How to Write Electron Configurations

The configuration is the chemistry

Almost every chemical property an element has — its bonding, its color, its magnetism, where it sits in the periodic table — traces back to how its electrons are arranged in orbitals. Once you know the electron configuration, you can predict whether an element will lose or gain electrons, what oxidation states are accessible, why iron is paramagnetic and zinc isn’t, and why neon refuses to react with anything.

Carbon’s configuration is 1s² 2s² 2p² — two electrons in the 1s, two in the 2s, two in the 2p. That’s six electrons total, matching its atomic number, and the four electrons in the n=2 shell explain why carbon makes four bonds. The whole framework is bookkeeping, but it’s the bookkeeping that runs the rest of inorganic chemistry.

Three rules you fill orbitals by

Aufbau principle. Fill from lowest energy upward. The order:

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p

Notice 4s comes before 3d, and 6s before 4f and 5d. The “diagonal rule” diagram is a memory aid, but recognizing the pattern in the periodic table itself works better long-term: row determines n, block (s/p/d/f) determines the subshell, and column gives you the electron count within the block.

Pauli exclusion. No two electrons in the same orbital can have the same spin. Practically: each orbital holds at most 2 electrons. Subshell capacities: s holds 2, p holds 6, d holds 10, f holds 14.

Hund’s rule. When orbitals have the same energy (the three p orbitals, the five d orbitals), put one electron in each before pairing any up. Pairing electrons costs repulsion energy; Hund’s rule minimizes that cost. All unpaired electrons in a degenerate set have parallel spins.

How to write one

1. Count electrons. Z for a neutral atom, Z − charge for a cation, Z + |charge| for an anion.

2. Fill in Aufbau order until you’ve placed all of them. Keep a running tally.

3. Write each occupied subshell with its electron count as a superscript: 1s² 2s² 2p⁶ 3s¹.

4. Optional shorthand. Replace the noble-gas core with [NobleGas]. Sodium full: 1s² 2s² 2p⁶ 3s¹. Shorthand: [Ne] 3s¹.

Worked examples

Oxygen (Z = 8). 1s² (2 placed), 2s² (4), 2p⁴ (8). 1s² 2s² 2p⁴, or [He] 2s² 2p⁴. Six valence electrons in the n=2 shell — that’s why oxygen wants two more to complete its octet.

Iron (Z = 26). Through argon takes 18 electrons. Then 4s² (20), then 3d⁶ (26). [Ar] 4s² 3d⁶ (or equivalently [Ar] 3d⁶ 4s² — both notations are accepted; reordering by principal quantum number is common in textbooks because it groups the n=3 electrons together).

Chromium (Z = 24) — exception. You’d predict [Ar] 4s² 3d⁴. The actual configuration is [Ar] 4s¹ 3d⁵. A half-filled d subshell is unusually stable thanks to exchange energy (the QM bonus for having parallel spins across degenerate orbitals). One electron promotes from 4s to 3d to grab that stability.

Copper (Z = 29) — exception. Predicted: [Ar] 4s² 3d⁹. Actual: [Ar] 4s¹ 3d¹⁰. A fully filled d subshell wins for the same reason. Other 4d-block exceptions — Mo, Pd, Ag, Au — follow analogous patterns; lanthanides and actinides have even more.

Barium (Z = 56). Through xenon takes 54. Then 6s² (56). [Xe] 6s². Group 2, two valence electrons, behaves like the heavier alkaline earth it is.

Cations: 4s leaves before 3d

Here’s the wrinkle that gets missed. The Aufbau order says 4s fills before 3d when you build up the atom. But once the atom is built, 3d has actually dropped below 4s in energy in the populated atom (the orbital energies depend on which orbitals are filled). When you ionize a transition metal, electrons leave from the highest principal quantum number first — so 4s electrons leave before 3d electrons.

• Fe (Z = 26): [Ar] 4s² 3d⁶
• Fe²⁺: [Ar] 3d⁶ (remove both 4s)
• Fe³⁺: [Ar] 3d⁵ (remove both 4s and one 3d — and notice 3d⁵ is the half-filled stability that makes Fe³⁺ common)
• Cu (Z = 29): [Ar] 4s¹ 3d¹⁰
• Cu⁺: [Ar] 3d¹⁰ (remove the lone 4s)
• Cu²⁺: [Ar] 3d⁹

For main-group ions: Na⁺ has 10 electrons, configuration [Ne] (same as neon). Cl⁻ has 18, configuration [Ar]. Iso-electronic with the nearest noble gas — that’s the whole reason these ions form.

Valence electrons

For main-group elements (groups 1–2 and 13–18), valence count is the group number’s last digit (groups 1–2 give 1–2; group 13 gives 3, 14 gives 4, …, 18 gives 8). They live in the highest-n s and p subshells.

Transition metals are messier — the (n−1)d electrons can also participate in bonding, which is why Mn ranges from +2 to +7 oxidation states. Counting “valence” for a transition metal usually means the ns electrons plus chemically accessible (n−1)d electrons.

Traps to watch for

Forgetting the d-block exceptions. Cr and Cu in row 4 are the canonical ones every general chemistry course tests. [Ar] 4s¹ 3d⁵ and [Ar] 4s¹ 3d¹⁰. Memorize the reasoning (half-filled and fully-filled d stability) so you can also predict Mo and Ag.

Removing d electrons before s for transition metal cations. Fe → Fe²⁺ removes the two 4s, not two of the 3d. Same for every first-row transition metal cation.

Confusing filling order with energy order. During build-up, 4s fills first because it’s lower at that point. After build-up, 3d sits lower. This is why ionization removes the 4s electrons first — they’re now the higher-energy ones. Both statements are correct; they describe different stages.

Exceeding subshell capacity. s holds 2, p holds 6, d holds 10, f holds 14. Writing 2p⁷ or 3d¹¹ is always wrong — recount your electrons.

Forgetting to adjust for charge. Na⁺ has 10 electrons, not 11. F⁻ has 10, not 9. Always set the electron count to Z minus charge before you start filling.

Use the Electron Configuration Calculator to generate full and shorthand configurations, orbital diagrams, and valence counts for any element or ion.