Electron Configuration of the 3d Series
Simple Explanation
The 3d series (scandium to zinc) fills its 3d subshell after the 4s subshell — but because the two subshells are so close in energy, a few elements (notably chromium and copper) have irregular configurations that don't follow the simple filling pattern.
Why Do We Need It?
The exact electron configuration — and especially how easily 4s and 3d electrons are lost — explains almost everything distinctive about transition metals: their variable oxidation states, their coloured compounds, and their catalytic behaviour.
Why Does This Work?
The 4s and 3d subshells are unusually close in energy, so the normal "fill lowest energy first" rule (4s before 3d) sometimes gets disrupted by the extra stability of a half-filled or completely filled 3d subshell — this is exactly why chromium adopts [Ar]3d⁵4s¹ (a half-full 3d) instead of the "expected" [Ar]3d⁴4s², and copper adopts [Ar]3d¹⁰4s¹ (a full 3d) instead of [Ar]3d⁹4s².
Real-Life Example
Iron's electron configuration and rust
Iron commonly forms both Fe²⁺ and Fe³⁺ compounds — for example, in the two different iron oxides that make up rust.
Iron's configuration, [Ar]3d⁶4s², loses its two 4s electrons easily to form Fe²⁺, and can lose one additional 3d electron to form the also-stable Fe³⁺ — this flexibility in how many electrons are lost is a direct consequence of its electron configuration.
Practice
What is unusual about the electron configuration of chromium (Cr, atomic number 24)?
MediumCommon mistake
Assuming EVERY transition element follows the simple 'fill 4s then 3d' pattern without exception — chromium and copper are well-known, important exceptions caused by the extra stability of half-filled and fully-filled 3d subshells.
Quick Review
- The 3d series fills the 3d subshell after 4s, with some irregularities.
- Cr: [Ar]3d⁵4s¹ and Cu: [Ar]3d¹⁰4s¹ are the classic exceptions.
- Electron configuration explains transition metals' variable oxidation states and other properties.