Why doesn't Helium behave as an ideal gas? I am a bit confused (might be due to some conceptual misunderstanding) as to why doesn't Helium behave as an ideal gas (it shows a deviation from the $pV$ vs $p$ graph)? (Given the fact that it is highly inert therefore there is no attraction between the molecules.)
Here is the graph for reference :

 A: A substance being inert does not mean that it doesn't have forces of attraction or repulsion, or even that the collisions undergone by gas molecules are elastic. It is true that helium gas is quite inert, but there still are Van der Waals' interactions present in helium, moreover the gas particles are not exactly point sized which the kinetic theory assumes gases to be. Hence there is deviation from ideal gas behaviour.
A: Helium may have negligible interactions, but it still has a non-zero volume.   As you compress the pressure, the space available for the atoms to occupy does not decrease linearly.
An ideal gas would have components that occupy zero volume.  So each component would have collisions only with the sides of the container.
The molecules of a real gas occupy a volume. This volume reduces the effective volume of the container, which means collisions happen more often and the pressure is greater than would be expected from an ideal gas.  The smaller the molecule/atom/component, the lower the slope on the graph.
