The elements in the periodic table are organized into vertical columns known as groups. As one moves down a group, atoms follow a clear trend in their ability to accept a new electron. This tendency, known as electron affinity, generally decreases as the atomic number increases. Understanding this decrease requires examining the fundamental forces within the atom that change with increasing size.
What Electron Affinity Measures
Electron affinity (EA) is a quantitative measure of an atom’s desire to acquire an additional electron. It is defined as the energy change that occurs when a neutral gaseous atom gains one electron to form a negative ion. This process is typically exothermic for most elements, meaning energy is released because the resulting negative ion is more stable.
When discussing the magnitude of the attraction, chemists typically report electron affinity as a positive value corresponding to the amount of energy released. A higher electron affinity value signifies a stronger attraction for the incoming electron and a greater amount of energy released. The general trend of decreasing electron affinity down a group means that atoms at the bottom release less energy, indicating a weaker attraction for the new electron. Elements in Group 2 (alkaline earth metals) and Group 18 (noble gases) are common exceptions because their stable electron configurations make accepting an electron energetically unfavorable.
The Role of Atomic Radius
The most significant factor causing electron affinity to decrease down a group is the continuous increase in atomic radius. As you descend a column, each element adds a new principal energy level, effectively adding a new, complete shell of electrons. This shell addition dramatically increases the distance between the positive nucleus and the outermost valence shell, where the new electron settles.
The attractive force between the positive nucleus and the incoming electron is governed by distance; as distance increases, the attractive force weakens rapidly. In larger atoms at the bottom of the group, the nucleus is farther away from the valence shell. This greater separation results in a diminished pull on the added electron, making the process of gaining an electron less energetically favorable. Consequently, less energy is released when the electron is added, corresponding to the observed decrease in electron affinity.
The electron cloud of a larger atom is more diffuse, meaning the negative charge of the added electron is spread over a greater volume. This reduces the concentration of charge near the nucleus, lessening the strength of the electron-nucleus attraction. This geometric effect of increasing size is the primary reason why the ability to attract and bind an external electron diminishes down a group.
Understanding Electron Shielding
Complementary to the effect of increasing size is the phenomenon of electron shielding, also called screening. As new electron shells are added down a group, the electrons in these inner, filled shells come between the nucleus and the incoming electron. These inner electrons, being negatively charged, repel the incoming valence electron, effectively shielding it from the full attractive force of the positive nucleus.
While the total number of protons (the nuclear charge, Z) increases down a group, the repulsive effect of the growing number of core electrons increases almost proportionally. This shielding reduces the net positive charge that the incoming electron “feels,” a concept known as the Effective Nuclear Charge (\(Z_{eff}\)). The \(Z_{eff}\) experienced by the valence shell is significantly lower than the actual nuclear charge.
This reduction in the effective positive pull makes the atom less willing to accept and hold onto the new electron. The combined effects of increased distance and increased repulsion from inner electrons destabilize the newly formed negative ion. Atoms lower in the group experience a weaker net attraction for an extra electron, leading to the characteristic decrease in electron affinity.