Bonding in the Natural Environment
Why Bonding Determines Properties
The physical properties of a substance — its melting point, hardness, electrical conductivity and whether it dissolves in water — all depend on its type of bonding and structure. This is one of the most powerful ideas in chemistry: if you know how the atoms are held together, you can predict how the material will behave. This link matters, so every property you state must be explained in terms of the forces involved.
Metallic Bonding
In a metal, atoms are arranged closely together in a giant metallic structure. The metal atoms lose their outer electrons, leaving a lattice of positively charged metal ions surrounded by a ‘sea’ of delocalised electrons. Strong electrostatic attraction between the positive ions and the negative electrons holds the whole structure together.
This model explains several key properties. Metals conduct electricity because the delocalised electrons can drift along the structure when a voltage is applied. They conduct heat well because those same mobile electrons transfer kinetic energy quickly. Metals are malleable and ductile — the layers of ions can slide over each other without breaking, because the delocalised electrons can reorganise to maintain attraction throughout. High melting and boiling points arise from the strength of the metal–electron attraction.
Metallic properties explained: conducts electricity (delocalised electrons carry charge), malleable (ions slide over each other), high melting point (strong attraction between ions and electron sea).
Giant Ionic Structures
Ionic compounds such as sodium chloride form a giant ionic lattice. Many strong electrostatic forces act in all directions between the positive and negative ions. This gives ionic compounds high melting and boiling points because a great deal of energy is needed to break all these forces.
Ionic compounds do not conduct electricity when solid because the ions are locked in fixed positions. When melted or dissolved in water, the ions become free to move and carry charge — so they conduct electricity in these states. Many ionic compounds dissolve in water because water molecules can surround and pull away individual ions from the lattice. Ionic solids are often brittle: applying force shifts layers, so like charges align and repel, shattering the crystal.
Ionic compounds: high melting points (many strong electrostatic forces), conduct when molten or dissolved (ions free to move), do not conduct as solids (ions fixed in the lattice). A common exam error is saying solid ionic compounds ‘have no ions’ — they do; the ions just cannot move.
Simple Molecular Substances
Simple molecular substances (such as water, carbon dioxide and methane) have strong covalent bonds within each molecule but weak intermolecular forces between molecules. When the substance melts or boils, it is the weak intermolecular forces that are broken — not the covalent bonds. This means simple molecular substances have low melting and boiling points and are often liquids or gases at room temperature.
Simple molecular substances do not conduct electricity in any state because they have no free electrons or ions. A very common exam mistake is to say the covalent bonds are weak — the bonds within the molecules are strong; it is the forces between molecules that are weak.
Larger molecules have more electrons and so form stronger intermolecular forces. This means they need more energy to overcome those forces and boil. For example, methane (\(CH_4\)) has a much lower boiling point than octane (\(C_8H_{18}\)) even though both are alkanes with similar covalent bonds within each molecule. The difference in boiling point reflects the difference in intermolecular force strength, not bond strength within the molecule.
Giant Covalent Structures
Giant covalent structures contain billions of atoms bonded together by strong covalent bonds in a continuous network. They have very high melting points because many strong covalent bonds must be broken. Carbon forms several giant covalent allotropes.
Carbon can form four covalent bonds. In diamond, each carbon bonds to four others in a rigid tetrahedral arrangement, making diamond extremely hard and a non-conductor (no free electrons — all four bonding electrons are used in covalent bonds). In graphite, each carbon bonds to three others in flat hexagonal layers; one electron per carbon atom is delocalised within each layer, so graphite conducts electricity. The layers are held together by weak forces, allowing them to slide, making graphite soft and slippery. Graphene is a single layer of graphite — one atom thick, very strong, and an excellent conductor with potential uses in electronics and composite materials.
Silicon dioxide (found in sand and quartz) is another giant covalent structure. Like diamond, it has very high melting and boiling points and does not conduct electricity. Each silicon atom is bonded to four oxygen atoms in a network that extends in all directions.