Electron Affinity

Electron affinity measures the ease with which an atom gains an electron, indicating the energy change when an electron is added to a gaseous atom. A more negative electron affinity value means the atom more readily accepts an electron and releases energy in an exothermic process. This property helps predict an element's chemical reactivity and tendency to form anions.

Although fluorine is smaller and has greater nuclear attraction, the incoming electron experiences significant repulsion from electrons already present in fluorine's small n=2 shell. Chlorine's incoming electron enters the larger n=3 shell, experiencing reduced electron-electron repulsions. This makes chlorine more favorable for electron acceptance, resulting in a more negative electron affinity value of −348.6 kJ/mol compared to fluorine's −322 kJ/mol.

Moving across a period from left to right, electron affinities generally become more negative as effective nuclear charge increases. Halogens have the most negative electron affinities because the incoming electron helps achieve stable noble gas configurations. This trend reflects increasing attraction between atoms and added electrons across the period.

Noble gases have completely filled electron shells, making them highly stable. When an electron is added to a noble gas, it must occupy a higher principal energy level, which is energetically unfavorable. This requires energy input rather than releasing it, resulting in positive electron affinity values that reflect the instability of the resulting anion.

Group 15 elements like phosphorus have half-filled p-subshells. When an electron is added, it must pair with an existing p-orbital electron, increasing electron-electron repulsions and making the process energetically unfavorable. This results in less negative electron affinities for group 15 compared to group 14 elements, disrupting the expected periodic trend.

Moving down a group, atomic size increases as electrons occupy higher principal quantum numbers. Incoming electrons experience less nuclear attraction due to increased distance and shielding effects, leading to less negative electron affinities. However, the second element in each group often has the most negative electron affinity due to the small size of the n=2 shell.

Group 2 elements have filled ns subshells, so the next electron must enter the higher-energy np subshell. This energetically unfavorable placement results in electron affinities that are either positive or less exothermic than the periodic trend would predict, disrupting the expected pattern across the periodic table.