AQA A-Level Chemistry: Periodicity (7405)

This guide teaches section 3.2.1 Periodicity of the AQA AS and A-level Chemistry specification (7404/7405), version 1.2, July 2026, for AS and A-level exams June 2016 onwards. It covers both sub-sections: 3.2.1.1 Classification (s, p, d and f blocks) and 3.2.1.2 Physical properties of Period 3 elements (atomic radius, first ionisation energy and melting point from Na to Ar). All of it is AS content, so nothing here is A-level only. The specification places 3.2.1 on AS Paper 1 (inorganic sections 3.2.1 to 3.2.3) and on A-level Paper 1 (all of section 3.2), and A-level Paper 3 can test any content.

For quick recall, use the periodicity revision notes. To test yourself, use the periodicity practice questions. The course hub is AQA A-Level Chemistry, and the printable checklist lists every specification point. If you want to find your weak spots first, try one of the free 10-minute diagnostics.

This topic leans on two earlier ones. Electron configurations and the definition of first ionisation energy come from atomic structure, and the four types of crystal structure come from bonding. The reactions of Period 3 elements and their oxides are a separate topic (3.2.4, A-level only) and are not taught here.

What this unit covers

Spec point	What you must be able to do	Status
3.2.1.1 Classification	Classify an element as s, p, d or f block from its position in the Periodic Table, which is set by its proton number	AS and A-level
3.2.1.2 Atomic radius	Describe and explain the trend in atomic radius across Period 3	AS and A-level
3.2.1.2 First ionisation energy	Describe and explain the trend in first ionisation energy from Na to Ar	AS and A-level
3.2.1.2 Melting point	Explain the melting points of Na to Ar in terms of their structure and bonding	AS and A-level

The specification lists no required practical in section 3.2.1.

3.2.1.1 Classifying elements into blocks

The Periodic Table is arranged in order of proton number (atomic number, Z). Proton number fixes the number of electrons in a neutral atom, and so fixes the electron configuration. That is why proton number decides where an element sits.

An element’s block is named after the sub-shell that holds its highest-energy electron, the one added last when you build the configuration in order of increasing energy.

Block	Highest-energy electron in	Position in the Periodic Table
s block	an s sub-shell	Groups 1 and 2, plus hydrogen and helium
p block	a p sub-shell	the six groups on the right, Group 3 to Group 0 (13 to 18)
d block	a d sub-shell	the central block between Group 2 and Group 3
f block	an f sub-shell	the two rows usually printed below the main table (lanthanides and actinides)

Helium is the one to watch. It sits with the noble gases in Group 0, but its configuration is 1s², so its highest-energy electron is in an s sub-shell. It is an s-block element.

Chromium (3d⁵ 4s¹) and copper (3d¹⁰ 4s¹) have unusual configurations, but both are still d-block elements. Block follows position, and position follows proton number.

Worked example 1: blocks from proton number

Classify the elements with Z = 22, 31 and 38.

Z = 22: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d² 4s²
        last electron added goes into 3d          -> d block (titanium)
Z = 31: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p¹
        last electron added goes into 4p          -> p block (gallium)
Z = 38: [Kr] 5s²
        last electron added goes into 5s          -> s block (strontium)

Notice that Z = 31 has a full 3d sub-shell, but the electron added last is in 4p. Do not classify by the sub-shell that is written last in the line; classify by the one filled last in energy order.

Period 3 at a glance

Every Period 3 element has its outer electrons in the third shell. Moving from Na to Ar adds one proton and one electron each time.

Element	Z	Outer electrons	Block	Structure	Bonding or forces that must be overcome to melt
Na	11	3s¹	s	giant metallic lattice	metallic bonding
Mg	12	3s²	s	giant metallic lattice	metallic bonding
Al	13	3s² 3p¹	p	giant metallic lattice	metallic bonding
Si	14	3s² 3p²	p	macromolecular (giant covalent)	covalent bonds
P	15	3s² 3p³	p	simple molecular, P₄	van der Waals forces
S	16	3s² 3p⁴	p	simple molecular, S₈	van der Waals forces
Cl	17	3s² 3p⁵	p	simple molecular, Cl₂	van der Waals forces
Ar	18	3s² 3p⁶	p	single atoms	van der Waals forces

3.2.1.2 Trend in atomic radius

Trend: atomic radius decreases across Period 3 from sodium to chlorine.

Explanation, in three linked points:

1. Nuclear charge increases: each element has one more proton than the one before.
2. The outer electrons are all in the same (third) shell, so shielding by inner electrons stays about the same.
3. The outer electrons are therefore attracted more strongly to the nucleus and are pulled closer in.

Worked example 2: comparing two radii

Explain why a sulfur atom is smaller than a magnesium atom.

• Sulfur has 16 protons and magnesium has 12, so sulfur has the greater nuclear charge.
• In both atoms the outer electrons are in the third shell, with similar shielding from the inner 1s² 2s² 2p⁶ electrons.
• So the outer electrons in sulfur are attracted more strongly and are held closer to the nucleus.

An answer that only says “sulfur has more protons” is incomplete: it does not say why extra shielding does not cancel the extra protons.

3.2.1.2 Trend in first ionisation energy

First ionisation energy is the energy needed to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions, for example:

Na(g) → Na⁺(g) + e⁻

Trend: first ionisation energy generally increases from Na to Ar. Sodium has the lowest value in the period and argon the highest.

Explanation: the same three factors as the radius trend. Nuclear charge rises, shielding stays about the same because the electron is removed from the same shell, and the atom gets smaller. The outer electron is held more strongly, so more energy is needed to remove it.

The general rise has two small drops: aluminium is lower than magnesium, and sulfur is lower than phosphorus. These are evidence for sub-shells and orbitals, and they are taught (with practice) in the atomic structure study guide and the atomic structure revision notes. In a periodicity answer you should still name both drops when you describe the trend.

Worked example 3: explaining a comparison

Explain why phosphorus has a higher first ionisation energy than magnesium.

Step 1  Nuclear charge:  P has 15 protons, Mg has 12, so P has greater nuclear charge.
Step 2  Shielding:       both outer electrons are in the third shell, so shielding is similar.
Step 3  Distance:        P has a smaller atomic radius.
Step 4  Conclusion:      the outer electron in P is attracted more strongly,
                         so more energy is needed to remove it.

3.2.1.2 Melting points and structure

Melting points rise from sodium to a peak at silicon, then fall sharply. The pattern only makes sense through structure and bonding.

Sodium, magnesium and aluminium: metallic bonding

These three elements are giant metallic lattices. Metallic bonding is the attraction between delocalised electrons and positive ions. Melting needs a lot of energy to weaken these strong attractions.

From Na to Al the melting point increases, because:

• the charge on the metal ion increases: Na⁺, Mg²⁺, Al³⁺
• the number of delocalised electrons per atom increases: 1, 2, 3
• the ions get smaller

So the attraction between the positive ions and the delocalised electrons gets stronger, and more energy is needed to overcome it.

Silicon: macromolecular

Silicon is macromolecular (giant covalent), with each atom covalently bonded to four others in a structure like diamond. To melt it you must break many strong covalent bonds. That needs a great deal of energy, so silicon has the highest melting point in Period 3.

Phosphorus, sulfur and chlorine: simple molecular

These elements are made of small molecules: P₄, S₈ and Cl₂. Melting does not break the covalent bonds inside the molecules. It only overcomes the weak van der Waals forces (induced dipole–dipole forces) between molecules, so the melting points are low.

The strength of van der Waals forces increases with the number of electrons in the molecule, because a larger electron cloud is more easily distorted to form temporary dipoles. So:

S₈  : 8 × 16 = 128 electrons   strongest van der Waals forces   highest melting point of the three
P₄  : 4 × 15 =  60 electrons
Cl₂ : 2 × 17 =  34 electrons   weakest                          lowest melting point of the three

Argon: single atoms

Argon exists as separate atoms. With only 18 electrons per particle, its van der Waals forces are the weakest of all, so argon has the lowest melting point in Period 3.

Worked example 4: one metal against another

Explain why aluminium has a higher melting point than sodium.

• Both are giant metallic lattices held together by metallic bonding.
• Aluminium forms Al³⁺ ions and contributes three delocalised electrons per atom; sodium forms Na⁺ ions and contributes one.
• Al³⁺ is also smaller than Na⁺.
• So the attraction between the ions and the delocalised electrons is stronger in aluminium, and more energy is needed to overcome it.

Worked example 5: crossing a structure boundary

Explain why silicon has a much higher melting point than phosphorus.

• Silicon is macromolecular: melting breaks many strong covalent bonds.
• Phosphorus is simple molecular (P₄): melting only overcomes weak van der Waals forces between molecules; the P–P covalent bonds do not break.
• Far less energy is needed to overcome the van der Waals forces, so phosphorus melts at a much lower temperature.

When the two elements have different structures, your answer must name both structures and say what is broken or overcome in each.

Sketching and reading the trends

You may be asked to sketch a trend or to pick out an element from a graph. Use these shapes:

• Atomic radius: a steady fall from Na to Cl.
• First ionisation energy: a general rise from Na to Ar, with small drops at Al and S, and Ar the highest.
• Melting point: a rise from Na to Al, a sharp peak at Si, a large drop to P, a smaller rise to S, then lower values for Cl and Ar, with Ar the lowest.

On an unlabelled melting point chart, find silicon first (the tallest bar) and count from there.

Common errors

• Explaining the radius trend with “more electrons, so more repulsion”. The extra electrons go into the same shell; the deciding factor is the rising nuclear charge with similar shielding.
• Saying “more shells” across a period. The number of occupied shells is the same for every Period 3 element.
• Writing “covalent bonds break” when phosphorus, sulfur or chlorine melts. Only van der Waals forces are overcome.
• Calling silicon “molecular” or “simple covalent”. It is macromolecular (giant covalent).
• Describing metallic bonding as attraction between atoms, or between ions and ions. It is between positive ions and delocalised electrons.
• Classifying helium as p block because it sits in Group 0.
• Classifying an element such as gallium as d block because the 3d sub-shell appears in its configuration.

Where to go next

Condense this into the revision notes, then test it with the practice questions. For paper structure, see AQA A-Level Chemistry exam preparation.

Official syllabus

AQA, AS and A-level Chemistry specification (7404/7405), version 1.2, July 2026, for AS and A-level exams June 2016 onwards: section 3.2.1 Periodicity (3.2.1.1 Classification and 3.2.1.2 Physical properties of Period 3 elements).