Why is the first energy of ionization of oxygen lesser than that of nitrogen? The following question arises from a question I found in my book.


  Experimentally it has been determined that the value of the first
  energy of ionization of oxygen is lesser than the first energy of
  ionization of nitrogen. Select the statement which explains this:
  
  
  The higher electronegativity of oxygen
  The lesser electron affinity of nitrogen
  The lesser radius of oxygen
  The higher stability of nitrogen
  The higher electron affinity of nitrogen
  


According to my book the answer is the fourth option, but I couldn't find a way to justify this. What sort of explanation can do this?. 

My guess is that nitrogen forms a triple bond with itself which is higher than the oxygen and this makes the molecule more stable but I don't know if this can be translated into a higher energy of ionization nitrogen rather than the oxygen which does have a sigma and pi bonds. Am I right? 

Can someone please illustrate this or bring some sort of diagram with some values to justify this?. It would help a lot if an answer could help me why should I discard the second option?. 

Doesn't nitrogen has smaller electron affinity than the oxygen? Wouldn't it mean that having a smaller electron affinity cause it to be easier to strip out an electron from its outer shell rather than the opposite as the question indicates?
Two electrons of the oxygen molecule are located on a $\pi $* level. They are relatively easy to ionize. It is not the case in $N_2$ : all electrons are located in binding orbitals.