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The atomic radius of an element is the distance from the nucleus to the outermost (valence) electrons. This radius shifts in predictable ways across the periodic table, driven by the interplay between the nucleus’s positive charge and the surrounding electrons.
Electrons orbit the nucleus in discrete energy levels. Each level contains subshells—s, p, d, f—that can hold a fixed number of orbitals, and consequently a fixed maximum number of electrons. When a subshell fills, additional electrons must occupy orbitals in the next higher energy level. The higher the energy level, the farther its electrons sit from the nucleus.
Moving left to right across a main‑group period, atomic radii steadily decrease while the valence‑electron count rises. The reason is a rising net nuclear charge that pulls the existing valence electrons closer, without adding a new energy level. The slight deviation observed in transition metals stems from their partially filled d subshells, which reduce the effective pull on outer electrons.
Shielding occurs when inner electrons partially neutralize the nucleus’s positive charge. The remaining “effective” nuclear charge is what the valence electrons feel. Across a period, the number of inner electrons stays constant while the nuclear charge increases, so the effective charge grows, tightening the electron cloud and shrinking the radius.
Descending a group, the valence electrons occupy successively higher energy levels. Even though the valence‑electron count stays the same, the additional shells push the outer electrons farther from the nucleus. The increased number of protons is counterbalanced by the extra shielding from inner electrons, resulting in a net increase in atomic radius.
