Why do some gases not obey Boyle's Law at any pressure?

In the case of NO2, if you count up the valence electrons, you will realize that it has an odd number of electrons.  The presence of an unpaired electron makes it particularly reactive, and in fact, it always exists in an equilibrium with the gas N2O4.  When you decrease the volume or increase the pressure of the gas, NO2 is more likely to interact with itself and shift the equilibrium in favor of N2O4.  Doing the reverse has the opposite effect, and N2O4 tends to fall apart into two molecules of NO2.

NO2 essentially violates the notion that ideal gas molecules do not interact with each other at every temperature and pressure.  It's not so much that NO2 isn't ideal, because no gas is ideal.  It just doesn't approximate an ideal gas at normal temperatures and pressures like many other gases do.  It is too reactive.

To my knowledge, NF2 does not exist naturally, but if it did, it would probably have the same problem of reacting with itself.  It would have an odd number of electrons and would probably tend to dimerize.

When I read the title I said NO2 fits that criteria.  I am not sure what NF2 is (N2F4? fluorohydrazine?)
Anyway NO2 exists as a monomer dimer equilibrium at temperatures near RT:
 N2O4  ? 2NO2
Kp  = [NO2]^2/N2O4]
If you apply pressure to this system the equlibrium shifts to the left (Le Chateliers) that decreases the volume more than that would be expected from Boyle's Law   P1/V1 = P2/V2
Another system that would do this is acetic acid which is also a monomer ? dimer eq > 120°C    (check this temp).
Aha! Greenwood and Earnshaw p 804 (1st ed) to the rescue:
N2F4: most intriguing property is an ability to dissociate at RT and above to give the free radical NF2.   At 150°C the equilibrium constant for the dissociation N2F4 ? 2NF2 is K-0.03 atm...  

I have learn some new chemistry today!  cheers.