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.


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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.

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    $\begingroup$ As more shells are added, electrons are shielded from the nucleus, and become freer to wander -- i.e., behave as metals. Note even the chalcogens, starting with oxygen, wind up as semimetals by the time one reaches polonium. See sciencenotes.org/printable-periodic-table $\endgroup$ Commented Apr 11, 2016 at 4:03
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    $\begingroup$ And indeed, grey tin is a semi-metal. Liquid silicon and germanium are metallic. Every time you start thinking that the 'periodic' table should be nice and predictable you come up against an 'anomaly' that seems weird. $\endgroup$
    – Jon Custer
    Commented Apr 11, 2016 at 16:01
  • $\begingroup$ Thanks Jon, I know about liquid Si and Ge. If I am not mistaken, the melting of those elements going along with the change in bonding means there is a "collapse" like that in water (ice) when it melts. So Si and Ge, like water, contract upon melting. Tin undegoes this contraction while still solid. $\endgroup$ Commented Apr 11, 2016 at 16:19

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