How to Write an Electron Configuration

An electron configuration is a list of where an atom keeps its electrons. Writing one is a mechanical procedure with about four rules and twenty exceptions, and once you can do it you can read most of the periodic table's structure straight off the page — why the table has the block widths it has, why group 2 exists, why the lanthanides are a separate strip at the bottom.

This page is the how-to. It assumes you know what a shell and a subshell are; if not, the orbitals guide covers that first.

The notation

A configuration is a string of terms like 1s² 2s² 2p⁶. Each term has three parts:

• the principal quantum number, n — the shell, counting outward from 1;
• the subshell letter — s, p, d or f, which describes the shape of the orbitals;
• a superscript — how many electrons are in that subshell.

Each subshell holds a fixed maximum, and the numbers are worth memorising because everything else follows from them: s holds 2, p holds 6, d holds 10, f holds 14. Those are the widths of the four blocks of the periodic table, which is not a coincidence — the s-block is two columns wide, the p-block six, the d-block ten and the f-block fourteen.

The superscripts of a neutral atom always add up to the atomic number. That is the single best check on a configuration you have just written: if the total is not Z, something is wrong.

Rule one: fill in order of energy, not in order of shell

The rule that catches everyone is that subshells do not fill in the order 1, 2, 3, 4. They fill in order of increasing energy, and from the fourth shell onward those two orders diverge.

The order is set by the Madelung rule: subshells fill in order of increasing n + ℓ, and where two subshells tie, the one with the smaller n goes first. Taking ℓ as 0 for s, 1 for p, 2 for d and 3 for f, that produces:

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

The famous consequence is that 4s fills before 3d. Potassium and calcium put their outer electrons into 4s while 3d is still empty, and only then does the first transition series begin filling 3d. The same thing happens one row down with 5s before 4d, and again with 6s before 4f and 5d.

The usual way to reproduce this without memorising it is the diagonal diagram: write the subshells in a grid with one row per shell, then read the diagonals from top right to bottom left. It is a mnemonic for the n + ℓ rule, not a separate fact.

Rule two: stop when you run out of electrons

Take the atomic number, walk the filling order, and put electrons in until they are gone.

Iron, Z = 26:

1. 1s² — 2 used, 24 left
2. 2s² — 4 used, 22 left
3. 2p⁶ — 10 used, 16 left
4. 3s² — 12 used, 14 left
5. 3p⁶ — 18 used, 8 left
6. 4s² — 20 used, 6 left
7. 3d⁶ — all 26 used

So iron is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶. Check: 2 + 2 + 6 + 2 + 6 + 2 + 6 = 26. Correct.

Rule three: two ways to write the same answer

Iron's configuration is also written 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s², with 3d before 4s. This is the same configuration, reordered by shell rather than by filling energy. Both appear in textbooks and both are accepted; the shell-ordered form has the advantage that it groups each shell together, which makes the outer shell easier to spot.

Which you should write depends on what your course wants. If nobody has told you, write it in shell order and say that is what you have done. The pages on this site print both.

Rule four: abbreviate with the previous noble gas

Writing out all 86 electrons of radon before getting to the interesting part of a heavy element is a waste of everyone's time, so the inner shells are replaced by the symbol of the last noble gas before the element, in square brackets.

Iron becomes [Ar] 3d⁶ 4s², because argon is element 18 and accounts for the first eighteen electrons exactly. Silver becomes [Kr] 4d¹⁰ 5s¹. Sodium becomes [Ne] 3s¹.

The abbreviation is not just shorthand. What is left outside the bracket is the part of the atom that does chemistry, so the shorthand quietly makes the chemically relevant electrons the only ones you see. Every element in group 1 ends in s¹ outside its bracket; every halogen ends in s² p⁵. That pattern is the periodic table's whole logic in one line.

The twenty exceptions

Twenty elements do not have the ground-state configuration the filling order predicts, and they are not arbitrary. A half-filled or completely filled d subshell is unusually stable, so where promoting one s electron into d achieves that, the atom does it.

Chromium is the standard example. Aufbau predicts [Ar] 3d⁴ 4s². The actual ground state is [Ar] 3d⁵ 4s¹ — one electron moved from 4s to 3d, buying a half-filled d subshell. Copper does the same to reach a full one: [Ar] 3d¹⁰ 4s¹ rather than 3d⁹ 4s².

The others are molybdenum, niobium, ruthenium, rhodium, palladium, silver, lanthanum, cerium, gadolinium, platinum, gold, actinium, thorium, protactinium, uranium, neptunium, curium and lawrencium. Palladium is the most extreme: it empties its 5s subshell completely and sits at [Kr] 4d¹⁰.

For an exam, chromium and copper are the two you are expected to know. The rest are worth recognising as a class — "the ones that steal an s electron to complete a d or f subshell" — rather than memorising individually.

Ions: take from the outermost shell, not the last one filled

This is the other place people reliably lose marks. When a transition metal forms a cation, it loses its 4s electrons before its 3d electrons — even though 4s filled first.

Iron is [Ar] 3d⁶ 4s². Fe²⁺ is [Ar] 3d⁶, not [Ar] 3d⁴ 4s². Fe³⁺ is [Ar] 3d⁵, which is a half-filled d subshell, and part of why iron(III) is so common.

The rule to carry away is: fill by energy, empty by distance. Electrons go in lowest-energy first and come off outermost first, and for the fourth row those are not the same subshell.

For a main-group element there is no such conflict. Sodium loses its single 3s electron to become [Ne]; chlorine gains one to become [Ar]. Both end up with a noble gas configuration, which is the whole reason they react at all.

Checking your answer

Three checks, in order of how often they catch something:

1. Do the superscripts add to Z? For an ion, to Z minus the charge.
2. Does the last term match the element's block? A d-block element must end in a d subshell in filling order; a p-block element in a p subshell.
3. Is it one of the twenty? If your element is chromium, copper, or anything in the silver and gold columns, check the exception list before trusting the Aufbau answer.

The electron configuration calculator on this site does all three, and flags which of the twenty exceptions it has applied rather than silently returning a different answer from the one your working produced.

One last habit worth building: after writing a configuration, look at where the element sits in the table and check that the two agree. The block tells you the last subshell, the period tells you the shell number, and the position within the block tells you how many electrons are in it. Selenium is in period 4, in the p-block, four columns into it — so it ends in 4p⁴, and it does. A configuration that fails that check is wrong before you have counted anything.