Metallic and Covalent Bonding
Metallic bonding creates a fascinating structure called the 'sea of electrons'. Metal atoms release their outer electrons, which become delocalised and can move freely throughout the metal. These mobile electrons create strong attractions with the positive metal centres, giving metals their unique properties.
This electron mobility explains why metals are so useful. They're ductile and malleable because layers can slide over each other without breaking bonds. They conduct heat and electricity brilliantly because those free electrons can transfer energy and charge throughout the structure.
Covalent bonding works completely differently - non-metal atoms share electrons to fill their outer shells. This sharing creates some incredible structures like diamond, where each carbon atom bonds to four others, making it incredibly hard. Graphite is fascinating because each carbon only bonds to three others, leaving spare electrons that can conduct electricity.
Giant covalent structures like diamond and silicon dioxide have sky-high melting points because breaking all those strong bonds requires enormous energy. However, they typically can't conduct electricity (except graphite) because there are no free electrons or ions to carry charge.
Remember: Covalent bonding is all about sharing - atoms share electrons rather than transferring them completely like in ionic bonding.