Introduction

Electron configuration patterns across Periods 2 and 3 reveal predictable arrangements of electrons in sub-shells, enabling classification of elements into s-, p-, and d-blocks based on their highest-energy electrons.

Electron Configuration Across Periods 2 and 3

The arrangement of electrons in atoms is fundamental to understanding periodicity. As proton number increases across a period, electrons fill sub-shells in a specific sequence determined by increasing energy levels. This produces repeating patterns that form the basis for chemical periodicity.

When discussing electron configuration, the term electron configuration refers to the ordered distribution of electrons among energy levels and sub-shells according to quantum rules. The periodic table reflects this structure directly, meaning its layout is a visual map of the order in which electrons occupy orbitals.

Electron configurations across Periods 2 and 3 follow a consistent pattern:

• Period 2 elements fill the 2s and then the 2p sub-shells.

• Period 3 elements fill the 3s and then the 3p sub-shells.

• Sub-shells fill in order of increasing energy, which broadly follows the sequence 1s → 2s → 2p → 3s → 3p → 4s → 3d.

Diagram showing the relative energies of s, p, d and f sub-shells, reinforcing the electron filling order. Includes shells beyond Periods 2 and 3 but follows the same pattern. Source

The Aufbau Principle, Pauli Exclusion Rule and Hund’s Rule

Although these principles are not named explicitly in the OCR specification, they underpin the arrangement of electrons and therefore support understanding of sub-shell filling patterns:

• Electrons occupy the lowest-energy available orbital first.

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

• Within a set of degenerate (equal-energy) orbitals, electrons fill singly before pairing.

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Sub-shell Structure and Filling Patterns

Across Periods 2 and 3, sub-shells follow predictable filling orders because:

• Increasing nuclear charge across a period draws electrons closer, influencing energetic ordering.

• The relative energies of s and p sub-shells remain consistent across these periods, leading to repeated sub-shell filling patterns.

Key Features of Sub-shell Filling

• s-sub-shells contain a single orbital and hold up to two electrons.

• p-sub-shells contain three orbitals and hold up to six electrons.

• Each successive element increases proton number by one and introduces one additional electron.

Figure showing s, p and d orbital shapes, reinforcing the number of orbitals in each sub-shell. Includes d orbitals, which extend beyond Periods 2 and 3 but support understanding of orbital structure. Source

These patterns explain why Periods 2 and 3 have eight elements each: 2 electrons in the s-sub-shell and 6 in the p-sub-shell.

Block Classification of the Periodic Table

The periodic table is divided into blocks based on the type of sub-shell that receives the highest-energy electron. This classification reinforces periodicity by grouping elements with similar electronic structures and therefore similar chemical behaviours.

The presence of this definition aligns with OCR expectations while supporting understanding of periodic arrangement.

The s-Block

The s-block includes Groups 1 and 2, plus hydrogen and helium. The distinguishing feature is that the highest-energy electron occupies an s-orbital:

• Group 1 elements have an outer configuration ending in s¹.

• Group 2 elements have an outer configuration ending in s².

These configurations explain the strong reactive trends in these groups, although the details of reactivity fall outside this subsubtopic.

The p-Block

The p-block spans Groups 13–18. Here the highest-energy electron occupies a p-orbital. Configurations end in p¹ to p⁶, depending on the group.
The p-block contains a wide range of element types, including:

• Non-metals (e.g., oxygen, nitrogen)

• Metalloids (e.g., silicon)

• Some metals (e.g., aluminium)

These variations arise because p-orbitals allow greater diversity in bonding and structure.

The d-Block

The d-block contains the transition metals, which begin in Period 4. The highest-energy electron generally occupies a d-orbital, although the 4s orbital fills before the 3d orbital due to relative energy differences.

Even though the specification for this subsubtopic focuses on Periods 2 and 3, including the d-block in classification is essential, as OCR requires block terminology to be understood in full.

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Periodic table showing s-, p-, d- and f-block regions, illustrating how highest-energy electrons define periodic table classification. Includes lanthanoids and actinoids, which extend beyond Periods 2 and 3. Source

Relationship Between Electron Configuration and Periodicity

Periodicity arises from recurring electron configuration patterns. As elements across a period fill equivalent sub-shell types, predictable trends emerge in:

• Atomic radius

• Ionisation energy

• Electronegativity

• Chemical reactivity

While detailed trends belong to other subsubtopics, their origin lies in the electron configuration patterns described here.

Key Observations

• Elements with similar outer electron configurations appear in the same group.

• Repetition of s- and p-sub-shell filling across adjacent periods creates repeating chemical behaviour.

• Classification into s-, p-, and d-blocks provides a structural way to interpret these patterns.

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