ELECTRON CONFIGURATION
In lessons 2.1 and 2.2 we talked about shells and valence electrons. Now we go one level deeper: a precise address system that tells you exactly where every single electron in an atom lives — not just which shell, but which subshell within that shell.
A RECURRING THEME IN CHEMISTRY
Systems in nature tend toward their lowest possible energy state. A ball rolls downhill, a stretched spring contracts, and electrons are no different — given a choice, they always occupy the lowest energy level available before moving to a higher one. This single principle explains almost everything about how electron configurations are built.
ELECTRON SPIN
Electrons have an intrinsic quantum property called spin. Despite the name, electrons are not literally spinning — spin is a fundamental property of the particle itself, like charge or mass. What matters is that spin can only ever take one of two values, which we call spin-up (↑) and spin-down (↓).
This two-value property has a profound consequence: the Pauli exclusion principle states that no two electrons in the same atom can be in exactly the same quantum state. Since spin is the only thing that can differ between two electrons in the same orbital, each orbital can hold at most two electrons — one spin-up and one spin-down.
This is why every subshell capacity is an even number, and why the orbital box diagrams you will see below always show pairs of ↑↓ arrows — each pair represents one full orbital.
SUBSHELLS
Each electron shell is divided into subshells — regions of space with different shapes and different energies. Subshells are labelled with a number (the shell they belong to) and a letter (the type of subshell):
// INTERACTIVE — click a row · toggle 2D / 3D · select individual orbitals
Each orbital can hold exactly 2 electrons. So an s subshell with 1 orbital holds 2, a p subshell with 3 orbitals holds 6, and so on.
ORBITALS AS PROBABILITY CLOUDS
So where does the spherical shape of an s orbital actually come from? An electron doesn't sit still — it moves continuously around the nucleus. Because of quantum mechanics, we can't know exactly where it is at any moment, only the probability of finding it in a given region. The orbital shape is simply a map of that probability — denser where the electron spends more time. Use the animation below to build that intuition.
// INTERACTIVE — ORBITAL FORMATION ANIMATION
THE AUFBAU PRINCIPLE
"Aufbau" is German for building up. The principle is simple: electrons fill the lowest available energy subshell first, then the next, and so on. You build an atom's configuration by adding electrons one at a time into the cheapest available seat. Based on that, can you guess the order we should fill orbitals in? Give it a try in the diagram below.
CHALLENGE
Click each subshell in the correct filling order — lowest energy first.
The filling order is not simply 1s → 2s → 2p → 3s → 3p → 3d, because the 4s subshell is actually slightly lower in energy than 3d. The correct full filling order is:
FILLING ORDER
Notice that 4s comes before 3d because it's lower in energy.
There is one more rule to know: when filling a subshell that has multiple orbitals — like the three orbitals in a p subshell — electrons spread out into empty orbitals first, one per orbital, before any pairing occurs. Only once every orbital in the subshell has one electron do they start doubling up. This is called Hund's rule, and it happens because electrons repel each other — occupying separate orbitals keeps them further apart and lowers the overall energy.
EXERCISE — HUND'S RULE
Place electrons into the 2p subshell — fill empty orbitals before pairing.
Carbon— place 2 electrons into the 2p subshell
0 / 2 electrons placed — click an orbital to add an electron
UP NEXT
Now that we know the rules for filling orbitals, lesson 2.4 covers how to write and read the notation — the precise shorthand chemists use to record every electron in an atom. We'll also look at the s/p/d blocks and noble gas abbreviated form.