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The randomness in gases is due to the flow of molecules in the space in which the gas has been kept and due to which we see the randomness in other words we call it disordered behaviour. This randomness is due to the motion in the molecules of gases which comes from the molecular kinetic energy.

The kinetic theory of gases was a theory that was developed in various stages over the past few centuries. Some date its origins to the time of the ancient Greeks. Since then, a number of scientists have made contributions to the theory. In this explanation, we will introduce the kinetic theory of gases and describe and explain its key principles and assumptions.

## The kinetic theory of gases: Definition

The kinetic theory of gases explains the macroscopic properties of gas — such as its pressure and its temperature—in terms of the behavior of the molecules that make it up. In physics, we study the behavior of gases, which are governed by a system of equations called gas laws.

The **kinetic theory of gases** is a physical theory of the thermodynamic behavior of gases.

## Kinetic theory of gases principles:

To understand the properties of a gas we need to have a standard method of measuring the volume of the gases to do so we would like to introduce the concept of the mole. A mole is one of the seven SI base units of the physical measuring system.

One **mole** is the **number of atoms** in a 12 g sample of** carbon-12**.

Now we have another important question: “How many **atoms** or molecules are there in a **mole**?”

The answer to this question has been determined experimentally, which has a value of

$$N_A\;=\;6.02\;\times\;10^{23}\;\mathrm{mol^{-1}},$$

where** mol ^{−1} **represents the inverse mole or

**“per mole,”**and

**mol**is the abbreviation for

**mole**.

This number \(N_A\)** **has been named **Avogadro’s number **after Italian scientist **Amedeo Avogadro (1776–1856), **who suggested that all **gases **occupy the **same volume **under the **same conditions **of **temperature **and **pressure **and contain the **same number **of **atoms or molecules.**

Now from Avogadro’s number, we can determine that the number of **moles **\(n\)* *which constitutes any gas is equal to the

**ratio**of the

**number of molecules**\(N\) in the sample to the number of

**molecules**\(N_A\)

**in 1 mol**as follows:

$$n=\frac{N}{N_A}.$$

It has been observed that if we put **1 mol** of **different gases** in containers of **identical volume **and at the **same temperature their** **pressure **will also be **equal, **thus gases at lower densities follow the **ideal gas equation or the perfect gas law.**

$$pV\;={\;nRT}.$$

Where p is the absolute (not gauge) pressure, n is the number of moles of gas present, and T is the temperature in kelvins. The symbol R is a constant called the universal gas constant that has the same value for all gases.

The value is the gas constant \(R\;=\;8.314\mathrm{J/(mol K)}\).

It is assumed and observed that at** lower density** this law holds true for** any specific gas** because under this condition **real gases** tend to behave like **ideal gases**. This gives us another important equation that has been derived from equation 1 and can be written in terms of** Boltzmann's constant **\(k_B\), which is defined as:

$$\\\begin{array}{rcl}k\;=\frac R{N_A}\\k&=&\frac{8.31{\frac{\mathrm J}{\mathrm{mol}\;\mathrm K}}\;}{6.023\;\mathrm x\;10^{23}\;\frac1{\mathrm{mol}}}\\k&=&1.38\;\mathrm x\;10^{-23}\;\;{\frac{\mathrm J}{\mathrm K}}\end{array}$$

Boltzmann constant (*k*) = universal gas constant (R) divided by Avogadro's number (N_{A} ) gives us the value of the Boltzmann constant.

The equation above enables us to arrive at the following equation,

$$R\;=\;kN_A.$$

With the equation

$$n=\frac{N}{{N}_A},$$

we can write the following equation,

$$nR\;=\;Nk.$$

Substituting the above equation into the ideal gas equation we get,

$$pV\;=\;NkT.$$

Note the difference between the two expressions shown above, for the ideal gas law equation or the perfect gas law, involves the number of moles \(n\), and the equation above involves the number of molecules \(N\).)

## Kinetic theory of gases assumptions:

Just like any physical theory, there are a number of assumptions that underlie the kinetic theory of gases which must be understood in order to get a comprehensive view of the kinetic theory of gases. Here are the assumptions of the kinetic theory of gases:

- The kinetic theory of gases models gases as consisting of an ensemble of a large number of very small particles. The gas molecules are assumed to be so small that the sum of the volumes of the particles is negligible compared to the volume that the gas as a whole occupies.
- The kinetic theory of gases is a statistical description of gases. We assume that we are justified in using a statistical treatment because the number of particles is so large. We often refer to this assumption as the thermodynamic limit.
- The particles that make up the gas are moving very quickly and are constantly bumping into each other. We say that these collisions are perfectly elastic which means that the total amount of kinetic energy of the particles before and after each collision is conserved and the gas particles can be considered to be rigid spheres.
- Interactions between the gas particles are negligible. The only forces experienced by the particles are due to their collisions with one another.

If we want to model more specific gas systems we can apply more specific assumptions such as keeping certain state variables of the gaseous system constant.

**At constant temperature (Isothermal):**

Suppose also that we allow the gas to expand from an initial volume** V**** _{i}** to a final volume

**V**

_{f }while we keep the temperature

**T**of the gas constant. At a constant temperature, such a process is called an

**isothermal expansion**(and the reverse is called an

**isothermal compression**). The equation shown below represents an ideal gas in an isothermal process.

$$W\;=\;nRT\;ln(\frac{V_f}{V_i})$$

The symbol **\(\ln\)** denotes the natural logarithm, having a base **\(e\).**

**At constant volume and constant pressure (isobaric):**

Work done by gas can be a constant-volume process and a constant-pressure process. If the volume of the gas is constant, then,

$$W\;=\;p\triangle V.$$

In a constant volume scenario, the change in volume is 0 hence, the work done is zero:

$$W\;=\;0$$

(at constant volume). So the work done is zero. The volume changes while the pressure \(P\) of the gas is held constant,

$$\begin{array}{rcl}W&=&p(V_f-\;V_i\;)\;\\W&=&p\triangle V\end{array}$$

## Kinetic theory of gases: formula

Here we provide a comprehensive list of the formulae that you might need to know for your studies of the topic of the kinetic theory of gases.

$$n\;=\;\frac N{N_A}$$

$$pV\;=nRT$$

$$R\;=\;kN_A$$

$$nR\;=\;Nk$$

$$pV\;=\;NkT$$

$$W\;=\;nRT\;ln(\frac{V_f}{V_i})$$

$$W\;=\;p\triangle V$$

Where \(N\) is the number of gas particles, \(P\) is the pressure of the gas, \(V\) is the volume of the gas, \(T\) is the temperature of the gas, \(n\) is the number of moles of gas, \(\Delta\;V\) is the change in volume of the gas, \(N_A\) is Avagadro's number, \(R\) is the gas constant and \(k\) is Boltzmann's constant.

## The Kinetic Theory of Gases – Key takeaways

- Gas is the third state of matter which has the capability to flow.
- The behavior of gases is governed and explained by the kinetic theory of gases, which takes into consideration the temperature pressure, and volume.
- The number of molecules or atoms present in one mole of any gas is explained by Avogadro’s number \(N_A\), which has a value of
**6.023×10**^{23}. - The constant \(R\)
**8.314 J/(mol K).** - There is no such thing as an ideal gas, it's an assumption that is taken to simplify and compare our analysis.
- Another constant is the
**Boltzmann constant**\(k_B\)**1.38×10**^{-23}J/K. - Work done by any gas is equal to the pressure times the change in volume, and it can be measured in two different situations which are at constant volume and at a constant temperature.

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##### Frequently Asked Questions about Kinetic Theory of Gases

What is the kinetic theory of gases?

The kinetic theory of gases is an important historic theory of the macroscopic properties of gases, and how these gases change under changes in the state variables that describe gases including volume, pressure and temperature. The kinetic theory of gases models gases as consisting of a large ensemble of identical rigid spheres that occupy a volume that is negligible compared to the volume that the gas is occupying as a whole.

What does kinetic theory of gases explain?

The kinetic theory of gases explains how gases change as their macroscopic properties such as volume, temperature and pressure change. The kinetic theory of gases explains that gases consist of a large number of small fast-moving particles.

What are the main assumptions of the kinetic theory of gases?

The main assumptions of the kinetic theory if gases are the following:

- The volume of the particles of a gas are negligible compared to the volume that the gas occupies as a whole.
- The collisions between gas particles are perfectly elastic, meaning that the total kinetic energy of the colliding balls before the collision is equal to the total kinetic energy of the gas particles after the collision. The spheres act as rigid spheres.the gas particles are the
- The only forces that act between the gas particles are due to the collisions between the gas particles. All other interactions between gas particles are negligible.
- Gases are made up of a very large number of particles such that gases can be treated statistically and the statistical properties of gases can be determined. This is known as the thermodynamic limit.

What are the main principles of the kinetic theory of gases?

The main principle of the kinetic theory of gases can be summarized as a simple description of gases as consisting of a large ensemble of fast-moving elastically colliding rigid spheres of negligible volume compared to the volume occupied by the gas as whole, which frequently collide with one another, exchanging momentum and kinetic energy as they do.

What is the equation of kinetic theory of gases?

Perhaps the most important equation of the kinetic theory of gases is the equation which states that the product of the pressure and volume of an ideal gas is equal to the product of the number of moles that make up the gas, the thermodynamic temperature of the gas and the universal gas constant. This equation is important because it relates the number of particles in the gas to the macroscopic variables of the gas, providing a link between the statistical properties of a gas and its thermodynamic properties.

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