Introduction

The arrangement of electrons within atoms follows defined principles that ensure electrons occupy the lowest-energy positions available while maintaining maximum stability. This subtopic explains how orbitals fill in increasing energy order and how Hund’s rule determines electron distribution within degenerate orbitals.

The Concept of Orbital Filling

Electron configuration describes how electrons are arranged within atomic orbitals, following a predictable sequence based on increasing energy. Understanding this order is essential for explaining patterns in the Periodic Table and the reactivity of different elements.

Increasing Energy Order from 1s to 4p

Electrons fill orbitals in a sequence determined by orbital energy, with lower-energy orbitals filling before higher-energy ones. This sequence applies to all atoms and ensures that the most stable configuration is achieved.

The increasing energy order of sub-shells needed for OCR is 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p.

This orbital energy diagram shows the relative energies of the 1s, 2s, 2p, 3s, 3p, 4s, 3d and 4p sub-shells, with energy increasing up the vertical axis. It illustrates that the 4s orbital lies slightly lower in energy than the 3d orbitals, explaining why 4s fills before 3d. The diagram also includes electrons in each orbital, which exceeds OCR requirements but helps students relate energy ordering to occupied orbitals. Source

The Energy Levels and Sub-shell Structure

Each principal energy level contains sub-shells (s, p, d, f), each with a specific number of orbitals. Understanding how these sub-shells relate to electron filling provides structure to electron configuration.

Orbitals within the same sub-shell (for example, the three p orbitals) have the same energy and are known as degenerate orbitals. The way electrons occupy these affects atomic behaviour and magnetic properties.

Between major energy levels, there is overlap, meaning energy values do not increase evenly. This is why the 4s orbital fills before the 3d orbital, even though 4s is in a higher principal energy level.

Electron Spin and Pairing Principles

Electrons possess a property called spin, described as either ‘up’ or ‘down’. When two electrons occupy the same orbital, they must have opposite spins to minimise repulsion.

Orbital filling rules use spin to maintain stable arrangements and minimise electron–electron repulsion. These rules underpin the formation of the electron configurations of atoms.

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Hund’s Rule and Its Importance in Electron Filling

Hund’s rule is fundamental when filling orbitals of equal energy within the same sub-shell. The rule ensures electrons remain unpaired where possible to maintain a lower-energy configuration.

Hund’s rule applies only to degenerate orbitals, such as the three p orbitals or five d orbitals. It ensures that electron repulsion is minimised by spreading electrons out. As a result, atoms with half-filled or fully filled sub-shells often display increased stability.

Applying Hund’s Rule to Sub-shells

When filling p or d sub-shells, electrons follow a specific distribution pattern:

• Electrons occupy each orbital singly before pairing begins.

• All single electrons in degenerate orbitals have parallel spins.

• Pairing occurs only once all orbitals contain one electron.

This leads to characteristic patterns in electron configuration diagrams. For example:

• The p sub-shell (three orbitals) can hold up to six electrons.

• According to Hund’s rule:

  • The first three electrons occupy different p orbitals singly.

  • The next three electrons pair up to complete the sub-shell.

Following this pattern helps predict magnetic behaviour—species with unpaired electrons often show paramagnetism.

This diagram shows the three 2p orbitals for boron, carbon, nitrogen and oxygen, with arrows representing electron spins. It demonstrates how each orbital is singly occupied with parallel spins before any pairing occurs, directly illustrating Hund’s rule. The inclusion of specific elements goes beyond OCR requirements but helps students see how the rule functions in real electron configurations. Source

The Aufbau Principle and Filling Order

Hund’s rule works alongside the general principle governing how orbitals fill: the Aufbau principle, which states that electrons occupy the lowest-energy orbital available first.

The combination of Hund’s rule and the Aufbau principle produces the characteristic electron configurations used throughout A-Level Chemistry.

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Key Rules for Determining Electron Configuration

To use orbital filling order and Hund’s rule effectively, remember the core principles that govern electron arrangement:

Essential Orbital Filling Rules

• Lower-energy orbitals fill before higher-energy ones.

• The 4s orbital fills before the 3d orbital for neutral atoms.

• Each orbital holds a maximum of two electrons with opposite spins.

• Orbitals in the same sub-shell are filled singly before pairing (Hund’s rule).

• Electrons remain unpaired for as long as possible to minimise repulsion.

• Electron configurations follow the pattern dictated by energy order from 1s to 4p.

Benefits of Understanding Filling Patterns

Understanding electron filling allows students to:

• Predict stable electron configurations.

• Recognise unusual configurations such as those involving transition metals.

• Explain magnetic and chemical properties of atoms.

• Interpret patterns in the Periodic Table linked to sub-shell structure.