Earth Materials. John O'Brien
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from each atom.
A relevant nonmineral example is oxygen gas (O2) molecules. Each oxygen atom from column 16 (group VIA) requires two electrons to achieve a stable electron configuration in its highest principal quantum level. Since both oxygen atoms have an equally large electron affinity and electronegativity, they tend to share two electrons in order to achieve the stable electron configuration. This sharing is modeled as interpenetration or overlapping of the two electron clouds (Figure 2.12). Interpenetration of electron clouds due to the sharing of valence electrons forms a strong covalent or electron‐sharing bond. Because the bonds are localized in the region where the electrons are “shared” each atom has a larger probability of electrons in the area of the bond than it does elsewhere in its electron cloud. This causes each atom to become electrically polarized with a more negative charge in the vicinity of the bond and a less negative charge away from the bond. Polarization of atoms during covalent bonding is accentuated when covalent bonds form between different atoms with different electronegativities. This causes the electrons to be more tightly held by the more electronegative atom which in turn distorts the shape of the atoms so that they cannot be as effectively modeled as spheres in contact.
Other diatomic gases with covalent bonding mechanisms similar to oxygen include the column 17 (group VIIA) gases chlorine (Cl2), fluorine (F2), and iodine (I2) in which single electrons are shared between the two atoms to achieve a stable electron configuration. Another gas that possesses covalent bonds is nitrogen (N2) from column 15 (group V) where three electrons from each atom are shared to achieve a stable electron configuration. Nitrogen is the most abundant gas (>79% of the total) in Earth's lower atmosphere. In part because the two atoms in nitrogen and oxygen gas are held together by strong electron‐sharing bonds that yield stable electron configurations, these two molecules are the most abundant constituents of Earth's lower atmosphere.
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:
1 Hard and brittle at room temperature.
2 Insoluble in polar substances such as water.
3 Crystallize from melts.
4 Moderate to high melting temperatures.
5 Absorb very little light, producing transparent to translucent minerals with light colors and vitreous to sub‐vitreous lusters in macroscopic crystals.
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.
Figure 2.14 A model of metallic bonds with delocalized electrons (dark red) surrounding positive charge centers that consist of tightly held lower energy electrons (light red dots) surrounding individual nuclei (blue).
Minerals containing metallic bonds are generally characterized by the following features:
1 Fairly soft to moderately hard minerals.
2 Deform plastically; malleable and ductile.
3 Excellent electrical and thermal conductors.
4 Frequently high specific gravity.
5 Excellent absorbers and reflectors of light; so are commonly opaque with a metallic luster in macroscopic crystals.
2.3.5 Transitional (hybrid) bonds
Transitional or hybrid bonds display combinations of ionic, covalent and/or metallic bond behavior. Some transitional bonds can be modeled as ionic–covalent transitional, others as ionic–metallic or covalent–metallic transitional. A detailed discussion of all the possibilities is beyond the scope of this book, but because most bonds in Earth materials are transitional, it is a subject worthy of mention. The following discussion also serves to illustrate once again the enigmatic behavior of which electrons are capable.
Earlier in this chapter, we defined electronegativity in relation to the periodic table. Linus Pauling developed the concept of electronegativity (En) to help model transitional ionic–covalent bonds. In models of such bonds, electrons are partially transferred from the more metallic, more electropositive element to the less metallic, more electronegative element to produce a degree of ionization and electrostatic attraction typical of ionic bonding. At the same time, the electrons are partially shared between the two elements to produce a degree of electron sharing associated with