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VCE Chemistry

Periodic Table Trends: Electronic Configuration, Atomic Radius, Ionisation Energy and More

Published 2026-10-05 · Updated 2026-10-05

Periodic table trends are among the most reliably tested topics in the VCAA Chemistry Examination, yet they are also among the most commonly mishandled. Students often memorise the direction of a trend without understanding why it occurs, which means a single unfamiliar question phrasing can throw them completely. VCAA examiners are very good at asking questions that require reasoning, not just recall. If you understand core charge and shielding, every trend in this topic follows logically from the same two ideas.

What the syllabus asks

This concept covers:

The idea, explained

How the periodic table is organised

Elements are arranged in order of increasing atomic number (Z). Each new period (row) begins when a new principal energy level starts to fill. Each group (column) contains elements that share the same valence electron configuration, which is why elements in the same group behave similarly. The table is also divided into four blocks, named after the subshell being filled: the s-block (groups 1 and 2), the p-block (groups 13 to 18), the d-block (groups 3 to 12, the transition metals), and the f-block (lanthanides and actinides).

Shell and subshell configurations

Electrons occupy shells (principal energy levels, numbered 1, 2, 3 ...) which are subdivided into subshells (s, p, d, f). The maximum electron capacity of each subshell is: s = 2, p = 6, d = 10, f = 14. Electrons fill subshells from lowest energy upward, following the Aufbau order:

1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p

Note that 4s fills before 3d because it sits at slightly lower energy in a neutral atom. Within a subshell, electrons occupy separate orbitals singly before pairing (Hund's rule), and each orbital holds a maximum of two electrons with opposite spins (Pauli exclusion principle).

Some worked examples:

Core charge and shielding: the engine behind every trend

Core charge is the effective nuclear charge felt by a valence electron, approximated as the number of protons minus the number of inner-shell electrons. Across a period, protons are added to the same shell, so inner-shell shielding stays roughly constant and core charge increases. Down a group, each new period adds a new inner shell, so shielding increases and the valence electrons feel a weaker effective attraction despite the higher proton count.

Atomic radius

Across a period, atomic radius decreases because the rising core charge pulls valence electrons closer to the nucleus. Down a group, atomic radius increases because each new shell sits further from the nucleus.

First ionisation energy

First ionisation energy is the energy required to remove the first electron from a gaseous atom. Across a period it generally increases because a higher core charge holds valence electrons more tightly. Down a group it decreases because the valence electron is further away and more shielded.

Electronegativity

Electronegativity is the tendency of a bonded atom to attract shared electrons. It increases across a period and decreases down a group, for the same core charge and shielding reasons. Fluorine is the most electronegative element.

Metallic and non-metallic character

Metallic character increases toward the lower-left of the table, where atoms lose electrons readily (low ionisation energy, low electronegativity). Non-metallic character increases toward the upper-right, where atoms gain electrons readily (high electronegativity, high ionisation energy).

Reactivity

For alkali metals (Group 1), reactivity increases down the group. Each successive metal loses its single valence electron more easily because ionisation energy decreases. Lithium reacts gently with water, sodium reacts rapidly, potassium ignites immediately, and rubidium and caesium react explosively.

For halogens (Group 17), reactivity decreases down the group. Each successive halogen gains an electron less readily because of the larger atomic radius and greater shielding. The displacement order is F2 > Cl2 > Br2 > I2.

What the exam asks

VCAA examiners commonly ask students to:

For example: Compare the atomic radius, first ionisation energy and electronegativity of Cl and Br. Both are in Group 17. Moving down from Period 3 (Cl) to Period 4 (Br), a new shell is added, so Br has a larger atomic radius. The larger atom with greater shielding means the valence electron is easier to remove, so Cl has a higher first ionisation energy. By the same reasoning, Cl also has a higher electronegativity. One core argument, three answers.

Common mistakes

If you want to work through these ideas with immediate feedback and exam-style questions tailored to the VCAA Chemistry Examination, Avocado is an AI-native tutoring service built specifically for VCE students. Start with the The Periodic Table: Electronic Configuration and Trends lesson to practise writing configurations and applying trends until they become second nature.