Why does ionization energy increase regularly across the periodic table from group 1A to group 8A, whereas electron affinity increases irregularly from group 1A to group 7A and then falls dramatically for group 8A?

Key Concepts

Ionization Energy

Ionization energy is the energy required to remove an electron from an atom in its gaseous state. It generally increases across a period from left to right due to increasing nuclear charge, which attracts electrons more strongly, making them harder to remove. Additionally, as atomic size decreases, the outer electrons are held more tightly, contributing to the trend.

Electron Affinity

Electron affinity refers to the energy change that occurs when an electron is added to a neutral atom in the gas phase. This property does not follow a simple trend across the periodic table; while it generally becomes more negative (indicating a greater tendency to gain electrons) from group 1A to group 7A, the dramatic drop in group 8A occurs because noble gases have a complete valence shell and do not favorably accept additional electrons.

Periodic Trends

Periodic trends are patterns observed in the properties of elements across the periodic table, influenced by atomic structure. These trends include ionization energy, electron affinity, atomic radius, and electronegativity. Understanding these trends helps explain why certain properties change in a predictable manner across periods and groups, reflecting the underlying principles of atomic interactions and electron configurations.