Gases, as well as mixture of gases, behave as seen above, with regard to changes in temperature, pressure or both. The following assumptions can be made on the basis of the kinetic theory of gases:
- A molecule of a gas means the smallest unit of a gas that can exist independently. It may consist of one atom, as in the case of noble gases; two in the case of oxygen, hydrogen, carbon monoxide, nitric oxide etc; three in the case of ozone, carbon dioxide, sulphur dioxide etc; four in the case of ammonia phosgene, phosphine (PH3) etc; and five in the case of methane, and so on.
A gas is made up of molecules which are far apart from each other in space. The total volume of molecules is negligible as compared to the total volume of the gas i.e., the total space in which the molecules move (Fig. 15.3 (a & b)).
- Every molecule in a gas is in a state of constant straight-line motion. Molecules very frequently collide with one another, and against the sides of the container (Fig. 15.4). It is assumed that the molecules do not lose energy during these collisions. Hence the sum total amount of the kinetic energy of the gas is not affected because of these collisions.
- Since the molecules are far apart they have a tremendous velocity. Hence the effect of intermolecular attractions becomes negligible.
- The average kinetic energy of a gas depends on the absolute (Kelvin) temperature. Higher the absolute temperature, greater is the average kinetic energy.
- At a given temperature and pressure all gases have the same kinetic energy.
