AQA GCSE Chemistry 8462: Chemical bonds and ionic compounds – Study Guide

This guide teaches two subsections of the AQA GCSE Chemistry (8462) specification, for teaching from September 2016 onwards and exams in 2018 onwards (version 1.1): 4.2.1.1 Chemical bonds and 4.2.1.3 Ionic compounds. Both sit in topic 2, Bonding, structure, and the properties of matter, which is assessed on Paper 1 (1 hour 45 minutes, Foundation and Higher Tier). Nothing in these two subsections is marked (HT only), so all of it can be tested on either tier. There is no required practical for this content.

How electrons move to make ions (4.2.1.2, dot and cross diagrams and ion charges from group number) has its own pages: the chemical bonds and ionic bonding study guide, its revision notes and its practice questions. This guide goes further into what those pages only touch: explaining every bond type with electrostatic forces, and the giant ionic lattice itself. For quick recall use the chemical bonds and ionic compounds revision notes, and test yourself with the chemical bonds and ionic compounds practice questions. The course hub is AQA GCSE Chemistry, and the printable checklist lists every specification point.

What this unit covers

Spec point	What you must be able to do	Tier
4.2.1.1 Chemical bonds	Name the three strong bonds and the particles in each; say which kinds of substance have each bond	Both
4.2.1.1 Chemical bonds	Explain bonding in terms of electrostatic forces and the transfer or sharing of electrons	Both
4.2.1.3 Ionic compounds	Describe an ionic compound as a giant structure of ions held by electrostatic forces acting in all directions	Both
4.2.1.3 Ionic compounds	Deduce that a compound is ionic from a diagram of its structure	Both
4.2.1.3 Ionic compounds	Describe the limitations of dot and cross, ball and stick, 2D and 3D diagrams of a giant ionic structure	Both
4.2.1.3 Ionic compounds	Work out the empirical formula of an ionic compound from a model or diagram	Both

The specification expects you to be familiar with the structure of sodium chloride. You do not need to know the structures of other ionic compounds. If a question uses another compound, it will give you the diagram.

4.2.1.1 The three strong chemical bonds

There are three types of strong chemical bond: ionic, covalent and metallic. Each is described by the particles involved.

Bond	Particles	Found in
Ionic	Oppositely charged ions	Compounds formed from metals combined with non-metals
Covalent	Atoms that share pairs of electrons	Most non-metallic elements, and compounds of non-metals
Metallic	Atoms that share delocalised electrons	Metallic elements and alloys

Read the “found in” column carefully.

• Ionic bonding needs a compound. A single element cannot be ionic.
• Covalent bonding covers both elements (such as chlorine, Cl₂) and compounds (such as water, H₂O). The word “most” matters: the Group 0 noble gases do not easily form molecules, because their atoms already have stable arrangements of electrons.
• Metallic bonding covers pure metals and alloys. Brass is a mixture of metals, but its bonding is still metallic.

Explaining bonding with electrostatic forces

The specification asks you to explain chemical bonding in terms of electrostatic forces and the transfer or sharing of electrons. Every strong bond is an attraction between opposite charges. What changes is where the charges come from.

• Ionic: electrons are transferred from metal atoms to non-metal atoms. The metal atoms become positive ions and the non-metal atoms become negative ions. The bond is the strong electrostatic attraction between the oppositely charged ions.
• Covalent: two atoms share a pair of electrons. The shared pair is negative. It is attracted to the positive nuclei of both atoms, which holds the two atoms together.
• Metallic: the outer-shell electrons of metal atoms are delocalised, free to move through the whole structure. The sharing of these delocalised electrons gives strong metallic bonds: the negative delocalised electrons attract the positively charged metal centres they move among.

Worked example 1: classify and explain

For each substance, name the type of bonding and explain how it holds the particles together: potassium bromide, hydrogen chloride, copper.

Potassium bromide
  K is a metal, Br is a non-metal       -> compound of metal + non-metal -> IONIC
  Electrons transfer from K to Br       -> K+ ions and Br- ions
  Strong electrostatic attraction between the oppositely charged ions

Hydrogen chloride
  H and Cl are both non-metals          -> compound of non-metals -> COVALENT
  Atoms share a pair of electrons
  The shared pair is attracted to the positive nuclei of both atoms

Copper
  A metallic element                    -> METALLIC
  Outer electrons are delocalised and shared through the structure
  Electrostatic attraction between metal centres and delocalised electrons

Each answer has three parts: the elements, the bond type, then how electrons are transferred or shared and what attracts what. “They share electrons” alone does not explain the bond.

Worked example 2: is it ever “none of these”?

Classify: argon, bronze (an alloy of copper and tin), sulfur dioxide, magnesium fluoride.

Substance	Elements	Bonding
Argon	One noble gas	No strong bonds; atoms stay separate
Bronze	Two metals (alloy)	Metallic
Sulfur dioxide	Two non-metals	Covalent
Magnesium fluoride	Metal + non-metal	Ionic

Bronze contains two different elements, but that does not make it ionic. Ionic bonding needs a metal and a non-metal.

4.2.1.3 Ionic compounds: the giant lattice

An ionic compound is a giant structure of ions. It is held together by strong electrostatic forces of attraction between oppositely charged ions. These forces act in all directions in the lattice, and this is called ionic bonding.

Two ideas here trip people up.

1. There are no molecules. Sodium chloride is not made of separate “NaCl” pairs. Each ion is attracted to every oppositely charged ion around it, not to one partner.
2. “Giant” means the pattern repeats in three dimensions through the whole crystal. The formula NaCl tells you the ratio of ions, 1 : 1, not the size of a unit.

The structure of sodium chloride

In solid sodium chloride:

• sodium ions (Na⁺) and chloride ions (Cl⁻) alternate in a regular cubic pattern
• each Na⁺ ion is surrounded by six Cl⁻ ions, and each Cl⁻ ion by six Na⁺ ions
• the chloride ions are larger than the sodium ions
• the pattern continues in every direction, so there are equal numbers of each ion.

Worked example 3: deducing that a structure is ionic

Four diagrams are described. Which one shows an ionic compound?

• A Separate pairs of identical atoms, each pair joined by a single line.
• B One kind of positive particle in rows, with “e⁻” symbols scattered between them.
• C A regular grid of two kinds of particle, one labelled + and the other −, alternating, with the grid cut off at the edges of the drawing.
• D One kind of atom, every atom joined to four others by lines.

Answer: C. It shows two kinds of charged particle (positive and negative ions) in a regular repeating lattice that continues past the edge of the drawing (a giant structure). A shows small molecules (covalent). B shows positive centres among delocalised electrons (metallic): there are no negative ions. D shows one element joined by covalent bonds in a giant structure.

Useful clues for “ionic” in any diagram:

• charges (+ and −, or 2+ and 2−) on the particles
• square brackets with a charge around each ion in a dot and cross diagram
• a key naming a metal ion and a non-metal ion
• a regular, alternating pattern with no separate groups of particles.

Limitations of the diagrams

You must be able to describe the limitations of using dot and cross, ball and stick, two-dimensional and three-dimensional diagrams to represent a giant ionic structure. Each one shows some features well and hides others.

Diagram	What it shows well	Limitations
Dot and cross	Which electrons transfer, and the charge on each ion	Shows one pair of ions, not the giant lattice; gives no idea of the 3D arrangement or the relative sizes of the ions; does not show the forces acting in all directions
Ball and stick	The 3D arrangement and the alternating pattern	Sticks suggest physical bonds between particular pairs, when the forces act in all directions; ions are drawn far apart, when really they are packed closely; ion sizes may not be to scale; only a tiny part of the lattice
2D diagram	The alternating pattern and the ratio of ions	Shows only one layer, so you cannot see the ions above and below; does not show the 3D shape or the forces in all directions
3D space-filling	The relative sizes of the ions and how closely they pack	You cannot see the ions inside, so the arrangement is hard to follow; the forces are not shown; only a tiny part of the lattice

One limitation is shared by every model: a real crystal contains an enormous number of ions, so any drawing shows only a small fragment of the lattice.

Worked example 4: empirical formula from a 2D diagram

The diagram shows part of one layer of lithium oxide. L = lithium ion, O = oxide ion. Work out the empirical formula.

L  O  L  L  O  L
L  O  L  L  O  L
L  O  L  L  O  L

Count L:   4 per row x 3 rows = 12
Count O:   2 per row x 3 rows = 6
Ratio L : O = 12 : 6 = 2 : 1           (divide both by 6)
Empirical formula: Li2O  (written Li₂O)
Check charges: 2 x (+1) + 1 x (-2) = 0  -> neutral, so the ratio makes sense

The empirical formula is the simplest whole-number ratio of the ions. Write the metal first and use subscripts only for numbers greater than 1. The charge check is a useful habit, but it only works when you know the charges. The specification limits charges from group number to Groups 1, 2, 6 and 7 (lithium is in Group 1 and oxygen in Group 6).

Worked example 5: watch the edges of a 3D model

A model of sodium chloride is a 3 × 3 × 3 cube of 27 ions, with Na⁺ ions at the eight corners. A student counts 14 Na⁺ and 13 Cl⁻ and writes Na₁₄Cl₁₃. What has gone wrong?

Ions are alternating, so corners, edges and faces do not split evenly
in a small block: 14 of the corner type, 13 of the other.
In the real lattice the pattern carries on in every direction,
so the numbers of Na+ and Cl- are equal.
Check: Na+ is 1+, Cl- is 1-, so 1 : 1 is needed for a neutral compound.
Empirical formula: NaCl

A small model has edges that matter; a real crystal is so large that they do not. When a count gives an odd ratio, look at the repeating pattern and check the charges.

Common errors

• Calling an alloy ionic because it contains two elements. Alloys are metallic.
• Saying ionic compounds are made of “molecules of NaCl”. They are giant lattices.
• Writing “the electrons are shared” for a covalent bond without saying the shared pair is attracted to both nuclei.
• Saying the forces in an ionic lattice act “between each pair of ions”. They act in all directions.
• Giving a limitation that is just “it’s not realistic”. Name the specific feature: sticks, gaps, size, layer or fragment.
• Copying raw counts as a formula (Li₁₂O₆) instead of the simplest ratio (Li₂O).
• Treating a diagram with positive particles and electrons as ionic. Ionic diagrams have negative ions too.

Next steps

• Chemical bonds and ionic compounds revision notes
• Chemical bonds and ionic compounds practice questions
• AQA GCSE Chemistry course hub and printable checklist
• All free 10-minute diagnostics

Official syllabus

AQA GCSE Chemistry (8462) specification, for teaching from September 2016 onwards, exams in 2018 onwards, version 1.1 (4 October 2019), published by AQA – sections 4.2.1.1 Chemical bonds and 4.2.1.3 Ionic compounds.