LITMIR.BIZ

Figure 2.13 (a) Covalent bonding (double lines) in a carbon tetrahedron with the central carbon atom bonded to four carbon atoms that occupy the corners of a tetrahedron (dashed lines). (b) A larger scale diamond structure with multiple carbon tetrahedra.

      Source: Courtesy of Steve Dutch.

The best known mineral with covalent bonding is diamond, which is composed of carbon (C). Because carbon is a column 14 atom, it must either lose four electrons or gain four electrons to achieve a stable electron configuration. In diamond, each carbon atom in the structure is bonded to four nearest neighbor carbon atoms that share with it one of their electrons (Figure 2.13). In this way, each carbon atom attracts four additional electrons, one from each of its neighbors, to achieve the stable noble electron configuration. The long‐range crystal structure of diamond is a pattern of carbon atoms in which every carbon atom is covalently bonded to four other carbon atoms.

      Covalently bonded minerals are generally characterized by the following:

      2.3.4 Metallic bonds

When metallic atoms bond with other metallic atoms, a metallic bond is formed. Because very metallic atoms have low first ionization energies, are highly electropositive and possess low electronegativities they do not tend to hold their valence electrons strongly. In such situations, each atom releases valence electrons to achieve a stable electron configuration. The positions of the valence electrons fluctuate or migrate between atoms. Metallic bonding is difficult to model, but is usually portrayed as positively charged partial atoms (nuclei plus the strongly held inner electrons) in a matrix or “gas” of “delocalized” valence electrons that are only temporarily associated with individual atoms (Figure 2.14). The weak attractive forces between positive partial atoms and valence electrons bond the atoms together. Unlike the strong electron‐sharing bonds of covalently bonded substances, or the frequently strong electrostatic bonds of ionically bonded substances, metallic bonds are rather weak, less permanent and easily broken and reformed. Because the valence electrons are not strongly held by any of the partial atoms, they are easily moved in response to stress or in response to an electric field or thermal gradient.

      Excellent examples of metallic bonding exist in the native metals such as native gold (Au), native silver (Ag), and native copper (Cu). Such materials are excellent conductors of electricity and heat. When materials with metallic bonds are subjected to an electric potential or field, delocalized electrons flow toward the positive anode, which creates and maintains a strong electric current. Similarly, when a thermal gradient exists, thermal vibrations are transferred by delocalized electrons, making such materials excellent heat conductors. When metals are stressed, the weakly held electrons tend to flow, which helps to explain the ductile behavior that characterizes native copper, silver, gold, and other metallically bonded substances.