Why doesn't NF₃ undergo hydrolysis?
Nitrogen trifluoride is extremely inert to hydrolysis, attack by base or acids or oxidation because of the fact that there is no d-orbitals in Nitrogen trifluoride and fluorine is far more electronegative than nitrogen. However, in the presence of base or acids, nitrogen trifluoride can however undergo hydrolysis because there is now excess of nucleophile which can attack nitrogen trifluoride.

There is a old paper which discuss the hydrolysis of nitrogen trifluoride in presence of acids and base.


  In the presence of aqueous base, however, slow hydrolysis occurs at
  100°C yielding nitrite and fluoride. 
  
  $$\ce{NF3 + 4 OH- -&gt;[100℃] NO2-  + 3F- + 2H2O}$$


It further says:


  This behavior differs sharply from that of nitrogen trichloride, 
  which is known to give ammonia and hypochlorite under similar con-
  ditions. The latter products are readily explained in terms of 
  nucleophilic attack directed at the chlorine atoms, a mechanism which 
  appears reasonable in view of the fact that the electronegativities 
  of $\ce{N}$ and $\ce{Cl}$ are very nearly the same and that the
  halogen may easily  expand its valence shell. Obviously, these
  considerations cannot be  applied to the nitrogen trifluoride molecule
  since fluorine is considerably more electronegative than nitrogen and
  it has no available  d orbitals. Although the nitrogen atom also has
  no free orbitals,  the relatively low electron density would at least
  offer less resistance to the approach of a nucleophile .
  
  
  
  The proposed intermediate $\ce{HONF2}$ would be expected to be
  unstable with  respect to the loss of $\ce{HF}$, as is apparently the
  case with the unknown  perfluoro alcohols . 
  
  Two other reaction modes worth considering include the formation of 
  an intermediate activated complex with water.