- 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 5: Polarity and Basic Lewis Structures
In the last section, we only saw simple diatomic Lewis structures where two of the same atom were bound together. However, most molecules are made up more than two atoms, and two atoms not of the same type of element will share electrons, too. Also, as the molecules and their Lewis structures get more complicated, the sharing of electrons between the atoms can become uneven, leading to "polar covalent bonds."
In practice, electrons are shared equally only in covalent bonds between identical atoms, such as in diatomic molecules like Cl2, O2, or N2. In covalent bonds between two different atoms, the more electronegative atom pulls the electron density of the electron cloud toward itself more than the less electronegative atom. Electron density is just another way of describing areas of space where the electrons in bonds and lone pairs are likely to be. It is like when one of the teams in the tug-of-war is stronger than the other and ends up with more of the rope. For example, when a chlorine atom is bound to a hydrogen atom, as in hydrogen chloride or hydrochloric acid (HCl), the electrons are a little more likely to be found near the chlorine than near the hydrogen, because chlorine is more electronegative. The chlorine portion of the molecule ends up with a slight negative charge as a result, and the hydrogen end is left with a slight positive charge. (Figure 5-8)
Figure 5-8. Electron Clouds and Lewis Structures
The H2 and the HCl molecules are represented here showing the electron density of their electron clouds as they covalently share electrons. Note how the very electronegative chlorine atom in HCl helps to tilt the cloud of electrons on the molecule towards its side, making the chlorine end a bit more negative than the hydrogen end, which ends up a bit more positive, and thus polar. By contrast, the hydrogen molecule has a perfectly symmetrical cloud from two hydrogen atoms evenly sharing electrons, and the molecule is therefore non-polar.
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Because of this uneven electron cloud, and the partial positive and negative charges around an unsymmetrical molecule like hydrogen chloride, the molecule is said to be polar. Once again, polar refers to an uneven distribution of electrons and can apply to an individual bond or an entire molecule. So, the H-Cl bond is a polar covalent bond and the HCl molecule is polar. We will see later in this unit that the polarity of a molecule is one of the key factors in determining many of its properties.
Most bonds can therefore be said to lie on a continuum from pure covalent bonds (like fluorine (F2)), in which both atoms have identical electronegativity and share the electrons equally, through polar covalent bonds (like hydrogen chloride (HCl)), in which the atoms' electronegativities are different and electrons are shared unequally between them, to ionic bonds (like sodium chloride (NaCl)), in which the difference in electronegativity is so great that electrons have been completely transferred from one of the atoms to the other; that is, one tug-of-war team is so much stronger they end up with all of the rope.
We can describe molecules using the Lewis dot symbolism we saw above for individual atoms and diatomic molecules. However, we want to learn more about the bonds in larger molecules and how that will affect the three-dimensional shapes of these molecules. In order to do that, we need to look at Lewis structures for more complicated molecules. The four examples in Figure 5-9 are common molecules that will let us see the common patterns of Lewis structure and bonding.
Let's start with the most common molecule on planet Earth: water (Figure 5-9A). With a formula of H2O, three atoms need to share electrons to reach their octets. The hydrogen atoms are always limited to at most one bond, since they can only hold at most two electrons, and one bond represents two electrons. The oxygen makes one covalent bond to each hydrogen atom, and it has two lone pairs as well. These two covalent bonds give four electrons towards the octet, and the two lone pairs give the other four for a total of eight. This matches perfectly with the pattern that was in the table in the last section, where it showed that oxygen likes to make two bonds and have two lone pairs to reach an octet. Each of these bonds between two atoms where just two electrons are being shared is called a "single bond."
Figure 5-9. Basic Lewis Structures
The Lewis structures for A. water, B. ammonia, C. formaldehyde, and D. acetylene. Water and ammonia show examples of single bonds, formaldehyde and acetylene have a double and triple bond, respectively.
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Another molecule that is closely related to water is ammonia (NH3, Figure 5-9B). Ammonia is a common household cleaner and a compound produced in small quantities in the kidneys to control the level of acidity in urine before being excreted by the body. Nitrogen follows its pattern here again to reach an octet: It likes to make three single bonds and have one lone pair for a total of eight electrons.
It can also happen that two atoms share more than one pair of electrons, a situation referred to as a "multiple bond." For example, let's consider the molecule formaldehyde (CH2O), which was used historically to preserve biological samples in jars (Figure 5-9C). For carbon to reach an octet, it wants to make four covalent bonds, but with only three atoms present, it is going to have to make more than one covalent bond to one of the other atoms. Specifically, two pairs of electrons are being shared between the carbon and oxygen. When two pairs of electrons are shared in this way, it is called a "double bond."
It is also possible for atoms to share three pairs of electrons, a situation called a "triple bond." An example of this is the molecule acetylene, used in oxy-acetylene welding torches (Figure 5-9D). Again, note that the octet rule is satisfied. The carbon in acetylene has eight electrons just as it did in formaldehyde, but this time it accomplishes this by making a single bond to the hydrogen and a triple bond to the other carbon, for the same total of four covalent bonds.