BAND THEORY OF METALS
As described in the previous section, most metals crystallize in close-packed structures.
The ability of metals to conduct electricity and heat must result from strong electronic
interactions of an atom with its 8 to 12 nearest neighbors. This might be surprising at
first if we recall that each Group IA and Group IIA metal atom has only one or two
valence electrons available for bonding. This is too few to participate in bonds localized
between it and each of its nearest neighbors.
Bonding in metals is called metallic bonding.It results from the electrical attractions
among positively charged metal ions and mobile, delocalized electrons belonging to the
crystal as a whole. The properties associated with metals—metallic luster, high thermal
and electrical conductivity, and so on—can be explained by the band theoryof metals,
which we now describe.
The overlap interaction of two atomic orbitals, say the 3sorbitals of two sodium atoms,
produces two molecular orbitals, one bonding orbital and one antibonding orbital (Chapter
9). If Natomic orbitals interact, Nmolecular orbitals are formed. In a single metallic13-17
Figure 13-32 Portions of the atomic arrangements in three covalent solids. (a) Diamond.
Each C is bonded tetrahedrally to four others through sp^3 - sp^3 -bonds. (1.54 Å).
(b) Graphite. C atoms are linked in planes by sp^2 - sp^2 -bonds (1.42 Å). The crystal is soft,
owing to the weakness of the attractions between planes (3.40 Å). Electrons move freely
through the -bonding network in these planes, but they do not jump between planes easily.
(c) Quartz (SiO 2 ). Each Si atom (gray) is bonded tetrahedrally to four O atoms (red).528 CHAPTER 13: Liquids and Solids
Metals can be formed into many
shapes because of their malleability
and ductility.
(a) Diamond (b) Graphite (c) A natural quartz crystal