Validity of van der Waals equation Van der Waals  had stated that at high pressure, pressure exerted by ideal gas on the walls of container is more than real gas.

However in this graph, one can clearly see that, the at high pressure volume of real gas tends to be more than volume of ideal gas. Hence real gas should be exerting more pressure on the walls of the container, as it occupies more volume. So why does this not hold true?
It's because in ideal gases there are no intermolecular forces whereas in a real gas the force with which it collides is reduced due to intermolecular forces.Now for the volume occupying part ideal gas particles are point sizes but in a real gas the molecules do occupy some volume.That's why some corrections were made to volume and pressure in the ideal gas equation.
Note that the above were assumptions of the Kinetic theory of gases.