Covalent Bonds

Only five elements have atoms with their valence p-subshells filled; these are the inert gases in the far right column of the periodic table. Their lack of chemical reactivity is explained by their stable electron configurations.

All other chemical elements need to lose or gain electrons to achieve electronic stability. Table 1 shows the stable electron configurations for the elements in the first three rows of the periodic table.

Most atoms achieve a stable number of valence electrons by sharing electrons with other atoms. Begin with fluorine, element 9, which has the electron configuration 1 s² 2 s² 2 p⁵. The orbitals of the valance electron shell are 2 s² 2 p⁵, and these 7 valence electrons can be portrayed in a diagram devised by the American chemist Gilbert Lewis (1875–1946). In Figure 1 , a Lewis diagram shows each valence electron as a single dot.

Figure 1 The valence electrons.

In Figure 2 , two fluorine atoms can each fill their valence orbitals with 8 electrons if they approach each other to share their single electrons.

Figure 2 Fluorine atoms sharing two electrons.

You can determine the number of valence electrons for the light elements by counting the columns from the left. (See Figure 3 .)

Figure 3 Counting valence electrons.

Phosphorus has 5 valence electrons, and chlorine has 7, so their isolated atoms have Lewis configurations, as shown in Figure 4 .

Figure 4 The valence electrons of phosphorus and chlorine.

Phosphorus must combine with 3 chlorines to complete its valence shell. (See Figure 5 .)

Figure 5 Six electrons shared between phosphorus and chlorine.

Study Figure 5 carefully. First, see that each atom is now surrounded by a full shell of 8 valence electrons. Of the 26 valence electrons, 6 are shared, and 20 are unshared. For the 6 that are shared to form the covalent bonds, the phosphorus atom contributed 3, and each of the chlorines contributed 1. The resulting PCl₃ molecule is usually drawn as shown in Figure 6 .

Figure 6 Three covalent bonds.

Each of the three lines represents the shared pair of electrons in a covalent bond. Some textbooks omit the nonbonding electrons for simplicity. Because the hydrogen atom has its single 1 s orbital completed with only 2 electrons, the hydrogen chloride molecule is drawn as shown in Figure 7 .

Figure 7 The hydrogen chloride molecule.

In some molecules, bonded atoms share more than 2 electrons, as in ethylene (C₂H₄), where the 2 carbons share 4 electrons. (See Figure 8 .)

Figure 8 A double bond between two carbon atoms.

Notice that each carbon achieves 8 electrons by this sharing. Because each shared pair constitutes a single covalent bond, the 2 shared pairs are called a double bond. The structure on the right side of Figure 8 shows this double bond of 4 shared electrons.

There are even triple bonds of 6 shared electrons, as in the nitrogen molecule. In N₂, each nitrogen atom contributes 5 valence electrons. Of the 10 electrons shown in Figure 9 , 4 are nonbonding, and 6 comprise the triple bond holding the nitrogen atoms together.

Figure 9 A triple bond between two nitrogen atoms.

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