Structure of Metals

Picture a metal as a neat stack of positive ions sitting in a pool of electrons that don't belong to any one atom. Each metal atom gives up its valence electrons to this shared pool, and the resulting cations are held in place by their attraction to that electron sea, not by a bond to any particular neighbor. Because the attraction has no direction, the cations can slide past each other and still be surrounded by the same sea. That's why hammering copper reshapes it instead of breaking it: push the ions into new positions and the electron sea just flows around them, keeping every cation bonded. Try the same shove on an ionic crystal and you shift ions of like charge next to each other, and the repulsion splits the crystal apart.

That same mobility explains conductivity. The delocalized electrons are barely held by any single nucleus, so when you connect a wire to a battery, the electric field pushes them to drift steadily toward the positive terminal, carrying charge through the whole piece of metal. Nothing about the lattice structure needs to change for this to happen, which is also why metals conduct in both solid and liquid form.

This sets metallic bonding apart from covalent bonding, which pairs two specific electrons between two specific nuclei in a fixed direction. Metallic bonding has no such pairing or direction: every electron in the sea interacts with every cation in the lattice at once. Heating a metal toward its melting point adds energy that lets ions vibrate and eventually break free of their fixed positions, but the electron sea stays intact, which is why molten metals still conduct and still resist shattering the way ionic melts can.