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Metal - Crystallography Of Metals

bonds valence cubic centered

Metals usually differ from nonmetals by their excellent thermal and electrical conductivities, and by their great mechanical strengths and ductilities. These properties follow directly from the nonlocalized electronic bonds in these materials. The electrons in metals are mobile; in a true metal, there are no underlying directed bonds.

With the exception of manganese and uranium, all true metals have one of the following crystal structures: body-centered cubic (sodium, potassium, molybdenum), iron face-centered cubic (copper, silver, gold), iron close-packed hexagonal (beryllium, magnesium, zirconium).

The origins of metallic behavior may be understood by considering the first and simplest of these three structures. There are eight nearest neighbors in a body-centered cubic structure. The number of next nearest atoms is six. The one valence electron of a body-centered cubic element like sodium clearly cannot furnish 14 or even eight covalent bonds with its neighbors. Thus, the single valence electron is shared.

The elements on the left-hand side of the periodic table readily pool their valence electrons, as they have low ionization potentials. Their large de-localization energies result in net binding. As one moves to the right of Group 1 in the periodic table, the metallic properties of the elements become weaker, and the tendency to form covalent bonds increases. As a result, thermal and electrical conductivities diminish, densities decrease, and the materials become hard, but brittle.

Carbon in Group 14, for example, does not allow its valence electrons to escape, but readily shares them with four neighbors. Graphitic carbon is made up of well separated layer planes with high conductivities along the planes but weak conductivities at right angles to these planes; consequently graphite is a two dimensional metal. In diamond, the electron bonds are tetrahedral and highly directed; this has the effect of making diamond brittle. Silicon, germanium, and grey tin also have diamond-like structures, and their bonding is largely covalent.

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