Kinetic theory

In physical chemistry, the kinetic theory of gases is a theory that explains the macroscopic[?] properties of gases by consideration of their composition at a molecular level.

Table of contents 1 Postulates 2 Pressure 3 Temperature 4 External links

Postulates The fundamental principles of the kinetic theory are given in the form of several postulates:

• Gases are composed of molecules in constant random, motion. The moving particles constantly collide with each other and with the walls of the container.
• The collisions between gas molecules are elastic.
• The total volume of the gas molecules is neglible compared to the volume of the container.
• The forces of attraction between the molecules are negligible.

The above postulates accurately describe the behavior of ideal gases. Real gases[?] approach ideality under conditions of low density and high temperature.

Pressure Pressure is explained by the kinetic theory as arising from the force exerted by collisions of gas molecules with the walls of the container. The derivation of the mathematical expression for pressure is given below:

Consider a gas with N molecules, each of mass m, enclosed in a cuboidal container of volume V. Suppose that a gas molecule collides with a wall of the container which is perpendicular to the x co-ordinate axis. Then the momentum lost by the particle and gained by the wall is given by

<math>2mv_x</math>

where v_x is the x-component of the initial velocity of the particle.
Now, force is the rate of change of momentum. The particle under consideration impacts with the wall once every 2l/v_x time units, where l is the length of the container. Therefore the force due to this particle is

<math>mv_x *v_x \over l</math>

and the total force on the wall is

<math>m\sum_j v_{jx}^2 \over l</math>

where the summation is over all the gas molecules in the container. Since the particles are moving randomly in all directions, and since v² = v_x² + v_y² + v_z² for each particle, the expression for the total force becomes

<math>m\sum_j v_j^2 \over 3l</math>

This can be written as

<math>Nmv_{rms}^2 \over 3l</math>

where v_rms is the root mean square velocity of the gas. Therefore, pressure, the force per unit area, equals

<math>Nmv_{rms}^2 \over 3Al</math>

where A is the area of the wall. Thus, we have the following expression for the pressure

<math>P=Nmv_{rms}^2 \over 3V</math>

This result is interesting and significant because it relates pressure, a macroscopic property, to the average (translational) kinetic energy per molecule (1/2 mv_rms²), which is a microscopic property.

Temperature The above equation tells us that the product of pressure and volume is proportional to the average molecular kinetic energy. Further, the ideal gas equation tells us that this product is proportional to the absolute temperature. Putting the two together, we arrive at one important result of the kinetic theory: average molecular kinetic energy is proportional to the absolute temperature. The constant of proportionality is 3/2 times Boltzmann's constant, which is the ratio of the gas constant to Avogadro's number. This result is related to the equipartition theorem[?].

See also:

• Maxwell-Boltzmann distribution

External links

• http://www.phys.virginia.edu/CLASSES/252/kinetic_theory Kinetic theory of gases - A brief review (http://www.phys.virginia.edu/CLASSES/252/kinetic_theory)
• http://www.ucdsb.on.ca/tiss/stretton/chem1/gases9 Introduction to the kinetic molecular theory of gases (http://www.ucdsb.on.ca/tiss/stretton/chem1/gases9)