I know that lead is classified as a metal, from its appearance, nonexistent band gap, and its position on the periodic table... etc. However, all of the elements above it (carbon, silicon, germanium, and tin) usually form covalent network bonding with four surrounding atoms. Why is lead suddenly different, engaging in delocalized metallic bonding when it has enough electrons? Or is it different? Lead has a low melting point for a metal (600 K), and a lower melting point than, say, copper.
As you proceed down Group 14 the metallic bonding structure tends to become more stable than the tetrahedrally bonded structure we see in diamond, Si and Ge. Tin is actually a transitional case. At typical room temperature it is metallic (white tin) but upon chilling it transforms to the tetrahedral allotrope (gray tin). By the time we get to lead we have fully gone over to the metallic phase.
Because the tetrahedral bonding structure requires each atom to contribute four valence electrons, this tetrahedral --> metallic transition may be seen as a form of the "inert pair" effect that emerges in heavy atoms.