Assumptions of the Kinetic Theory of Gases

Kinetic theory does not attempt to track the exact motion of every individual molecule in a gas -- with roughly $10^{23}$ molecules in even a small sample, this would be both impossible and pointless. Instead, it makes a handful of simplifying, idealised assumptions about how molecules behave on average, and then derives the gas's bulk properties purely by averaging over an enormous number of them. For an IDEAL gas, these assumptions are:

1. A gas consists of a very large number of identical molecules, in a state of ceaseless, random motion, moving in every direction with a wide range of speeds. This random motion is what is meant by the thermal motion of a gas.

2. The size of a molecule is negligible compared to the average distance between molecules. A molecule is treated, for most purposes, as a point mass occupying essentially no volume of its own -- the actual volume occupied by the gas is overwhelmingly empty space between widely separated molecules.

3. Molecules exert no force on one another except during a collision. Between collisions, a molecule is assumed to move in a straight line at constant velocity (Newton's first law), completely unaffected by any other molecule, however close it might be -- there is no long-range attraction or repulsion assumed in the ideal-gas model.

4. Collisions -- between two molecules, or between a molecule and the container wall -- are perfectly elastic, and of negligible duration compared to the time spent travelling freely between collisions. Kinetic energy is exchanged during a collision but never permanently lost to any internal, non-mechanical process.

5. Molecules obey Newton's laws of motion at every instant, including during the brief instant of a collision itself.

6. Gravity is neglected for the molecules' own motion -- their kinetic energies are assumed to be far too large, and the molecules themselves far too light, for the weak pull of gravity on an individual molecule to matter over the short distances between collisions.

These assumptions together are what gives an ideal gas its name: a mathematically idealised gas that no real gas matches perfectly, but that most real gases (particularly at low pressure and moderately high temperature, where molecules genuinely are far apart) approximate closely enough for the theory's predictions to match experiment very well. …