Bonding and Structure: Ionic, Covalent, Metallic Bonding and the Four Crystal Types

This guide covers sections 3.1.3.1 to 3.1.3.4, the bonding-types and crystal-structure half of the Bonding section (3.1.3) in OxfordAQA International AS and A-level Chemistry 9620. It follows Amount of substance (3.1.2) and covers how atoms and ions are held together, and how that bonding determines a substance’s macroscopic properties.

Ionic bonding

You need to understand ionic bonding as electrostatic attraction between oppositely charged ions arranged in a lattice. This includes knowing the formulae of named compound ions – sulfate, hydroxide, nitrate, carbonate and ammonium – and being able to predict the charge on a simple ion from an element’s position in the periodic table, then use that to construct formulae for ionic compounds.

Covalent and dative covalent bonds

You need to understand a single covalent bond as containing one shared pair of electrons, with multiple bonds (double, triple) containing multiple shared pairs. A co-ordinate, or dative covalent, bond is distinct: it also contains a shared pair of electrons, but both electrons in that pair are supplied by a single atom rather than one from each. You need to be able to represent a covalent bond using a line, and a co-ordinate bond specifically using an arrow, pointing from the atom donating both electrons to the atom receiving them.

Metallic bonding

You need to understand metallic bonding as attraction between delocalised electrons and positive ions arranged in a lattice. The “sea of delocalised electrons” model is what explains many of a metal’s characteristic physical properties, covered next.

Predicting ionic charge and constructing formulae

Predicting an ion’s charge from periodic table position follows a pattern: Group 1 elements typically form 1+ ions, Group 2 elements typically form 2+ ions, Group 6 elements typically form 2− ions, and Group 7 elements typically form 1− ions, all reflecting the number of electrons lost or gained to reach a full outer shell. Combined with the named compound-ion formulae above, this lets you construct formulae for a wide range of ionic compounds: balancing the total positive and negative charge to zero, using brackets around a polyatomic ion whenever more than one is needed (for example, Mg(OH)₂, since magnesium’s 2+ charge requires two 1− hydroxide ions to balance it).

Energy changes associated with changes of state

You need to be able to explain the energy changes associated with changes of state in terms of the bonding or intermolecular forces being overcome. Melting and boiling always require energy input, because separating particles further apart means working against the attractive forces holding them together – whichever type of bonding or intermolecular force is present in that particular structure. Critically, when a molecular substance melts or boils, it is the weak intermolecular forces between molecules that break, not the strong covalent bonds within each molecule – a distinction worth stating explicitly in any answer, since conflating the two is a common source of lost marks. By contrast, melting a giant covalent structure like diamond genuinely does require breaking covalent bonds, which is why its melting point is so much higher than a molecular substance’s.

The four crystal structure types

The specification names four types of crystal structure – ionic, metallic, macromolecular (giant covalent), and molecular – and requires knowledge of six named example structures spanning all four types: diamond, graphite, ice, iodine, magnesium, and sodium chloride. You need to be able to relate melting point and conductivity to the type of structure and bonding present, explain the energy changes associated with changes of state, and draw diagrams representing these structures with a specified number of particles.

Crystal type	Named example(s)	Typical melting point	Conducts electricity?
Ionic	Sodium chloride	High	Only when molten or dissolved
Metallic	Magnesium	High	Yes, always
Macromolecular (giant covalent)	Diamond, graphite	Very high	Diamond: no. Graphite: yes (in one plane)
Molecular	Ice, iodine	Low	No

Diamond and graphite are worth studying as a contrasting pair even though both are giant covalent structures of carbon: diamond’s four-bonds-per-carbon tetrahedral network makes it an electrical insulator and extremely hard, while graphite’s layered structure (three bonds per carbon, with one delocalised electron per atom free to move within each layer) makes it a conductor along the layers and soft, because the layers themselves are only held together by weak forces and can slide over one another.

How to approach it

Learn the six named crystal structures with an actual sketch of each, not just a verbal description – questions frequently ask you to draw a structure showing a specified number of particles, and a genuinely practised diagram is far faster to produce accurately under exam conditions than one reconstructed from memory of a description. When explaining melting point differences, connect your answer explicitly to the type and strength of the forces being broken (ionic lattice electrostatic attraction, metallic delocalised-electron attraction, or covalent bonds within a giant structure, versus only weak intermolecular forces in a molecular crystal) rather than a vague answer about “strong bonds” – examiners specifically look for the correct force being named. For conductivity questions, remember the two-part condition for ionic conduction (molten or dissolved, not solid) and the graphite exception among covalent structures (conducts in-plane only, due to its one delocalised electron per carbon atom), since these are the details most often oversimplified in a rushed answer.

Official syllabus

OxfordAQA, International AS and A-level Chemistry (9620) specification, Version 5.3, for International AS and A-level exams May/June 2020 onwards: official specification PDF, sections 3.1.3.1-3.1.3.4 “Bonding”. Verified 2026-09-02.