Giant Covalent Structures

From covalent bonds to a giant network

A covalent bond forms when atoms share a pair of electrons. The bond is the strong electrostatic attraction between that negatively charged shared pair and the positively charged nuclei of the bonded atoms.

In a simple molecule, covalent bonding stops after a small, fixed group of atoms has formed. In a giant covalent structure, it does not produce separate, identifiable molecules. Instead, covalent bonds link a vast number of atoms into one continuous, repeating network, or lattice. The word giant describes the extent of this network rather than the size of an individual atom.

A diagram can show only a tiny fragment of such a structure. Lines represent covalent bonds, but the model is not a collection of balls and sticks in reality. A two-dimensional page can also hide bonds that point towards the viewer, away from the viewer or into neighbouring parts of a three-dimensional lattice.

Diamond: bonding in three dimensions

Diamond is an allotrope of carbon. Each carbon atom forms four covalent bonds to four other carbon atoms. The four bonds point towards the corners of a tetrahedron and connect the atoms into a rigid network extending in all three dimensions.

There is no boundary at which one diamond molecule ends and another begins: the bonding continues throughout the crystal. Every outer electron of each carbon is used in bonding, so diamond has no delocalised electrons available to move through the structure.

Graphite: strong layers with weak forces between them

Graphite is another allotrope of carbon, but its atoms have a different arrangement. Each carbon atom forms three covalent bonds to three other carbon atoms. This produces flat sheets of joined hexagonal rings.

The covalent bonding continues across each sheet, so every sheet is part of the giant structure. However, there are no covalent bonds joining one sheet to the next. Only weak forces act between the layers. It is therefore important to distinguish strong covalent bonds within a layer from weak forces between layers.

Each carbon contributes three of its four outer electrons to covalent bonds. The remaining electron is delocalised and can move along a layer. This makes graphite structurally different from diamond even though both substances contain carbon alone.

Silicon dioxide: a network containing two elements

Giant covalent structures are not limited to elemental carbon. In silicon dioxide, SiO₂, each silicon atom is covalently bonded to four oxygen atoms, while each oxygen atom is bonded to two silicon atoms. The alternating connections continue in three dimensions.

The formula SiO₂ expresses the simplest ratio in the lattice: there are two oxygen atoms for every silicon atom overall. It does not mean that solid silicon dioxide consists of separate SiO₂ molecules.

Silicon dioxide and diamond are similar because both contain an extended, three-dimensional arrangement of strong covalent bonds. They differ in composition and connectivity: diamond contains only carbon, with each carbon bonded to four carbons; silicon dioxide contains silicon and oxygen, with four oxygen neighbours around each silicon and two silicon neighbours around each oxygen.

Comparing the arrangements

All three examples contain strong covalent bonds and enormous networks, but their geometry is not identical:

• Diamond: four covalent bonds per carbon; a rigid, three-dimensional tetrahedral network.
• Graphite: three covalent bonds per carbon; two-dimensional hexagonal layers, weak forces between layers and delocalised electrons.
• Silicon dioxide: a three-dimensional network in which each silicon bonds to four oxygens and each oxygen bonds to two silicons.

These arrangements help explain familiar differences. Many strong covalent bonds must be broken to melt a giant covalent substance, so very high temperatures are generally required. Diamond is hard because its strong bonds form a rigid network in every direction. Graphite is soft because its layers can slide over one another, although the covalent bonds within each layer remain strong. Graphite can conduct electricity because it contains mobile delocalised electrons; diamond cannot because all four outer electrons on each carbon are held in bonds.