Diamond & Graphite: Giant Covalent Structures

• Giant covalent structures are those with a three-dimensional regular arrangement of atoms or molecules joined by strong covalent bonds throughout the entire structure.
• They are also sometimes described as macromolecules

Giant Covalent Structure (Diamond)

• Diamond and graphite are giant covalent structures made of carbon.
• Diamond and graphite are allotropes of carbon element; allotropes are the same element in a different form.
• In diamond, each carbon atom forms covalent bonds with 4 other carbon atoms and are arranged tetrahedrally forming a giant structure.
• Diamonds have a very high melting point because a lot of energy is needed to break the strong covalent bonds between C atoms
• Diamonds are very hard, and cannot be easily scratched
• Diamonds do not conduct electricity because they do not have free moving ions or electrons.
• Diamonds exist as colorless glittering crystals and are used in jewelry and cutting tools.

Giant Covalent Structure (Graphite)

• Graphite is a black shiny solid.
• Each carbon atom in graphite is joined to three other carbon atoms covalently, arranged in hexagons.
• As each carbon only forms 3 covalent bonds with 3 other C atoms, each carbon has one free moving/ delocalized electron, meaning it can conduct electricity and can be used as an electrode.
• The strong covalent bonding between C atoms means that breaking the bonds requires a lot of energy
• Hexagonal rings of carbon atoms are arranged in layers but the bonding between the layers is weak.
• The layers can slide over each other if a force is applied.
• Graphite has a slippery feel and can easily scratched therefore it is used as a lubricant and in pencil leads.

Giant Covalent Structure (Silicon Dioxide)

• Silicon (IV) oxide (SiO₂) is also referred to as silicon dioxide.
• The silicon (IV) oxide structure is similar to diamond’s tetrahedral structure.
• Each silicon atom is bonded to four oxygen atoms and each oxygen atom is bonded to two silicon atoms.
• Silicon dioxide is used to make glass.
• Silicon (IV) oxide has a high melting and boiling point due to the significant amount of energy required to break the strong covalent bonds between atoms.
• Silicon (IV) oxide does not conduct electricity because it has no free-moving electrons.

Metallic Bonding and Metal Lattice Structure

• Metal atoms are packed closely together in a regular arrangement. Valence electrons tend to move away from their atoms
• The structure of a metal can be described as a giant lattice of positive metal ions arranged regularly in a 'sea of electrons'.
• Electrons are free, mobile or delocalized.
• Metallic bonding results from the strong electrostatic attraction between the positive ions and negatively charged electrons.
• Metals are good conductors of electricity.
• When a voltage is applied across a piece of metal, the delocalized electrons move to the positive end and a current flows through the metal.
• In a metallic bond the attractive forces between the metal ions and the electrons act equally in any direction.
• Positive ions in a metal are arranged regularly in layers.
• When a force is applied, the layers can slide over each other and new bonds can easily form, thus altering the metal's shape.

Properties of Metals

• Metals generally have a high melting point and boiling point. This is because a lot of energy is needed to overcome the strong electrostatic attraction between the positive metal ions and negative electrons.

Physical properties of metals and nonmetals

• Metals are good conductors of electricity and heat, are malleable and ductile, have high melting and boiling points (except for group one metals and mercury), and have a shiny, lustrous appearance.
• Non-metals are poor conductors of electricity and heat, tend to be brittle, have low melting and boiling points (except for carbon and silicon), and have a dull appearance.