Electron Configuration Guide

Where Electrons Live — and Why It Matters for Every Chemical Reaction

An electron configuration describes how an atom's electrons are distributed across energy levels and orbital types. Every chemical property — reactivity, color in compounds, magnetic behavior, oxidation state — flows from electron configuration. Knowing where electrons are tells you what reactions an element will participate in and how strongly.

The configuration is written as a series of notations like 1s²2s²2p⁶3s²3p⁴ (sulfur), where the number is the principal energy level, the letter is the orbital type, and the superscript is the number of electrons in that subshell.

The Three Rules: Aufbau, Hund's Rule, Pauli

Three rules govern electron placement. Knowing all three is required to write any configuration correctly.

1. Aufbau Principle: Fill Lowest Energy First

Electrons fill orbitals from lowest to highest energy. The energy order isn't simply 1s, 2s, 2p, 3s, 3p, 3d... — the d subshells are higher in energy than the next s subshell.

The actual filling order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p

Memory trick: The diagonal arrow mnemonic (Madelung rule) generates this order. Or simply: 4s fills before 3d; 5s before 4d; 6s before 4f and 5d.

Orbital capacities:

• s subshell: 1 orbital × 2 electrons = 2 electrons max
• p subshell: 3 orbitals × 2 electrons = 6 electrons max
• d subshell: 5 orbitals × 2 electrons = 10 electrons max
• f subshell: 7 orbitals × 2 electrons = 14 electrons max

2. Pauli Exclusion Principle: Two Electrons Per Orbital, Opposite Spins

No two electrons in the same atom can have the same set of four quantum numbers. Practically: each orbital holds at most 2 electrons, and those 2 must have opposite spins (one spin-up ↑, one spin-down ↓).

What this means for writing configurations: An s subshell fills to 2 electrons before you move to the next; a p subshell fills through all three orbitals (giving 6 electrons total) before moving on.

3. Hund's Rule: Maximize Unpaired Electrons Before Pairing

When filling orbitals of equal energy (like the three p orbitals or five d orbitals), put one electron in each orbital before pairing any.

Carbon (6 electrons): 1s²2s²2p²

The 2 electrons in the 2p subshell go into separate p orbitals (both spin-up), not paired in the same orbital:

• Correct: 2p↑__ 2p↑__ 2p__ (both unpaired)
• Wrong: 2p↑↓ 2p__ 2p__ (forced pairing violates Hund's rule)

This matters for magnetic properties: carbon with 2 unpaired electrons is paramagnetic (weakly attracted to magnets). If electrons were forced to pair, the material would be diamagnetic.

Writing Configurations: Step by Step

Chlorine (Cl, atomic number 17 — 17 electrons):

Fill in order:

• 1s: 2 electrons used (2 total, 15 remaining)
• 2s: 2 electrons (4 total, 13 remaining)
• 2p: 6 electrons (10 total, 7 remaining)
• 3s: 2 electrons (12 total, 5 remaining)
• 3p: 5 electrons (17 total, 0 remaining)

Full configuration: 1s²2s²2p⁶3s²3p⁵

Note the 3p⁵ — one electron short of a complete p subshell (which would be 3p⁶). This is why chlorine is highly reactive: it needs one more electron to reach noble gas configuration (like argon, which has 3p⁶). Chlorine readily gains that electron, forming Cl⁻.

Noble gas shorthand notation: Instead of writing 1s²2s²2p⁶3s²3p⁵, represent the previous noble gas in brackets: Argon (Ar) is 1s²2s²2p⁶3s²3p⁶. So chlorine = [Ne] 3s²3p⁵ where [Ne] = 1s²2s²2p⁶.

The Notable Exceptions: Chromium and Copper

The textbook filling order predicts incorrect configurations for several transition metals. The two you must know are chromium (Cr) and copper (Cu).

Chromium (Cr, Z = 24)

Predicted by aufbau: [Ar] 3d⁴ 4s² Actual configuration: [Ar] 3d⁵ 4s¹

An electron moves from 4s to 3d to create a half-filled d subshell (d⁵). Half-filled subshells (and fully-filled subshells) have extra stability due to exchange energy — the symmetric arrangement of electrons with parallel spins lowers the total energy.

The d⁵ configuration has one electron in each of the five d orbitals, all with the same spin (Hund's rule applied to all five orbitals). This maximizes exchange energy and creates unusual magnetic properties — Cr³⁺ ions give many chromium compounds their characteristic green or purple colors.

Copper (Cu, Z = 29)

Predicted by aufbau: [Ar] 3d⁹ 4s² Actual configuration: [Ar] 3d¹⁰ 4s¹

A fully-filled d subshell (d¹⁰) is even more stable than a half-filled one. One electron moves from 4s to complete the d subshell. This explains copper's excellent electrical conductivity — the single loosely-held 4s electron is the charge carrier in copper metal.

Other exceptions (less commonly tested):

• Molybdenum (Mo): [Kr] 4d⁵ 5s¹ (like Cr)
• Silver (Ag): [Kr] 4d¹⁰ 5s¹ (like Cu)
• Gold (Au): [Xe] 4f¹⁴ 5d¹⁰ 6s¹ (like Cu, with a full d¹⁰)
• Tungsten (W), Palladium (Pd), and others have various exceptions driven by stability of filled or half-filled d or f subshells

Configurations for Ions

Ions form by gaining or losing electrons. Metal cations lose electrons from the highest principal quantum number first (not necessarily the highest energy orbital).

Iron (Fe, Z = 26): [Ar] 3d⁶ 4s²

Fe²⁺: Remove 2 electrons from 4s first: [Ar] 3d⁶ (not [Ar] 3d⁴ 4s²) Fe³⁺: Remove the 4s electrons plus 1 d electron: [Ar] 3d⁵

The Fe³⁺ configuration has a half-filled d subshell (3d⁵) — which is why Fe³⁺ is particularly stable and commonly formed.

Nonmetal anions gain electrons: O²⁻ (oxide ion): O = 1s²2s²2p⁴, adds 2 electrons to 2p: 1s²2s²2p⁶ — same as neon configuration (isoelectronic with Ne).

Paramagnetism and Diamagnetism

Atoms with unpaired electrons are paramagnetic — weakly attracted to magnetic fields. Atoms with all electrons paired are diamagnetic — weakly repelled.

Counting unpaired electrons:

• Ne: 1s²2s²2p⁶ — all paired → diamagnetic
• O: 1s²2s²2p⁴ → 2p has 4 electrons in 3 orbitals: two orbitals get one each (unpaired), one gets two (paired) → 2 unpaired electrons → paramagnetic
• Fe³⁺: [Ar] 3d⁵ → 5 electrons in 5 d orbitals, each orbital with one electron → 5 unpaired electrons → strongly paramagnetic

This is why iron oxide (Fe₂O₃) is the basis for magnetic data storage — the unpaired electrons in Fe³⁺ create a permanent magnetic moment that can be aligned.