What is Electron Configuration?
Electron configuration describes the distribution of electrons in an atom's orbitals. It shows how electrons are arranged in shells and subshells around the nucleus. Understanding electron configuration is fundamental to chemistry because it determines an element's chemical properties, reactivity, and bonding behavior. The configuration follows specific rules including the Aufbau principle, Pauli exclusion principle, and Hund's rule, which govern how electrons fill available orbitals.
How to Write Electron Configuration
Electron configuration is written using orbital notation with superscripts indicating the number of electrons. The general format is: energy level (1, 2, 3...) + sublevel type (s, p, d, f) + number of electrons as superscript. For example, oxygen (atomic number 8) has 8 electrons distributed as: 1s² 2s² 2p⁴. The superscripts add up to the total number of electrons. Each s orbital holds 2 electrons maximum, p orbitals hold 6, d orbitals hold 10, and f orbitals hold 14 electrons.
The Aufbau Principle Explained
The Aufbau principle states that electrons fill orbitals starting from the lowest available energy level before moving to higher levels. The filling order follows a specific sequence: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p. This order can be remembered using the diagonal rule or n+l rule. Notice that 4s fills before 3d even though 3d is in a lower shell—this is because 4s has lower energy. Understanding this principle is crucial for correctly determining electron configurations for all elements.
Noble Gas Notation (Abbreviated Configuration)
Noble gas notation is a shorthand method for writing electron configurations using the previous noble gas as a starting point. Instead of writing the full configuration, we use the noble gas symbol in brackets followed by the remaining electrons. For example, sodium (Na, 11 electrons) has full configuration 1s² 2s² 2p⁶ 3s¹, but abbreviated as [Ne] 3s¹. This method is especially useful for elements with many electrons, making configurations easier to write and compare. The noble gases used as reference points are He, Ne, Ar, Kr, Xe, and Rn.
Valence Electrons and Chemical Properties
Valence electrons are the electrons in the outermost shell of an atom. These electrons determine an element's chemical reactivity and bonding behavior. Elements in the same group of the periodic table have the same number of valence electrons, which explains why they have similar chemical properties. For example, all alkali metals (Group 1) have one valence electron and are highly reactive. Noble gases have full valence shells (8 electrons except helium with 2), making them very stable and unreactive. Understanding valence electrons helps predict how elements will bond and form compounds.
Orbital Diagrams and Electron Spin
Orbital diagrams provide a visual representation of electron configuration, showing individual orbitals as boxes and electrons as arrows. Arrows pointing up (↑) and down (↓) represent electron spin. According to Hund's rule, when filling orbitals of equal energy (like the three 2p orbitals), electrons occupy empty orbitals singly with parallel spins before pairing up. This minimizes electron-electron repulsion. The Pauli exclusion principle requires that paired electrons in the same orbital must have opposite spins. Orbital diagrams help visualize these concepts and predict magnetic properties of atoms.
Exceptions to Standard Configuration
Some elements have electron configurations that deviate from the expected pattern. Chromium (Cr) and copper (Cu) are classic examples. Instead of [Ar] 4s² 3d⁴ and [Ar] 4s² 3d⁹, they have [Ar] 4s¹ 3d⁵ and [Ar] 4s¹ 3d¹⁰ respectively. These exceptions occur because half-filled and fully-filled d subshells provide extra stability. Similar exceptions occur in heavier transition metals and lanthanides. These irregularities reflect the complex interplay of electron-electron repulsions and orbital energies.
Electron Configuration of Ions
When atoms form ions by gaining or losing electrons, their electron configurations change. For cations (positive ions), electrons are removed from the highest energy orbital first—often the outermost s orbital, not the d orbital that filled last. For example, Fe²⁺ is [Ar] 3d⁶ (not [Ar] 4s² 3d⁴) because 4s electrons are lost first. For anions (negative ions), electrons are added following the normal filling order. Chlorine (Cl) gains one electron to become Cl⁻ with configuration [Ne] 3s² 3p⁶, which is the same as argon. Understanding ion configurations is essential for predicting ionic charge and chemical behavior.
Quantum Numbers and Electron Address
Each electron in an atom has a unique set of four quantum numbers that describe its "address": principal quantum number (n) indicates energy level/shell, azimuthal quantum number (l) indicates orbital shape (s, p, d, f), magnetic quantum number (mₗ) indicates orbital orientation, and spin quantum number (mₛ) indicates electron spin direction. These quantum numbers arise from solving the Schrödinger equation and provide a complete description of an electron's state. The Pauli exclusion principle states no two electrons can have identical sets of all four quantum numbers.
Using This Calculator
Our electron configuration calculator makes finding configurations effortless. Simply enter any element by name (like "iron"), symbol ("Fe"), or atomic number (26) to instantly get the full electron configuration, abbreviated noble gas notation, orbital diagram, number of valence electrons, and quantum number information. The calculator handles all 118 elements including transition metals with their exceptions. Use ion mode to calculate configurations for charged atoms, or compare mode to see configurations side-by-side. Perfect for chemistry homework, exam preparation, and quick reference in lab work.