0.6 Covalent bonding and electron pair sharing  (Page 7/8)

Looking at the data, it is very clear that the bonding between the carbon atoms in these three molecules must be very different. The bond strength increases when there are fewer H atoms, and the bond length gets progressively shorter with fewer H atoms. These observations are reinforced by looking at carbon-carbonbond strengths and lengths in other C _n H _2n+2 molecules, like propane and butane. In all of these molecules, the bond strengths and lengths are comparable to those in ethane.

The bond in ethene is about one and a half times stronger than the bond in ethane; this suggests that there might be a second bond between the two carbon atoms. This could be formed from the two unpaired and unshared electrons in the ethene structure above. Similarly, the bond in acetylene is about two and a half times stronger than the bond in ethane, so we can imagine that this results from the sharing of three pairs of electrons between the two carbon atoms. These assumptions produce the following Lewis structures:

These structures appear to make sense from two regards. First, these structures would explain the trend in carbon-carbon bond strengths as arising from the increasing number of shared pairs of electrons. Second, each carbon atom has a complete octet of electrons. Our model now reveals that there is a “double bond” in ethane and a “triple bond” in acetylene.

We thus extend our model of valence shell electron pair sharing to conclude that carbon atoms can bond by sharing one, two, or three pairs of electrons as needed to complete an octet of electrons, and that the strength of the bond is greater when more pairs of electrons are shared. Moreover, the data above tell us that the carbon-carbon bond in acetylene is shorter than that in ethene, which is shorter than that in ethane. We conclude that triple bonds are shorter than double bonds, which are shorter than single bonds.

Observation 5: hydrocarbon compounds containing nitrogen, oxygen, and the halogens

Many compounds composed primarily of carbon and hydrogen also contain some oxygen or nitrogen, or one or more of the halogens. These include the majority of biomolecules, such as amino acids and proteins. This means that it would be very desirable to extend our understanding of bonding by developing Lewis structures for these types of molecules.

Recall that a nitrogen atom has a valence of 3 and has five valence electrons. In our notation, we could draw a structure in which each of the five electrons appears separately in a ring, similar to what we drew for C. However, our new understanding of covalent bonding tells us that unpaired electrons on atoms are shared with other atoms to form bonds. Having all five valence electrons in an N atom unpaired would imply that an N atom would generally form five bonds to pair its five valence electrons. Since the valence is actually 3, our notation should have three unpaired electrons. One possibility looks like: