Chemistry: Properties of materials

Ionic compounds are formed between metals and non-metals. Electrons are transferred from the metal atoms to the non-metal atoms, forming positive ions (cations) and negative ions (anions). These oppositely charged ions are strongly attracted to each other by electrostatic forces. They arrange themselves into a regular, repeating three-dimensional structure called a giant ionic lattice. A huge amount of energy is needed to overcome these strong forces and break down the lattice, which is why ionic compounds have very high melting and boiling points. They are usually crystalline solids at room temperature. They cannot conduct electricity when solid because the ions are fixed in position. However, when molten or dissolved in water, the ions are free to move and can carry an electrical charge.

Simple molecular substances are formed between non-metal atoms, which share electrons to form strong covalent bonds. These substances exist as discrete (separate) small molecules, such as water (H₂O), ammonia (NH₃), and methane (CH₄). While the covalent bonds *within* each molecule are very strong, the forces of attraction *between* the molecules are very weak. These are known as intermolecular forces. When a simple molecular substance melts or boils, it is these weak intermolecular forces that are overcome, not the strong covalent bonds. Because little energy is needed to overcome them, simple molecular substances have low melting and boiling points. They are often gases or liquids at room temperature. They do not have any free electrons or ions, so they do not conduct electricity in any state.

Diamond is an allotrope of carbon, meaning it's a different structural form of the same element. In diamond, every carbon atom is joined to four other carbon atoms by strong covalent bonds. This creates a rigid, tetrahedral network that extends in three dimensions to form a single giant molecule, known as a giant covalent structure or macromolecular structure. To melt or break diamond, you would need to break these numerous strong covalent bonds throughout the structure. This requires a vast amount of energy, giving diamond an extremely high melting point (around 4000°C) and making it the hardest known natural substance. As all four of each carbon's outer electrons are used in bonding, there are no free-moving electrons or ions. Therefore, diamond does not conduct electricity.

Graphite is another allotrope of carbon. In graphite, each carbon atom is covalently bonded to three other carbon atoms, forming hexagonal rings arranged in flat layers. The covalent bonds within the layers are very strong. However, there are only weak intermolecular forces between the layers. This structure explains graphite's properties: it is soft and slippery because the layers can easily slide over each other. This makes it useful as a lubricant and in pencil 'lead'. Crucially, because each carbon atom only uses three of its four outer electrons for bonding, one electron per carbon atom is delocalised. These delocalised electrons are free to move along the layers, allowing graphite to conduct electricity. Graphene is simply a single layer of graphite. It is incredibly strong for its mass, transparent, and an even better electrical conductor than graphite.

Metals have a unique type of bonding called metallic bonding. They form a giant structure consisting of a regular, closely packed lattice of positive metal ions. The outer shell electrons from each metal atom are delocalised, meaning they are free to move throughout the entire structure. This creates a 'sea' of delocalised electrons surrounding the positive ions. The metallic bond is the strong electrostatic attraction between the positive ions and the sea of delocalised electrons. This strong attraction is why most metals have high melting points and are strong. The delocalised electrons also explain why metals are excellent conductors of electricity and heat. Furthermore, metals are malleable (can be hammered into shape) and ductile (can be drawn into wires) because the layers of ions can slide over one another without breaking the metallic bond, as the delocalised electrons can move with them.