Forget Kinetic Theory

The kinetic theory of gases describes a gas as a large number of small particles (atoms or molecules), all of which are in constant, random motion. The rapidly moving particles constantly collide with each other and with the walls of the container. Kinetic theory explains macroscopic properties of gases, such as pressure, temperature, viscosity, thermal conductivity, and volume, by considering their molecular composition and motion. This theory considers that the number of molecules contained in an ideal gas sample is very large, and motion in all directions is identical. As such the average velocities of all gas molecules at an instant is zero.

However if the number of molecules is small, the case is not so. Consider a gas sample with only five molecules of an ideal gas, with speeds $$1m/s$$, $$2m/s$$, $$3m/s$$, $$4m/s$$ and $$5m/s$$. Then, the magnitude of average velocity of particles at this instant is

$$Image Credit : Wikipedia$$

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