Ideal and Real Gases

Properties of Ideal and Real Gases

Ideal Gas and Real Gas: How Are They Different?

Ideal gas is a theoretical gas that behaves in a very ideal manner. No gas which exists in the environment acts perfectly as an ideal gas, though some do come really close to it under certain conditions of temperature and pressure. Since ideal gas is a theoretical gas, there are some assumptions that are made for its behaviour, which can be considered as properties of ideal gas

• Ideal gas follows the gas laws at all temperatures and pressures.

• The gas molecules are solid, spherical particles, which do attract or repel each other.

• The motion of these molecules is completely random.

• The particles travel in a straight line until they hit another particle or the wall of the container.

• Collision of these particles with each other or the wall of the container is completely elastic. I.e. no energy is lost in the collision and therefore the kinetic energy of the gas remains constant.

• The average distance between the particles is much larger than the size of the molecules.

• Kinetic energy of each particle can vary but the average kinetic energy of all particles remains constant.

Any gas that exists in reality is a real gas. Some gases come close to ideal gas behaviour under some conditions, but can never be perfectly ideal. This is because:

• Gases are made up of matter - atoms and molecules - which will always have some forces of attraction or repulsion between them.
• Atoms and molecules can never be point-sized as they will always occupy some volume.
• Under some conditions of high pressure or low temperature, the size of atoms/molecules can no longer be considered negligible as compared to distance between the particles.

Generally, the behaviour of gases tends towards ideal gas behaviour at high temperatures and low pressures. Gases generally deviate from ideal gas behaviour at low temperatures or high pressures.

Ideal Gas vs Real Gas: Similarities, Differences, and Examples

Although ideal gas and real gas seem to be completely different, let's not forget, they're both gases and have some similarities:

We can tabulate all the differences between ideal and real gases like so:

Examples of Ideal gases are - Hydrogen, Oxygen, Nitrogen and noble gases like Helium and Neon. These gases show behaviour that is very close to that of ideal gases at conditions of standard temperature and pressure (STP).

All gases found in the environment are examples of real gases. Even Hydrogen, Oxygen, and Nitrogen behave like real gases at conditions of low temperatures and high pressures - That is why it is possible to liquify them.

Deviation of Gases from Ideal Gas Behaviour

Conditions of Ideal Gas Behaviour

Imagine there is a container of large volume containing a real gas in it. The temperature inside the container is high.

Fig. 1: Large Container with Real Gas at High Temperature | Kanishk Singh StudySmarter Originals

By “large volume", we mean that the size of the particle, and the average distance between particles is negligible compared to the size of the container. And by “high temperature”, we mean that the kinetic energy of the particles is high enough so that intermolecular attraction/repulsion is negligible (the long red arrows are indicators of high kinetic energy of particles). The gas inside the container has the right conditions for it to behave like an ideal gas.

Therefore, the pressure of this gas can be determined by the ideal gas equation -

$$P\cdot V=n\cdot R\cdot T$$

Rearranging:

$$P=\frac{n\cdot R\cdot T}{V}$$

Deviation from Ideal Gas Behaviour due to Low Temperature

Now, let’s cool down the same container to a very low temperature. Volume of the container is the same.

Fig. 2: Large Container with Real Gas at Low Temperature | Kanishk Singh StudySmarter Originals

Looking at the equation for pressure from ideal gas equation:

$$P=\frac{n\cdot R\cdot T}{V}$$

Temperature is in the numerator. Therefore, if we reduce the temperature, pressure should reduce. Let’s compare this with what is happening inside the container.

As the temperature of the gas is reduced, gas particles now have very low kinetic energy (the small blue arrows are indicators of the low kinetic energy possessed by the particles). This will reduce the speed at which these particles hit the wall of the container, reducing the pressure. But, reducing the temperature had another effect - Since the particles don’t have much kinetic energy, the intermolecular forces of attraction or repulsion between the particles are not negligible anymore. This further reduces the speed at which the particles collide with the container. Due to this, the pressure drops more than what is predicted by the ideal gas equation. The behaviour of the gas can no longer be accurately predicted by the ideal gas equation, and therefore, the gas is said to be deviated from ideal gas behaviour.

Deviation from Ideal Gas Behaviour due to High Pressure

Now, let's take another container with the same amount of the same gas. Temperature is set to high, same as before. But this time, the volume of the container is small, and by "small" we mean that the size of the molecules are not negligible compared to the size of the container.

Fig. 3: Small Container with Real Gas at High Temperature Kanishk Singh StudySmarter Originals

Looking at the equation of P from the ideal gas equation:

$$P=\frac{n\cdot R\cdot T}{V}$$

We can see that volume is in the denominator. So when we decreased the volume, the value of the overall expression increased, and therefore pressure increased.

Since the amount of gas is the same as before and volume is smaller, the molecules are more densely packed. Therefore, the size and the average distance between the molecules is not negligible compared to the size of the container. This means that the space to move around for the particles is even lesser and the number of molecular collisions is higher. Due to this, the particles collide with the walls of the container with more vigour. This results in pressure increasing more than what was predicted by the ideal gas equation. The behaviour of the gas can no longer be predicted accurately by the ideal gas law, and the gas is said to be deviated from ideal gas behaviour.

Therefore, you learnt that under conditions of low temperature or high pressure, the behaviour of real gases tends to deviate from ideal gas behaviour.