Periodic Trends: Atomic and Ionic Radii, Ionization Enthalpy

Atomic and ionic radii. On descending either group, one full principal quantum shell is added at every step, so both the atomic radius and the ionic radius increase steadily and substantially down the group: for Group 1, $\text{Li} < \text{Na} < \text{K} < \text{Rb} < \text{Cs}$; for Group 2, $\text{Be} < \text{Mg} < \text{Ca} < \text{Sr} < \text{Ba}$. The increase in principal quantum number (and the resulting extra shell of inner, screening electrons) outweighs the modest increase in nuclear charge, so the outermost electron sits, on average, farther from the nucleus at every step down the group.

Comparing across the two groups in the same period, a Group 2 atom is always distinctly smaller than the Group 1 atom immediately before it (for example, $\text{Na}$, $186\ \text{pm}$, versus $\text{Mg}$, $160\ \text{pm}$), because the Group 2 atom has one extra proton in the nucleus (higher effective nuclear charge) while both outer electrons are added to the very same shell, pulling the whole electron cloud in more tightly.

Ionization enthalpy. Descending either group, the first ionization enthalpy falls steadily, because the outermost electron is progressively farther from the nucleus and more effectively shielded by the growing core of inner electrons, so it takes less energy to remove: $\text{Li}\,(520) > \text{Na}\,(496) > \text{K}\,(419) > \text{Rb}\,(403) > \text{Cs}\,(376\ \text{kJ mol}^{-1})$, and similarly down Group 2. This falling ionization enthalpy is exactly why chemical reactivity increases down each group -- it becomes progressively easier for the atom to lose its valence electron(s) and be oxidized. …