Electron arrangement determines many chemical properties. AP Chemistry describes electron locations using energy levels (shells), sublevels (subshells), and orbitals, along with basic filling rules that produce an atom’s lowest-energy (ground-state) electron configuration.

Shells: principal energy levels

Electrons in atoms occupy shells, which are broad energy levels centered around the nucleus. Shell number increases as electrons, on average, occupy regions farther from the nucleus and higher in energy.

Shells help explain why elements in the same period fill the same overall energy level: moving across a period adds electrons to the same principal shell until it is energetically favourable to begin a new shell.

This capacity is a useful upper limit, but actual filling depends on subshell energies within and between shells.

Subshells: s, p, d, f

Within each shell, electrons are grouped into subshells (also called sublevels). Subshells differ in shape and energy and are labelled s, p, d, f.

Key subshell facts used in AP Chemistry:

• s subshell: 1 orbital, holds up to 2 electrons

• p subshell: 3 orbitals, holds up to 6 electrons

• d subshell: 5 orbitals, holds up to 10 electrons

• f subshell: 7 orbitals, holds up to 14 electrons

• Each shell contains specific subshells:

  • $n=1$: s only

  • $n=2$: s and p

  • $n=3$: s, p, and d

  • $n\ge 4$: s, p, d, and f (as applicable)

Subshell labels are combined with the shell number (for example, 2p means the p subshell in the second shell).

This diagonal-rule (Aufbau) chart summarizes the relative filling order of atomic subshells (e.g., $1s \rightarrow 2s \rightarrow 2p \rightarrow 3s \rightarrow 3p \rightarrow 4s \rightarrow 3d \dots$). Following the arrows gives the standard sequence used to build ground-state electron configurations from lowest to higher orbital energy. Source

Orbitals: where electrons are most likely found

A subshell is made of individual orbitals.

The figure shows the three p orbitals ($p_x$, $p_y$, $p_z$) as distinct, equal-energy orbitals oriented along perpendicular axes. Seeing them side-by-side helps explain why a p subshell contains three orbitals (and therefore can accommodate up to 6 electrons, two per orbital with opposite spins). Source

An orbital is not a fixed path; it is a region in space where an electron is likely to be found.

Counting orbitals links directly to electron capacity:

• s: 1 orbital → max 2 e⁻

• p: 3 orbitals → max 6 e⁻

• d: 5 orbitals → max 10 e⁻

• f: 7 orbitals → max 14 e⁻

Orbital filling rules for ground-state configurations

A ground-state electron configuration describes the arrangement of electrons that minimises the atom’s energy, consistent with electrostatic attractions and electron–electron repulsions.

Two principles govern how electrons fill orbitals within a subshell.

In addition, electrons spread out among equal-energy orbitals before pairing.

These orbital filling diagrams use boxes (orbitals) and arrows (electron spins) to show how electrons are distributed within subshells. The examples make Hund’s rule visible (single occupancy with parallel spins across degenerate orbitals before pairing) while also enforcing the Pauli exclusion principle (paired electrons in one orbital must have opposite spins). Source

Practical implications for AP Chemistry:

• Within a subshell, fill one electron per orbital first (Hund), then begin pairing.

• Never exceed two electrons per orbital (Pauli).

• Electron configurations are commonly written using subshell notation with superscripts (for example, 1s² means two electrons in the 1s orbital).