⚛️ Electron Configuration Generator – Orbital Diagrams & Notation
The Electron Configuration Generator instantly produces the complete electron configuration for any of the 118 elements — as a standard orbital notation string, a noble-gas shorthand, and a colour-coded box diagram showing individual spin states. Just type a symbol like Fe or a full name like Iron, optionally specify an ion charge, and choose your preferred display format.
📘 What Is Electron Configuration?
Electron configuration describes how an atom's electrons are distributed across its atomic orbitals. Each orbital is characterised by a principal quantum number (n), a subshell type (s, p, d, or f), and the number of electrons it contains. The configuration determines nearly every chemical property of an element — its reactivity, oxidation states, magnetic behaviour, and position in the periodic table.
⚙️ The Three Principles Behind the Tool
Aufbau Principle
Electrons fill orbitals from lowest to highest energy. The filling order follows the (n + l) rule: subshells with lower (n + l) values fill first, and when equal, the lower n fills first.
Hund's Rule
Within a subshell, electrons occupy separate orbitals with parallel spins before any pairing occurs. This minimises electron–electron repulsion and gives the atom its lowest energy state.
Pauli Exclusion Principle
No two electrons in the same atom can have identical quantum numbers. In practice this means each orbital holds at most two electrons, which must have opposite spins (↑↓).
🔢 Aufbau Filling Order
Electrons are added to subshells in the following sequence (the diagonal mnemonic):
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Each subshell can hold a fixed maximum: s = 2, p = 6, d = 10, f = 14 electrons.
📐 Worked Example – Iron (Fe, Z = 26)
Iron has 26 electrons. Following the Aufbau sequence:
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶ (full notation)
[Ar] 3d⁶ 4s² (noble gas shorthand)
Box diagram for 3d⁶ (5 orbitals, Hund's rule):
3d: [↑↓] [↑ ] [↑ ] [↑ ] [↑ ]The four singly occupied 3d orbitals make iron paramagnetic — it is attracted to magnets because of its unpaired electrons.
⚡ Transition Metal Exceptions
Several transition metals deviate from pure Aufbau predictions because half-filled or completely filled d subshells are energetically favourable:
| Element | Expected (Aufbau) | Actual Configuration | Reason |
|---|---|---|---|
| Cr (Z=24) | [Ar] 3d⁴ 4s² | [Ar] 3d⁵ 4s¹ | Half-filled 3d⁵ stability |
| Cu (Z=29) | [Ar] 3d⁹ 4s² | [Ar] 3d¹⁰ 4s¹ | Full 3d¹⁰ stability |
| Mo (Z=42) | [Kr] 4d⁴ 5s² | [Kr] 4d⁵ 5s¹ | Half-filled 4d⁵ stability |
| Au (Z=79) | [Xe] 4f¹⁴ 5d⁹ 6s² | [Xe] 4f¹⁴ 5d¹⁰ 6s¹ | Full 5d¹⁰ stability |
🔬 Noble Gas Shorthand
To save writing, chemists replace the filled inner-core configuration with the symbol of the preceding noble gas in square brackets. Only the valence-region electrons — the ones that participate in bonding — are written out explicitly:
Na (Z=11): 1s² 2s² 2p⁶ 3s¹ → [Ne] 3s¹
Ca (Z=20): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² → [Ar] 4s²
Fe (Z=26): ... → [Ar] 3d⁶ 4s²⚗️ Ions and Electron Configuration
When an atom loses electrons (forms a cation), electrons are removed starting from the highest principal quantum number shell. For transition metals, the outer ns electrons are removed before the (n−1)d electrons. When an atom gains electrons (forms an anion), electrons are added by continuing the Aufbau sequence. For example:
Fe²⁺: [Ar] 3d⁶ (remove both 4s electrons from Fe)
Fe³⁺: [Ar] 3d⁵ (remove one more from 3d)
O²⁻: 1s² 2s² 2p⁶ (add 2 electrons to O's 2p subshell)🎯 Orbital Colour Code
The box diagram uses colour-coded borders to distinguish orbital types at a glance: s (blue), p (green), d (orange), f (purple). Each box represents one individual orbital. An upward arrow (↑) means spin-up, a downward arrow (↓) means spin-down, and an empty box means the orbital is unoccupied.
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