Electron Configuration Builder

What is electron configuration?

Every neutral atom has as many electrons as protons. Those electrons do not sit randomly: they occupy quantised energy levels called shells, which contain subshells (s, p, d, f), each of which contains one or more orbitals. The configuration tells you exactly how many electrons are in each subshell and therefore predicts chemical behaviour — which electrons are available for bonding, which elements behave similarly, and why the periodic table has its shape.

This Electron Configuration Builder lets you select any of the 118 confirmed elements, see the accepted ground-state configuration, view orbital box diagrams with correct Hund and Pauli filling, and watch electrons placed step by step. Everything runs in your browser with no sign-up and no data sent to a server.

How electron configuration works

Three rules govern how electrons fill orbitals in neutral atoms. First, the Aufbau principle directs electrons into orbitals in order of increasing energy: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, and so on. Second, Hund's rule requires that when a subshell has several equivalent orbitals (such as the three 2p orbitals), each orbital gets one electron before any is paired. Third, the Pauli exclusion principle limits each orbital to two electrons with opposite spins.

Shell distribution is derived by adding the electrons in all subshells that share the same principal quantum number n. Carbon has 2 electrons in shell 1 (1s²) and 4 in shell 2 (2s² 2p²), giving K = 2 and L = 4. The tool calculates this automatically from the configuration string rather than using hard-coded tables.

Electron shells, subshells and orbitals

What are electron shells?

A shell is a principal energy level identified by quantum number n (1, 2, 3…). The first shell (K) holds up to 2 electrons, the second (L) up to 8, the third (M) up to 18, and so on, following 2n². In this tool, shell distribution counts how many electrons actually occupy each n for the selected element.

What are subshells?

Each shell contains subshells labelled s, p, d or f. An s subshell has 1 orbital (max 2 e⁻), p has 3 orbitals (max 6 e⁻), d has 5 orbitals (max 10 e⁻), and f has 7 orbitals (max 14 e⁻). The subshell letter tells you the orbital shape in quantum mechanics; here we show orbital boxes, not circular paths.

What is an orbital?

An orbital is a region of space that holds at most two electrons. Orbital box diagrams represent each orbital as a box; arrows (↑ ↓) show electrons and their spins.

What are s, p, d and f subshells?

What is the s subshell?

Hydrogen and helium occupy the 1s subshell. Because there is only one orbital, s electrons have no degeneracy and Hund's rule does not change the diagram. Helium completes 1s², the first noble gas configuration.

What is the p subshell?

Carbon, nitrogen and oxygen illustrate p filling clearly. Carbon 2p² shows two unpaired electrons; nitrogen 2p³ shows three unpaired electrons (half-filled subshell); oxygen 2p⁴ shows the first p orbital paired with four electrons total across three boxes.

What is the d subshell?

Transition metals derive their colour, magnetism and variable oxidation states from partially filled d subshells. Iron's [Ar] 3d⁶ 4s² configuration places six d electrons that participate in bonding and redox chemistry.

What is the f subshell?

Lanthanum and actinium start the f blocks; cerium's [Xe] 4f¹ 5d¹ 6s² shows that f and d filling can compete at the beginning of the series. The tool uses the accepted configurations from the dataset for every lanthanide and actinide.

Aufbau principle, Hund's rule and Pauli exclusion

Aufbau principle

The Aufbau principle describes the general tendency of electrons to occupy lower-energy orbitals before higher-energy ones. The canonical filling sequence is 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p. The builder highlights which subshells are occupied for your selected element and animates electrons in this order.

Hund's rule

When degenerate orbitals of equal energy are available — the three 2p orbitals, for example — electrons occupy separate orbitals with parallel spins before any pairing occurs. This maximises total spin and lowers exchange energy. In the orbital diagram, nitrogen's 2p³ appears as three separate ↑ arrows, not one paired box and one empty box.

Pauli exclusion principle

No two electrons in an atom can have the same set of four quantum numbers. Practically, each orbital holds at most two electrons and they must have opposite spins (↑↓). The visualizer never shows two parallel spins in the same box.

Shell distribution explained

Shell distribution counts how many electrons occupy each principal energy level n. It is not the same as a Bohr planetary model — electrons are not orbiting in simple circles. Instead, the shell view summarises total electron count per n derived from the configuration.

For sodium ([Ne] 3s¹), shell 1 holds 2 electrons, shell 2 holds 8, and shell 3 holds 1. For argon ([Ne] 3s² 3p⁶), shell 3 holds 8 electrons, completing the octet. The tool recalculates shell totals instantly when you change elements using the parsed subshell string, so the numbers always match the displayed configuration.

How to write electron configuration

Write subshells in order of filling. Each label has three parts: the principal number (1, 2, 3…), the subshell letter (s, p, d, f), and a superscript count of electrons. Example: 2p⁴ means four electrons in the 2p subshell.

Noble gas notation shortens the string by replacing inner electrons with the previous noble gas in brackets. Sodium is [Ne] 3s¹ instead of 1s² 2s² 2p⁶ 3s¹. Switch between full and noble gas views in the tool using the configuration tabs.

Electron configuration examples

• Hydrogen (1): 1s¹ — one electron in the lowest orbital.
• Helium (2): 1s² — first shell complete.
• Carbon (6): 1s² 2s² 2p² — two unpaired p electrons explain carbon's bonding versatility.
• Oxygen (8): 1s² 2s² 2p⁴ — six valence electrons (2s² 2p⁴).
• Sodium (11): [Ne] 3s¹ — one valence electron, typical alkali metal.
• Chlorine (17): [Ne] 3s² 3p⁵ — one electron short of a full p subshell.
• Iron (26): [Ar] 3d⁶ 4s² — transition metal with partially filled d orbitals.

Transition metal exceptions: chromium and copper

Half-filled and fully filled d subshells gain extra stability from electron exchange energy. Chromium promotes one 4s electron to 3d to achieve 3d⁵ 4s¹. Copper promotes one 4s electron to reach 3d¹⁰ 4s¹. Similar exceptions appear for molybdenum, silver, gold and several others. This builder uses the accepted configurations stored in the ToolAdda element dataset — not a naive algorithm that would misreport these elements.

4s and 3d: which fills first?

4s is normally filled before 3d when building the atom from scratch, but the 3d subshell lies lower in energy once the atom is formed. That is why transition-metal chemistry is dominated by d electrons even though 4s fills first on the way up.

Common mistakes

• Pairing electrons in p orbitals before each orbital has one electron (violates Hund's rule).
• Putting two electrons with the same spin in one orbital (violates Pauli exclusion).
• Using the simplified Aufbau prediction for Cr and Cu instead of the established ground state.
• Confusing shell capacity (2n²) with the actual number of electrons in a shell for a specific element.
• Drawing Bohr-style circular orbits and calling them orbitals — orbital box diagrams are the standard teaching representation.

Frequently asked questions

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