Shells, Subshells, and Electron Configuration

Electrons avoid crowding into one level because they repel each other and because each shell only offers so much room at a given energy. Mutual repulsion pushes electrons into the lowest available energy states, but the Pauli exclusion principle caps how many can occupy each orbital, so they spread out across shells (n = 1, 2, 3...) and subshells within those shells. Lower n means lower energy and a smaller, tighter region around the nucleus; higher n means more room but higher energy.

Subshells differ in shape and capacity. The s subshell is a single sphere-shaped orbital holding 2 electrons. The p subshell has three dumbbell-shaped orbitals along different axes, holding 6 electrons total. The d subshell has five orbitals, holding 10 electrons. Capacity always follows 2(2l+1), where l = 0 for s, 1 for p, 2 for d. To build iron's ground state (26 electrons), fill in order of increasing energy: 1s2 2s2 2p6 3s2 3p6 4s2 3d6. Notice 4s fills before 3d despite 3d belonging to a lower principal shell, because 4s is briefly lower in energy for the neutral atom.

Ionization reverses this logic. Once 3d starts filling, its orbitals drop below 4s in energy for that atom, so removing electrons empties 4s first: Fe3+ is 1s2 2s2 2p6 3s2 3p6 3d5, not a configuration missing 3d electrons. This connects directly to excited states, where a photon bumps one electron to a higher empty orbital temporarily, then photon emission when it falls back, the basis of atomic spectra covered next.