AQA GCSE Chemistry: Chemical Bonds and Ionic Bonding — Revision Notes

Condensed for the final weeks. For the full explanation, use the Chemical Bonds and Ionic Bonding study guide.

The three bond types (4.2.1.1)

Bond type	Particles involved	Occurs in
Ionic	Oppositely charged ions	Metal + non-metal compounds
Covalent	Atoms sharing pairs of electrons	Non-metallic elements; compounds of non-metals
Metallic	Atoms sharing delocalised electrons	Metallic elements and alloys

Learn these as a matched set from the start: metal + non-metal signals ionic; non-metal + non-metal signals covalent; metal alone (or an alloy) signals metallic. A full-mark answer justifies the identification from the elements’ positions in the periodic table, not just names the bond type. All three types are explained by electrostatic forces and the transfer or sharing of electrons.

Ionic bonding by electron transfer (4.2.1.2)

When a metal reacts with a non-metal, outer-shell electrons transfer: the metal atom loses electrons → positive ion; the non-metal atom gains electrons → negative ion.

Group	Electrons lost/gained	Ion charge
Group 1 (metal)	Loses 1	1+
Group 2 (metal)	Loses 2	2+
Group 6 (non-metal)	Gains 2	2−
Group 7 (non-metal)	Gains 1	1−

In every case, the resulting ion has the electronic structure of the nearest noble gas — this is the rule to state explicitly, not just the final charge.

Dot and cross diagrams

Convention: dots = electrons from one atom; crosses = electrons from the other; the usual convention is to show only outer (valence) shell electrons, though showing all occupied shells is not wrong; enclose each finished ion in square brackets with its charge. A diagram that omits the brackets/charge loses marks even if the electron count is right.

Worked example: magnesium oxide

Magnesium (Group 2):  2 outer electrons; loses both
                       -> Mg2+, electronic structure of neon
Oxygen (Group 6):      6 outer electrons; gains 2 (from magnesium)
                       -> O2-, electronic structure of neon
Result:                oppositely charged Mg2+ and O2- ions held
                       together by strong electrostatic forces of
                       attraction = the ionic bond in MgO

State both ions’ resulting electronic structures and name the noble gas each now matches — this demonstrates full understanding, not just the final formula.

Worked example: sodium chloride, and reading the pattern

Sodium (Group 1):     1 outer electron; loses it
                       -> Na+, electronic structure of neon
Chlorine (Group 7):   7 outer electrons; gains 1 (from sodium)
                       -> Cl-, electronic structure of argon
Result:                Na+ and Cl- attract electrostatically,
                       forming the ionic bond in NaCl

Notice the pattern across both worked examples: a Group 1 or 2 metal always loses electrons equal to its group number, and a Group 6 or 7 non-metal always gains electrons equal to (8 minus its group number). Practising this rule with several different element pairs — not just memorising two fixed examples — is what lets you apply it confidently to an unfamiliar compound in an exam question.

Why ionic compounds behave the way they do

Although 4.2.1.1 and 4.2.1.2 focus on bond formation, it is worth connecting this to the properties ionic compounds show as a result, since exam questions often link the two. The strong electrostatic forces of attraction between oppositely charged ions extend in all directions through a giant ionic lattice, which is why ionic compounds have high melting and boiling points — a large amount of energy is needed to overcome many strong ionic bonds throughout the structure, not just one. This lattice structure, and the properties it produces, is developed further elsewhere in Topic 2, but recognising the link between “how the bond forms” (this sub-topic) and “why the resulting solid behaves as it does” strengthens answers that ask you to explain a property in terms of bonding.

Key terms

Ion — an atom or group of atoms that has lost or gained electrons, giving it an overall charge. Electron transfer — movement of electrons from one atom to another; the mechanism of ionic bonding. Electrostatic force of attraction — the specific named force holding oppositely charged ions together (say this exact phrase, not “they stick together”). Noble gas electronic structure — a full outer shell; the stable arrangement ions form towards. Dot and cross diagram — shows only outer-shell electrons, dots for one atom, crosses for the other.

Common mistakes

• Naming a bond type without justifying it from the elements involved.
• Omitting the square brackets and charge around a finished ion in a dot and cross diagram.
• Forgetting square brackets and the ion’s charge in the finished diagram.
• Applying “charge = group number” outside Groups 1, 2, 6 and 7 — the specification only requires (and the simple rule only reliably applies to) these four groups.
• Describing ionic bond formation as one step (“electrons move”) instead of the required two-step sequence: transfer, then electrostatic attraction between the ions formed.

Quick self-test

• State which bond type forms between a metal and a non-metal, and justify why.
• Work out the ion charge formed by a Group 2 metal and a Group 6 non-metal.
• List the two conventions (dots/crosses; outer shell only) required in a correct dot and cross diagram.
• Explain, in two steps, how sodium chloride’s ionic bond forms.
• State the specific named force holding ions together in an ionic compound.

Related resources

• Chemical Bonds and Ionic Bonding study guide
• Chemical Bonds and Ionic Bonding practice questions

Official syllabus

AQA GCSE Chemistry (8462) specification, Version 1.1, exams June 2018 onwards — aqa.org.uk/8462.