ELECTRON CONFIGURATION

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.

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):

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.

For most of general chemistry — and all elements up through krypton (Z=36) — you only need s, p, and d. The f subshell only comes into play for the lanthanides and actinides far down the periodic table.

"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.

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:

Notice that 4s comes before 3d. This is why potassium (Z=19) fills its 4s subshell before touching the 3d, giving it the configuration [Ar] 4s¹ rather than [Ar] 3d¹.

An electron configuration lists each occupied subshell in order, with a superscript showing how many electrons it contains. The format is:

For example, carbon (Z=6) has 6 electrons to place. The first 2 go into 1s, the next 2 fill 2s, and the final 2 begin filling 2p:

Step through the explorer below to see how configurations build up across the first 36 elements.

The shape of the periodic table is not arbitrary — it directly reflects which subshell is being filled as you move across each period. This divides the table into named blocks:

The block an element belongs to tells you which subshell its highest-energy electrons occupy. This is directly related to its chemical behaviour — the valence electrons we studied in lesson 2.2 are always the outermost s and p electrons (for s and p block elements), which is why Group number predicts valence electron count so cleanly.

Writing out the full configuration for every element gets tedious fast. Because noble gases have completely filled, stable electron arrangements, chemists use them as shorthand. Instead of writing everything from 1s onward, you write the previous noble gas in square brackets, then list only the electrons added after it.

The abbreviated form is what you'll most often see in textbooks and exams. It strips away the stable, chemically inert core and highlights only the electrons that actually participate in bonding — exactly the valence electrons we studied in lesson 2.2.