If an atom were a building, the electron configuration would be its floor plan — it tells you exactly which rooms (orbitals) the electrons occupy and how many are in each. Electrons do not just swirl randomly around the nucleus; they fill specific energy levels in a precise order, much like water filling containers from the lowest point upward. Hydrogen's single electron sits in the 1s orbital (written as 1s¹). Carbon has six electrons arranged as 1s² 2s² 2p², and iron has 26 electrons filling orbitals all the way up to 3d⁶. This notation might look like algebra, but it is really a powerful address system for electrons.
The orbitals come in four types — s, p, d, and f — each with a distinctive shape. The s orbital is spherical, p orbitals are dumbbell-shaped, d orbitals look like four-leaf clovers, and f orbitals are even more complex. Electrons fill these orbitals following three key rules: the Aufbau principle (fill lowest energy first), the Pauli exclusion principle (maximum two electrons per orbital, with opposite spins), and Hund's rule (spread out before pairing up in equal-energy orbitals). These rules explain why the periodic table has its distinctive shape — the s block is 2 columns wide, the p block is 6 columns, the d block is 10, and the f block is 14.
Why does this matter in real life? Electron configuration directly explains an element's chemical behavior. Elements in the same group of the periodic table have the same outer (valence) electron configuration, which is why they react similarly. Sodium (1s² 2s² 2p⁶ 3s¹) and potassium (1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹) both have one valence electron, so both are reactive alkali metals. The brilliant colors of transition metal compounds — the blue of copper sulfate, the purple of potassium permanganate — arise from d-electron transitions between energy levels. Understanding electron configuration is like having the blueprint for all of chemistry.