Theory of Gases

It is possible to explain the three macroscopic properties of a gas in terms of the microscopic nature of the atoms and molecules that make up the gas by using the kinetic theory of gases. Most of the time, the physical properties of solids and liquids can be described in terms of their size and shape as well as their mass and volume, among other things. When it comes to gases, on the other hand, there is no definitive shape or size, and neither mass nor volume can be measured in any meaningful way. The Kinetic Theory of Gases is useful in this situation and can be used to solve it.

Thanks to the kinetic theory of gases, the physical properties of any gas can be defined in terms of three measurable macroscopic properties, which can be expressed as a generalisation. The pressure, volume, and temperature of the container in which the gas is being stored or present are all important considerations. This concept will be discussed in greater depth later on.

What is Kinetic Theory of Gases and How Does It Work?

The kinetic theory of gases is a theoretical model that describes the molecular composition of a gas in terms of a large number of submicroscopic particles, which include atoms and molecules. It is a branch of physics that was developed in the 1960s. Furthermore, the theory explains that gas pressure is created as a result of particles colliding with one another and the container’s walls. As well as properties such as temperature, volume, and pressure, kinetic theory of gases defines transport properties such as viscosity and thermal conductivity as well as mass diffusivity. Fundamentally, it explains all of the characteristics that are associated with the microscopic phenomenon.

Because it assists in developing a correlation between macroscopic properties and microscopic phenomena, the theory has significant practical value. In layman’s terms, the kinetic theory of gases also aids in the investigation of the action of molecules on one another. In general, the molecules of gases are constantly in motion, and they have a tendency to collide with one another as well as with the walls of containers. The model also aids in the understanding of related phenomena such as Brownian motion, which are discussed further below.

Kinetic Theory of Gases Assumptions

Molecular kinetic theory of gases considers individual atoms or molecules in a gas as constantly moving point masses separated by vast distances between them and capable of undergoing perfectly elastic collisions. The following are the ramifications of these assumptions:

i) Particles

A gas is a large number of atoms or molecules gathered together in a small amount of room.

ii) Point Masses

The atoms or molecules that make up the gas are extremely small particles with a very small mass, similar to a point (dot) on a piece of paper.

iii) Negligible Volume Particles

Particles are typically spaced so far apart that the inter-particle distance between them is much greater than the particle size, and there is a significant amount of free unoccupied space in the container. The volume of the particle is insignificant when compared to the volume of the container (zero volume).

iv) Nil Force of Interaction

Particles are completely independent of one another. These individuals do not have any (attractive or repulsive) interactions with one another.

v) Particles in Motion

The particles are always in motion, no matter what. A lack of interactions and the availability of free space cause the particles to move randomly in all directions except in a straight line, as shown in the diagram.

1. vi) Volume of Gas

In order to accommodate the motion of gas particles, regardless of how small or large the container is, the total volume of the container must be considered when calculating the volume of gases contained within.

vi) Mean Free Path

When a particle meets another particle, this is the average distance travelled by the particle.

vii) Kinetic Energy of the Particle

Because the particles are constantly in motion, they have an average kinetic energy that is proportional to the temperature of the gas in which they are contained.

viii) Constancy of Energy / Momentum

It is possible for moving particles to collide with other moving particles or containers. However, the collisions are perfectly elastic in nature. Collisions have no effect on the energy or momentum of a particle involved.

ix)Pressure of Gas

An upward force is exerted on the container’s walls by the collision of the particles against the walls of the container. Pressure is defined as force per unit area. Accordingly, the pressure of gas is proportional to the number of particles colliding (the frequency of collisions) in a unit of time per unit area on the container’s wall.

Kinetic Theory of Gases Postulates

The postulates of the kinetic theory of gases are useful in deducing the macroscopic properties of a gas from its microscopic properties, and vice versa.

• Gases are made up of a large number of tiny particles that form a gaseous state (atoms and molecules). When compared to the distance between the particles, the size of these particles is extremely small. The individual particle’s size is considered insignificant, and the vast majority of the volume occupied by the gas is empty space. These molecules are in a state of constant random motion, which causes them to collide with one another and with the walls of the container as they move around. When gas molecules collide with the walls of a container, the molecules impart a small amount of momentum to the walls of the container. In essence, this results in the generation of a force that can be measured. Consequently, the pressure is defined as the product of dividing the force by the area.
• Collisions between molecules and walls are perfectly elastic because the molecules and walls are made of elastic material. This means that when the molecules collide, they don’t lose any kinetic energy in the process. Molecules never slow down and will continue to move at their current speed.
• The average kinetic energy of the gas particles varies as a function of the ambient temperature. In other words, the higher the temperature, the greater the average kinetic energy of the gas.   Except during collisions, the molecules do not exert any force of attraction or repulsion on one another.

Conclusion:

The kinetic theory of gases describes the relationship between the macroscopic properties of a gas, such as temperature, pressure, and volume, and the microscopic properties of a gas, such as speed, momentum, and position. atoms and molecules are constantly moving at random speeds, colliding with one another and the walls of the container that contains the gas. It is because of this motion that physical properties such as heat and pressure are produced. For the purposes of this article, let us look into the kinetic theory of gases in greater depth.