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

A comparison of bond lengths is consistent with our reasoning: the single $CO$ bond in ethanol is 148 pm, whereas the double bond in $C{O}_2$ is 116.

Knowing that oxygen atoms can double-bond, we can easily account for the structure of formaldehyde. The strengthof the $CO$ bond in $H_{2}CO$ is comparable to that in $C{O}_2$ , consistent with the Lewis structure here .

What about nitrogen atoms? We can compare the strength of the $CN$ bond in $HCN$ , 880 kJ, to that in methyl amine, 290 kJ. This dramaticdisparity again suggests the possibility of multiple bonding, and an appropriate Lewis structure for $HCN$ is shown here .

We can conclude that oxygen and nitrogen atoms, like carbon atoms, are capable of multiple bonding.Furthermore, our observations of oxygen and nitrogen reinforce our earlier deduction that multiple bonds are stronger than singlebonds, and their bond lengths are shorter.

As our final examples in this section, we consider molecules in which oxygen atoms are bonded to oxygenatoms. Oxygen-oxygen bonds appear primarily in two types of molecules. The first is simply the oxygen diatomic molecule, $O_2$ , and the second are the peroxides, typified by hydrogen peroxide, $H_2{O}_2$ . In a comparison of bond energies, we find that the strength of theOO bond in $O_2$ is 499 kJ whereas the strength of the OO bond in $H_2{O}_2$ is 142 kJ. This is easily understood in a comparison of the Lewis structures of these molecules, showing that the peroxide bond is asingle bond, whereas the $O_2$ bond is a double bond, shown here .

We conclude that an oxygen atom can satisfy its valence of 2 by forming two single bonds or by forming onedouble bond. In both cases, we can understand the stability of the resulting molecules by in terms of an octet of valenceelectrons.

Interpretation of lewis structures

Before further developing our model of chemical bonding based on Lewis structures, we pause to considerthe interpretation and importance of these structures. It is worth recalling that we have developed our model based on observations ofthe numbers of bonds formed by individual atoms and the number of valence electrons in each atom. In general, these structures areuseful for predicting whether a molecule is expected to be stable under normal conditions. If we cannot draw a Lewis structure inwhich each carbon, oxygen, nitrogen, or halogen has an octet of valence electrons, then the corresponding molecule probably is notstable. Consideration of bond strengths and bond lengths enhances the model by revealing the presence of double and triple bonds inthe Lewis structures of some molecules.

At this point, however, we have observed no information regarding the geometries of molecules. For example, wehave not considered the angles measured between bonds in molecules. Consequently, the Lewis structure model of chemical bonding doesnot at this level predict or interpret these bond angles. (This will be considered here .) Therefore, although the Lewis structure of methane is drawn as shown here .

This does not imply that methane is a flat molecule, or that the angles between $CH$ bonds in methane is 90°. Rather, the structure simply reveals that the carbon atom has a complete octet of valence electrons in amethane molecule, that all bonds are single bonds, and that there are no non-bonding electrons. Similarly, one can write the Lewisstructure for a water molecule in two apparently different ways, shown here .