Kinetic Theory Of Gases

Kinetic theory of gases explains the three macroscopic properties of a gas in terms of the microscopic nature of atoms and molecules making up the gas. Usually, physical properties of solids and liquids can be described by their, size, shape, mass, volume etc. However, when we talk about gases, they have no definite shape, size while mass and volume are not directly measurable. The Kinetic theory of gases is useful and can be applied in this case.

With the help of the kinetic theory of gases, physical properties of any gas cane defined generally in terms of three measurable macroscopic properties of pressure, volume and temperature of the container where the gas is stored or present. We will learn about this concept in detail below.

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What is Kinetic Theory of Gases?

Kinetic theory of gases is a theoretical model that describes the molecular composition of the gas in terms of a large number of submicroscopic particles which include atoms and molecules. Further, the theory explains that gas pressure arises due to particles colliding with each other and the walls of the container. Kinetic theory of gases also defines properties such as temperature, viscosity, and thermal conductivity and all these properties are related to the microscopic phenomenon.

The significance of the theory is that it helps in developing a correlation between the macroscopic properties and the microscopic phenomenon. In simple terms the kinetic theory of gases also helps us study the action of the molecules. Generally, the molecules of gases are always in motion and they tend to collide with each other and the walls of the containers.

Kinetic Theory of Gases Assumptions

Kinetic theory of gases considers the atoms or molecules of a gas as a constantly moving point masses, with huge inter-particle distance and may undergo perfectly elastic collisions. Implications of these assumptions are –

i) Particles

Gas is a collection of a large number of atoms or molecules.

ii) Point Masses

Atoms or molecules making up the gas are very small particles like a point(dot) on a paper with a small mass.

iii) Negligible Volume Particles

Particles are generally far apart such that their inter-particle distance is much larger than the particle size and there is large free unoccupied space in the container. Compared to the volume of the container, the volume of the particle is negligible (zero volume).

iv) Nil Force of Interaction

Particles are independent. They do not have any (attractive or repulsive) interactions among them.

v) Particles in Motion

The particles are always in constant motion. Because of lack of interactions and the free space available, the particles randomly move in all directions but in a straight line.

vi) Volume of Gas

Because of motion, gas particles, occupy the total volume of the container whether it is small or big and hence the volume of the container to be treated as the volume of the gases.

vi) Mean Free Path

This is the average distance a particle travels to meet another particle.

vii) Kinetic Energy of the Particle

Since the particles are always in motion, they have average kinetic energy proportional to the temperature of the gas.

viii) Constancy of Energy / Momentum

Moving particles may collide with other particle or container. But the collisions are perfectly elastic. Collisions do not change the energy or momentum of the particle.

ix) Pressure of Gas

Collision of the particles on the walls of the container exerts force on the walls of the container. Force per unit area is the pressure. The pressure of the gas is thus proportional to the number of particles colliding (frequency of collisions) in unit time per unit area on the wall of the container.

Kinetic Theory of Gases Postulates

The kinetic theory of gas postulates are useful in the understanding of the macroscopic properties from the microscopic properties.

A) Understanding Gas Laws of Ideal Gases:

i) Pressure α Amount or Number of Particles at Constant Volume:

Collision of the particles on the walls of the container creates pressure. Larger the number of the particle (amount) of the gas, more the number of particles colliding with the walls of the container.

At constant temperature and volume, larger the amount (or the number of particles) of the gas, higher the pressure.

ii) Avogadro’s Law – N α V at Constant Pressure:

When there is a greater number of particles it increases the collisions and the pressure. If the pressure is to remain constant, the number of collisions can be reduced only by increasing the volume.

At constant pressure, the volume is proportional to the amount of gas.

ii) Boyle’s Law – Pressure \(\alpha \frac{1}{v}\) at Constant Temperature:

At a constant temperature, the kinetic energy of particles remains the same. If the volume is reduced to a constant temperature, then the number of particles in unit volume or area increases. If there is an increased number of particles in the unit area then it increases the frequency of collisions per unit area.

At constant temperature, smaller the volume of the container, larger the pressure.

ii) Amonton’s Law: P α T at Constant Volume:

The kinetic energy of the particle increases with temperature. When the volume is constant, with increased energy, particles move fast and increase the frequency of collisions per unit time on the walls of the container and hence the pressure.

At constant volume, the higher the temperature higher will be the pressure of the gas.

iv) Charles’s Law – V α T at Constant Pressure:

Change of temperature changes proportionately the pressure. If the pressure also has to remain constant, then the number of collisions has to be changed proportionately. At constant pressure and constant amount of substance, collisions can be changed only by changing the area or volume.

At constant pressure, volume changes proportionally to temperature.

v) Graham Law of Diffusion – \(v\,\,\alpha \sqrt{\frac{1}{M}}\)

Two gases with molecular weights M1 and M2 will have the same kinetic energy at the same temperature. Then,

\(\frac{1}{2}{{M}_{1}}v_{1}^{2}=\frac{1}{2}{{M}_{2}}v_{2}^{2}\) or \({{M}_{1}}v_{1}^{2}={{M}_{2}}v_{2}^{2}\) or \(\frac{v_{1}^{2}}{v_{2}^{2}}=\frac{{{M}_{2}}}{{{M}_{1}}}\) or \(\frac{{{v}_{1}}}{{{v}_{2}}}=\sqrt{\frac{{{M}_{2}}}{{{M}_{1}}}}\)

The velocity of the molecules are inversely proportional to their molecular weights.

Read more about the different laws by visiting the links given below;

B) Understanding Non-ideal Gas Behaviour:

All the gas molecules obey the ideal gas laws only under special conditions of low pressures and high temperatures. The deviations of the real gases, from the ideal gas behaviour, is traced to mainly to wrong or incorrect assumptions in the postulates.

They are,

  • The particles are point charges and have no volume. Then, it should be possible to compress the gases to zero volume. But, gases cannot be compressed to zero volume indicates that particles do have volume though small and cannot be neglected.
  • Particles are independent and do not interact. Particles do interact depending upon their nature. The interactions affect the pressure of the gas. Volume and the interactions differ from gas to gas. Many gas laws have been developed for the real gases incorporating correction factors in the pressure and volume of the gases.
  • Particular Collisions are not elastic. Particular collisions are not elastic and they exchange energy. The particles hence do not have the same energy and have a distribution of energy.

C) Maxwell – Boltzmann Molecular Distribution of Energy and Velocity:

The kinetic theory of gas postulates predicted the particles as always in motion and that they have kinetic energy proportional, to the temperature of the gas.

This concept was used by Maxwell – Boltzmann, to find the distribution of gaseous particles between energy zero to infinity and calculate the most probable, average and root mean square velocity of the particles.