Electron Configuration Calculator
Select any element and optional ion charge to get its complete ground-state electron configuration, noble-gas shorthand notation, valence electron count, and the electron distribution across shells. All 118 elements are supported, including transition metals with known Aufbau exceptions such as chromium, copper, molybdenum, silver, palladium, gold, and platinum. Cations and anions are handled using standard ionization rules.
What is electron configuration?
Electron configuration is a notation that describes how electrons are distributed among the atomic orbitals of an atom. Each electron occupies a specific subshell defined by a principal quantum number (n = 1, 2, 3 ...) and an azimuthal quantum number (l = 0 for s, 1 for p, 2 for d, 3 for f). The notation lists each occupied subshell followed by a superscript showing how many electrons it contains, for example 1s2 2s2 2p6. Electron configuration determines an element's chemical behaviour, its position on the periodic table, the types of bonds it forms, and its magnetic properties.
The three rules for filling electrons
Three fundamental principles govern how electrons fill orbitals. The Aufbau principle states that electrons enter orbitals of lowest energy first, filling them in the order 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, and so on (the diagonal or Madelung rule). The Pauli exclusion principle states that each orbital holds at most two electrons, and they must have opposite spins (spin-up and spin-down). Hund's rule states that when multiple orbitals of equal energy (degenerate orbitals) are available, electrons occupy each one singly before any pairing occurs, and the unpaired electrons align their spins in the same direction. Together these three rules determine the ground-state configuration of any element.
Noble-gas shorthand notation
Writing out the full configuration for a heavy element is cumbersome. The noble-gas shorthand replaces the completely filled inner core with the symbol of the preceding noble gas in square brackets, then lists only the valence-shell orbitals. For example, the full configuration of iron is 1s2 2s2 2p6 3s2 3p6 3d6 4s2. Because the first 18 electrons match the configuration of argon, this is abbreviated as [Ar] 3d6 4s2. This makes it easy to compare the outer-electron arrangements that drive chemical reactivity.
Aufbau exceptions and why they occur
A handful of elements, most notably chromium (Z=24) and copper (Z=29), and their heavier analogues molybdenum, silver, gold, and palladium, deviate from the expected Aufbau filling order. For chromium, the predicted configuration would be [Ar] 3d4 4s2, but the observed configuration is [Ar] 3d5 4s1. For copper it is [Ar] 3d10 4s1 rather than [Ar] 3d9 4s2. These deviations arise because exactly half-filled (d5) and fully filled (d10) d subshells have extra exchange-energy stabilisation that outweighs the small energy cost of promoting one electron from the 4s into the 3d. Similar reasoning applies to palladium ([Kr] 4d10 with no 5s electrons), gold ([Xe] 4f14 5d10 6s1), and platinum ([Xe] 4f14 5d9 6s1). These exceptions must be memorised because simple energy-diagram rules do not predict them reliably.
Ions: cations and anions
When an atom gains or loses electrons it becomes an ion. A cation has fewer electrons than protons (positive charge); an anion has more electrons than protons (negative charge). To write an ion's configuration, start from the neutral atom's ground-state configuration and then remove or add electrons. For main-group elements, remove electrons from the outermost shell and add them to the next available orbital. For transition metals the rule is more nuanced: when forming a cation, electrons are removed from the highest principal quantum number shell first, and specifically from the ns subshell before the (n-1)d subshell, even though 4s fills before 3d in the neutral atom. So Fe2+ loses both 4s electrons and becomes [Ar] 3d6, not [Ar] 3d4 4s2.
Subshell capacities and Aufbau filling order
| Filling order | Subshell | Max electrons | Running total |
|---|---|---|---|
| 1 | 1s | 2 | 2 |
| 2 | 2s | 2 | 4 |
| 3 | 2p | 6 | 10 |
| 4 | 3s | 2 | 12 |
| 5 | 3p | 6 | 18 |
| 6 | 4s | 2 | 20 |
| 7 | 3d | 10 | 30 |
| 8 | 4p | 6 | 36 |
| 9 | 5s | 2 | 38 |
| 10 | 4d | 10 | 48 |
| 11 | 5p | 6 | 54 |
| 12 | 6s | 2 | 56 |
| 13 | 4f | 14 | 70 |
| 14 | 5d | 10 | 80 |
| 15 | 6p | 6 | 86 |
| 16 | 7s | 2 | 88 |
| 17 | 5f | 14 | 102 |
| 18 | 6d | 10 | 112 |
| 19 | 7p | 6 | 118 |
Each subshell holds a maximum of 2(2l+1) electrons. Fill in the order shown.
Frequently asked questions
What is the electron configuration of iron (Fe)?
The full electron configuration of iron (Z=26) is 1s2 2s2 2p6 3s2 3p6 3d6 4s2. The noble-gas shorthand is [Ar] 3d6 4s2. Iron has 8 valence electrons. As Fe2+, both 4s electrons are removed first, giving [Ar] 3d6. As Fe3+, one additional 3d electron is removed, giving [Ar] 3d5.
Why is chromium an Aufbau exception?
Chromium (Z=24) is expected to be [Ar] 3d4 4s2 based on simple Aufbau order, but spectroscopy shows the actual ground state is [Ar] 3d5 4s1. The half-filled 3d subshell (five unpaired electrons) has extra exchange-energy stability, and the energy gained by achieving this configuration outweighs the small cost of moving one electron from 4s to 3d. Copper (Z=29) is similar: it achieves a full 3d10 by promoting one 4s electron, giving [Ar] 3d10 4s1 instead of [Ar] 3d9 4s2.
How do I write the configuration of a cation?
Start with the neutral atom's ground-state configuration, then remove electrons beginning with the outermost principal shell. For transition metals specifically, remove s electrons before d electrons. For example, Mn2+ (Z=25, charge +2) starts from [Ar] 3d5 4s2 and loses both 4s electrons to become [Ar] 3d5. The key is that 4s electrons are higher in energy in the presence of the 3d electrons, so they are ionised first even though they fill first in the neutral atom.
What are valence electrons and why do they matter?
Valence electrons are the electrons in the outermost shell (highest principal quantum number n) of an atom. They are the ones that participate in chemical bonding, determine an element's reactivity, and govern what kinds of ions or molecules it can form. Main-group elements have 1 to 8 valence electrons matching their group number in the periodic table. For example, oxygen (group 16) has 6 valence electrons and typically forms 2 bonds, while sodium (group 1) has 1 valence electron and readily loses it to form Na+.
What is the difference between the full configuration and noble-gas shorthand?
The full electron configuration lists every subshell from 1s onward, for example 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5 for bromine. The noble-gas shorthand replaces the completely filled inner core (which matches the configuration of the preceding noble gas) with that noble gas symbol in brackets, leaving only the outer electrons: [Ar] 3d10 4s2 4p5 for bromine. Both are equivalent; the shorthand simply saves writing and makes the valence electrons immediately visible.
How many electrons can each subshell hold?
s subshells hold up to 2 electrons (1 orbital x 2 spins), p subshells hold up to 6 (3 orbitals x 2), d subshells hold up to 10 (5 orbitals x 2), and f subshells hold up to 14 (7 orbitals x 2). This follows from the general formula 2(2l+1), where l is the azimuthal quantum number (0 for s, 1 for p, 2 for d, 3 for f).