3.7 Giant Covalent Structures

A free GCSE Chemistry lesson for AQA (8462) from OBA School: the structures of diamond, graphite and silicon dioxide, and how their bonding explains their properties. Watch the video, read the explanation and worked examples, then try the revision questions. You can work through the full lesson, with every video, practice quizzes and printable worksheets with mark schemes, in our free online course.

What you will learn

• Recognise giant covalent structures from diagrams
• Explain why substances with giant covalent structures have very high melting points
• Describe and explain the properties of diamond and graphite in terms of their structure and bonding
• Explain why graphite conducts electricity and is soft, and why diamond does not conduct and is hard
• Describe the structure and properties of silicon dioxide

Key words

• Giant covalent structure: A huge 3D network of atoms all joined by strong covalent bonds. Also called a macromolecule.
• Delocalised electron: An electron that is not attached to one atom or one bond and can move through the structure.
• Layer: A flat sheet of atoms. In graphite each layer is made of hexagonal rings of carbon atoms.
• Allotrope: A different form of the same element in the same state, e.g. diamond and graphite are both pure carbon.

What is a giant covalent structure?

In a giant covalent structure, every atom is joined to other atoms by strong covalent bonds, in a network that keeps going. There are no separate molecules, so there are no weak intermolecular forces to overcome.

Key fact: To melt a giant covalent substance you must break very many strong covalent bonds. This needs a huge amount of energy, so the melting and boiling points are very high.

• very high melting and boiling points; Why very many strong covalent bonds must be broken
• insoluble in water; Why water cannot break the strong covalent bonds of the network
• hard (except graphite); Why every atom is held firmly in a rigid 3D network
• do not conduct electricity (except graphite); Why all outer electrons are used in bonds, and there are no ions

Examples you need to know: diamond and graphite (both forms of carbon) and silicon dioxide (silica, the main compound in sand).

Diamond

• very hard (used in cutting and drilling tools); Explanation each carbon atom forms 4 strong covalent bonds in a rigid 3D network
• very high melting point; Explanation very many strong covalent bonds must be broken
• does not conduct electricity; Explanation all 4 outer electrons of each carbon are used in bonds, so there are no delocalised electrons or ions

Graphite

• soft and slippery (used in pencils and as a lubricant); Explanation each carbon forms 3 bonds within its layer; there are no covalent bonds between layers, only weak forces, so the layers can slide over each other
• very high melting point; Explanation very many strong covalent bonds within the layers must be broken
• conducts electricity and heat (used for electrodes); Explanation each carbon has one outer electron not used in bonding; these delocalised electrons can move through the layers and carry charge

The same delocalised electrons make graphite a good thermal conductor: when one end is heated, the electrons gain energy and move faster, carrying the energy quickly along the layers.

Link: Graphite is like a metal in one way: it has delocalised electrons. This is why it conducts electricity, even though it is a non-metal.

Silicon dioxide

In silicon dioxide, each silicon atom is covalently bonded to 4 oxygen atoms, and each oxygen atom is bonded to 2 silicon atoms. The network gives silicon dioxide a high melting point (about 1700 °C) and makes it hard. It has no delocalised electrons or ions, so it does not conduct electricity.

Worked example: why the formula is SiO₂

Each Si has 4 bonds to O; each O has 2 bonds to Si.
Number of Si–O bonds = 4 × (number of Si) = 2 × (number of O)
So there are twice as many O atoms as Si atoms.

Answer: ratio Si : O = 1 : 2, so the formula is SiO₂

Exam tip: In exam answers about melting giant covalent substances, write “covalent bonds must be broken”. Do not mention intermolecular forces: there are none to overcome in diamond or silicon dioxide.

Common mistakes

• Saying graphite conducts because of free ions. It has delocalised electrons.
• Saying the covalent bonds between graphite layers are weak. There are no covalent bonds between layers, only weak forces.
• Saying diamond is hard because the bonds are “hard to break” without mentioning that each atom has 4 bonds in a rigid 3D network.
• Writing “intermolecular forces break” when diamond or silicon dioxide melts. Covalent bonds must be broken.
• Thinking silicon dioxide is made of SiO₂ molecules. It is a giant structure; SiO₂ is the ratio.

Revision questions

Try these, then check your answers below each one.

Question 1. How many covalent bonds does each carbon atom form in diamond?

Answer: 4

Working: Each carbon atom forms 4 covalent bonds in a rigid 3D network.

Question 2. How many covalent bonds does each carbon atom form in graphite?

Answer: 3

Working: Each carbon atom forms 3 covalent bonds within its layer; the fourth electron is delocalised.

Question 3. Why does graphite conduct electricity? (Choose from: Each carbon atom has one delocalised electron that can move through the layers, Graphite contains ions that can move, Graphite is a metal, The layers can slide over each other.)

Answer: Each carbon atom has one delocalised electron that can move through the layers

Working: Each carbon forms only 3 bonds, so one electron per atom is delocalised. These electrons move through the structure and carry charge.

Related lessons

• Previous lesson: 3.6 Structure of Simple Molecules
• Next lesson: 3.8 Fullerenes and Graphene
• In this chapter: 3.1 States of Matter
• In this chapter: 3.2 Atoms and Ions
• In this chapter: 3.3 Ionic Bonding

Continue in the full course

This lesson is section 3.7 (AGCCH0307) of our free GCSE Chemistry course for AQA. In the full course you get every video for this section, a self-marking practice quiz, a printable worksheet with a mark scheme, chapter tests, and your progress saved. Go to the full GCSE Chemistry course.