By Sean Lancaster
Updated Mar 24, 2022
Understanding an atom’s structure involves examining its nucleus and the surrounding electron orbitals. Electrons occupy concentric energy shells around the nucleus; the closer a shell is to the nucleus, the lower its energy. The outermost electrons—those in the valence orbitals—participate in chemical bonding. These orbitals are typically s and p types for covalent bonds, and as you move down a period, additional d orbitals become available.
Electrons fill orbitals from the lowest to the highest energy, with each orbital holding a maximum of two electrons. When an orbital is doubly occupied, its energy rises relative to a singly occupied orbital.
Determine the element’s total number of electrons. This equals its atomic number.
Write the full electron configuration, filling orbitals in the order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s. Remember: s orbitals hold 2 electrons, p orbitals hold 6, and d orbitals hold 10.
Identify the last s or p orbital filled; these contain the valence electrons. For example, silicon (atomic number 14) has the configuration 1s² 2s² 2p⁶ 3s² 3p². Thus, its valence electrons reside in the 3s and 3p orbitals—four in total.
