1.4.3 Covalent structures and bonding

A covalent bond is formed when two atoms share two electrons, through overlap and merging of two electron orbitals, one from each atom (Figure 12c). Crystals containing covalent bonds tend to have more complex structures than those of ionic or metallic structures. Covalent bonding requires the precise overlap of electron orbitals, so if an atom forms several covalent bonds, these are usually constrained to specific directions. As covalent bonds are directional, unlike metallic or ionic bonds, this places additional constraints on the arrangements of atoms within such a crystal. One result is that covalent structures tend to be more open - and hence have lower densities than metallic or ionic structures.

Diamond is an example of a covalently bonded solid. In this form of carbon, each atom is covalently bonded to four other carbon atoms, arranged at the corners of a tetrahedron (Figure 16a). The resulting structure, which has a repeating cubic shape, is illustrated in Figure 16b. The structure contains much more unoccupied space than close-packed metal structures.

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Figure 16 The diamond structure: (a) the tetrahedral arrangement of covalent bonds around a carbon atom; (b) the arrangement of carbon atoms and bonds in the unit cell (see Section 2.6); the unit cell is shown by black lines. (c) Phase diagram illustrating the stability fields for graphite and diamond. Temperature and pressure conditions within the Earth are such that diamond tends to form only at depths greater than 150 km.

Another form of solid carbon with covalent bonding is graphite. Unlike those in diamond, the carbon atoms in graphite are covalently bonded to three neighbours in the same plane (Figure 17a), producing a strong sheet of carbon atoms. However, each carbon atom has one extra electron available for bonding that forms very weak bonds, which serve to keep the carbon sheets together (Figure 17b).

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Figure 17 The graphite structure: (a) the triangular arrangement of covalent bonds around a carbon atom in the same plane; (b) part of the three-dimensional structure of graphite, showing strongly bonded hexagonal sheets of carbon atoms, connected by weak bonds. (c) Specimen of graphite (3.5 cm across).

Diamond and graphite have the same chemical composition (pure carbon) but different crystal structures. They are known as polymorphs of carbon. Diamond is formed under higher pressures than graphite (Figure 16c), and is less stable than graphite at the surface of the Earth. However, because of the strong bonding, it is very difficult to break down the diamond structure, so diamonds (fortunately) will not spontaneously transform into graphite! Note that not only is diamond harder than graphite, but it is also denser, as predicted by its structure (Table 3).

Minerals are never chemically pure; they always contain some foreign atoms. These impurity atoms may simply squeeze into the interstices. Another possibility is that certain elements may be able to directly replace (substitute for) the normal atoms in the ideal structure - although, for a comfortable fit, the substituting element must have a similar size and charge to the original atom. This phenomenon is called ionic substitution. An example is the substitution of Fe²⁺ for Mg²⁺, or Mg²⁺ for Fe²⁺, which occurs in the mineral olivine (Section 3.3.1).