Giant covalent structures
Diamond, graphite and silicon(IV) oxide — structure, bonding and properties.
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Learning objectives
What you need to be able to do
Teacher-mapped phrasing — check against the official Cambridge syllabus for exact wording.
- 2.6.1Describe the giant covalent structures of diamond and graphite and relate them to their properties and uses.
7 minute read
Giant covalent structures
Some covalent substances do not form small, separate molecules at all — instead, every atom is covalently bonded to its neighbours in one enormous, continuous lattice. These are giant covalent structures, and their properties are completely different from simple molecular substances, even though both involve covalent bonding.
Diamond
In diamond, each carbon atom forms four strong covalent bonds to four other carbon atoms, arranged in a rigid three-dimensional lattice. This makes diamond extremely hard, with a very high melting point — a huge number of strong covalent bonds must be broken to melt it. Because every outer electron is used in a bond, there are no free electrons or ions, so diamond does not conduct electricity. It is used in cutting tools and drill tips because of its hardness.
Graphite
In graphite, each carbon atom bonds to only three others, forming flat hexagonal layers. This leaves one delocalised electron per atom, free to move along the layers — which is why graphite conducts electricity, unusually for a non-metal. The layers themselves are held together by only weak forces, so they can slide over each other easily, making graphite soft and useful as a lubricant and in pencil "lead".
Silicon(IV) oxide
Silicon(IV) oxide (silica, found in sand) has a giant covalent structure similar to diamond, with each silicon atom bonded to four oxygen atoms. Like diamond, it is hard with a very high melting point, because many strong covalent bonds must be broken.
Think of it like this
Diamond is like a climbing frame welded rigidly at every joint in three dimensions — nothing can move without breaking a weld. Graphite is like stacked sheets of chicken wire, each sheet strongly welded internally but merely resting on the sheet below, so the sheets slide apart easily even though the wire within each one is just as strong.
Common misconceptions
- Assuming all covalent substances have low melting points, based on simple molecules like water. Giant covalent structures like diamond have extremely high melting points, because breaking the structure means breaking many strong covalent bonds, not just weak intermolecular forces.
- Thinking graphite conducts because of ions, like a metal or an electrolyte. It conducts because of delocalised electrons, similar in principle to metallic bonding, even though graphite is not a metal.
- Believing diamond and graphite, both pure carbon, should have identical properties. Their very different properties come entirely from their different bonding arrangements (four bonds vs three bonds per atom).
In the exam
- A "compare diamond and graphite" question should be answered in matched pairs — structure of diamond then its property, structure of graphite then its property — rather than describing all of one substance and then all of the other.
- The reason graphite conducts is specifically the delocalised electron per carbon atom — say "delocalised electrons are free to move", not just "graphite has free electrons".