Electron Affinity is the energy associated with the addition of an electon to a gaseous atom.
Example:
| Cl(g) + e- → Cl-(g) | E.A. = -349 kJ/mol |
Notice the sign on the energy is negative. This is because energy is usually released in this process, as opposed to ionization energy, which requires energy. A more negative electron affinity corresponds to a greater attraction for an electron. (An unbound electron has an energy of zero.)
Trends:
As with ionization energy, there are two rules that govern the periodic trends of electron affinities:
Electron affinity becomes less negative down a group.
As the principal quantum number increases, the size of the orbital increases and the affinity for the electron is less. The change is small and there are many exceptions.
Electron affinity decreases or increases across a period depending on electronic configuration.
This occurs because of the same subshell rule that governs ionization energies.
Example:
Since a half-filled "p" subshell is more stable, carbon has a greater affinity for an electron than nitrogen.
Obviously, the halogens, which are one electron away from a noble gas electron configuration, have high affinities for electrons:
(More negative energy = greater affinity)
| Element | Electron Affinity |
|---|---|
| I | -295.2 kJ/mole |
| Br | -324.5 kJ/mole |
| Cl | -348.7 kJ/mole |
| F* | -327.8 kJ/mole |
*Fluorine's electron affinity is smaller than chlorine's because of the higher electron - electron repulsions in the smaller 2p orbital compared to the larger 3p orbital of chlorine.
