[Tom]

At very high pressures, the volume of most real gases is greater than that predicted by the ideal gas law.  What would be the best explanation for this?
   

[integral0]

At very high pressures, the ideal gas law can no longer descibe accurately the state of those "real gases."  Because, in a real gas, the volume of the gas particles at high pressures needs to be taken into account.  In reality, those particles are not independent of each other i.e. they really do effect each other (attractive forces).  Therefore, in order to account for the change of volume observed at high pressures, one will need to utilize the van der Waals Equation.

[Mitch]

gas molecules have volume too. The volume of the molecules does not need to be taken into account when working at room temperature. But when you increase the pressure by a lot you will notice the volume predicted by the ideal gas law is no longer valid and this was because you forgot to include the volume of the gas molecules themselves.

[AWK]

The volume of the molecules does not need to be taken into account when working at room temperature
better: when working far from critical point.

[Tom]

I understand what you mean. So we know that the molecules of real gas occupy space, and its safe to say the kinetic energy of a real gas is greater than that of an ideal gas, and furthermore the molecules of a real gas are influenced by the interparticle attractions.

[Donaldson Tan]

You cannot compare ideal gas and real gas in this manner.

There is no intermolecular attraction between ideal gas molecules. It's not as if the intermolecular attraction between the molecules has been overcame by the kinetic energy of the gas molecules.

The volume of ideal gas molecules is always negligible compare to the volume of the system. However, in a real system, high pressure translate to small volume of the system. Hence, the total volume of the particles isn't negligible pon comparision to the volume of the empty space in the system.