Electron Configuration Calculator

Determine electron configurations for elements and ions. Visualize orbital filling order and electron distribution across energy levels for chemistry education.

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Atomic Number (Z)

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Full Configuration

Noble Gas Notation

About Electron Configuration Calculator

Quantum Mechanical Foundation

Electron configuration theory emerged from early quantum mechanics, developed by pioneers like Niels Bohr (1913), Wolfgang Pauli (1925), and Friedrich Hund (1927). Their work revolutionized our understanding of atomic structure and chemical bonding, establishing the quantum mechanical model of the atom.

Noble Gas Notation

Noble Gas notation is a shorthand way to write electron configurations using the symbol of the nearest noble gas with fewer electrons than the element in question. This method simplifies writing long configurations while emphasizing the valence electrons that are most important for chemical behavior.

For example:

Na (Z=11): [Ne]3s¹ instead of 1s²2s²2p⁶3s¹
Fe (Z=26): [Ar]4s²3d⁶ instead of 1s²2s²2p⁶3s²3p⁶4s²3d⁶
Au (Z=79): [Xe]6s¹4f¹⁴5d¹⁰ instead of writing out all 79 electrons

Fundamental Principles

n = principal quantum number (1,2,3,...)
l = angular momentum (s,p,d,f)
ml = magnetic quantum number
ms = spin quantum number (±½)

Aufbau Principle: Electrons fill orbitals in order of increasing energy, following the n + l rule. This ensures the most stable electron configurations
Pauli Exclusion: No two electrons in an atom can have identical quantum numbers. Each orbital can hold maximum two electrons with opposite spins
Hund's Rule: Electrons in orbitals of equal energy (degenerate) will occupy separate orbitals and maintain parallel spins before pairing
Electronic Stability: Atoms prefer complete or half-filled subshells due to symmetric electron distribution and enhanced stability

Orbital Structure

s Orbitals:
- Perfect spherical shape with uniform electron density in all directions
- Single orbital representing the simplest atomic orbital shape
- Can hold exactly two electrons with opposite spins (↑↓)
- Equal probability of finding electrons anywhere at a given radius from nucleus
p Orbitals:
- Figure-eight or dumbbell shape with two lobes separated by a nodal plane
- Three perpendicular orbitals (px, py, pz) forming 90° angles
- Each orbital can hold two electrons, total six for complete p subshell
- Responsible for directional bonding in molecules along x, y, z axes
d Orbitals:
- Five different orbital shapes: dxy, dxz, dyz, dx²-y², dz²
- Complex spatial arrangements with four lobes each (except dz²)
- Critical for transition metal chemistry and colored compounds
- Involved in crystal field splitting and ligand interactions
f Orbitals:
- Seven different orbital shapes with complex angular distributions
- Responsible for unique magnetic and spectroscopic properties
- Found in rare earth elements and actinides
- Key to understanding inner transition metal chemistry

Aufbau Filling Order

Energy Level Sequence:
- Complete sequence: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f
- n + l rule determines orbital energy: lower sum fills first, then lower n if sums equal
- Exceptions occur due to electron-electron repulsion and nuclear charge effects
- Diagonal rule visualized through the Madelung energy diagram
Common Exceptions:
- Chromium [Ar]4s¹3d⁵: Enhanced stability from half-filled 3d and singly-filled 4s
- Copper [Ar]4s¹3d¹⁰: Extra stability from completely filled 3d sublevel
- Half-filled subshells provide extra stability due to symmetric electron distribution
- d-s electron shifting occurs in many transition metals due to orbital energy overlap

Chemical Properties

Periodic Trends:
- Atomic size decreases across periods, increases down groups due to electron shielding
- Ionization energy increases across periods, decreases down groups with atomic size
- Electron affinity generally increases across periods with some exceptions
- Electronegativity follows diagonal trend, highest in upper right of periodic table
Valence Electrons:
- Determine primary bonding behavior and chemical reactivity patterns
- Define possible oxidation states based on electron loss/gain potential
- Influence molecular geometry through electron pair repulsion
- Control electrical conductivity in metals and semiconductors

Applications

Chemistry:
- Chemical bonding: Predicts bond types, strengths, and molecular geometries
- Molecular orbitals: Forms basis for understanding complex molecular interactions
- Spectroscopy: Explains atomic spectra and energy level transitions
- Reaction mechanisms: Helps predict reactive sites and electron transfer processes
Materials Science:
- Band theory: Explains conductivity in metals, semiconductors, and insulators
- Semiconductor properties: Crucial for designing electronic devices and solar cells
- Magnetic materials: Determines ferromagnetic, paramagnetic, and diamagnetic behavior
- Crystal field theory: Explains color and magnetic properties of transition metal compounds

Learn More

Electron Configuration Atomic Orbitals Aufbau Principle