Electron Configuration Calculator

Frequently Asked Questions

What is electron configuration?

Electron configuration is the distribution of electrons in an atom across orbitals (subshells), written as 1s2 2s2 2p6... etc. Each subshell designation is: principal quantum number n, subshell letter (s, p, d, f), and superscript electron count. The configuration determines chemical reactivity and bonding behavior.

What is the Aufbau principle?

Aufbau (German for "building up") is the principle that electrons fill orbitals from lowest to highest energy. The filling order is: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p. Note that 4s fills before 3d because at typical atomic sizes, 4s is slightly lower in energy.

Why do chromium and copper have unusual electron configurations?

Chromium (Z=24) is [Ar] 3d5 4s1 instead of the expected [Ar] 3d4 4s2, and copper (Z=29) is [Ar] 3d10 4s1 instead of [Ar] 3d9 4s2. Both exceptions arise because half-filled (d5) and fully filled (d10) d subshells have extra stability from exchange energy. The energy gained by achieving these configurations exceeds the cost of leaving 4s with one electron.

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How This Calculator Works

The Electron Configuration Calculator builds the ground-state electron configuration for any element (Z = 1 to 118) by filling subshells in Aufbau order. It handles known exceptions like chromium and copper where one electron is promoted to half-fill or fully fill a d subshell. For ions, it removes electrons from the highest-energy subshell first (for cations) or adds to the next available subshell (for anions).

Enter an atomic number between 1 and 118, or one of 30 supported symbols (H, He, Li, Be, B, C, N, O, F, Ne, Na, Mg, Al, Si, P, S, Cl, Ar, K, Ca, Fe, Co, Ni, Cu, Zn, Br, Kr, Ag, I, Au). For other elements, use the atomic number.

The Aufbau Principle

The Aufbau principle (German for "building up") states that electrons fill the lowest available energy subshells first. The filling order by energy is: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p. Notice that 4s fills before 3d because at that stage 4s has lower energy, even though n=4 is a higher principal quantum number than n=3.

Subshell capacities: s holds 2 electrons, p holds 6, d holds 10, f holds 14. These limits come from the number of magnetic quantum number values and the Pauli exclusion principle (each orbital holds at most 2 electrons with opposite spins).

Hund's Rule and Pauli Exclusion

Within a subshell, electrons occupy separate orbitals with parallel spins before pairing (Hund's rule). This minimizes electron-electron repulsion. The Pauli exclusion principle requires that no two electrons in the same atom share all four quantum numbers - two electrons in the same orbital must have opposite spins.

Orbital Notation Explained

Each term in the configuration is written as nLx where n is the principal quantum number, L is the subshell letter (s, p, d, f), and x is the number of electrons in that subshell. For example, 3d10 means the 3d subshell is completely filled with 10 electrons. Noble gas shorthand replaces the filled inner shells with the symbol of the nearest noble gas in brackets, such as [Ar] for the first 18 electrons.

Exceptions to the Aufbau Principle

Several transition metals deviate from simple Aufbau filling. The most commonly tested exceptions are:

Chromium (Z=24): [Ar] 3d5 4s1 instead of [Ar] 3d4 4s2 - half-filled d subshell is extra stable

Copper (Z=29): [Ar] 3d10 4s1 instead of [Ar] 3d9 4s2 - fully filled d subshell is extra stable

Molybdenum (Z=42), silver (Z=47), and gold (Z=79) follow analogous patterns in the 4d and 5d series

Valence Electrons and Chemical Reactivity

Valence electrons are those in the outermost principal quantum number shell (and the outermost d subshell for transition metals). They determine how an element bonds and reacts. Main-group elements in the same column share the same valence electron count, explaining the periodic trends in reactivity. Noble gases with 8 valence electrons (helium has 2) are chemically inert. Transition metals can use both their ns and (n-1)d electrons in bonding, giving them variable oxidation states.

Electron Configuration Calculator

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