Why are van der Waals' forces between gas molecules strong near its critical temperature? My textbook says


  …easily liquefiable gases i.e. with higher critical temperatures are readily adsorbed as van der Waals' forces are stronger near the critical temperatures.


What I do understand is that at a temperature higher than the critical temperature $(T_\mathrm c)$ the gas molecules have higher kinetic energy which is why they're unable to form strong bonds. 

But, what if the temperature is lower than the critical temperature? Or is my reasoning incorrect?
Critical temperature is  kind of measure of the strength of intermolecular van der Waals force.

Comparison of actual and critical temperature is kind of comparison of average kinetic energy of molecules with energy needed to break the intermolecular bonds.

So temperature itself does not say anything about the bonding strength. It only indirectly determine probability these bonds get broken by thermal molecular motion.