Period 3 Trends: Atomic Radius and Ionisation Energy Revision Notes | A Level AQA Chemistry

1. Why Period 3 Trends Matter

Picture the eight elements from sodium to argon sitting in a row. Across that row, the nuclear charge climbs from +11 to +18 — yet the number of inner shielding electrons stays almost the same. That single tension drives every trend you need to explain in this topic.

AQA expects you to explain atomic radius, first ionisation energy, and melting point trends — and to link each one back to structure and bonding. These three trends are deeply connected, so understanding the underlying logic once lets you answer very different exam questions with confidence.

2. Atomic Radius Across Period 3

Atomic radius decreases steadily from Na to Cl across Period 3 (argon is a noble gas with no standard covalent radius, so it is typically excluded from this comparison). The reason is straightforward once you see the pattern: each element adds one more proton to the nucleus and one more electron to the same third shell.

The inner shells — the first and second — provide shielding that barely changes across the period. So the outer electrons feel a progressively stronger effective nuclear charge as you move from Na to Cl. A stronger pull drags those outer electrons closer to the nucleus, shrinking the atom.

The chart confirms a clear, steady decrease. Notice there are no anomalies here — unlike first ionisation energy, atomic radius decreases smoothly because the same mechanism (increasing effective nuclear charge, constant shielding) operates without interruption across the period.

3. First Ionisation Energy: The General Trend and Two Anomalies

First ionisation energy (IE₁) generally increases across Period 3 for the same reason atomic radius decreases: rising effective nuclear charge makes outer electrons harder to remove. However, there are two dips in the trend — at aluminium (Al) and at sulphur (S) — and these are the points examiners love to probe.

The line rises overall but dips at Al and again at S. Each dip has a precise structural explanation rooted in electron configuration — and AQA mark schemes reward that precision.

3a. The Anomaly at Aluminium (Al)

Magnesium's electron configuration ends in $3s^2$. Aluminium's ends in $3s^2 3p^1$. The outer electron in Al occupies a 3p orbital, which is at a slightly higher energy level than the 3s orbital and is therefore further from the nucleus on average. It also experiences additional shielding from the filled 3s sub-shell beneath it.

The result: Al's outer electron is easier to remove than Mg's, despite Al having a higher nuclear charge. IE₁ drops from 738 kJ mol⁻¹ (Mg) to 578 kJ mol⁻¹ (Al). This is the key exam point — the 3p electron is both further away and more shielded than the 3s electron it replaces.

3b. The Anomaly at Sulphur (S)

Phosphorus has the configuration $[Ne] 3s^2 3p^3$. Each of the three 3p electrons occupies a separate 3p orbital — each is unpaired, occupying a separate 3p orbital with the same spin direction (Hund's rule of maximum multiplicity). Sulphur adds a fourth 3p electron, which must pair up with one of the existing 3p electrons in the same orbital.

That paired electron experiences extra electron–electron repulsion from its orbital partner, making it easier to remove. IE₁ therefore dips from 1012 kJ mol⁻¹ (P) to 1000 kJ mol⁻¹ (S). The key phrase for the mark scheme: the paired electron in S experiences greater inter-electron repulsion, lowering the energy needed to remove it.

4. Melting Points Across Period 3: Structure and Bonding

Melting point trends across Period 3 are more complex than atomic radius or ionisation energy, because the type of structure and bonding changes dramatically across the period. You cannot explain melting points with a single mechanism — you need to identify the structure first, then apply the appropriate bonding argument.

Period 3 elements fall into four structural categories: metallic lattices (Na, Mg, Al), a giant covalent lattice (Si), simple molecular structures (P₄, S₈, Cl₂), and a monatomic gas (Ar). Each category has a characteristic melting point range and a distinct reason for it.

4a. Metallic Lattices: Na, Mg, Al

Sodium, magnesium, and aluminium all form metallic lattices — a regular arrangement of positive metal ions surrounded by a delocalised sea of electrons. Melting a metal requires breaking these electrostatic attractions between the ions and the delocalised electrons.

Melting point rises from Na to Al because the metallic bonding strengthens significantly. Sodium contributes one delocalised electron per atom, magnesium contributes two, and aluminium contributes three. More delocalised electrons mean stronger electrostatic attraction and a higher charge density on the ion (Al³⁺ is smaller and more highly charged than Na⁺). More energy is needed to overcome these stronger forces, so melting point increases: Na (98 °C) → Mg (650 °C) → Al (660 °C).

4b. Giant Covalent Lattice: Si

Silicon has the highest melting point in Period 3 (1414 °C) because it forms a giant covalent lattice — every Si atom is covalently bonded to four others in a continuous three-dimensional network. Melting requires breaking a very large number of strong, directional covalent bonds simultaneously, which demands an enormous amount of energy.

4c. Simple Molecular Structures: P₄, S₈, Cl₂, Ar

Phosphorus (P₄), sulphur (S₈), chlorine (Cl₂), and argon (Ar) all exist as discrete molecules or atoms. The covalent bonds within each molecule are strong, but the forces between molecules are weak van der Waals (London dispersion) forces. Melting only requires overcoming these weak intermolecular forces — the covalent bonds themselves remain intact.

The melting point order within this group reflects the size of the molecules: S₈ (113 °C) > P₄ (44 °C) > Cl₂ (−101 °C) > Ar (−189 °C). Larger molecules have more electrons and a greater surface area for contact, generating stronger van der Waals forces and therefore higher melting points. Argon, as a monatomic noble gas, has the weakest forces of all and the lowest melting point.

Trend questions in AQA A Level chemistry almost always carry 'explain' as the command word. That means stating what happens is never enough — you must give the reason using precise scientific language. The mark scheme will have specific phrases, and vague answers lose marks even when the conclusion is right.

5. Exam Technique: Answering Trend Questions at A Level

For atomic radius and ionisation energy questions, build your answer in three layers: (1) state what changes across the period (nuclear charge increases, shielding stays approximately constant), (2) state the consequence (effective nuclear charge increases), and (3) link to the property (electrons are held more tightly / harder to remove). For melting point questions, always identify the structure and bonding type before explaining the energy needed to overcome it.

Before moving on, pause and connect the structural model to the trend: metallic bonding explains the high melting points of Na to Al, giant covalent bonding explains the exceptionally high melting point of silicon, and simple molecular or monatomic structures explain the much lower melting points of the non-metals and argon.