulated the Maxwell–Boltzmann distribution.
At the beginning of the 20th century, atoms were considered by many physicists to be purely hypothetical constructs, rather than real objects. An important turning point was Albert Einstein's (1905) and Marian Smoluchowski's (1906) papers on Brownian motion, which succeeded in making certain accurate quantitative predictions based on the kinetic theory.
Following the development of the Boltzmann equation, a framework for its use in developing transport equations was developed independently by David Enskog and Sydney Chapman in 1917 and 1916.
Assumptions
The application of kinetic theory to ideal gases makes the following assumptions:
- The gas consists of very small particles. This smallness of their size is such that the sum of the volume of the individual gas molecules is negligible compared to the volume of the container of the gas. This is equivalent to stating that the average distance separating the gas particles is large compared to their size, and that the elapsed time of a collision between particles and the container's wall is negligible when compared to the time between successive collisions.
- The number of particles is so large that a statistical treatment of the problem is well justified. This assumption is sometimes referred to as the thermodynamic limit.
- The rapidly moving particles constantly collide among themselves and with the walls of the container. All these collisions are perfectly elastic, which means the molecules are perfect hard spheres.
- Except during collisions, the interactions among molecules are negligible. They exert no other forces on one another.
