Question

According to kinetic theory of gases, the collisions between molecules of a gas (a) occur in a zig-zag path (b) occur in a straight line (c) change velocity and energy (d) result in settling down of molecules.

Short answer

The correct answer is (c) change velocity and energy.

Step-by-step solution

Understanding Kinetic Theory of Gases

The Kinetic Theory of Gases describes the behavior of molecules in a gas. It is based on the concept that gas molecules are in constant, random motion and collide with each other and the walls of their container.

Analyzing the Options

Option (a) suggests a zig-zag path which refers to the overall movement of gas molecules after multiple collisions, not the individual collisions themselves. Option (b) indicates that collisions occur in a straight line, which is true for the movement between collisions. Option (c) states that collisions change velocity and energy, which is accurate as collisions can transfer energy between molecules, and result in a change in the direction and speed of the molecules. Option (d) implies that collisions cause molecules to settle down, this contradicts the theory that gas molecules are in perpetual motion.

Selecting the Correct Answer

Evaluate the options given in light of the Kinetic Theory of Gases. Only option (c) is consistent with the theory as it correctly describes the effect of collisions on the molecules of a gas.

Key concepts

Behavior of Gas Molecules

Under the umbrella of the Kinetic Theory of Gases, we delve into the intriguing world of microscopic participants: gas molecules. Their behavior is anything but stagnant; they're perpetual motion incarnate. These minuscule entities whiz around their container at rapid velocities, embodying the quintessence of freedom as no forces of attraction significantly impede their journey.

Yet, these molecular sprinters are not running a lonely marathon. Envision a bustling city: the molecules traverse similar paths with encounters at every turn. While these meetings might sound orderly, they're the complete opposite—there's no predicting where a molecule might jet off to next. This spontaneous itinerary is a product of countless, random collisions. Each collision, much like a cue ball strikes in a game of pool, results in a change in direction and possibly speed, varying the molecules’ velocity vectors in a fashion encompassing all possible directions.

Understanding this fundamental aspect of molecular behavior provides insight into gas properties such as pressure, which is fundamentally the collective push of these microscopic particles on their container's walls. It's a dance of dynamics, with every molecule contributing to the grand performance of phenomena we observe on a macroscopic scale.

Molecular Collisions

Now, let's narrow our focus to those pivotal instances—the molecular collisions. When gas molecules clash, it's akin to billiard balls colliding on a pool table; the impacts change their trajectories and can redistribute kinetic energy between them. These interactions are perfectly elastic, meaning no energy is lost to the surroundings; it's merely exchanged or transformed among the participants.

Imagine this in practice: when two molecules bump into each other, they may recoil with altered speeds and divergent paths. This constant reshuffling is responsible for the dispersion of energy within the gas. It's like a never-ending game where each collision can either speed or slow a molecule down, depending on the relative velocities and angles of impact. It's this aspect of molecular collisions that lays the foundation for the thermal equilibrium in a substance—as the faster molecules give up energy to the slower ones, a sort of 'thermal democracy' is achieved.

This concept also sheds light on why temperature is a measure of the average kinetic energy of the molecules in a substance. When you heat a gas, you accelerate these collisions—like turning up the tempo of our billiard ball dance—causing an increase in the average speed of the molecules and thus, the temperature.

Random Motion of Particles

The words 'random motion' evoke an image of chaos, and rightly so when it comes to gas molecules. This randomness is a pivotal cornerstone of the Kinetic Theory. The particles' paths are unpredictable; their motion is not defined by a clear pattern but rather a series of erratic twists and turns. Importantly, this random motion isn't just a side note—it's critical to understanding diffusion, where molecules spread from areas of higher concentration to lower concentration, scenting the air from a sprayed perfume.

But what enables this randomness? A mix of rapid speeds and frequent collisions ensures that gas molecules visit virtually every possible location in the available space, moving in straight lines between collisions only to be sent careening in new directions time and again. In a closed container, the randomness is also reflected in how gas pressure remains uniform throughout—the molecules are equally likely to strike any part of the walls.

In essence, the random motion encapsulates the constant, haphazard zig-zag journey gas molecules undergo—without ever settling down, unless cooled to a standstill. It's this frenetic character that underlies the familiar phenomena of pressure, temperature, and diffusion, painting a vivid portrait of the microscopic kinetic frenzy that propels our macroscopic world.