Electron configuration: watch every electron find its orbital

How to use this simulation

1. Just watch first: iron’s 26 electrons drop into the orbital boxes one at a time, and the note under the stage names the rule each one is obeying.
2. Move the Z slider to any element from 1 to 36, or tap a quick example: Na, Cl, K, Cr, Fe, Cu, or the ions Na⁺, Cl⁻, Fe²⁺ and Fe³⁺.
3. Switch the view to “Energy ladder”: the 4s rung sits below 3d, which is exactly why 4s fills first.
4. Pause and tap the diagram to add one electron at a time, so you can write the configuration in your notebook alongside it.
5. Compare the readings panel: full configuration, noble-gas shorthand, shells K L M N, outer-shell and unpaired electrons, and period · group · block.

Bus seats, hostel rooms and an address for every electron

Think about getting on a bus. People take the best empty seats first. If one of two seats in a pair is taken, most newcomers would rather sit alone in another empty pair, and only share when there is no empty pair left. Electrons behave in almost exactly the same way inside an atom.

Now picture a hostel. Each floor is a shell, each wing on a floor is a subshell, and each room is an orbital. A room holds at most two people, and if two share, one sleeps with their head to the north and the other to the south. Nobody wants to climb stairs, so the ground floor fills first. Writing down who lives where is exactly what an electron configuration does.

Why bother? Because the arrangement explains nearly everything about an element: why sodium fizzes in water, why neon refuses to react with anything, why iron sticks to a magnet and why copper conducts electricity so well. The outermost electrons in particular decide how many bonds an atom makes and where it sits in the periodic table.

Starting from zero: what is an electron configuration?

At the centre of an atom is the nucleus, made of protons and neutrons. Around it move the negatively charged electrons. In a neutral atom the number of electrons equals the number of protons, which is the atomic number Z. Iron has Z = 26, so a neutral iron atom has 26 electrons.

Those electrons are not scattered at random. Each one has a definite energy, and electrons with similar energies group together in levels. The electron configuration of an atom is the list of which shells, subshells and orbitals its electrons occupy, written according to a fixed set of rules.

There are two common ways to write it. Early on you meet the shell form, for example sodium: 2, 8, 1, meaning two electrons in the K shell, eight in L and one in M. Later you meet the subshell form: 1s² 2s² 2p⁶ 3s¹. Here 3s¹ means one electron in the s subshell of the third shell.

Read each term as: number (the shell, n), letter (the subshell s, p, d or f), and a small superscript (how many electrons are in it). Add up the superscripts and you must get the total number of electrons. Checking that sum at the end of every answer is the single best habit you can build in this chapter.

Shells K, L, M, N and the 2n² rule

Moving outwards from the nucleus, the main shells are called K, L, M and N, with principal quantum numbers n = 1, 2, 3 and 4. The further out a shell is, the higher its energy and the larger it is. Electrons in the K shell, closest to the nucleus, are held most tightly.

The largest number of electrons a shell can hold is 2n². For n = 3 (the M shell) that is 18; for n = 4 (the N shell) it is 32. The reason is simple: shell n contains n² orbitals, and each orbital holds two electrons.

Be careful, though: 2n² is a maximum, not a promise. Potassium (Z = 19) is 2, 8, 8, 1. Even though the M shell has room for 18, the nineteenth electron goes into the N shell after M has only 8. The reason lies in the energy order of subshells, which the Aufbau principle below explains. A handy rule for early classes is that the outermost shell never holds more than 8 electrons.

Subshells s, p, d, f and orbitals

Each main shell is split into smaller groups called subshells, labelled s, p, d and f. Shell n has exactly n subshells: K has only 1s; L has 2s and 2p; M has 3s, 3p and 3d; N has 4s, 4p, 4d and 4f.

Inside each subshell are orbitals. An orbital is a three-dimensional region around the nucleus where an electron is most likely to be found. It is not a neat circular track like the Bohr model draws, but more like a cloud. s orbitals are spheres, p orbitals look like dumbbells, and most d orbitals look like four-leaf clovers.

An s subshell has 1 orbital, p has 3, d has 5 and f has 7. Each orbital holds at most two electrons, so s holds 2, p holds 6, d holds 10 and f holds 14. Every small box in the simulation is one orbital, which is why 3d shows five boxes side by side.

The four quantum numbers, in plain words

A complete address for one electron needs four numbers, the way a postal address needs a country, city, street and house number. These are the quantum numbers.

The principal quantum number n says which shell (n = 1, 2, 3 …). The azimuthal (angular momentum) quantum number l says which subshell (l = 0 to n − 1; 0 is s, 1 is p, 2 is d, 3 is f). The magnetic quantum number m says which orbital inside the subshell (m runs from −l to +l, giving 2l + 1 values). The spin quantum number s says which way the electron spins (+½ or −½).

Everything else in this chapter follows from these four. For n = 2, l can be 0 or 1, so there are 2s and 2p. For l = 1, m = −1, 0, +1, so there are three p orbitals. Each takes a spin of +½ and −½, so p holds six electrons.

The Aufbau principle and the (n + l) rule

Aufbau is German for “building up”. The Aufbau principle says electrons enter the lowest-energy orbital available first, and only move on to a higher-energy orbital once the lower ones are full. It is like pouring water into a jug: the bottom fills first.

But how do you know which subshell is lower in energy? Use the (n + l) rule. Add n and l for each subshell. The one with the smaller n + l has lower energy and fills first. If two subshells have the same n + l, the one with the smaller n fills first.

The most famous result is that 4s fills before 3d. For 4s, n + l = 4; for 3d, n + l = 5. So even though 4s belongs to a higher shell, it has lower energy than 3d. And 3d and 4p both have n + l = 5, but 3d has the smaller n, so 3d fills before 4p.

The easiest way to remember the order is the diagonal diagram. Write 1s; 2s 2p; 3s 3p 3d; 4s 4p 4d 4f in rows, then draw arrows from top right to bottom left. The order in which the arrows pass through the subshells is the filling order. The simulation draws this chart on the right of the orbital boxes.

Pauli exclusion principle: two per room, facing opposite ways

Wolfgang Pauli stated that no two electrons in the same atom can have all four quantum numbers the same. At least one must differ.

The direct consequence: an orbital holds at most two electrons, and those two must have opposite spins. In one orbital n, l and m are already identical, so the only thing left to tell them apart is spin: one is +½ (↑), the other −½ (↓). A third electron would have no spin value left to take.

Watch the simulation: no box ever holds two ↑ arrows, and no box ever takes three arrows. The moment a second electron drops into a box, the note under the stage names the Pauli principle.

Hund's rule: sit alone first, then pair up

Orbitals in the same subshell have the same energy; the three 2p orbitals, for example. So where does a new electron go? Hund's rule of maximum multiplicity says: in orbitals of equal energy, electrons first go in one at a time with parallel spins. Only when every orbital has one electron do they begin to pair.

The reason is repulsion. Electrons are all negative and push each other away. Staying in separate orbitals keeps them further apart, which lowers the energy and makes the atom more stable. Parallel spins also let the electrons exchange places, which lowers the energy a little further.

Take nitrogen: 1s² 2s² 2p³. Its three 2p electrons sit alone in three different orbitals (↑ ↑ ↑), so nitrogen has 3 unpaired electrons. In oxygen the fourth 2p electron pairs up in the first orbital (↑↓ ↑ ↑), leaving 2 unpaired. In the simulation the amber arrows are unpaired and the coloured ones are paired.

The exceptions: why chromium and copper break the pattern

Following Aufbau blindly, chromium (Z = 24) should be [Ar] 3d⁴ 4s². In reality it is [Ar] 3d⁵ 4s¹: one 4s electron moves into 3d. In the same way copper (Z = 29) “should” be [Ar] 3d⁹ 4s², but is actually [Ar] 3d¹⁰ 4s¹.

The reason is the extra stability of half-filled (d⁵) and completely filled (d¹⁰) subshells. When a d subshell is half full, each of the five orbitals holds one electron with the same spin, so the charge is spread evenly all round the nucleus (symmetry). Electrons with parallel spins can also swap places, which releases exchange energy; the more parallel pairs, the larger it is.

Because 4s and 3d are so close in energy, moving one electron to reach a half-filled or full d subshell is worth it. In chromium this raises the unpaired electrons from 4 to 6. Copper ends with just 1 unpaired electron, in 4s.

Choose Cr or Cu in the simulation. After all the electrons are in, you will see one red arrow hop from 4s into 3d. Between Z = 1 and 36 these are the only two exceptions, which is why they turn up in exams again and again.

Configurations of ions: 4s empties first

An atom that loses electrons becomes a cation (positive ion); one that gains electrons becomes an anion (negative ion). To write an ion’s configuration, write the neutral atom first, then remove or add electrons.

Na⁺ has 10 electrons: 1s² 2s² 2p⁶, which is [Ne], the same stable octet as neon. Cl⁻ has 18 electrons: 1s² 2s² 2p⁶ 3s² 3p⁶, which is [Ar], just like argon.

For transition metals there is one rule that matters: when a cation forms, electrons leave the outermost shell first, meaning 4s before 3d. 4s fills first and empties first. So Fe²⁺ is [Ar] 3d⁶, and Fe³⁺ is [Ar] 3d⁵. Do not write [Ar] 3d⁴ 4s² by taking electrons out of 3d.

Why? While electrons are being added, 4s is lower in energy. But once 3d holds electrons, the 3d energy drops below 4s. At that point 4s (n = 4) is the outermost subshell, furthest from the nucleus, so its electrons are the easiest to remove.

From configuration to position in the periodic table

Know the configuration and you can place an element in the periodic table without memorising the table at all.

The period is the highest shell number (largest n). The block is the subshell that receives the last electron: s, p, d or f. The group (IUPAC 1–18) comes from the outer electrons: for the s block it equals the number of outer s electrons; for the p block it is 10 + outer s and p electrons (that is, 12 + the p electrons); for the d block it is the outer s electrons plus the (n − 1)d electrons.

Example: chlorine is 1s² 2s² 2p⁶ 3s² 3p⁵. The largest n is 3, so it is in period 3. The last electron goes into 3p, so it is p-block, group 17. For iron the largest n is 4 and the last electron enters 3d, so it is d-block, group 2 + 6 = 8.

Valence electrons are the electrons in the outermost shell; they are the ones that take part in bonding. The valency (how many bonds an atom forms) is usually that number if it is 1 to 4, or 8 minus that number if it is 5 to 7. So sodium and chlorine both have valency 1, oxygen 2 and nitrogen 3. A full outer shell of 8 (2 for helium) makes a noble gas.

The stability of half-filled and filled subshells shows up across the table too. Nitrogen’s 2p³ is half full, so nitrogen has a higher first ionisation energy than its neighbour oxygen. Noble gases have every subshell full, so they hardly react at all.

Key terms at a glance

All the vocabulary of this chapter in one table. Reading just this the night before a test goes a long way.

Try it yourself in the simulation

Work through these in order. Before each one, predict the answer and write it down, then check it in the simulation.

Solved problems

You can check every answer in the simulation. The first problem uses the simulation’s starting atom, iron.

Electron configuration chart for elements 1 to 36

Every row below is generated by the same engine the simulation uses. Look out for the two exceptions, Cr (24) and Cu (29). The shorthand is written in shell order, 3d before 4s.

Mistakes almost everyone makes

These cost the most marks. The correct version is next to each one.

• Writing potassium as 2, 8, 9. It is 2, 8, 8, 1, because the 19th electron goes into 4s, not 3d.
• Writing Cr as [Ar] 3d⁴ 4s² and Cu as [Ar] 3d⁹ 4s². They are 3d⁵ 4s¹ and 3d¹⁰ 4s¹.
• Taking electrons out of 3d for Fe²⁺. 4s empties first, so Fe²⁺ = [Ar] 3d⁶.
• Forgetting Hund and drawing 2p³ as ↑↓ ↑ _. It is ↑ ↑ ↑.
• Drawing two ↑ arrows in one orbital. By Pauli, a pair is always ↑↓.
• Not checking that the superscripts add up. They must total Z (or Z minus the charge for an ion).
• Mixing up orbital and subshell. 3d is one subshell containing five orbitals.
• Calling a Bohr-model circular orbit an orbital. An orbital is a probability cloud, not a fixed path.

Electron configuration in real life

Firework colours: sodium burns yellow, copper green-blue and strontium red. Heat pushes electrons up to higher levels, and as they drop back they give out light of a particular colour. That colour depends on the energy gaps between levels, which come straight from the electron configuration. Yellow sodium street lamps glow for the same reason.

Magnets: iron, cobalt and nickel have many unpaired d electrons. When their spins line up, countless tiny magnets add up to a big one. From the magnet on your fridge to the generators in a power station, unpaired electrons are doing the work.

Blood: haemoglobin carries oxygen using Fe²⁺ ions. If Fe²⁺ turns into Fe³⁺, haemoglobin can no longer carry oxygen. Why Fe³⁺ forms so easily is explained by the stability of 3d⁵.

Phones and computers: silicon has four outer electrons. Adding a trace of phosphorus (five outer electrons) or boron (three) turns it into a semiconductor, the material every chip is made from. And neon signs glow because neon’s full configuration keeps it unreactive, lighting up only when electricity excites its electrons.

Exam corner

Electron configuration appears in almost every chemistry syllabus, from GCSE and middle-school science to IB, AP and A-level chemistry, and it underpins the periodic table, bonding and transition-metal chapters. The same question types come up everywhere.

• “Write the full and shorthand configuration of X.” Always check that the superscripts add up to Z.
• “Which fills first, 4s or 3d?” 4s. “Which empties first when a transition-metal ion forms?” Also 4s.
• “Explain why Cr is 3d⁵ 4s¹.” Mention the extra stability of a half-filled subshell: symmetry and exchange energy.
• “Draw the orbital (box) diagram.” Label each group of boxes with its subshell, and follow Hund before pairing.
• “How many unpaired electrons, and is it paramagnetic?” Count lone arrows; any unpaired electron means paramagnetic.
• “Deduce the period, group and block.” Largest n, the subshell of the last electron, and the outer electrons.
• “Maximum electrons in shell n = 3?” 2n² = 18. “Orbitals when l = 2?” 2l + 1 = 5.

One-screen revision

Read this list before a test and the whole chapter comes back.

• Electron configuration = how electrons are arranged in shells, subshells and orbitals.
• Shells K, L, M, N; at most 2n² electrons; n² orbitals.
• Subshells s, p, d, f; 1, 3, 5, 7 orbitals; 2, 6, 10, 14 electrons.
• Filling order: 1s 2s 2p 3s 3p 4s 3d 4p; lower n + l first, and if equal, lower n first.
• Pauli: at most 2 per orbital, with opposite spins.
• Hund: one each with parallel spins, then pair.
• Exceptions: Cr = [Ar] 3d⁵ 4s¹, Cu = [Ar] 3d¹⁰ 4s¹.
• Cations lose 4s before 3d: Fe²⁺ = [Ar] 3d⁶, Fe³⁺ = [Ar] 3d⁵.
• Period = largest n; block = subshell of the last electron; group from the outer electrons.

Frequently asked questions