MCQ Detailed View

The Kinetic Molecular Theory (KMT) explains the behavior of gases based on the motion of their molecules. One of its postulates states that gas molecules move continuously in straight-line paths until they collide with other molecules or the walls of the container. These collisions are perfectly elastic, meaning there is no loss of total kinetic energy, although energy may be redistributed among colliding molecules.

In an ideal gas, molecules are assumed to have negligible volume compared to the space they occupy, and intermolecular forces are considered absent. Because of this, gas molecules travel freely and independently in straight lines. When two molecules collide, the direction of motion changes, but between collisions, each molecule continues moving straight.

The path of a gas molecule between two successive collisions is known as the mean free path. In real gases, intermolecular forces exist, and high pressure or low temperature reduces the mean free path, leading to more frequent collisions. However, the fundamental assumption of the theory still holds that motion between collisions is straight.

The other options do not represent gas behavior in KMT. Molecules do not follow parabolic or hyperbolic paths because no continuous external force acts on them in a uniform way. They also do not move in a zig-zag manner on their own; the zig-zag effect is only the overall impression caused by frequent collisions, but between collisions the motion remains straight.

This principle helps explain gas pressure, diffusion, and effusion. The random straight-line motion of gas molecules and their elastic collisions account for the uniform distribution of gases in a container, as well as their ability to expand and fill any available volume.