Diamond

Summary: Giant covalent allotrope of carbon in which every atom bonds tetrahedrally to four others — the hardest natural substance, with a very high melting point and no electrical conductivity; used in cutting tools. Tags: igcse chemistry Created: 2026-07-18

Diamond is a giant covalent structure of the element carbon and, together with Graphite, one of the two allotropes of carbon required at IGCSE. Each carbon atom is joined by strong covalent bonds to four other carbon atoms arranged tetrahedrally, producing a rigid, continuous three-dimensional giant lattice with no separate molecules — the whole crystal is effectively one giant molecule. Because every one of carbon’s four outer electrons is locked into a covalent bond, there are no free or delocalised electrons, so diamond does not conduct electricity, in sharp contrast to graphite. Breaking the network requires breaking huge numbers of strong covalent bonds, which makes diamond the hardest naturally occurring substance and gives it a very high melting point. These properties directly explain its main practical use in cutting tools and drill bits, and the same tetrahedral network is mirrored in Silicon(IV) Oxide, which behaves very similarly.

Structure

• A giant covalent structure (giant lattice): strong covalent bonding extends in all directions throughout the crystal.
• Each carbon atom is covalently bonded to 4 other carbon atoms.
• The four bonds point to the corners of a tetrahedron (bond angle 109.5°).
• The result is a rigid, interlocking 3D network — there are no layers, no weak regions, and no separate molecules.
• All four outer electrons of every carbon atom are used in bonding, so none are delocalised.

Structure–Property–Use Table

Why Diamond Is Used for Cutting Tools

A cutting or drilling tool must be harder than the material it cuts and must not wear away or melt from frictional heat. Diamond’s rigid tetrahedral network of strong covalent bonds makes it harder than any rock, metal, or glass it is used on, and its very high melting point means frictional heating does not damage it. Diamond-tipped drills and saws are therefore used for cutting stone, concrete, glass, and other hard materials.

Comparison with Graphite and Silicon(IV) Oxide

Diamond and graphite are allotropes — the same element with different structures and therefore dramatically different properties. Silicon(IV) oxide is a different compound entirely, but its diamond-like tetrahedral structure gives it diamond-like properties. See Giant Structures for the full comparison of giant ionic, covalent, and metallic lattices.

Hardness of Diamond

Diamond is the hardest known natural substance. This extreme hardness is due to its giant covalent structure: each carbon atom is bonded to four other carbon atoms by strong covalent bonds in a rigid tetrahedral 3D network. To scratch or cut diamond, these strong covalent bonds must be broken, which requires enormous energy. This is why diamond is used in cutting tools and drill bits for drilling through rock and other hard materials.

Exam point: Diamond is hard because each C atom is held by 4 strong covalent bonds in a tetrahedral network — no weak planes to break along.

Cutting Tools and Diamond

Diamond is used in industrial cutting tools, drill bits, and grinding wheels because it is the hardest known natural material. The giant covalent structure — each carbon atom tetrahedrally bonded to 4 others by strong covalent bonds — means enormous force is needed to break the structure. Diamond-tipped tools can cut through rock, concrete, glass, and other very hard materials.

Exam point: Diamond’s use in cutting tools follows from its structure: 4 strong covalent bonds per C atom in a rigid 3D tetrahedral lattice = extreme hardness.

Sources

• Cambridge IGCSE Chemistry 0620 Syllabus (2023-2025) — Section 2.6 Giant covalent structures, Core points 1–2, Cambridge Assessment International Education

Common Misconceptions