OxfordAQA A-Level Chemistry: Periodicity (9620)

Here you will learn section 3.2.1 Periodicity from the OxfordAQA International AS and A-level Chemistry (9620) specification, Version 5.3, for International AS and A-level exams May/June 2020 onwards. The section has two points, 3.2.1.1 (Classification) plus 3.2.1.2 (the Physical properties of Period 3 elements), and this guide teaches both. Everything here is International AS content, so nothing is International A-level only. The specification examines 3.2.1 in Unit 1 (Inorganic 1 and Physical 1); Unit 5 (Practical and synoptic), taken at International A2, draws on all content.

For a condensed version, open the periodicity revision notes. To test yourself, work through the periodicity practice questions. Other 9620 topics sit on the OxfordAQA A-Level Chemistry hub, and you can track each point with the printable checklist. Not sure where you stand? A free 10-minute diagnostic will show you.

Background you need: sub-shells and ionisation energy from atomic structure, crystal structures from bonding and structure, and van der Waals forces from intermolecular forces.

What this unit covers

Point	You should be able to	Stage
3.2.1.1 Classification	Say whether an element belongs to the s, p, d or f block, using its place in the table (set by proton number)	International AS
3.2.1.2 Atomic radius	Account for the way atoms shrink across Na–Ar	International AS
3.2.1.2 First ionisation energy	Account for the way first ionisation energy changes across Na–Ar, including the two dips	International AS
3.2.1.2 Melting point	Link each melting point, Na–Ar, to the element’s structure and bonding	International AS

The specification’s list of required practical activities has no entry for 3.2.1. Its overview for the section mentions the history of the Periodic Table, but the specification says overviews are not directly assessed.

3.2.1.1 Classification into blocks

Elements are placed in the Periodic Table in order of proton number. The proton number tells you how many electrons a neutral atom has, and those electrons fill sub-shells in a fixed energy order:

1s  2s  2p  3s  3p  4s  3d  4p

So proton number fixes position, and position fixes the block: the type of sub-shell (s, p, d or f) that holds the element’s highest-energy electron.

Block	Columns or rows	Sub-shell being filled
s	Groups 1 and 2, with hydrogen and helium	s
p	Groups 3 to 0 (13 to 18), apart from helium	p
d	the ten columns between Group 2 and Group 3	d
f	the lanthanide and actinide rows, printed under the main table	f

Worked example 1: proton number to block

Classify the elements with proton numbers 3, 25 and 35.

Z = 3   1s² 2s¹                                   last electron in 2s  -> s block (lithium)
Z = 25  1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵ 4s²               last electron in 3d  -> d block (manganese)
Z = 35  1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁵          last electron in 4p  -> p block (bromine)

For manganese, 4s fills before 3d, so the last electron added goes into 3d. For bromine, 3d is already full and filling has moved on to 4p.

Worked example 2: position to block and configuration

An element sits in Period 4, Group 5 (15). State its block and write its electron configuration.

1. Groups 3 to 0 make up the p block, so the element is p block.
2. Period 4 means its outer shell is the fourth shell, and Group 5 means five outer electrons: 4s² 4p³.
3. Before 4p fills, 3d must be full. The configuration is 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p³ (Z = 33, arsenic).

Any element in the lanthanide or actinide rows is f block, even though the 9620 configuration work stops at Z = 36. You classify it from its position alone.

Period 3 data

Every element from sodium to argon has its outer electrons in the third shell. Each step to the right adds one proton to the nucleus and one electron to that shell. The values below are rounded from standard data tables (phosphorus as white P₄; sulfur as S₈).

Element	Z	Block	First ionisation energy / kJ mol⁻¹	Melting point / K	Structure
Na	11	s	496	371	giant metallic
Mg	12	s	738	923	giant metallic
Al	13	p	578	933	giant metallic
Si	14	p	787	1687	macromolecular (giant covalent)
P	15	p	1012	317	molecular, P₄
S	16	p	1000	388	molecular, S₈
Cl	17	p	1251	172	molecular, Cl₂
Ar	18	p	1521	84	separate atoms

3.2.1.2 Atomic radius

Pattern: atoms get smaller from sodium to chlorine.

Why, as a chain of three ideas:

• Nuclear charge rises. There is one extra proton at each step.
• Shielding hardly changes. The new electron goes into the third shell, the same shell as the other outer electrons. The inner electrons (1s² 2s² 2p⁶) are the same ten in every case.
• Attraction wins. With more charge pulling and similar shielding, the outer electrons are drawn nearer to the nucleus.

Worked example 3: magnesium against silicon

Explain why a silicon atom has a smaller radius than a magnesium atom.

• Silicon has 14 protons; magnesium has 12. Silicon’s nuclear charge is larger.
• Both atoms keep their outer electrons in the third shell, behind the same ten inner electrons, so the shielding is about the same.
• The silicon nucleus therefore attracts its outer electrons more strongly and holds them closer.

3.2.1.2 First ionisation energy

First ionisation energy: the energy required for a mole of gaseous atoms to lose a mole of electrons, leaving a mole of gaseous 1+ ions. For silicon:

Si(g) → Si⁺(g) + e⁻

Pattern: a general rise from 496 kJ mol⁻¹ (Na) to 1521 kJ mol⁻¹ (Ar), the highest in the period. Two values break the rise: aluminium is below magnesium, and sulfur is below phosphorus.

The general rise comes from the same three ideas as the radius trend. Nuclear charge increases, the electron always leaves the third shell so shielding stays similar, and the atom is smaller. The outer electron is held more tightly, so more energy is needed to remove it.

The two drops are evidence for sub-shells (see atomic structure). A full account of the trend names them and gives a reason for each.

Worked example 4: reading the data

Use the data table to explain (a) why silicon’s first ionisation energy is higher than aluminium’s, and (b) why aluminium’s is lower than magnesium’s.

(a) Si 787, Al 578: Si is higher by 209 kJ mol⁻¹.
    Both lose an electron from the 3p sub-shell.
    Si has one more proton (14 v 13), similar shielding and a smaller radius,
    so its 3p electron is held more strongly.

(b) Mg 738, Al 578: Al is lower by 160 kJ mol⁻¹.
    Mg loses a 3s electron; Al loses its single 3p electron.
    The 3p sub-shell is higher in energy than 3s, and the 3p electron is
    partly shielded by the 3s electrons, so it needs less energy to remove.

The drop from phosphorus (1012) to sulfur (1000) is small. In phosphorus each 3p orbital holds one electron. In sulfur one 3p orbital holds a pair, and the two electrons in it repel each other, so one of them is a little easier to remove.

3.2.1.2 Melting point

Across Period 3 the structure changes from metallic to macromolecular to molecular to single atoms. Melting point depends on what must be overcome, so always name the structure first.

Shape of the pattern: up from Na to Al, a sharp peak at Si (1687 K), a collapse at P, a rise to S, then down through Cl to Ar (84 K), the lowest in the period.

Sodium, magnesium and aluminium

These are giant metallic lattices: positive ions held together by attraction to delocalised electrons. From Na to Al:

• the ions carry a larger charge: 1+, 2+, 3+
• each atom gives more electrons to the delocalised sea: one, two, three
• the ions are smaller

Each of these changes makes the metallic bonding stronger, so it takes more energy to melt the metal. The data show a large rise from Na to Mg (552 K) and a much smaller one from Mg to Al (10 K), but the direction is the same.

Silicon

Silicon is macromolecular (giant covalent): every atom is joined to four neighbours by covalent bonds, in a three-dimensional network like diamond. Melting means breaking a great many strong covalent bonds, so silicon has the highest melting point in the period, 754 K above aluminium.

Phosphorus, sulfur and chlorine

These are molecular: P₄, S₈ and Cl₂. When they melt, the covalent bonds inside each molecule stay intact. Melting just has to overcome the van der Waals forces between molecules, which are weak, so these melting points are low.

The more electrons a molecule has, the stronger its van der Waals forces: a bigger electron cloud distorts more readily, so temporary dipoles form more easily and are larger.

Argon

Argon is made of separate atoms. Each has only 18 electrons, so the van der Waals forces between atoms are the weakest in the period, and argon’s melting point is the lowest.

Worked example 5: sulfur against phosphorus

Explain why sulfur has a higher melting point than phosphorus.

1. Structures: both are molecular. Sulfur is S₈; phosphorus is P₄.
2. What melting overcomes: the van der Waals forces holding the molecules together. The S–S and P–P bonds inside the molecules survive.
3. Comparison: an S₈ molecule has 8 × 16 = 128 electrons (Mr 256.8). A P₄ molecule has 4 × 15 = 60 electrons (Mr 124.0). The larger electron cloud in S₈ gives stronger van der Waals forces.
4. Link to energy: more energy is needed to separate S₈ molecules, so sulfur melts at a higher temperature (388 K against 317 K in the table).

Worked example 6: crossing from metal to molecule

Explain why aluminium has a much higher melting point than chlorine.

• Aluminium is a giant metallic lattice. Melting has to overcome the strong attraction between Al³⁺ ions and the delocalised electrons throughout the lattice.
• Chlorine is molecular (Cl₂). To melt it you separate whole Cl₂ molecules by overcoming weak van der Waals forces; the Cl–Cl bond itself is untouched.
• Far more energy is needed for aluminium, so its melting point is far higher (933 K against 172 K).

Reading Period 3 graphs

Questions often give a bar chart or line graph with the elements unlabelled. Find a landmark first:

• Melting point: silicon is the tallest bar; argon is the shortest.
• First ionisation energy: argon is the highest; sodium is the lowest; the two small dips sit at aluminium and sulfur.
• Atomic radius: a steady fall, with sodium the largest.

Then count along from the landmark in order of proton number.

Common errors

• Claiming the number of shells increases across Period 3. All eight elements use the third shell for their outer electrons.
• Explaining the radius trend with nuclear charge alone. Similar shielding is a separate point and needs to be stated.
• Writing that covalent bonds break when P₄, S₈ or Cl₂ melts.
• Calling silicon “molecular”, or crediting it with strong intermolecular forces.
• Describing metallic bonding as attraction between atoms. It is between positive ions and delocalised electrons.
• Leaving out the state symbols, or using X(s), in a first ionisation energy equation.

Next steps

Next, shrink all this down with the revision notes and check your answers against the practice questions.

Official syllabus

Source document: OxfordAQA, International AS and A-level Chemistry (9620) specification, Version 5.3, for International AS and A-level exams May/June 2020 onwards. Section 3.2.1 Periodicity (International AS): point 3.2.1.1, Classification; point 3.2.1.2, on the Physical properties of Period 3 elements.