Question: I was tasked with figuring out whether carbon or nitrogen has a more negative electron affinity value. I initially picked nitrogen, just because nitrogen has a higher $Z\_\mathrm{eff}$, creating a larger attraction between electrons and protons, decreasing the radius, causing a higher ionization energy, and therefore decreasing the electron affinity value, but I was actually wrong, and the solutions manual explains it as this:

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> "As you go from C to N across the Periodic Table you would normally expect N to have the more negative electron affinity. However, N has a half-filled p subshell, which lends it extra stability; therefore, it is harder to add an electron."
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Are there any major exceptions to the rules when comparing electron affinity? I'm hesitant to use nitrogen as an exception, because I don't know how far it extends. If nitrogen has a more positive EA than carbon, does that also extend to boron, aluminium, or phosphorus? 

I later found that this also applies when comparing silicon and phosphorus. The explanation given was the same. 

What exceptions should be noted when comparing electron affinities? Are there any at all? And how far does the exception with atoms with half-filled p subshells extend?

Answer: 
This exception rule is actually orbital filling rule. For two electrons to be in same orbital they need to have different spins (Pauli exclusion principal). This electron pairing requires additional energy and thus it is easier to add electrons if there are free orbitals. When element has a half-filled p sublevel all 3 orbitals have one electron and pairing takes place (difference between energy levels of 2p and 3s is greater than electron pairing energy).

Electron pairing effects have significant impact to physical properties of coordination complexes (like color and magnetic properties).
I say the reason why the electron affinity of fluorine is not as negative as chlorine and that of O is not as negative as S is because of the electron repulsions in the small compact atoms keep the added electrons from being tightly bound as we might expect. . .