Certainly! Electron affinity refers to the energy change that occurs when a neutral atom in the gas phase gains an electron to form a negatively charged ion. The trends you mentioned—decreasing electron affinity down a group and increasing electron affinity across a period—can be explained by considering two main factors: atomic size and effective nuclear charge.
Atomic Size: As you move down a group in the periodic table, the atomic size or radius generally increases. This is because each successive element has an additional electron shell, leading to increased distance between the nucleus and the outermost electrons. With a larger atomic size, the outermost electrons are further away from the nucleus, resulting in a weaker attractive force between the nucleus and the incoming electron. Therefore, it becomes easier to add an electron, and the electron affinity decreases down a group.
Effective Nuclear Charge: When moving across a period from left to right, the number of protons in the nucleus increases, resulting in a higher effective nuclear charge. The effective nuclear charge is the net positive charge experienced by the outermost electrons after accounting for shielding effects by inner electrons. The increased effective nuclear charge pulls the outermost electrons closer to the nucleus, reducing the atomic size.
As the atomic size decreases across a period, the attraction between the nucleus and an incoming electron becomes stronger. Consequently, it requires more energy to add an electron, and the electron affinity increases from left to right across the periodic table.
It's important to note that while these trends generally hold true, there can be some exceptions or irregularities due to specific electronic configurations or other factors. However, the explanations provided above offer a broad understanding of the observed trends in electron affinity.