Giant covalent structures and metallic bonding
Chemistry · GCSE · The School
Some covalent substances form giant structures, not small molecules
Water and carbon dioxide are small covalent molecules. Each one is a handful of atoms, with weak forces between separate molecules. Diamond, graphite and silicon dioxide are different. Every atom is bonded to its neighbours in one giant lattice. There are no separate molecules at all. That one difference explains why these three behave nothing like water or carbon dioxide.
Diamond: every carbon bonded to four others
Each carbon atom in diamond forms four strong bonds. Every one links to four other carbons, in a rigid 3D lattice. Breaking diamond means breaking many of these bonds at once. That is why it is so hard, with such a high melting point. It does not conduct electricity, either. Every outer electron is locked into a bond, so none are free to move.
Graphite: layers that slide over each other
Each carbon in graphite bonds to only three others. This forms flat layers, shaped like honeycomb. Weak forces hold the layers together. So the layers slide over each other easily. That is why graphite is soft, and used as a lubricant and in pencils. Unlike diamond, graphite DOES conduct electricity. Each carbon has one spare electron, and it moves freely along the layers.
Silicon dioxide: diamond’s structure, a different element
Silicon dioxide is found in sand. It has a giant structure much like diamond. Each silicon atom bonds to four oxygen atoms, in one rigid lattice. Like diamond, it has a very high melting point. It does not conduct electricity either. Comparing the two side by side checks the pattern is really understood, not just remembered for one substance.
Metallic bonding: a lattice of ions in a sea of electrons
A metal is a lattice of positive metal ions. Every atom gives up its outer electrons. Those electrons become delocalised: free to move through the whole structure, not stuck to any one ion. A strong pull forms between the positive ions and this shared "sea" of electrons. That pull is what holds the lattice together, and it is called metallic bonding.
Explaining metals’ properties from the same picture
Metals conduct electricity because the delocalised electrons carry charge freely through the lattice. Metals are malleable, too: the layers of ions can slide over each other without breaking the bonding. The sea of electrons just flows around them, holding the new shape together. Both properties come from the SAME feature. Electrons that belong to the whole lattice, not to one atom.