Electron Configuration: Rules, Notation, and Practice for High School

Understanding where electrons are located in an atom is one of the most important ideas in chemistry. Electron configuration describes the arrangement of electrons in an atom's orbitals, and it explains why elements behave the way they do, why the periodic table is organized the way it is, and how atoms form bonds. This guide covers the rules, notation, and practice you need for high school chemistry.

What electron configuration tells us

Every atom has a nucleus containing protons and neutrons, surrounded by electrons. The electrons are not randomly scattered; they occupy specific energy levels and orbitals. An electron configuration is a written description of which orbitals are occupied and how many electrons are in each one. Knowing the electron configuration of an element lets you predict its chemical properties, the types of bonds it forms, and its position on the periodic table.

Energy levels, subshells, and orbitals

Electrons are organized into energy levels (also called shells), numbered 1, 2, 3, 4, and so on. Each energy level contains one or more subshells, labeled s, p, d, and f. Each subshell contains a specific number of orbitals, and each orbital can hold a maximum of 2 electrons. The s subshell has 1 orbital (2 electrons max). The p subshell has 3 orbitals (6 electrons max). The d subshell has 5 orbitals (10 electrons max). The f subshell has 7 orbitals (14 electrons max).

The first energy level has only an s subshell (1s). The second level has s and p subshells (2s, 2p). The third level has s, p, and d subshells (3s, 3p, 3d). The fourth level has s, p, d, and f subshells (4s, 4p, 4d, 4f). As you go to higher energy levels, more subshells become available, and the total number of electrons that can be accommodated increases.

The Aufbau principle

The Aufbau principle (from the German word for "building up") states that electrons fill orbitals starting from the lowest energy level and working upward. The order in which orbitals fill is not simply 1s, 2s, 2p, 3s, 3p, 3d. Because of the way energy levels overlap, the actual filling order is: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.

A helpful memory tool is the diagonal rule diagram. Write the subshells in rows: 1s on the first row, 2s 2p on the second, 3s 3p 3d on the third, and so on. Then draw diagonal arrows from the top right to the bottom left. Following the arrows gives you the correct filling order. Many periodic tables also include this information.

The Pauli exclusion principle

The Pauli exclusion principle states that no two electrons in the same atom can have the same set of four quantum numbers. In practical terms, this means each orbital can hold at most 2 electrons, and those 2 electrons must have opposite spins (one spin-up and one spin-down). This is why the s subshell holds 2 electrons, the p subshell holds 6, the d subshell holds 10, and the f subshell holds 14.

Hund's rule

Hund's rule states that when electrons fill orbitals of equal energy (such as the three p orbitals or the five d orbitals), one electron enters each orbital before any orbital gets a second electron. All of the singly occupied orbitals will have electrons with the same spin. Think of it like passengers on a bus: everyone takes their own seat before anyone has to share.

This rule minimizes electron-electron repulsion and results in a more stable arrangement. It is especially important when writing orbital diagrams, where you draw boxes for orbitals and arrows for electrons.

How to write electron configurations

The standard notation lists each occupied subshell followed by a superscript indicating the number of electrons in that subshell. For example, hydrogen has 1 electron in the 1s orbital: 1s¹. Helium has 2 electrons: 1s². Lithium has 3 electrons: 1s² 2s¹. Carbon has 6 electrons: 1s² 2s² 2p². Neon has 10 electrons: 1s² 2s² 2p⁶.

Continuing through the first 20 elements: sodium (Na, 11 electrons) is 1s² 2s² 2p⁶ 3s¹. Magnesium (Mg, 12) is 1s² 2s² 2p⁶ 3s². Aluminum (Al, 13) is 1s² 2s² 2p⁶ 3s² 3p¹. Silicon (Si, 14) is 1s² 2s² 2p⁶ 3s² 3p². Phosphorus (P, 15) is 1s² 2s² 2p⁶ 3s² 3p³. Sulfur (S, 16) is 1s² 2s² 2p⁶ 3s² 3p⁴. Chlorine (Cl, 17) is 1s² 2s² 2p⁶ 3s² 3p⁵. Argon (Ar, 18) is 1s² 2s² 2p⁶ 3s² 3p⁶. Potassium (K, 19) is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹. Calcium (Ca, 20) is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s².

Noble gas shorthand notation

Writing out the full configuration for heavier elements gets long. A shortcut is to use the symbol of the nearest noble gas in brackets to represent the core electrons. For example, sodium's full configuration is 1s² 2s² 2p⁶ 3s¹. The core electrons (1s² 2s² 2p⁶) match neon's configuration, so you can write sodium as [Ne] 3s¹. Iron (Fe, 26 electrons) is [Ar] 3d⁶ 4s². This shorthand makes it easier to focus on the valence electrons that determine chemical behavior.

Exceptions: chromium and copper

Most elements follow the Aufbau principle perfectly, but a few exceptions exist because half-filled and fully filled d subshells are especially stable. Chromium (Cr, 24 electrons) is expected to be [Ar] 3d⁴ 4s², but its actual configuration is [Ar] 3d⁵ 4s¹. One electron moves from the 4s to the 3d subshell to create a half-filled d subshell. Similarly, copper (Cu, 29 electrons) is expected to be [Ar] 3d⁹ 4s², but its actual configuration is [Ar] 3d¹⁰ 4s¹, achieving a fully filled d subshell.

These exceptions occur because the energy difference between the 4s and 3d orbitals is small, and the extra stability of a half-filled or fully filled subshell tips the balance. You should memorize these two common exceptions for your chemistry course.

Connection to the periodic table

The periodic table is organized by electron configuration. Elements in the same group (column) have the same number of valence electrons and similar configurations in their outermost shell. The s-block (Groups 1 and 2) contains elements whose last electron enters an s orbital. 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.

This is why elements in the same group have similar chemical properties. Lithium, sodium, and potassium all have one electron in their outermost s orbital, which is why they are all reactive metals that form +1 ions. Understanding electron configuration turns the periodic table from a chart you memorize into a map you can read.

Orbital diagrams

An orbital diagram is a visual representation of electron configuration. Each orbital is drawn as a box, and electrons are drawn as arrows pointing up or down to represent their spin. Orbital diagrams make it easy to see how Hund's rule and the Pauli exclusion principle work in practice. For example, the orbital diagram for nitrogen (1s² 2s² 2p³) shows the 1s box with two arrows (up and down), the 2s box with two arrows, and three 2p boxes each with one up arrow.

Drawing orbital diagrams is excellent practice for understanding electron configuration. You can use tools like MyChemLab AI to check your orbital diagrams and electron configurations interactively, getting immediate feedback as you work through the elements.

Key takeaways

Electron configuration describes how electrons are distributed among an atom's orbitals. Three rules govern the filling order: the Aufbau principle (fill lowest energy first), the Pauli exclusion principle (maximum 2 electrons per orbital with opposite spins), and Hund's rule (spread electrons across equal-energy orbitals before pairing). Noble gas shorthand simplifies notation for heavier elements. A few elements like chromium and copper are exceptions due to the stability of half-filled and fully filled d subshells. Mastering electron configuration connects atomic structure to the periodic table and to chemical bonding.