Orbital / Electron Configuration Calculator

Electron configuration shows how an atom’s electrons are distributed across its subshells. The ground-state configuration follows three rules: the Aufbau principle (fill lowest energy subshells first), Pauli exclusion (max two electrons per orbital, opposite spins), and Hund’s rule (fill degenerate orbitals singly before pairing).

Subshell capacities: s holds 2, p holds 6, d holds 10, f holds 14. Each subshell is labeled by principal quantum number n and letter (1s, 2s, 2p, 3s, 3p, 3d, etc.).

The Madelung filling order, by the (n+l) rule with smaller n breaking ties, runs: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p

Notice 4s fills before 3d, even though n=3 < n=4. The reason: 4s sits slightly lower in energy for neutral atoms despite the higher principal quantum number.

Eight known exceptions exist within elements 1-54. The famous two, chromium and copper, are the half-filled and fully-filled d shells everyone learns. The other six are subtler: in the 4d row the 4d and 5s levels sit so close together that the ordering is decided by electron-electron repulsion rather than by any tidy stability rule.

Technetium (43) is the one that catches people out. It sits between Mo and Ru, both exceptions, and follows the rule anyway at [Kr] 4d⁵ 5s². Half-filled stability is a tendency, not a law.

The noble gas shorthand replaces the inner core with the symbol of the previous noble gas in brackets. Carbon’s full configuration 1s² 2s² 2p² becomes [He] 2s² 2p², where [He] stands for the same 1s² found in helium.

Valence electrons are the outermost-shell electrons, which determine chemical bonding. For main-group elements they are the s and p electrons of the highest n. For transition metals it is conventionally the (n)s plus (n-1)d electrons.

Unpaired electrons and magnetism

Hund’s rule says a set of equal-energy orbitals fills singly before any orbital doubles up, so a p³ subshell holds three unpaired electrons and a p⁴ holds only two. Count the unpaired electrons and you know the magnetic behavior:

• One or more unpaired: paramagnetic, weakly pulled into a magnetic field
• All paired: diamagnetic, weakly pushed out

Oxygen is the textbook case. Liquid oxygen sticks to the poles of a magnet because each O atom carries two unpaired 2p electrons, and O₂ keeps two unpaired electrons in its molecular orbitals. Nitrogen, one place to the left with three unpaired electrons in the atom, pairs all of them up in N₂ and pours straight through.

Ions

Pick a charge and the calculator adjusts the electron count. The removal order for cations trips up almost everyone: you take electrons from the highest principal quantum number first, which for a transition metal means the ns electrons go before the (n-1)d electrons, the reverse of the order they filled in.

Iron is [Ar] 3d⁶ 4s². Fe²⁺ is [Ar] 3d⁶, not [Ar] 3d⁴ 4s². Fe³⁺ then drops to [Ar] 3d⁵, a half-filled d shell, which is a large part of why iron(III) is so common in nature.

Anions simply keep filling in Madelung order, so O²⁻ has ten electrons and matches neon exactly. Two species with the same electron count are called isoelectronic, and the calculator names the match when there is one.