Electron Configuration Calculator
Look up the electron configuration of any element by name, symbol or atomic number. Shows shorthand and full notation, orbital filling diagram, valence electrons and block classification for all 118 elements.
Common elements:
Electron Configuration
Search for an element to see its configuration
Aufbau Filling Order
1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p
All 118 Elements
Complete database from hydrogen (Z=1) through oganesson (Z=118) with correct configurations including all exceptions like chromium and copper.
Orbital Diagrams
Visual box diagrams showing electron spin arrows in each orbital. See exactly how electrons fill according to Hund's rule and the Pauli exclusion principle.
Properties at a Glance
Block (s/p/d/f), period, group, valence electron count, unpaired electron count and both shorthand and full configuration displayed instantly.
Common Electron Configurations
| Element | Z | Configuration | Valence e- | Block |
|---|---|---|---|---|
| Hydrogen | 1 | 1s1 | 1 | s-block |
| Carbon | 6 | [He] 2s2 2p2 | 4 | p-block |
| Oxygen | 8 | [He] 2s2 2p4 | 6 | p-block |
| Sodium | 11 | [Ne] 3s1 | 1 | s-block |
| Chlorine | 17 | [Ne] 3s2 3p5 | 7 | p-block |
| Chromium | 24 | [Ar] 3d5 4s1 | 1 | d-block |
| Iron | 26 | [Ar] 3d6 4s2 | 2 | d-block |
| Copper | 29 | [Ar] 3d10 4s1 | 1 | d-block |
| Silver | 47 | [Kr] 4d10 5s1 | 1 | d-block |
| Gold | 79 | [Xe] 4f14 5d10 6s1 | 1 | d-block |
How to Find Electron Configuration
Electron Configuration: Complete Guide
What Is Electron Configuration?
Electron configuration describes how electrons are distributed among the orbitals of an atom. Each electron occupies a specific orbital defined by quantum numbers: the principal quantum number (n, the shell), the angular momentum number (l, the subshell type: s, p, d or f), the magnetic quantum number (ml, the specific orbital), and the spin quantum number (ms, up or down). The configuration determines nearly everything about how an element behaves chemically.
The Three Rules
Three rules govern how electrons fill orbitals:
- Aufbau principle: Electrons fill orbitals starting from the lowest available energy level. The filling order follows the diagonal rule: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, and so on. The 4s orbital fills before 3d because it has slightly lower energy in multi-electron atoms.
- Pauli exclusion principle: No two electrons in an atom can have the same four quantum numbers. This means each orbital holds at most 2 electrons, and they must have opposite spins (one up, one down).
- Hund's rule: When filling degenerate orbitals (orbitals with equal energy, like the three 2p orbitals), electrons spread out first with parallel spins before pairing up. This minimizes electron repulsion and stabilizes the atom.
Notable Exceptions
Most elements follow the standard Aufbau filling order, but several transition metals are exceptions. The most commonly tested ones are:
- Chromium (Z=24): Expected [Ar] 3d4 4s2, actual [Ar] 3d5 4s1. A half-filled d subshell is extra stable.
- Copper (Z=29): Expected [Ar] 3d9 4s2, actual [Ar] 3d10 4s1. A fully-filled d subshell is extra stable.
- Palladium (Z=46): Expected [Kr] 4d8 5s2, actual [Kr] 4d10. All 10 d electrons and no 5s electrons at all.
- Gold (Z=79): Expected [Xe] 4f14 5d9 6s2, actual [Xe] 4f14 5d10 6s1. Same half/full shell stabilization.
These exceptions happen because the energy difference between the ns and (n-1)d orbitals is very small in transition metals, and the extra stability of half-filled or fully-filled d subshells tips the balance.
Electron Configuration and the Periodic Table
The periodic table is organized by electron configuration. Each block corresponds to the type of orbital being filled: s-block (groups 1-2), p-block (groups 13-18), d-block (groups 3-12, the transition metals), and f-block (the lanthanides and actinides). Moving left to right across a period, you add one electron at a time. Moving down a group, you add a new principal energy level.
This is why elements in the same group have similar chemical properties. They have the same valence electron configuration, just in a higher shell. Sodium (3s1) and potassium (4s1) both have one valence electron and both form +1 ions.
Ions and Electron Configuration
When atoms form ions, the electron configuration changes. For cations (positive ions), remove electrons starting from the orbital with the highest principal quantum number, not the last orbital filled. For transition metals, this means removing s electrons before d electrons. Fe2+ is [Ar] 3d6, not [Ar] 3d4 4s2. For anions (negative ions), add electrons following the normal Aufbau filling order.
Why Electron Configuration Matters
Electron configuration predicts chemical behavior, bonding, reactivity, color, magnetism, and ionization energy. It explains why noble gases are inert (full outer shells), why alkali metals are highly reactive (one loosely held valence electron), why transition metals form colored compounds (partially filled d orbitals), and why iron is magnetic (four unpaired 3d electrons).