why is the pressure exerted by ideal gas on the walls of the container more than that of real gas I have learnt that ideal gases have no intermolecular forces. So even at high pressure the volume gets completely compressed. 
But in case of real gases,  we have both attractive and repulsive forces. At far distance,  attractive forces dominate. 
But at high pressure , when the particles are closer to each other,  the repulsive forces dominate over the attractive forces. 
So my question is,  if the repulsive forces dominate over the attractive forces at close range,  then real gas should exert more pressure on the walls of the container.  With this one more doubt arises,  how do gases liquify if repulsive forces are dominant?
For neutral gases, the intermolecular forces that we observe tend to be attractive. Significant repulsive forces don't really show up until your particles have charge. 

The other difference between ideal gases and real gases are that real gases take up space, but the volume deviation from the ideal gas law is usually swamped out by the attractive intermolecular forces. This causes a real gas to exert less pressure than an ideal gas, and also is the reason matter undergoes phase changes.