0.4 Covalent bonding and electron pair sharing  (Page 5/8)

These structures appear sensible from two regards. First, the trend in carbon-carbon bond strengths can beunderstood as arising from the increasing number of shared pairs of electrons. Second, each carbon atom has a complete octet ofelectrons. We refer to the two pairs of shared electrons in ethene as a double bond and the three shared pairs in acetylene as a triple bond .

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

Observation 3: compounds of nitrogen, oxygen, and the halogens

Many compounds composed primarily of carbon and hydrogen also contain some oxygen or nitrogen, or one or moreof the halogens. We thus seek to extend our understanding of bonding and stability by developing Lewis structures involvingthese atoms. Recall that a nitrogen atom has a valence of 3 and has five valence electrons. In our notation, we could draw a structurein which each of the five electrons appears separately in a ring, similar to what we drew for C. However, this would imply that anitrogen atom would generally form five bonds to pair its five valence electrons. Since the valence is actually 3, our notationshould reflect this. One possibility looks like this .

Note that this structure leaves three of the valence electrons "unpaired" and thus ready to join ina shared electron pair. The remaining two valence electrons are "paired," and this notation implies that they thereforeare not generally available for sharing in a covalent bond. This notation is consistent with the available data, i.e. five valence electrons and a valence of 3. Pairing the two non-bonding electrons seems reasonable in analogyto the fact that electrons are paired in forming covalent bonds.

Analogous structures can be drawn for oxygen, as well as for fluorine and the other halogens, as shown here .

With this notation in hand, we can now analyze structures for molecules including nitrogen, oxygen, and thehalogens. The hydrides are the easiest, shown here .

Note that the octet rule is clearly obeyed for oxygen, nitrogen, and the halogens.

At this point, it becomes very helpful to adopt one new convention: a pair of bonded electrons will now bemore easily represented in our Lewis structures by a straight line, rather than two dots. Double bonds and triple bonds are representedby double and triple straight lines between atoms. We will continue to show non-bonded electron pairs explicitly.

As before, when analyzing Lewis structures for larger molecules, we must already know which atoms are bonded towhich. For example, two very different compounds, ethanol and dimethyl ether, both have molecular formula $C_2{H}_{6}O$ . In ethanol, the two carbon atoms are bonded together and the oxygenatom is attached to one of the two carbons; the hydrogens are arranged to complete the valences of the carbons and the oxygenshown here .