Generate electron configurations, noble gas notation, and orbital diagrams for any element using the Aufbau principle and Hund's rule.
Enter an atomic number to determine its electron configuration using the Aufbau principle.
Enter a value from 1 (Hydrogen) to 118 (Oganesson)
Aufbau Principle: Electrons fill orbitals starting from the lowest energy level to higher energy levels following a specific order.
Exceptions: Some elements like Chromium (Cr) and Copper (Cu) have unusual configurations because half-filled and fully-filled d orbitals are more stable.
An electron configuration describes how electrons are distributed among atomic orbitals in an atom. It follows specific principles that govern how electrons fill available orbitals, determining an element's chemical properties and position in the periodic table.
Electrons fill orbitals from lowest to highest energy: 1s → 2s → 2p → 3s → 3p → 4s → 3d...
Each orbital holds maximum 2 electrons with opposite spins (↑↓).
Electrons occupy degenerate orbitals singly before pairing up (all spins parallel first).
Number (1, 2, 3...): Principal energy level (shell)
Letter (s, p, d, f): Orbital type (sublevel)
Superscript (², ⁶, ¹⁰): Number of electrons in that orbital
Electrons fill orbitals in order of increasing energy. The "diagonal rule" or Aufbau diagram helps visualize this sequence:
1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p
Notice: 4s fills before 3d, 5s before 4d, 6s before 4f and 5d!
s orbitals
1 orbital
Max: 2e⁻
p orbitals
3 orbitals
Max: 6e⁻
d orbitals
5 orbitals
Max: 10e⁻
f orbitals
7 orbitals
Max: 14e⁻
Step 1: Carbon has 6 electrons to distribute
Step 2: Fill orbitals following Aufbau principle
Result:
Or in noble gas notation: [He] 2s² 2p²
Iron has 26 electrons. Following the Aufbau order:
Full configuration:
Noble gas notation:
[Ar] represents 1s² 2s² 2p⁶ 3s² 3p⁶ (18 electrons)
Some elements have unexpected configurations because half-filled and fully-filled d orbitals are more stable than predicted by the Aufbau principle alone.
| Element | Expected | Actual | Reason |
|---|---|---|---|
| Cr (24) | [Ar] 4s² 3d⁴ | [Ar] 4s¹ 3d⁵ | Half-filled d orbital stability |
| Cu (29) | [Ar] 4s² 3d⁹ | [Ar] 4s¹ 3d¹⁰ | Fully-filled d orbital stability |
| Ag (47) | [Kr] 5s² 4d⁹ | [Kr] 5s¹ 4d¹⁰ | Fully-filled d orbital stability |
| Au (79) | [Xe] 6s² 4f¹⁴ 5d⁹ | [Xe] 6s¹ 4f¹⁴ 5d¹⁰ | Fully-filled d orbital stability |
Half-filled (d⁵) and fully-filled (d¹⁰) d subshells provide extra stability due to:
• Symmetry: Evenly distributed electron density
• Exchange Energy: Electrons with parallel spins have lower energy
• Minimized Repulsion: Balanced electron-electron interactions
Valence electrons (outermost electrons) determine how atoms bond. Elements in the same group have similar configurations, explaining periodic trends in reactivity.
Electrons transition between energy levels, absorbing or emitting specific wavelengths of light. This creates unique spectral "fingerprints" for element identification.
Unpaired electrons create magnetic moments. Knowing the electron configuration helps predict whether a substance is paramagnetic (attracted to magnets) or diamagnetic (repelled).
Atoms lose or gain electrons to achieve stable noble gas configurations. Understanding configurations predicts ionic charges (e.g., Na → Na⁺, Cl → Cl⁻).
The periodic table's structure directly reflects electron configurations. Groups share the same valence configuration, explaining similar chemical properties.
Electronic structure determines conductivity, color, hardness, and other material properties. Transition metals with d electrons show unique catalytic behavior.