Imagine a crowd at a concert. The fans (metal atoms) are packed closely together, and their phones (electrons) are being passed around overhead in every direction — no single phone belongs to any particular person anymore. That is essentially how a metallic bond works. Metal atoms release their outermost electrons into a communal pool, creating a lattice of positively charged ions bathed in a "sea" of delocalized electrons. Unlike covalent bonds where electrons are shared between specific atoms, or ionic bonds where electrons are transferred from one atom to another, metallic bonds involve electrons that belong to everyone and no one at once.
This sea of free-roaming electrons explains virtually every property we associate with metals. Metals conduct electricity brilliantly because these electrons can flow through the lattice when a voltage is applied — copper wiring in your house works thanks to metallic bonding. Metals conduct heat well for the same reason: electrons carry thermal energy rapidly from atom to atom. Metals are shiny because the free electrons absorb light and re-emit it, giving that characteristic metallic luster. And metals are malleable and ductile — you can hammer gold into sheets thinner than a human hair or draw it into wire — because when layers of ions slide over each other, the electron sea simply reshapes itself to maintain the bond. No bonds break, so the metal bends instead of shattering.
Metallic bonding varies in strength depending on how many electrons each atom contributes and how tightly the ions pack together. Sodium contributes one electron and melts at just 98°C; iron contributes two or three and melts at 1538°C; tungsten, with its many shared electrons and tight packing, holds the record among pure metals at 3422°C. This is why different metals suit different purposes — from the soft aluminum in soda cans to the unyielding tungsten in rocket nozzles.