Electron shells and orbitals are the starting point for understanding atomic structure, and they are needed again when we discuss periodic trends or the electron configuration of silicon. Electrons around the nucleus are not scattered at random: they sit in shells ordered by energy, and each shell is divided into orbitals.

Shells and their capacity

Shells are labeled K, L, M and so on outward from the nucleus, corresponding to principal quantum numbers n = 1, 2, 3. A shell can hold at most 2n² electrons: 2 in the K shell, 8 in the L shell and 18 in the M shell.

One point on energy: the closer a shell is to the nucleus, the lower its energy, because the electrons are held more strongly by the protons. Outer shells have higher energy and weaker attraction, so outer electrons are the ones most easily involved in reactions.

Orbital shapes and numbers

A shell can contain several kinds of orbitals. An s orbital is spherical, with one orbital per set; a p set contains three dumbbell-shaped orbitals; a d set contains five. Whatever its type, an orbital holds at most two electrons.

Rules for filling orbitals

1. Fill the lowest-energy orbital first, which is 1s.
2. Each orbital holds at most two electrons.
3. Orbitals of equal energy, such as the three 2p orbitals, each receive one electron before any of them is paired.

Helium: with atomic number 2, both electrons go into 1s. The K shell is full, giving a closed shell.

Oxygen: with atomic number 8, its configuration is 1s²2s²2p⁴. The 1s and 2s orbitals take two electrons each, and the remaining four go into 2p: one each in px, py and pz, then the fourth pairs up in px. Oxygen therefore has two unpaired electrons, and its L shell holds fewer than eight electrons, so it is an open shell and chemically reactive.