OCR A Level Chemistry A: Periodicity (H432)

This study guide teaches section 3.1.1, Periodicity, of the OCR A Level Chemistry A (H432) specification, Version 3.1 (May 2026), for first assessment in 2017. It opens Module 3, Periodic table and energy, and covers every lettered outcome, 3.1.1(a) to (g), with its additional guidance. H432 is a full A level, not split into AS and A2 here. According to the specification, components 01 (Periodic table, elements and physical chemistry) and 03 (Unified chemistry) examine Module 3. This section has no PAG link.

All H432 topics are linked from the OCR chemistry hub. Tick off each outcome on the printable checklist, and use the free diagnostics to find your weak spots.

Outcomes in 3.1.1

Outcome	Skill	Guidance notes
(a)(i)-(iii)	Describe the table: ordered by proton number, periods, groups	HSW: Periodic Law; new elements
(b)(i)-(ii)	Configurations across Periods 2 and 3; s-, p- and d-blocks	See 2.2.1(d)
(c)(i)-(ii)	Define first ionisation energy; explain trends; use successive values	M3.1; dips; Bohr model
(d)(i)-(ii)	Metallic bonding; giant metallic lattice	No packing details
(e)	Diamond, graphite, graphene, silicon	Graphene’s uses (HSW)
(f)	Properties of giant metallic and giant covalent lattices	Particles, forces, mobility
(g)	Melting points across Periods 2 and 3	M3.1; see 2.2.2(o)

You need the ideas from Electron structure and Bonding and structure before you start.

3.1.1(a): How the periodic table is arranged

1. Elements are placed in order of increasing atomic number, which is the number of protons in the nucleus.
2. A horizontal row is a period. Properties change in a pattern along each period, and the pattern repeats in the next: this is periodicity.
3. A vertical column is a group. Elements in a group have the same outer-shell electron arrangement, so they react in similar ways.

How the idea was accepted (HSW1, 7, 11)

In 1869 Dmitri Mendeleev ordered the known elements mainly by atomic mass but kept similar elements in the same column, leaving gaps and predicting the missing elements. Germanium, found in 1886, closely matched his prediction for the element below silicon, and successes like this won acceptance for the Periodic Law. Ordering by proton number later removed misfits such as tellurium and iodine.

New elements are added only after claims are checked independently. In November 2016 IUPAC approved the names nihonium, moscovium, tennessine and oganesson for elements 113, 115, 117 and 118, completing the seventh row.

3.1.1(b): Electron configurations and blocks

Across Periods 2 and 3 the outer sub-shells fill in the same order. Each period starts by adding an s electron to a new shell and ends with that shell’s s and p sub-shells full.

• Period 2: Li 2s¹, Be 2s², B 2s² 2p¹, C 2s² 2p², N 2s² 2p³, O 2s² 2p⁴, F 2s² 2p⁵, Ne 2s² 2p⁶ (each after 1s²).
• Period 3: Na 3s¹ through to Ar 3s² 3p⁶, in the same pattern (each after [Ne]).

Each element in Period 3 has the same outer arrangement as the element above it, with the principal quantum number raised by one. That is why groups behave alike.

The block an element belongs to is the type of sub-shell (s, p or d) occupied by its highest-energy electron. Groups 1 and 2 form the s-block, Groups 13 to 18 the p-block, and the transition-metal columns the d-block. Helium (1s²) is placed in Group 18 but is an s-block element.

Worked example 1: block, group and period from proton number

Classify elements with Z = 56, 40 and 51.

• Z = 56: [Xe] 6s². The last electron enters 6s, so s-block, Group 2, Period 6 (barium).
• Z = 40: [Kr] 4d² 5s². The last electron enters 4d, so d-block, Period 5 (zirconium).
• Z = 51: [Kr] 4d¹⁰ 5s² 5p³. The last electron enters 5p, so p-block, Group 15, Period 5 (antimony).

The period is the highest shell number in use; for s- and p-block elements, the group follows from the outer s and p electrons.

3.1.1(c): Ionisation energy

The first ionisation energy is the energy required when 1 mol of gaseous atoms of an element each lose one electron, giving 1 mol of gaseous 1+ ions. For carbon:

C(g) → C⁺(g) + e⁻

Show (g) on the atom and the ion. Successive ionisation energies remove further electrons one at a time; the third for carbon is C²⁺(g) → C³⁺(g) + e⁻. Only the first ionisation energy needs a formal definition.

Three factors set the attraction between the nucleus and the outer electron:

• nuclear charge: more protons attract the outer electron more strongly;
• atomic radius: a more distant outer electron is attracted less;
• shielding: inner-shell electrons repel the outer electron and reduce the attraction it feels.

Trend across Periods 2 and 3

Across a period, nuclear charge rises, electrons are added to the same shell so shielding stays about the same, and atomic radius falls. The outer electron is held more strongly, so first ionisation energy rises overall. In Period 2 it rises from 520 kJ mol⁻¹ for lithium to 2081 kJ mol⁻¹ for neon.

Two small falls break the rise: from Group 2 to Group 13 (Be to B, Mg to Al) and from Group 15 to Group 16 (N to O, P to S). The same explanations work in both periods.

Worked example 2: two dips in Period 2

Use these values (kJ mol⁻¹): Be 900, B 801, N 1402, O 1314. Explain each fall.

• Be 1s² 2s² against B 1s² 2s² 2p¹. Boron has one more proton, but its outer electron is in 2p, higher in energy than 2s. The first ionisation energy falls by 900 − 801 = 99 kJ mol⁻¹.
• N 1s² 2s² 2p³ against O 1s² 2s² 2p⁴. Nitrogen’s three 2p electrons are unpaired, one per orbital. Oxygen’s fourth 2p electron shares an orbital, so repulsion between the pair helps remove it. The value falls by 1402 − 1314 = 88 kJ mol⁻¹.

Trend down a group

Down a group, first ionisation energy decreases. Each element has an extra shell, so radius and shielding both increase. These outweigh the larger nuclear charge, and the outer electron is held less tightly. In Group 18, the values are Ne 2081, Ar 1521 and Kr 1351 kJ mol⁻¹.

Successive ionisation energies and shells

Each successive ionisation energy is larger, because each electron leaves an increasingly positive ion. A much larger rise shows that the next electron sits in an inner shell, nearer the nucleus and less shielded. Counting the electrons before each large jump gives the number in each shell (M3.1: read them from a table or a graph, often plotted on a log scale).

Worked example 3: oxygen’s eight ionisation energies

Use these values for oxygen (kJ mol⁻¹): 1314, 3388, 5301, 7469, 10990, 13327, 71330, 84078.

• Up to the 6th value, each ratio to the one before is between 1.2 and 2.6; then 71330 ÷ 13327 = 5.35.
• So six electrons come from the outer shell and two from an inner shell: shells hold 2, 6.
• Six outer electrons places oxygen in Group 16. In the same way, values of 801, 2427, 3660 and 25026 kJ mol⁻¹ jump after the 3rd (ratio 6.84): Group 13, boron.

Steady rises followed by sudden jumps were evidence for electrons in separate shells, which supported the Bohr model of the atom (HSW1, 2).

3.1.1(d): Metallic bonding

Metallic bonding is the strong electrostatic attraction between cations (positive ions) and delocalised electrons. Each atom releases its outer electrons into a “sea” shared by the whole structure, and the cations sit in a regular, repeating giant metallic lattice. All metals have this structure; packing patterns are not needed.

Worked example 4: lithium against beryllium

Lithium melts at 181 °C and beryllium at 1287 °C. Explain the difference of 1106 °C.

• Both are giant metallic. Lithium forms Li⁺ and gives one delocalised electron per atom; beryllium forms Be²⁺ and gives two.
• Be²⁺ is smaller and more highly charged, and there are more delocalised electrons.
• So the electrostatic attraction between cations and delocalised electrons is stronger in beryllium, and more energy is needed to break it.

3.1.1(e): Giant covalent lattices

A giant covalent lattice is a network of atoms bonded by strong covalent bonds. You need four examples.

• Diamond: every carbon is bonded to four others in a tetrahedral network. There are no free electrons.
• Graphite: each carbon is bonded to three others in flat hexagonal layers. The fourth outer electron from each atom is delocalised across its layer. The layers are held together by weak London forces.
• Graphene: a single layer of graphite, one atom thick. Each carbon is bonded to three others, with delocalised electrons across the sheet.
• Silicon: each silicon atom is bonded to four others in a tetrahedral network, like diamond.

Graphene (HSW1, 9)

Graphene’s strength comes from the strong covalent bonds across the whole sheet. It conducts electricity because the delocalised electrons can move along the sheet. Possible uses include flexible electronics, sensors and strong, lightweight composites, if it can be made cheaply at scale.

3.1.1(f): Physical properties of the two giant lattices

Property	Giant metallic	Giant covalent
Melting and boiling points	High; strong attraction between cations and delocalised electrons must be overcome	Very high; many strong covalent bonds must be broken
Solubility	Insoluble in water and in non-polar solvents (some metals react with water instead)	Insoluble; the covalent bonds are too strong to be broken by a solvent
Electrical conductivity	Conducts when solid and when molten; delocalised electrons move through the lattice	Diamond does not conduct: every outer electron is held in a covalent bond. Silicon conducts only weakly (a semiconductor). Graphite and graphene conduct: delocalised electrons move along the layer or sheet

Always name the particles, the bonds or forces between them, and whether anything can move.

3.1.1(g): Melting points across Periods 2 and 3

Melting points rise across the metals, peak at Group 14, then drop sharply to the simple molecular elements and the noble gas.

Period 2	Melting point / °C	Structure	Period 3	Melting point / °C	Structure
Li	181	giant metallic	Na	98	giant metallic
Be	1287	giant metallic	Mg	650	giant metallic
B	2076	giant covalent	Al	660	giant metallic
C	sublimes above 3500	giant covalent	Si	1414	giant covalent
N₂	−210	simple molecular	P₄	44	simple molecular
O₂	−219	simple molecular	S₈	115	simple molecular
F₂	−220	simple molecular	Cl₂	−102	simple molecular
Ne	−249	single atoms	Ar	−189	single atoms

• Metals: ion charge and the number of delocalised electrons rise, so metallic bonding gets stronger.
• Giant covalent: carbon (as diamond or graphite), boron and silicon have the highest values because melting breaks strong covalent bonds throughout the lattice.
• Simple molecular: melting only overcomes weak London forces between molecules; the covalent bonds inside them are not broken. In Period 3 the order S₈ > P₄ > Cl₂ > Ar follows the number of electrons per particle (128, 60, 34, 18): more electrons, stronger London forces.

Worked example 5: boron against nitrogen

Boron melts at 2076 °C and nitrogen at −210 °C, a gap of 2286 °C. Boron is giant covalent, so melting breaks many strong covalent bonds. Nitrogen is simple molecular, N₂; melting overcomes only weak London forces between molecules, and the strong N≡N triple bond stays intact.

Common errors

• Leaving out (g) in an ionisation equation, or writing an e⁻ gained rather than lost.
• Explaining the B dip with “a full 2s sub-shell is stable”; the reason is the higher energy of the 2p sub-shell.
• Saying shielding increases across a period.
• Counting the electrons after the big jump instead of before it.
• Writing that melting a simple molecular element breaks covalent bonds.
• Describing metallic bonding as attraction between atoms rather than cations and delocalised electrons.

Next steps

• Condensed tables and a self-test: Periodicity revision notes
• Original questions with marked answers: Periodicity practice questions
• Find gaps across the course: all free 10-minute diagnostics

Official syllabus

OCR A Level Chemistry A (H432) specification, Version 3.1 (May 2026), for first assessment in 2017, section 3.1.1 Periodicity, learning outcomes (a) to (g). Published by OCR.