Question

In the kinetic theory of gases, we try to relate macroscopic properties of a gas, like pressure, volume and temperature to its microscopic properties (properties of each atom or molecule) like particle speeds, momenta, kinetic energies, etc. Which of the following is/are incorrect assumption(s) for building the theory?

• A. Molecules will experience inter-molecular forces of electrostatic attraction or repulsion, because of valence electrons and nuclear protons
• B. A collision between two molecules, or a molecule and a wall preserves net kinetic energy and momentum
• C. At equilibrium (i.e., after the system has been left undisturbed for a long time), the number of molecules inside a particular $1c{m}^3$ will be the same for any other $1c{m}^3$, inside the containing volume.
• D. Temperature is a property of the average internal energy of all the molecules, not of one molecule.

Answer 1

The correct option is A

Molecules will experience inter-molecular forces of electrostatic attraction or repulsion, because of valence electrons and nuclear protons

The kinetic theory treats gases like a boxful of large number of restless atoms or molecules, which are bouncing off each other elastically, and continuously. It is a simplified picture, and produces correct predictions for only gases of simple molecules - monoatomic molecules that do not interact among each other in the gas medium, collide elastically (hence preserve kinetic energy), and have a homogenous (uniform) and isotropic (in no special direction) velocity distribution. The average kinetic energy of such a gas dictates the temperature of the gas. Kinteic theory does not account for inter-molecular interactions, however small or big. Thus, option (a) is not an assumption of the theory.