Giant Structures and Metallic Bonding
Some substances have structures so large and complex that they behave very differently from simple molecules - and understanding these will help you ace your chemistry exams!
Diamond is pure carbon where each atom bonds to four others, creating an incredibly strong 3D network. This gives diamond its sky-high melting point because you'd need massive amounts of energy to break all those strong covalent bonds. However, it can't conduct electricity because all electrons are tied up in bonding.
Graphite is also pure carbon, but each atom only bonds to three others, forming layers. The layers slide over each other easily (making graphite slippery), and it conducts electricity thanks to delocalised electrons that can move freely.
Metallic bonding creates a "sea of delocalised electrons" around positive metal ions. These free-moving electrons make metals brilliant conductors of both electricity and heat. The strong electrostatic attraction between the positive ions and electron sea explains why metals typically have high melting points and stay solid at room temperature.
Remember: The key to understanding conductivity is asking yourself - are there free-moving electrons or ions available?