Calculate electron configurations for all 118 elements with orbital diagrams, noble gas notation, ion configurations, and quantum numbers. Step-by-step solutions for GCSE and A-Level Chemistry.
Orbitals fill in order of increasing energy (n+l rule).
Every orbital in a subshell is singly occupied (with parallel spin) before any is doubly occupied.
No two electrons can have the same set of four quantum numbers. Max 2 electrons per orbital.
Shell 1: 2, Shell 2: 8, Shell 3: 18, Shell 4: 32.
s has 1 orbital (2e⁻), p has 3 (6e⁻), d has 5 (10e⁻), f has 7 (14e⁻).
n: principal (1,2,3...), l: angular (0=s,1=p,2=d,3=f), ml: magnetic, ms: spin.
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Electron configuration describes the distribution of electrons in an atom's orbitals. It tells you which energy levels (shells) and subshells (s, p, d, f) the electrons occupy.
Understanding electron configuration is fundamental to chemistry — it explains why elements have certain properties, why they form particular ions, and why the periodic table is arranged the way it is. At GCSE level, you need to know the shell notation (e.g., 2.8.1 for sodium) for the first 20 elements. At A-Level, you must use subshell notation (e.g., 1s² 2s² 2p&sup6; 3s¹), understand noble gas shorthand, and handle configuration exceptions.
The Aufbau principle (from German "aufbauen", meaning "to build up") states that electrons fill orbitals from lowest energy to highest energy. The filling order is determined by the n+l rule: orbitals with lower (n+l) values fill first, and when two orbitals have the same (n+l) value, the one with lower n fills first.
Filling Order:
1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p
Each subshell can hold a maximum number of electrons: s = 2, p = 6, d = 10, f = 14. This comes from the number of orbitals (1, 3, 5, 7) multiplied by 2 (since each orbital holds a maximum of 2 electrons with opposite spins).
Hund's Rule states that electrons fill orbitals within a subshell singly (with parallel spins) before pairing up. This minimises electron-electron repulsion. For example, carbon's 2p² configuration has each electron in a separate p orbital (↑ ↑ _) rather than paired in one (↑↓ _ _).
The Pauli Exclusion Principle states that no two electrons in an atom can have the same set of four quantum numbers. In practice, this means each orbital can hold at most 2 electrons, and they must have opposite spins (+½ and -½).
Noble gas notation is a shorthand where the inner core electrons are replaced by the symbol of the nearest noble gas in square brackets. This makes configurations much shorter and focuses attention on the valence electrons.
Worked example: Iron (Fe, Z=26)
Full: 1s² 2s² 2p&sup6; 3s² 3p&sup6; 3d&sup6; 4s²
Noble gas: [Ar] 3d&sup6; 4s²
The first 18 electrons match argon's configuration, so we write [Ar] and list the remaining 8 electrons.
About 20 elements have electron configurations that differ from what the Aufbau principle predicts. The most important exceptions for A-Level are chromium (Cr) and copper (Cu).
Expected: [Ar] 4s² 3d&sup4;
Actual: [Ar] 4s¹ 3d&sup5;
A half-filled d subshell (3d&sup5;) is extra stable due to exchange energy.
Expected: [Ar] 4s² 3d&sup9;
Actual: [Ar] 4s¹ 3d¹&sup0;
A fully-filled d subshell (3d¹&sup0;) is extra stable.
Other notable exceptions include palladium (Pd), which uniquely has zero 5s electrons with all 10 electrons in 4d, and several lanthanide/actinide elements where 5d or 6d electrons appear instead of the expected f electrons.
When atoms form ions, the electron configuration changes. The key rules are:
Common Trap: Transition Metal Ions
Fe²+ is [Ar] 3d&sup6; (remove 4s² first), NOT [Ar] 3d&sup4; 4s². Although 4s fills before 3d, when both are occupied, 4s becomes higher in energy and is removed first.
Each electron in an atom is described by four quantum numbers:
Determines the shell (energy level). Values: 1, 2, 3, 4...
Determines the subshell shape. 0=s, 1=p, 2=d, 3=f.
Determines orbital orientation. Range: -l to +l.
Determines spin direction. Either +½ (up) or -½ (down).
At GCSE level, you need to write electron configurations using shell notation for the first 20 elements (hydrogen to calcium). Electrons fill shells in order: shell 1 holds 2, shell 2 holds 8, shell 3 holds 8 (at GCSE level), and shell 4 begins filling.
Examples:
The number of electrons in the outer shell tells you the group the element is in (for main group elements). For example, sodium has 1 outer electron and is in Group 1.
At A-Level, you must use subshell notation (e.g., 1s² 2s² 2p&sup6; 3s¹) and understand:
The block of an element is determined by which subshell is being filled: s-block (groups 1-2), p-block (groups 13-18), d-block (groups 3-12), and f-block (lanthanides and actinides).
Electron configuration describes how electrons are distributed across an atom's energy levels and subshells. It uses notation like 1s² 2s² 2p⁶ to show which orbitals contain electrons.
The Aufbau principle states that electrons fill orbitals from lowest to highest energy. The order follows the n+l rule: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, and so on.
Chromium is [Ar] 4s¹ 3d⁵ instead of 4s² 3d⁴ because a half-filled d subshell gives extra stability through exchange energy.
The 4s orbital has a lower energy than 3d in neutral atoms (n+l = 4 vs 5). However, once 3d electrons are present, 4s becomes higher in energy, which is why 4s electrons are removed first in cations.
For cations, remove electrons from the highest-energy orbital. For transition metals, remove ns before (n-1)d. For anions, add electrons following the Aufbau principle.
Four numbers that describe each electron: n (shell), l (subshell shape), ml (orbital orientation), and ms (spin direction). Together they give a unique address for every electron.
Electrons fill each orbital in a subshell singly (with parallel spin) before any orbital gets a second electron. This minimises repulsion between electrons.
A shorthand that replaces core electrons with the noble gas symbol in brackets. For example, Na is [Ne] 3s¹ instead of 1s² 2s² 2p⁶ 3s¹.
Species that have the same number of electrons and therefore the same electron configuration. Na⁺, F⁻, and Ne are all isoelectronic with 10 electrons.
s holds 2, p holds 6, d holds 10, and f holds 14. This comes from the number of orbitals (1, 3, 5, 7) times 2 electrons per orbital.
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