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The Electron Configuration Calculator provides the electronic structure of elements from hydrogen (Z=1) through krypton (Z=36), showing how electrons fill atomic orbitals according to the Aufbau principle, Pauli exclusion principle, and Hund's rule. Understanding electron configurations is fundamental to predicting chemical bonding, reactivity, and the periodic behavior of elements.
Electron configurations are written by filling orbitals in order of increasing energy, following the n+l rule (Madelung's rule): fill orbitals with lower n+l values first, and for equal n+l, fill lower n first. The subshell filling order is: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p... The notation 1s^2 2s^2 2p^6 means 2 electrons in 1s, 2 in 2s, 6 in 2p (subshells s hold max 2, p hold max 6, d hold max 10, f hold max 14).
Valence electrons are those in the outermost shell (highest principal quantum number n). They determine chemical bonding: elements with 1-2 valence electrons tend to be metals that give up electrons; elements with 6-7 valence electrons tend to be nonmetals that gain electrons; elements with 8 valence electrons (noble gases) are chemically inert. Transition metals (d-block, Z=21-30) have partially filled d subshells and complex bonding behavior.
Notable exceptions to simple filling occur at chromium (Z=24) and copper (Z=29): instead of the expected [Ar] 3d^4 4s^2 and [Ar] 3d^9 4s^2, they adopt [Ar] 3d^5 4s^1 and [Ar] 3d^10 4s^1 because half-filled and completely-filled d subshells confer extra stability. These exceptions are important in chemistry and spectroscopy.
The Aufbau principle fills orbitals: 1s(2), 2s(2), 2p(6), 3s(2), 3p(6), 4s(2), 3d(10), 4p(6). For Z up to 36, valence electrons = electrons in the outermost s and p subshells (for main group) or s and d (for transition metals). Period = row in periodic table (determined by the highest n occupied). The calculator covers all 36 elements through krypton.
Configuration shorthand uses noble gas cores: [He]=1s^2, [Ne]=[He]2s^2 2p^6, [Ar]=[Ne]3s^2 3p^6. Elements in the same group have the same valence electron configuration. Group 1 (alkali metals): ns^1. Group 2 (alkaline earth): ns^2. Group 17 (halogens): ns^2 np^5. Group 18 (noble gases): ns^2 np^6. Transition metals have partially filled d subshells.
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Iron: [Ar] 3d^6 4s^2. It has 8 electrons beyond argon, with 2 valence 4s electrons. The 6 d electrons give iron its magnetic and catalytic properties.
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Chlorine: [Ne] 3s^2 3p^5. With 7 valence electrons, it needs one more to complete its octet — explaining its high electronegativity and reactivity.
The Aufbau (German for 'building up') principle states that electrons fill atomic orbitals starting with the lowest energy orbital and working upward. The filling order follows the n+l rule: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p. This predicts the periodic table structure.
The Pauli exclusion principle states that no two electrons in an atom can have identical quantum numbers. Each orbital holds at most two electrons, and they must have opposite spins (spin up and spin down). This principle explains why electrons do not all collapse into the lowest energy orbital and underlies the structure of all matter.
Hund's rule states that when filling degenerate (equal energy) orbitals, electrons fill each orbital singly before pairing up, and the unpaired electrons have parallel spins. For example, nitrogen (2p^3) has three unpaired electrons, one in each 2p orbital. This maximizes the total spin angular momentum and minimizes electron-electron repulsion.
A completely filled d subshell (3d^10) is particularly stable due to exchange energy. Copper adopts [Ar] 3d^10 4s^1 rather than [Ar] 3d^9 4s^2 because the energy gained from completing the d subshell outweighs the cost of moving an electron from 4s to 3d. Similarly, chromium adopts [Ar] 3d^5 4s^1 (half-filled d).
Valence electrons are the outermost electrons that participate in chemical bonding. For main group elements, they are in the highest-n s and p subshells. For transition metals, the d electrons also participate. The number of valence electrons determines the element's chemical behavior and determines the group number in the periodic table.
Noble gas cores abbreviate the inner electron configuration. [He] replaces 1s^2. [Ne] replaces 1s^2 2s^2 2p^6. [Ar] replaces the first 18 electrons. For example, iron is written [Ar] 3d^6 4s^2 instead of the full 1s^2 2s^2 2p^6 3s^2 3p^6 3d^6 4s^2. This emphasizes the chemically relevant outer electrons.
Elements in the same period fill the same principal quantum shells. Elements in the same group have the same valence electron configuration. The s-block (groups 1-2) fills s orbitals. The p-block (groups 13-18) fills p orbitals. The d-block (transition metals, groups 3-12) fills d orbitals. The f-block (lanthanides and actinides) fills f orbitals.
Electron spin is an intrinsic quantum property with two states: spin-up (+1/2) and spin-down (-1/2). Spin is not classical rotation but a fundamental quantum attribute. Each orbital holds at most two electrons with opposite spins. Spin determines magnetic properties: unpaired electrons give atoms paramagnetism; paired electrons give diamagnetism.
Ionization removes an electron from an atom. The first ionization energy is highest for noble gases (full shells), lowest for alkali metals (one easily removed valence electron). Elements just past a half-filled or fully-filled subshell are slightly easier to ionize because removing one electron restores a more stable configuration. This explains the slight dip in ionization energy from N to O.
Transition metals (Z=21-30, 39-48, 72-80, 104-112) have partially filled d subshells. Their configurations follow [noble gas] (n-1)d^x ns^2 (with exceptions). The partially filled d orbitals give transition metals variable oxidation states, colorful compounds, magnetic properties, and catalytic activity. Iron can be Fe^2+ or Fe^3+ because losing 2 or 3 electrons both give stable configurations.
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