While most covalent substances exist as discrete, simple molecules, certain elements and compounds form vast, continuous networks of atoms. In these giant covalent structures (or macromolecules), there are no individual molecules. Instead, every atom is linked to its neighbors by strong covalent bonds throughout the entire sample. The IGCSE 0654 Extended syllabus focuses on the two primary allotropes of carbon: diamond and graphite.

Diamond is one of the hardest known natural substances. Its properties are a direct result of its three-dimensional crystalline geometry.

Structure

In diamond, each carbon atom forms four strong covalent bonds with four other carbon atoms in a rigid, tetrahedral arrangement. This pattern repeats infinitely in all directions.

Properties

• Extreme Hardness: The three-dimensional network of strong bonds makes it incredibly difficult to displace atoms, giving diamond its famous durability.
• Very High Melting Point: To melt diamond, a vast number of strong covalent bonds must be broken simultaneously, which requires extreme thermal energy.
• Non-conductor: All of carbon's valence electrons are localized in covalent bonds. There are no free-moving (delocalised) electrons to carry an electrical charge.

Graphite is an allotrope of carbon with properties that contrast sharply with diamond, despite being composed of the same element.

Structure

In graphite, each carbon atom forms three strong covalent bonds with three other carbon atoms, creating flat, hexagonal layers.

• Bonding: Only three of carbon's four valence electrons are involved in covalent bonding.
• Delocalisation: The fourth electron from each atom becomes delocalised and is free to move between the layers.
• Intermolecular Forces: While the bonds within the layers are strong, the layers themselves are held together only by weak intermolecular forces.

Properties

• Soft and Slippery: Because the forces between the layers are weak, the layers can easily slide over one another. This makes graphite an excellent lubricant and the ideal material for pencil "lead."
• Electrical Conductivity: Graphite is a rare example of a non-metal that conducts electricity. The delocalised electrons are free to move throughout the structure, allowing it to function as an electrode.
• Very High Melting Point: Similar to diamond, melting graphite requires breaking the strong covalent bonds within the hexagonal layers, which requires significant energy.

The study of diamond and graphite illustrates the profound impact of atomic arrangement on macroscopic properties. Even though both substances consist entirely of carbon atoms, the difference between three-way and four-way bonding creates two of the most physically diverse materials in chemistry. Understanding these giant structures is key to grasping the limits of material strength and conductivity.