Minerals and the crystalline state
Minerals and the crystalline state

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Minerals and the crystalline state

4.3 Covalent structures and bonding

Covalent bonding is formed when two atoms share two electrons (one from each atom) through overlap and merging of two electron orbitals (Figure 18). Crystals containing covalent bonds tend to have more complex structures than those of ionic or metallic structures.

Figure 18 Schematic representation of covalent bonding.

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 do 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 19a). The resulting structure, which has a repeating cubic shape, is illustrated in Figure 19b. The structure contains much more unoccupied space than do close-packed metal structures.

Figure 19 The diamond structure: (a) tetrahedral arrangement of covalent bonds around a carbon atom; (b) arrangement of carbon atoms and bonds in the unit cell (Section 6.2) - the unit cell is shown by black lines.

Another form of solid carbon with covalent bonding is graphite (Figure 20). Unlike those in diamond, the carbon atoms in graphite are covalently bonded to three neighbours in the same plane (Figure 20a), 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 20b).

Figure 20 The graphite structure: (a) triangular arrangement of covalent bonds around a carbon atom in the same plane; (b) part of the 3D structure of graphite, showing strongly bonded hexagonal sheets of carbon atoms, connected by weak bonds; (c) specimen of graphite (3.5 cm across).
  • How do the crystal structures of diamond and graphite account for the differences in hardness between the two minerals?

  • Diamond has a three-dimensional bonding pattern, with identical bonding in all directions, and no 'weak' directions. Graphite has a mainly two-dimensional pattern, with sheets of C-C bonds. Bonds between the sheets are very weak, so sheets can easily slide past each other, explaining graphite's use as a lubricant and why it is soft enough to mark paper.

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 21), and is less stable than graphite at the surface of the Earth. Indeed, temperature and pressure conditions within the Earth are such that diamond tends to form only at depths greater than 150 km. 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).

Described image
Figure 21 Phase diagram illustrating the stability fields for graphite and diamond.

Table 3 Relative densities of various minerals and ice, and notes on their structure and bonding.

SubstanceRelative density at room conditions (compared with water = 1.0)Structure and bonding
Ice, H2O0.9open structure; covalent bonds plus weak bonds between H2O molecules
Graphite, C2.2open structure; covalent bonds plus weak bonds between layers
Feldspar, KAlSi3O82.5open structure; predominantly covalent bonds
Quartz, SiO22.7open structure; predominantly covalent bonds
Olivine, Mg2SiO4−Fe2SiO43.2-4.4structure based on close-packing, but with ionic and covalent bonds; density increases as Fe content increases
Diamond, C3.5structure based on close-packing, but with covalent bonds
Barite, BaSO44.5ionic bonds between barium and sulfate groups
Hematite (iron oxide), Fe2O35.3structure based on close-packing; ionic and metallic bonds
Galena (lead sulfide), PbS7.6structure based on close-packing; ionic and metallic bonds
Silver, Ag10.5close-packed structure; metallic bonds
Gold, Au19.3close-packed structure; metallic bonds

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.

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