How does the band theory of metallic bonding explain the luster of metals?

Answer 1

According to the band theory, any given metal atom has only a limited number of valence electrons with which to bond to all of its nearest neighbours. Extensive sharing of electrons among individual atoms is therefore required. This sharing of electrons is accomplished through the overlap of atomic orbitals of equivalent energy on the metal atoms that are immediately adjacent to one another. This overlap is delocalized throughout the entire metal sample to form extensive orbitals that span the entire solid rather than being part of individual atoms. Each of these orbitals lies at different energies because the atomic orbitals from which they were constructed were at different energies to begin with. The orbitals, equal in number to the individual atomic orbitals that have been combined, each hold two electrons, and are filled in order from the lowest to the highest energy until the number of available electrons has been used up. Groups of electrons are then said to reside in bands, which are collections of orbitals. Each band has a range of energy values that the electrons must possess to be part of that band; in some metals, there are energy gaps between bands, meaning that there are certain energies that the electrons cannot possess. The highest energy band in a metal is not filled with electrons because metals characteristically possess too few electrons to fill it. The high thermal electrical conductivities of metals is then explained by the notion that electrons may be promoted by absorption of thermal energy into these unfilled energy levels of the band.

Answer 2

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