Metallic bonding
The idea
Picture a metal as a lattice of positive ions sitting in a shared sea of electrons. Each metal atom releases its outer electrons into a common pool that belongs to the whole crystal rather than to any single atom, and the attraction between the fixed cations and this mobile electron sea is the metallic bond. You already know metals sit on the left of the periodic table and hold their valence electrons loosely — metallic bonding is what those loose electrons do when there is no nonmetal around to take them.
This one model explains the familiar metal package. Conductivity: the delocalized electrons drift the moment a voltage is applied, and they also ferry heat quickly. Malleability: because no electron is tied to a particular neighbor, planes of cations can slide past each other while the electron sea keeps gluing everything together — the crystal dents instead of shattering. Bond strength scales with the glue: more delocalized electrons per atom, higher cation charge, and smaller cations all mean stronger bonding and higher melting points.
Do not picture metals as ionic compounds with fixed partners or as molecules with fixed shared pairs. The defining feature is that the bonding electrons are delocalized — owned by the entire lattice — which is exactly what an ionic crystal lacks, and why salt shatters under a hammer while copper flattens.
Worked example
Sodium melts at about 98 °C while magnesium melts at about 650 °C, yet both are period 3 metals. Use the metallic bonding model to explain the gap, and explain why both metals still conduct electricity well.
- Count the glue each atom contributes: sodium (group 1) releases 1 electron per atom into the sea, leaving Na⁺ cations, while magnesium (group 2) releases 2 per atom, leaving Mg²⁺ cations.
- Compare the attractions: magnesium has doubly charged cations attracted to twice the density of delocalized electrons, and Mg²⁺ is also smaller than Na⁺, so the cations sit closer to the electron sea. Every factor strengthens magnesium's bonding.
- Connect strength to melting: melting requires loosening the cations from their lattice positions against that attraction, so magnesium's much stronger bonding demands far more energy — hence roughly 650 °C versus 98 °C.
- Address conductivity separately: both metals have mobile, delocalized electrons that drift under an applied voltage, so both conduct well. Conductivity needs mobile charge, not strong bonding, which is why soft, low-melting sodium is still an excellent conductor.
Answer. Magnesium melts far higher because each atom supplies two electrons to the sea and forms smaller, doubly charged Mg²⁺ cations, giving much stronger attractions; both metals conduct because their electron seas are mobile either way.
Check your understanding
- Why can a metal be hammered into a sheet while an ionic crystal with similar bond strength shatters?
- How does the electron-sea model explain why metals conduct heat as well as electricity?
- What should happen to melting points going down group 1, and which part of the model predicts it?
- How would you explain the difference between delocalized and shared-pair electrons to a friend who just learned covalent bonding?
Build the foundations first
Metallic bonding builds on these concepts. If any feel shaky, start there.