Electron affinity (EA) is a measurable property that describes an atom’s tendency to attract an extra electron. This property changes predictably across the periodic table, revealing fundamental truths about atomic structure. The observable pattern shows that this electron-gaining tendency generally increases as one moves from the left side of the table toward the right. Understanding this trend requires examining how the nucleus’s pull and the atom’s physical size evolve with increasing atomic number.
Defining Electron Affinity
Electron affinity is defined as the energy change that occurs when a neutral gaseous atom gains an electron to form a negative ion. This process is represented by the equation \(text{X}(text{g}) + text{e}^- to text{X}^-(text{g})\), and the energy is typically measured in kilojoules per mole (\(text{kJ/mol}\)). For most atoms, energy is released, resulting in a negative EA value (an exothermic process). A more negative electron affinity indicates a greater attraction for the incoming electron.
This concept differs from electronegativity, which is a relative, unitless scale measuring an atom’s ability to attract a shared pair of electrons within a chemical bond. Electron affinity is a quantitative measurement taken on an isolated atom, revealing its inherent tendency to form an anion.
Observing the Periodic Trend
Electron affinity values generally increase as one moves horizontally from left to right across any given period. This means that atoms on the far right side of the table (excluding the noble gases) are the most receptive to acquiring an additional electron. For instance, in Period 3, the electron affinity increases significantly from sodium (Na) to chlorine (Cl).
The halogens in Group 17 display the highest overall electron affinities because they are only one electron away from achieving the stable, full-shell configuration of a noble gas. Conversely, the alkali metals (Group 1), alkaline earth metals (Group 2), and noble gases (Group 18) show very low or even positive electron affinity values.
The Influence of Effective Nuclear Charge and Atomic Radius
The increasing electron affinity trend across a period is fundamentally driven by the interplay between the nucleus’s positive charge and the atom’s physical size. Moving from left to right, the atomic number increases, adding more protons to the nucleus and providing a stronger attractive pull on all electrons. Since the incoming electron is added to the same outermost principal quantum shell, the shielding effect from inner-shell electrons remains relatively constant.
Because the number of protons increases while shielding is constant, the effective nuclear charge (\(text{Z}_{text{eff}}\)) experienced by the valence electrons increases significantly. This stronger net positive charge is the primary engine driving higher electron affinity, allowing the nucleus to exert a greater attractive force on the new addition.
This strengthened nuclear pull also influences the atom’s physical size, pulling the electron cloud inward and leading to a progressive decrease in the atomic radius across the period. A smaller atomic radius places the newly added electron much closer to the positively charged nucleus. According to Coulomb’s Law, the attractive force is inversely proportional to the square of the distance between charges. The combination of higher \(text{Z}_{text{eff}}\) and reduced atomic radius results in the added electron experiencing a significantly greater net attraction, translating into a more favorable, exothermic process.
Understanding Key Exceptions to the Rule
Despite the strong general trend, certain elements deviate significantly due to their specific electron configurations. The most obvious exception is Group 18, the noble gases, which have electron affinities close to zero or positive. Their configuration is a stable, completely filled valence shell (\(text{ns}^2text{np}^6\)), and adding an electron requires placing it into the next, much higher, principal energy level. This energetically unfavorable process means they resist accepting an electron.
Another notable deviation occurs with the alkaline earth metals (Group 2), such as beryllium and magnesium, which also exhibit very low electron affinities. These atoms possess a stable, completely filled s-subshell (\(text{ns}^2\)). To accept an electron, the new particle must enter the higher-energy p-orbital, which is less attractive to the nucleus.
A distinct dip in the trend occurs at Group 15, including nitrogen and phosphorus. These elements have a half-filled p-subshell (\(text{np}^3\)), a configuration that confers extra stability. Adding an electron requires placing it into an orbital that already contains an electron, significantly increasing electron-electron repulsion. This repulsion makes the electron-gaining process less energetically favorable than it is for neighboring elements, resulting in a lower-than-expected electron affinity.