Giant covalent structures Revision Notes

Diamond is a giant covalent structure composed entirely of carbon atoms.

Each carbon atom forms four strong covalent bonds with four other carbon atoms arranged tetrahedrally.

This creates a rigid three-dimensional network extending throughout the crystal lattice, with bond angles of approximately 109.5° between bonds (tetrahedral geometry).

• Each carbon atom uses sp³ hybrid orbitals to form four sigma (σ) bonds.
• The strong covalent bonds require a large amount of energy to break, resulting in diamond's high melting point (about 3550°C).
• No free electrons are present, so diamond does not conduct electricity.

Graphite is another giant covalent structure of carbon atoms, but arranged differently from diamond.

Each carbon atom bonds covalently to three others in a planar hexagonal lattice, forming flat layers.

The fourth valence electron of each carbon atom is delocalised and free to move within the layers.

• The layers are held together by weak van der Waals forces, allowing them to slide over each other easily, making graphite soft and slippery.
• The delocalised electrons enable graphite to conduct electricity along the planes.
• The bond angles within the layers are approximately 120°, consistent with trigonal planar geometry (sp² hybridisation).

Diamond is extremely hard due to its strong covalent bonds in all directions, making it useful for cutting tools.

Graphite is soft and slippery because the layers can slide over each other due to weak interlayer forces.

Diamond has a very high melting point because breaking the extensive covalent network requires a lot of energy.

• Graphite conducts electricity due to delocalised electrons; diamond does not conduct electricity.
• Graphite's layered structure results in anisotropic properties, with conductivity and softness along layers but not between them.
• Both substances are insoluble in water and most solvents due to the strength of their covalent bonds.

Diamond's structure can be represented by a three-dimensional tetrahedral network of carbon atoms, each bonded to four others.

Graphite's structure is represented by layers of hexagonal rings of carbon atoms, with delocalised electrons above and below the planes.

• Empirical formula for both is C, as they are allotropes of carbon.
• No simple molecular formula exists due to the giant covalent nature.
• Structural diagrams show the difference between tetrahedral bonding in diamond and planar hexagonal bonding in graphite.

Diamond has a melting point around 3550°C, whereas graphite sublimes at about 3900°C.

Both have high melting points due to strong covalent bonds, but graphite's layers can separate more easily than breaking diamond's 3D network.

• Melting diamond requires breaking all covalent bonds in the 3D network.
• Graphite's melting involves overcoming weaker van der Waals forces between layers, but the covalent bonds within layers remain intact.
• This explains why graphite sublimes rather than melts under normal pressure.

Graphite is a giant covalent structure composed entirely of carbon atoms. Each carbon atom forms three strong covalent bonds with three other carbon atoms in the same plane, creating hexagonal rings arranged in layers.

Within each layer, the carbon atoms are bonded by strong covalent bonds, forming a two-dimensional hexagonal lattice. The fourth outer electron of each carbon atom is delocalised and free to move within the layers, contributing to electrical conductivity.

The layers are held together by weak van der Waals forces, allowing them to slide easily over each other. This layered structure is responsible for graphite's lubricating properties.

• Each carbon atom bonded to three others in hexagonal layers.
• Delocalised electrons present within layers enable electrical conductivity.
• Weak forces between layers allow layers to slide over each other.

Graphite's layered structure allows the layers to slide over one another easily, reducing friction between moving parts. This makes graphite an effective dry lubricant, especially in high-temperature or vacuum environments where liquid lubricants are unsuitable.

The presence of delocalised electrons within the layers enables graphite to conduct electricity. This property allows graphite to be used as electrodes in electrolysis and in batteries.

Graphite electrodes are used in electric arc furnaces for steel production because they can withstand high temperatures and conduct electricity efficiently.

• Layered structure enables sliding, reducing friction (lubricant).
• Delocalised electrons allow electrical conductivity (electrode).
• High melting point and chemical inertness make graphite suitable for high-temperature applications.

Diamond is another giant covalent structure composed of carbon atoms. Each carbon atom forms four strong covalent bonds with four other carbon atoms arranged tetrahedrally.

This three-dimensional network of strong covalent bonds extends throughout the entire crystal, making diamond extremely hard and rigid.

There are no free electrons in diamond, so it does not conduct electricity.

• Each carbon atom bonded to four others in a tetrahedral arrangement.
• Strong covalent bonds throughout a 3D network.
• No delocalised electrons; diamond is an electrical insulator.

The hardness of diamond arises from its strong covalent bonds in a rigid three-dimensional structure. This makes diamond the hardest naturally occurring substance.

Because of this extreme hardness, diamond is used in cutting, grinding, and drilling tools to cut or shape hard materials such as metals, glass, and stone.

Diamond-tipped cutting tools maintain a sharp edge for longer and resist wear, improving efficiency and precision in industrial applications.

• Strong covalent bonding gives diamond exceptional hardness.
• Used in cutting tools for hard materials due to durability and sharpness.
• Does not conduct electricity, so not used as an electrode.

The contrasting structures of graphite and diamond explain their different physical properties and applications.

Graphite's layered structure with delocalised electrons results in lubricating properties and electrical conductivity, making it suitable as a lubricant and electrode.

Diamond's rigid tetrahedral network of covalent bonds results in extreme hardness, making it ideal for cutting tools.

• Structure determines physical properties such as hardness and conductivity.
• Properties determine practical applications in industry and technology.

Silicon dioxide (SiO₂) is composed of silicon and oxygen atoms arranged in a continuous three-dimensional network. Each silicon atom forms covalent bonds with four oxygen atoms, arranged tetrahedrally around it. Each oxygen atom is covalently bonded to two silicon atoms, acting as a bridge between them. This arrangement repeats throughout the solid, creating a giant covalent lattice without discrete molecules.

• Silicon atoms are at the centre of tetrahedra.
• Each silicon atom bonds to four oxygen atoms.
• Each oxygen atom bonds to two silicon atoms.
• The structure extends in all directions forming a giant lattice.

The bonds between silicon and oxygen atoms in SiO₂ are strong covalent bonds formed by the sharing of electrons. Each silicon atom shares electrons with four oxygen atoms, and each oxygen atom shares electrons with two silicon atoms. The strength and number of these covalent bonds throughout the lattice give silicon dioxide its stability and rigidity.

• Strong covalent bonds link silicon and oxygen atoms.
• Each bond involves shared electron pairs.
• The network is continuous and extensive.
• No individual SiO₂ molecules exist; it is a giant covalent structure.

The giant covalent structure of silicon dioxide results in very high melting and boiling points because a large amount of energy is required to break the many strong covalent bonds throughout the lattice. Silicon dioxide is also very hard and brittle due to the rigidity of the covalent network. Additionally, it does not conduct electricity because there are no free electrons or ions to carry charge.

• High melting point due to strong covalent bonds.
• High boiling point for the same reason.
• Hard and brittle solid.
• Does not conduct electricity.

Silicon dioxide is insoluble in water because the strong covalent bonds in the giant lattice are not broken by interaction with water molecules. The energy required to break these bonds is much greater than the energy released by hydration. This chemical stability makes silicon dioxide resistant to many chemical reactions under normal conditions.

• Insoluble in water due to strong covalent bonds.
• Stable and chemically inert under normal conditions.
• Does not dissolve or ionise in water.

Silicon dioxide's giant covalent structure consists of a continuous network of silicon and oxygen atoms linked by strong covalent bonds. This structure explains its high melting and boiling points, hardness, insolubility in water, and electrical insulating properties. Understanding this structure is essential for explaining the physical and chemical behaviour of silicon dioxide in various contexts.

• Giant covalent lattice of Si and O atoms.
• Strong covalent bonds throughout the structure.
• High melting and boiling points.
• Hard, brittle, and insoluble in water.
• Electrical insulator.

Diamond is a form of carbon where each carbon atom is covalently bonded to four other carbon atoms in a three-dimensional tetrahedral lattice. This creates a giant covalent structure extending throughout the crystal.

Silicon dioxide (SiO₂), commonly found as quartz, consists of silicon atoms each covalently bonded to four oxygen atoms in a tetrahedral arrangement. Each oxygen atom bridges between two silicon atoms, forming a continuous giant covalent network.

Both structures are giant covalent lattices with strong covalent bonds extending in all directions, resulting in very stable and rigid frameworks.

• Diamond: each C atom bonded to 4 C atoms tetrahedrally.
• SiO₂: each Si atom bonded to 4 O atoms tetrahedrally; each O bridges two Si atoms.
• Both form extensive 3D networks of strong covalent bonds.

The tetrahedral bonding in both diamond and silicon dioxide means that atoms are held in fixed positions by strong covalent bonds in all directions. This rigidity makes the structures very hard and gives them high melting points.

Because the bonding involves shared pairs of electrons localized between atoms, there are no free electrons or ions to carry charge, making both substances electrical insulators.

The strong covalent bonds require large amounts of energy to break, explaining the high melting points observed for both materials.

• Strong covalent bonds in tetrahedral arrangement.
• High hardness due to rigid 3D lattice.
• No free electrons or ions → electrical insulators.
• High melting points due to energy needed to break bonds.

Diamond has an extremely high melting point around 3550 °C because breaking the giant covalent lattice requires breaking many strong C–C covalent bonds.

Silicon dioxide melts at about 1710 °C, also very high due to the strong Si–O covalent bonds in its lattice.

Both materials are very hard; diamond is the hardest known natural substance due to the strong and rigid C–C bonds, while silicon dioxide is also hard but less so than diamond because of differences in bond strength and lattice structure.

• Diamond melting point ≈ 3550 °C; SiO₂ melting point ≈ 1710 °C.
• Both are very hard solids; diamond is harder than SiO₂.
• High melting points arise from strong covalent bonds throughout the lattice.

Neither diamond nor silicon dioxide conducts electricity because their electrons are all involved in covalent bonds and are not free to move.

In diamond, all four valence electrons of carbon form covalent bonds, leaving no free electrons or ions.

In silicon dioxide, the electrons are shared in Si–O bonds, and there are no mobile charged particles to carry current.

• No free electrons or ions → poor electrical conductivity.
• Diamond and SiO₂ are electrical insulators.
• Covalent bonding localizes electrons, preventing charge flow.