The seating chart of an atom
Every atom heavier than hydrogen has more than one electron, and they have to go *somewhere*. The list of which orbitals an atom's electrons occupy is its electron configuration — essentially a seating chart for electrons. We write it as a short code: oxygen, with eight electrons, is written 1s² 2s² 2p⁴. Read it gently. Each cluster says *which orbital* (the 1s, the 2s, a 2p) and the little raised number says *how many electrons sit there*. Add up the superscripts — 2 + 2 + 4 — and you get eight, exactly oxygen's electron count. The whole code is just an inventory of where everyone is sitting.
Why should you care about a seating chart? Because almost everything an element does — whether it is a reactive metal or an inert gas, what it bonds to, what colour it is — is set by how its electrons are arranged, especially the outermost ones. Get the configuration right and you can often predict the chemistry before you ever touch the substance. So our job is simply to learn the rules that decide the seating. There are only three.
Rule one: lowest seats first
The first rule is pure common sense. Water settles to the lowest point it can reach; a marble rolls to the bottom of a bowl. Electrons do the same with energy: each new electron drops into the lowest-energy seat still available. Fill the cheapest seats first, then the next cheapest, and only move up when everything below is taken. This building-up-from-the-bottom procedure is the aufbau principle — *aufbau* is just German for *building up*. It is less a law of physics than a tidy instruction for filling seats in price order.
Rule two: two to a seat, and no two alike
If lowest-first were the only rule, every electron would pile into the 1s orbital and atoms would be dull. They do not, because of a strict bouncer at the door called the Pauli exclusion principle. It says: no two electrons in the same atom may have the identical set of four quantum numbers. Recall that an electron's address is three numbers for its orbital plus a fourth for its spin, up or down. Two electrons can share the same orbital — same first three numbers — only if they differ in the fourth. So an orbital holds at most two electrons, one spin-up and one spin-down, and not a single one more.
This one limit — two per orbital — is what gives the atom its layered structure. Because the lowest seats fill and lock at two each, later electrons are *forced* outward into roomier, higher shells. That outward pressure is why a sodium atom is bigger than a hydrogen atom, why the shells fill up in the patterns they do, and ultimately why the periodic table has the shape it has. Without Pauli's bouncer there would be no rows, no columns, no chemistry — just every electron crushed into the basement.
Rule three: spread out before pairing up
The third rule settles a tie. A subshell often has several orbitals at the same energy level — the three p dumbbells, for instance, all cost the same. So when electrons enter a set of equal-energy orbitals, which do they take? Hund's rule gives the answer: spread out first. Electrons go into separate orbitals singly, all with the same spin direction, before any of them doubles up to share. They behave like strangers boarding an empty bus, each taking their own row before anyone agrees to sit beside someone else.
There is a real reason for the bus-seat behaviour: electrons all carry negative charge and repel one another, so they are *more comfortable*, lower in energy, when they keep to separate orbitals rather than crowding two into one. Spreading out keeps them apart. The rule is not arbitrary etiquette — it is electrons doing what costs them the least energy, just like rule one.
Putting the three rules to work
- Count the electrons. The atom is neutral, so the number of electrons equals the number of protons — the element's atomic number. Carbon has 6.
- Fill lowest-energy seats first (aufbau). For carbon's six: two go into 1s, two into 2s, leaving two for the 2p subshell.
- Cap every orbital at two electrons with opposite spins (Pauli). So 1s and 2s are each full with their pair.
- Spread the last electrons across equal orbitals before pairing (Hund). Carbon's two 2p electrons go into two different p orbitals, same spin — not paired in one.
- Read off the configuration: carbon is 1s² 2s² 2p². The same four steps, repeated, generate every element in order.