Guide
Electron Configuration Rules Made Clear
Electron configuration is the recipe for how an atom's electrons are arranged among its orbitals, and it is the hidden engine behind almost every chemical property. Where the electrons go determines how an element bonds, what charge its ions take, and where it sits on the periodic table. Students often find configurations fiddly at first, but they are governed by just three rules, and once those click the whole system becomes predictable.
This guide lays out the rules in order, shows how to write both full and shorthand configurations, and — most usefully — explains how the periodic table itself is a configuration chart you can read directly. Get comfortable here and a lot of chemistry that once seemed arbitrary starts to make sense.
Orbitals, subshells, and shells
Electrons occupy regions called orbitals, grouped into subshells labelled s, p, d, and f, which in turn belong to numbered shells. An s subshell holds up to 2 electrons, a p subshell up to 6, a d subshell up to 10, and an f subshell up to 14. The shell number tells you the energy level, so 2s means the s subshell of the second shell.
A configuration is simply a list of which subshells are occupied and by how many electrons, written as the shell number, the subshell letter, and a superscript count. Carbon, with six electrons, is written 1s2 2s2 2p2: two electrons in the first shell's s subshell, two in the second shell's s subshell, and two in the second shell's p subshell.
Rule 1 — the Aufbau principle
The first rule, Aufbau (German for 'building up'), says electrons fill the lowest-energy orbitals first before moving to higher ones. This gives the standard filling order: 1s, then 2s, then 2p, then 3s, and so on. The order is not simply by shell number, because subshells of different shells overlap in energy — the 4s subshell fills before the 3d, which is why the transition metals appear where they do.
The famous diagonal filling diagram is just a memory aid for this order. But there is an easier way to recover it, covered below: the periodic table is laid out in exactly the Aufbau sequence, so you can read the filling order straight off the chart.
Rule 2 — the Pauli exclusion principle
The second rule limits how many electrons share an orbital. The Pauli exclusion principle states that a single orbital holds at most two electrons, and those two must have opposite spins. This is why each subshell has its fixed capacity: an s subshell has one orbital and holds 2, a p subshell has three orbitals and holds 6, a d subshell has five orbitals and holds 10.
In practice this rule is what stops you from piling all of an atom's electrons into the lowest orbital. Each orbital fills to its limit of two paired electrons, then the next electron must go elsewhere, which is what forces the layered structure of shells and gives atoms their size and their bonding behaviour.
Rule 3 — Hund's rule
The third rule governs how electrons distribute within a single subshell that has several orbitals. Hund's rule says electrons spread out singly across the orbitals of a subshell, all with parallel spins, before any orbital gets a second electron. Electrons are negatively charged and repel one another, so occupying separate orbitals keeps them apart and lowers the energy.
For nitrogen's three 2p electrons, this means one electron in each of the three p orbitals rather than two in one and one in another. Hund's rule matters for predicting an element's magnetic behaviour and the number of unpaired electrons available for bonding.
Shorthand notation with noble gases
Full configurations get long fast — gold's runs to many subshells — so chemists use a shorthand that replaces the inner electrons with the previous noble gas in brackets. Sodium's full configuration 1s2 2s2 3s1 becomes simply [Ne] 3s1, because neon accounts for the first ten electrons. Iron becomes [Ar] 3d6 4s2, and gold becomes [Xe] 4f14 5d10 6s1.
The shorthand is not just tidier; it highlights the part that matters. The electrons written after the noble-gas core are the outer, or valence, electrons that actually do the chemistry. Focusing on them is often all you need to predict how an element behaves.
Reading configurations off the table
Here is the shortcut that ties everything together: the periodic table is arranged in Aufbau order, so an element's position gives its configuration directly. The blocks name the subshell being filled — the s-block fills s subshells, the p-block fills p, the d-block fills d, the f-block fills f — and the period number gives the shell. Count across the block to get the supershell count.
So chlorine, in period 3 of the p-block, three across, ends in 3p5, giving [Ne] 3s2 3p5. You did not need the diagonal diagram at all; you read it from the map. Practising this on a few elements, then checking against their listed configurations, quickly turns the abstract rules into something you can do on sight.
Frequently asked questions
What are the three electron configuration rules?
The Aufbau principle (fill lowest-energy orbitals first), the Pauli exclusion principle (two electrons per orbital, opposite spins), and Hund's rule (spread electrons singly across a subshell's orbitals before pairing them).
Why does 4s fill before 3d?
Because subshells from different shells overlap in energy, and the 4s subshell is slightly lower in energy than 3d. The Aufbau order follows energy, not shell number, which is why the transition metals sit where they do.
What is noble-gas shorthand?
A compact notation that replaces an atom's inner electrons with the symbol of the previous noble gas in brackets. Sodium becomes [Ne] 3s1, showing only the valence electrons that drive its chemistry.
Can I get the configuration from the periodic table?
Yes. The table is laid out in filling order: the block gives the subshell (s, p, d, f), the period gives the shell number, and counting across the block gives the electron count. Chlorine in period 3, five into the p-block, ends in 3p5.