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? 
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1$\begingroup$ At a fixed pressure real gas occupy more volume, that is because of repulsion.. $\endgroup$– user43021Commented Sep 18, 2017 at 19:36
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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.
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$\begingroup$ However, real gas molecules have some volume. Thus when they come too close, they tend to repel each other and oppose the compression. This is also the chief reason why volume of the real gas is more than that of ideal gas at a particular pressure. Because as we compress the, the gas particles prevent each other from occupying the same volume. So if volume increases, it should be exerting more pressure on the walls than the ideal gas. This is something that does not happen in Ideal gases. So how do the real gases compensate this repulsion during change of state? $\endgroup$– JajatiCommented Sep 17, 2017 at 3:04
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$\begingroup$ The repulsion you're talking about is only significant at very very short distances (smaller than even the distance between atoms in a chemical bond). Any larger than this distance and you see attractive forces pull the particles together. This is why a real gas almost always occupies a smaller volume than an ideal gas. $\endgroup$ Commented Sep 17, 2017 at 5:16
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$\begingroup$ Sir, kindly have a look at the picture I have added to the question. Here you can see that the at high pressure, real gas tends to occupy more volume than ideal gas.If repulsion really exists at such less distances, how can you explain the increase in volume of the real gases at high pressure as given in the graph? $\endgroup$– JajatiCommented Sep 17, 2017 at 6:31
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1$\begingroup$ The particles have less distance to travel before they strike another particle, causing the real gas to occupy a larger volume than an ideal gas (at high pressures). At low pressures, the distance between particles is so large, the difference between particles with no volume and particles with volume is negligible. Attractive forces cause the volume to be smaller than ideal gases at low pressures because the particles are "sticky" and pull towards each other. Finally, you cannot apply kinetic molecular theory (how we can think about gases) to phase changes. It's a whole different set of rules. $\endgroup$ Commented Sep 18, 2017 at 23:27
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1$\begingroup$ That's almost correct. There are no significant repulsive forces that play a part in kinetic molecular theory. The gas takes up more room because each particle takes up space. $\endgroup$ Commented Sep 20, 2017 at 0:09