How Electron Configuration Works: A Step-by-Step Guide

An electron configuration is a shorthand list of where every electron sits in an atom. It tells you which orbitals are filled, how many electrons occupy each one, and therefore how the atom is likely to behave when it meets other atoms. Once you can write a configuration, many facts about the periodic table, such as why sodium is so reactive or why the noble gases are inert, follow almost automatically.

This guide explains the three rules that govern filling, the order in which orbitals fill, and the shorthand notation chemists use. Each rule is followed by a worked example so you can check the logic yourself.

What an orbital is

An orbital is a region of space around a nucleus where an electron is most likely to be found. It is not a fixed path like a planet orbit. Each orbital can hold at most two electrons, and those two must have opposite spin.

Orbitals are grouped into subshells labelled s, p, d and f. Each subshell contains a fixed number of orbitals:

Subshell	Orbitals in subshell	Maximum electrons
s	1	2
p	3	6
d	5	10
f	7	14

Every subshell belongs to a principal shell, labelled by a number n = 1, 2, 3 and so on. The principal shell n can hold at most 2n² electrons in total. Shell 1 holds 2, shell 2 holds 8 and shell 3 holds 18.

The three filling rules

Three rules decide where each electron goes. They work together, so it helps to learn them as a set.

1. The Aufbau principle

Aufbau is German for "building up." Electrons enter the lowest-energy orbital that is still available. Energy generally rises with the sum of the principal number and the angular momentum number, which is why 4s fills before 3d even though 3d has a higher principal number.

The usual filling order is:

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

A helpful memory aid is the diagonal rule: write the shells in rows, then follow arrows running diagonally from the top right to the bottom left. The arrows pass through the subshells in the order above.

2. The Pauli exclusion principle

No two electrons in an atom can share the same four quantum numbers. In practice this means an orbital holds at most two electrons, and when it holds two, their spins must point in opposite directions. Chemists write this as an up arrow and a down arrow in the same box.

3. Hund's rule

When several orbitals have the same energy, such as the three 2p orbitals, electrons spread out across them first, each with the same spin, before any orbital receives a second electron. This arrangement is the most stable because electrons that do not share an orbital repel each other less.

The consequence matters for chemistry. Carbon has two electrons in its 2p subshell, and Hund's rule places them in two different p orbitals with parallel spins. That pattern gives carbon two unpaired electrons in its ground state.

Writing the configuration

Chemists write the configuration by listing each occupied subshell, then placing its electron count as a superscript. The total of all superscripts equals the atomic number Z, which is the number of protons and, in a neutral atom, the number of electrons.

Worked examples:

Element	Z	Full configuration	Shorthand
Hydrogen	1	1s¹	1s¹
Carbon	6	1s² 2s² 2p²	1s² 2s² 2p²
Sodium	11	1s² 2s² 2p⁶ 3s¹	[Ne] 3s¹
Iron	26	1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶	[Ar] 4s² 3d⁶
Chlorine	17	1s² 2s² 2p⁶ 3s² 3p⁵	[Ne] 3s² 3p⁵

The shorthand uses a noble gas in square brackets to stand in for the inner, filled shells. [Ne] means 1s² 2s² 2p⁶, which is ten electrons. Only the outer electrons, called valence electrons, need to be written out in full, and they are the ones that take part in bonding.

Note on the iron entry

Notice that iron is written as 4s² 3d⁶ rather than 3d⁶ 4s². In neutral atoms the 4s orbital is lower in energy than 3d, but once electrons are removed the 3d orbital becomes the lower-energy one. Iron loses its 4s electrons before its 3d electrons when it forms the Fe²⁺ and Fe³⁺ ions, which is why iron chemistry is dominated by the 3d subshell.

Exceptions that show the rules are approximate

The Aufbau order is a good guide, not a law. Chromium and copper break it because a half-filled or completely filled d subshell is unusually stable.

Element	Predicted	Observed
Chromium (Z = 24)	[Ar] 4s² 3d⁴	[Ar] 4s¹ 3d⁵
Copper (Z = 29)	[Ar] 4s² 3d⁹	[Ar] 4s¹ 3d¹⁰

In both cases one electron moves from 4s into 3d. A d⁵ or d¹⁰ arrangement gains extra stability from exchange energy, the small energy benefit of having many parallel-spin electrons in separate orbitals. Textbooks list a handful of such exceptions, and they are a reminder that configurations come from measured spectra and calculations, not from a rule that never fails.

Why configurations predict chemistry

The outermost electrons determine how an atom reacts. Elements in the same column of the periodic table share the same valence configuration, which is why they behave alike. Group 1 elements such as lithium, sodium and potassium all end in ns¹ and readily lose that single electron to form +1 ions. Group 17 elements such as fluorine, chlorine and bromine all end in ns² np⁵ and readily gain one electron to complete a noble-gas shell.

The noble gases end in ns² np⁶, a full outer shell of eight, except helium with two. Their stability explains why they rarely react and why the octet rule is a useful first guide to bonding. Atoms tend to gain, lose or share electrons until they reach a configuration like this.

Common mistakes to avoid

Students often make the same few errors. The first is writing 3d before 4s for neutral atoms, which conflicts with the filling order. The second is forgetting that a subshell can hold at most the number of electrons shown in the table above. The third is confusing the shell number with the number of electrons; shell 3 is not the same as three electrons.

A good check is to add up your superscripts and compare the total with the atomic number from the periodic table. If they match, the configuration accounts for every electron. If they do not, look for a missing subshell or a superscript that exceeds its limit.

Using configurations with the periodic table

The periodic table is arranged so that its rows and blocks follow the configuration. The s block on the left has two columns, the p block on the right has six, and the d and f blocks sit in between or below. Reading a configuration from the table, or the reverse, becomes quick with practice.

Key takeaways

Electrons fill orbitals from lowest to highest energy, follow the Aufbau order with the 4s before 3d exception for neutral atoms, and pair up only after each orbital in a subshell holds one electron. These three rules let you write the configuration of almost any element in the first few rows of the table.

Electron configuration is the bridge between an element's position on the periodic table and its chemical behaviour. The valence electrons, which sit in the outermost shell, decide how an atom bonds, which is why elements in the same group act alike.

Use the shorthand with a noble gas core to save effort, and check your work by summing the superscripts against the atomic number. Remember the known exceptions such as chromium and copper, because they show that the rules are a model fitted to measured data.