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Covalent Bonds

A stable configuration is a completely filled s‐type subshell and a p‐type subshell. 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 s 2 2 s 2 2 p 5. The orbitals of the valance electron shell are 2 s 2 p, with two electrons in the 2 s and five electrons in the 2 p. These seven 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.

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

Figure 2. Fluorine atoms sharing two electrons.

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Count the electrons in the Lewis diagram; notice that there are 14 electrons, with each atom contributing 7. The two fluorine atoms form a stable F 2 molecule by sharing two electrons; this linkage is called a covalent bond.

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.

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

Figure 4. The valence electrons of phosphorus and chlorine.

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

Figure 5. Six electrons shared between phosphorus and chlorine.

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

Figure 6. Three covalent bonds.

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Each of the three lines represents the shared pair of electrons in a covalent bond. When lines are used to represent bonding pairs of electrons, the structure is often called a structural formula. Some textbooks omit the nonbonding electrons for simplicity.

Because the hydrogen atom has its single 1 s orbital completed with only two electrons, the hydrogen chloride molecule is drawn as shown in Figure 7.

Figure 7. The hydrogen chloride molecule.

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The hydrogen and chlorine atoms each donate one electron to the covalent bond. In the molecule, the hydrogen has completed its valence shell with two electrons, and the chlorine has a full shell with eight valence electrons.

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

Figure 8. A double bond between two carbon atoms.

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Notice that each carbon achieves eight electrons by this sharing. Because each shared pair constitutes a single covalent bond, the two shared pairs are called a double bond. The structure on the right side of Figure 8 shows this double bond of four shared electrons with two lines, and the left side of Figure

shows the double bond as two pairs of dots.

There are even triple bonds of six shared electrons, as in the nitrogen molecule. In N 2, each nitrogen atom contributes five valence electrons. Of the 10 electrons shown in Figure 9, four are nonbonding, and six comprise the triple bond holding the nitrogen atoms together.

Figure 9. A triple bond between two nitrogen atoms.

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  • Look at the periodic table and deduce the number of valence electrons for aluminum and oxygen from the positions of the columns for those two elements.
  • Draw a Lewis diagram representing the electron configuration of the hydrogen sulfide molecule, H 2S.

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