Electron configuration follows the rule that the principal energy shell being filled corresponds to the period number for the s- and p-blocks, but for the d-block the filled subshell's energy level is one less than the period number, and for the f-block it is two less than the period number. This arises because d and f orbitals of a given period are actually lower in energy than the s/p orbitals of that same period, causing them to "backfill" an inner shell even though the atom is being built up period by period. This distinction matters because valence electrons—those in the outermost energy shell, which determine chemical reactivity—are not necessarily the last-added or highest-energy electrons; in transition metals and lanthanoids/actinoids, the highest-energy electrons occupy an inner shell (d or f) while the outer s-shell retains only its own electrons as valence electrons. This is a foundational concept in atomic structure and quantum chemistry, explaining the layout of the periodic table's blocks (s, p, d, f) and the electron-filling order (Hund's/Aufbau-related ordering conventions) that underlies periodic trends and chemical bonding behavior.
Electron configuration follows the rule that the principal energy shell being filled corresponds to the period number for the s- and p-blocks, but for the d-block the filled subshell's energy level i…