- Online Text
- 1. Introduction
- 2. Valence Electron Patterns and Lewis Structures
- 3. Ionic Bonds
- 4. Covalent Bonds and the Octet Rule
- 5. Polarity and Basic Lewis Structures
- 6. Advanced Lewis Structures
- 7. VSEPR Theory
- 8. Hybrid Orbitals
- 9. Intermolecular Forces
- 10. Physical Properties of Molecules
- 11. Conclusion
- 12. Further Reading
- Unit Guide (PDF)
Section 4: Covalent Bonds and the Octet Rule
While ionic structures are very common, their Lewis structures are not very interesting or informative, because all that really matters are their charges. However, the second type of bonding, covalent bonding, relies heavily on Lewis structures to show the roles of the electrons in these bonds. These structures will then be able to help us understand more about the molecules, from their shapes to their physical properties.
What is a covalent bond? Unlike an ionic bond where atoms have very different electronegativies, and one atom gives up electrons to the other to form ions, covalent bonds usually form between two atoms that have similar electronegativities. These elements tend to be the non-metals, like those shown at the top of Figure 5-6. In a covalent bond, elements share pairs of electrons between them, trying to reach this magical number of eight valence electrons.
Figure 5-6. Valence Electrons and Patterns
The groups of the non-metals that often form covalent bonds are shown with their Lewis dot atoms. At the bottom of each column, there is a sample Lewis pattern to show how each of these elements makes covalent bonds trying to reach its octet of electrons through sharing.
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For covalent bonding, elements try to follow the octet rule. The octet rule states that elements will work to get a total of eight valence electrons through a combination of lone pairs of electrons on the atom, and covalent bonds whereby two elements share a pair of electrons. A halogen, like chlorine, has seven valence electrons and needs one more to reach its octet; therefore, it can make one covalent bond along with its three lone pairs. In that covalent bond, two electrons are shared, one from each atom on either side of the bond. Whereas carbon, which has four valence electrons, needs to make four covalent bonds to get a total of eight electrons. Each time a covalent bond is formed between two atoms, each atom brings one electron—but through sharing, both atoms get to "feel" like they actually have two electrons from that bond. All of these patterns are summarized for the main group elements in Figure 5-6. The bottom row of the figure shows the patterns of covalent bonds and lone pairs each atom wants to reach its octet. Because the noble gases already have eight valence electrons, in general, they do not make covalent or ionic compounds because they are already stable and have satisfied their octet. Note that hydrogen is special. Because it only has that lonely 1s orbital, it can only hold two electrons. Because of this, hydrogen atoms are said to have an "octet of two," for the octet rule.
To reiterate, in covalent bonds, both atomic nuclei are holding onto the same two electrons in the bond, so like the teams in a tug-of-war game holding onto the same rope they remain linked to each other. Figure 5-7 shows four covalent molecules, each of which is a diatomic molecule with two atoms of the same element sharing electrons. For each atom, you can count up its lone pair electrons and the electrons in the bonds attached to it. When you do this for the fluorine (F), oxygen (O), or nitrogen (N), you get a total of eight electrons around each atom, showing that these follow the octet rule. The hydrogen atoms have just the bond and no lone pairs, but that works out perfectly because that lone bond gets hydrogen the two electrons it needs for its octet of two.
Figure 5-7. Diatomic Covalent Molecules
The simplest covalently-bound molecules are diatomic molecules; these happen between only two atoms sharing at least two electrons. See how the patterns of Figure 5-6 are repeated here: Fluorines make one covalent bond, oxygens make two covalent bonds, and nitrogens make three covalent bonds. We can make more than one covalent bond to the same atom in order to share enough electrons to reach the octet. For example, the oxygen atom has two lone pairs and two bonds, each of which represents two electrons and four lone pair electrons. These, plus four covalently bonded electrons equals eight electrons to satisfy the octet rule.
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Additionally, each of the molecules in Figure 5-7 are two identical elements bound together, so they have identical electronegavities, which means they share the electrons evenly between them. This is referred to as a true or a pure covalent bond, and, as we will see in a later section, a symmetrical sharing of electrons leads to molecules that have a balanced electron cloud, also referred to as "non-polar" molecules.
For us, it is less important to be able at first to draw Lewis structures from scratch; however, it is very important for us to recognize a Lewis structure and then to use it to learn things about that molecule and its structure and properties. If we know the rules for making Lewis structures, we can tell if the ones we see are correct, which is the first important step in looking at molecules.