Why Do Real Gases Deviate From Ideal Behaviour?

In the last post we have discussed few postulates of Kinetic molecular theory of gases which explains the ideal behaviour of gases. Today we will try to find out where this theory went wrong and why real gases deviate from the Ideal behaviour?

We all know gases can be liquefied under pressure. You must have heard of Compressed Natural Gas (CNG) and Liquid Petroleum Gas (LPG) both are liquid form of methane gas which are stored at high pressure. If gases can be liquefied, it means there must be some forces working between them which hold them together and gas molecules also possess volume.

That means two postulates of Kinetic molecular theory of gases are wrong in which it says:

• Ideal gas molecules are so small that they occupy a negligible space
• There is no force of attraction between these molecules.

We have learnt that molecules exert pressure when they collide with the wall of container. But it is observed that at high pressure molecules come closer and the force of attraction between them starts working. When a molecule is about to collide with the wall, this force of attraction drags it back so it cannot collide with its full impact. That’s why at high pressure, pressure experienced by the walls of container is less than the expected. For this reason scientists added a correction term to the observed pressure to get the total pressure exerted by the molecules of gas.

P_Total= P_real + an²/V²

Where ‘a’ is the constant, n is the number of moles and V is the volume of the container and P_real is the observed pressure.

Forces between molecules

Like attractive force, repulsive force also comes in action at higher pressure when molecular distance decreases. This repulsive force prevents squashing of molecules so that each molecule maintains a territory. So the volume available for the motion of molecules would be less than the volume of the container because some of its space is already occupied by the molecules. That’s why we have to subtract a correction term from the volume of the container to get the actual volume.

V_remaining = V- nb

Where ‘b’ is the constant, ‘n’ is the number of moles and ‘V’ is the volume of the container.

After doing these corrections we get a new equation for real gases which are derived from ideal gas equation:

(P_real + an²/V²) (V- nb) = nRT

This equation is known as Vander waals equation and ‘a’ and ‘b’ are known as Vander waals constants and their values depend on the characteristics of gas. ‘a’ is the measure of intermolecular attractive forces of gas and it is independent of pressure and temperature.

Now you have learnt that why real gases deviate from ideal behaviour. If you remember, I have told you that molecules are like us. Molecules don't behaved ideally just like we don’t. But there may be some conditions when they are likely to behave ideally. Can you tell me what those conditions are? In the next post we will try to find out its answer.  

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