- 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 6: Advanced Lewis Structures
Sometimes, when molecules form where the normal patterns of bonding are not followed, or if the molecules are actually charged molecular ions, the atoms themselves can have a formal charge on them. This is because any atom brought some number of valence electrons into the molecule, and it wants to keep that same number of electrons close to it, to balance the charge of the positive protons in its nucleus. So, if the atom ends up with more or fewer electrons around it, it can end up with a negative or positive formal charge, respectively. The other thing about formal charges, is that if a molecule has formal charges, they all must add up to the total charge on a molecule. Therefore, any Lewis structure of a cation or an anion must have at least one formal charge on it.
Figure 5-10. Advanced Lewis Structures
Three examples of Lewis structures containing formal charges: A. cynanide ion, B. hydronium ion, C. nitrous oxide.
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In order to talk about formal charges, we have to first talk about what it means to "own" an electron. First of all, an atom owns all of the electrons in their lone pairs, because it isn't sharing any of those electrons. Secondly, an atom owns half of the electrons in each of its covalent bonds, because it is sharing those electrons with an atom that owns the other half. Let's look at a valid Lewis structure for the cyanide ion (CN-), a highly poisonous anion that is used in gold mining (Figure 5-10A). The carbon has a minus charge on it; that's the formal charge. Why does carbon have a negative formal charge on it? Well, the carbon owns the two electrons in its lone pair, and half of each of its three bonds, for a total of five electrons. But how many valence electrons did carbon bring to the molecule? Just four electrons. So, by owning one more electron than it normally has in its valence, it gets a -1 charge. Also, see that carbon is not making its typical four covalent bonds: When atoms break their normal octet rule patterns, they end up with formal charges.
For comparison, let's look at the hydronium ion (H3O+), which is the molecule responsible for acidity in water solutions (Figure 5-10B). Here, the oxygen atom owns the two electrons in its lone pair and half of the three covalent bonds, for a total of 5 electrons. However, oxygen normally has six valence electrons so it has an overall formal charge of plus one. Also, this charge correctly matches the total charge on the overall ion, and note that once again oxygen is breaking its normal pattern of two bonds and two lone pairs. Sometimes, a neutral molecule can have formal charges, such at nitrous oxide (N2O), which is also known as "laughing gas" and is used in as an anesthetic in dental surgeries. Note how both the nitrogen and the oxygen have formal charges, but their charges add up to zero, which is always the net charge on a neutral molecule. It is actually possible to draw another Lewis structure for nitrous oxide, where there is a double bond between both sets of atoms, rather than a triple bond and a single bond (Figure 5-10C). This pair of Lewis structures is called a pair of "resonance structures," which are discussed in the Resonance sidebar.
Figure 5-11. Sulfuryl chloride (SO2Cl2)
The sulfur in the center of the molecule is actually making six covalent bonds, but that is allowed for sulfur, which is a third-row element. Notice that the oxygen, which is a second row element, and the chlorine still obey the octet rule. One of the other ways that Lewis structures for real molecules can bend the octet rule is when they are molecules called "radicals" (See the Radicals sidebar).
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For second row elements, their valence electrons are in their 2s and 2p orbitals, which can hold the eight electrons for the octet rule. However, when constructing the periodic table in Unit 4, we saw that the third shell has one s orbital, three p orbitals, and five d orbitals. However, since the 3s and 3p orbitals are lower energy than the 3d orbitals, electrons tend to occupy the s and p orbitals first. Hence, atoms in the third row of the periodic table can follow the octet rule. However, they don't have to. Sometimes, though, third and higher row elements will adopt an expanded octet. When this happens, an atom can have more than eight electrons in its valence shell, meaning that, unlike second-row atoms, these elements can form more than four bonds. Because of the 5d orbitals, elements that expand their octet can have up to 18 electrons. For example, the compound sulfuryl chloride (SO2Cl2), which is used as a source of chlorine in chemical reactions, has an expanded octet. (Figure 5-11)