What is the Kinetic Theory of Gases?

The kinetic theory of gases is a simple, classical model of the thermodynamic behaviour of gases. It established many fundamental thermodynamic concepts and is also historically significant. 

As determined by the model, gas consists of many identical submicroscopic particles (atoms or molecules) moving at a steady, quick and random rate. They are assumed to be much smaller in size than the average particle distance. Random elastic collisions occur between the particles and the container’s enclosing walls. The basic version of the model describes the ideal gas and considers no other particle interactions.

Assumptions of the kinetic theory of gases

Kinetic theory of gases tends to consider atoms and molecules of any gas as consistently moving masses with huge interparticle distances and could undergo elastic collisions. Implications of these assumptions are as follows:

• Particles: Gas should be considered an extensive collection of atoms or molecules.
• Point masses: Atoms or molecules that form gas are microscopic particles.
• The volume of Particles is negligible: Particles are often far apart due to high inter-particle distance. Additionally, the inter-particle distance is comparatively larger than the size of the particle, due to which there is a large, free and unoccupied space in the container. In this regard, the volume of the particle is negligible compared to the volume of the container.
• Nil force of interaction: Particles are independent; thus, they neither have attractive interaction nor repulsive interaction with each other.
• Particles are constantly in motion: Particles are consistently in constant motion due to the lack of interactions and ample availability of free space. In this context, particles move randomly in all directions rather than flowing in a straight line.
• The volume of gas: Due to the constant motion, gas particles tend to occupy the total volume of the container irrespective of its size; hence, the volume of gas is considered equal to the volume of the container.
• Mean free path: The average distance that a particle travels for meeting another particle independently.
• The kinetic energy of particles: As the particles are in constant motion, they tend to have kinetic energy, which is proportional to the temperature of the gas.
• The constancy of the momentum and energy: Moving particles tend to collide with each other and the walls of the containers; however, the collisions are perfectly elastic. Due to this reason, collisions do not induce changes in the momentum and energy of the particles.
• The pressure of gas: The collision of particles in a container exerts a force on the walls of that container, and it is well established that force per unit of area is the pressure. In this regard, the pressure of the gas is directly proportional to the number of particles colliding in a specific unit of time and specific unit area.

Kinetic Theory of Ideal Gases

The atoms in an ideal gas do not exert forces on each other but instead clash with the container’s walls. Based on experiments, ideal gas law relates the pressure, temperature, volume, and the number of moles of an ideal gas:

PV = nRT,

In which R is a constant also known as the universal gas constant.

Comments:

1. Make sure that the same unit system is used to express all quantities!
2. The temperature, T, must be expressed in absolute degrees Kelvin.
3. n denotes the number of moles of the gas, which is defined as

n = mass of sample/Molecular mass of gas

Suspicions of Kinetic Theory of Gases 

Following are the active suspicions of the kinetic theory of gases:

• All gases are composed of particles continually and determinedly moving in weird ways.
• The partition between the atoms is a lot more prominent than the size of particles.
• The divider is viewed as tiny during the period between two particles and an atom.
• All impacts among atoms, particles, and dividers are considered flexible.
• Assuming a gas test is left for an adequate time frame, it goes to a consistent state in the long run. This is because the thickness of particles and the appropriation of atoms are free of position, distance, and time.

Understanding Non-ideal Gas Behaviour:

All the gas molecules obey the perfect gas laws only under special conditions of low pressures and high temperatures. The deviations of the important gases from the perfect gas behaviour are traced mainly to wrong or incorrect assumptions within the postulates.

• The particles are point charges and haven’t any volume: Then, it should be possible to compress the gases to zero volume. But, gases can’t be compressed to zero volume indicates that particles do have volume though small and can’t be neglected.
• Particles are independent and don’t interact: Particles do interact depending upon their nature. The interactions affect the pressure of the gas. Volume and, therefore, the interactions differ from gas to gas. Many gas laws are developed for the important gases incorporating correction factors within the pressure and volume of the gases.
• Particle collisions are not elastic: Particle collisions are elastic, and they exchange energy. Hence, the particles don’t have equivalent energy and have a distribution of energy.

Conclusion

Overall, the Kinetic theory of gases focuses on considering the gas molecules as particles so that their movement and volume can be calculated. As per this theory, the pressure on gas molecules increases when they collide with each other or the surface of a container. Using this theory, multiple properties of gases are evaluated and determined. The kinetic molecular theory of gases explained, in detail, how ideal gases behave as described by ideal gas laws.