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When nonmetallic atoms bond with other nonmetallic atoms they tend to form covalent bonds , also called electron‐sharing bonds. Because the elements involved are highly electronegative they each tend to attract electrons; neither gives them up easily. This is a little bit like a tug‐of‐war in which neither side can be moved so neither side ends up with sole possession of the electrons needed to achieve stable electron configurations. In simple models of covalent bonding, the atoms involved share valence (thus covalent) electrons (Figure 2.12). By sharing electrons, each atom gains the electrons necessary to achieve a more stable electron configuration in its highest principal quantum level.

Figure 2.11 Relationship between attractive and repulsive forces between ions produces a minimum net force when the near spherical surfaces of the ions are in contact.

Figure 2.12 Covalent bonding in oxygen (O₂) by the sharing of two electrons 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.