Why does chlorine have a higher electron affinity than fluorine?

For #F#, #DeltaH# #=# #-328*kJ*mol^(-1)#;

For #Cl#, #DeltaH# #=# #-349*kJ*mol^(-1)#.

A similar disparity in electron affinities exists for the 2 first Group VI elements, oxygen (#-142*kJ*mol^(-1)#), and sulfur (#-200*kJ*mol^(-1)#).

So why should the 1st row atoms have reduced (magnitude) electron affinity? Two factors are important, (i) nuclear charge, and (ii) shielding by other electrons. In the case of the second row atoms, these have increased nuclear charge, and it should be easier to add electrons.

Also in the case of the second row atoms, the electrons are entering orbitals that are more diffuse (it builds on an inner electron shell), and physically larger, so that electron-electron repulsion is reduced.

The 2 factors combine to give the unusual Periodic trend.

The electron affinity of an element is the energy released when one mole of the element in the gaseous state each gain an electron under standard conditions.

Down a group, electron affinity typically decreases. This is because the atomic radius increases down a group. The electron gained ends up in the outermost shell. It is farther away from the nucleus and thus feels a weaker attraction. Less energy is released.

However, there are exceptions. Fluorine, which is higher up the group then chlorine, has a lower electron affinity. This is because the electrons in the outermost shell of a fluorine atom are closer together. The electron gained also feels a great amount of repulsion from the electrons originally in the outermost shell. Energy is required to keep the gained electron in the shell, causing fluorine to have a smaller electron affinity than chlorine.