0.5 Covalent bonding and electron pair sharing  (Page 8/8)

However, it is very important to realize that these two structures are identical in the Lewis model, because both show that the oxygen atom has a complete octet of valence electrons,forms two single bonds with hydrogen atoms, and has two pairs of unshared electrons in its valence shell. In the same way, the twostructures for Freon 114 shown here are also identical .

These two drawings do not represent different structures or arrangements of the atoms in the bonds.

Finally, we must keep in mind that we have drawn Lewis structures strictly as a convenient tool for ourunderstanding of chemical bonding and molecular stability. It is based on commonly observed trends in valence, bonding, and bondstrengths. These structures must not be mistaken as observations themselves, however. As we encounter additional experimentalobservations, we must be prepared to adapt our Lewis structure model to fit these observations, but we must never adapt ourobservations to fit the Lewis model.

Extensions of the lewis structure model

With these thoughts in mind, we turn to a set of molecules which challenge the limits of the Lewis model indescribing molecular structures. First, we note that there are a variety of molecules for which atoms clearly must bond in such away as to have more than eight valence electrons. A conspicuous example is $S{F}_6$ , where the sulfur atom is bonded to six F atoms. As such, the S atommust have 12 valence shell electrons to form 6 covalent bonds. Similarly, the phosphorous atom in $P{\mathrm{Cl}}_5$ has 10 valence electrons in 5 covalent bonds, the Cl atom in $\mathrm{Cl}{F}_3$ has 10 valence electrons in 3 covalent bonds and two lone pairs. We also observe the interesting compounds of the noble gas atoms, e.g. $\mathrm{Xe}{O}_3$ , where noble gas atom begins with eight valence electrons evenbefore forming any bonds. In each of these cases, we note that the valence of the atoms S, P, Cl, and Xe are normally 2, 3, 1, and 0,yet more bonds than this are formed. In such cases, it is not possible to draw Lewis structures in which S, P, Cl, and Xe obeythe octet rule. We refer to these molecules as "expanded valence" molecules, meaning that the valence of the centralatom has expanded beyond the expected octet.

There are also a variety of molecules for which there are too few electrons to provide an octet for every atom. Most notably, Boron and Aluminum, from Group III, displaybonding behavior somewhat different than we have seen and thus less predictable from the model we have developed so far. These atomshave three valence shell electrons, so we might predict a valence of 5 on the basis of the octet rule. However, compounds in whichboron or aluminum atoms form five bonds are never observed, so we must conclude that simple predictions based on the octet rule arenot reliable for Group III.

Consider first boron trifluoride, $B{F}_3$ . The bonding here is relatively simple to model with a Lewis structure if we allow each valenceshell electron in the boron atom to be shared in a covalent bond with each fluorine atom.

Note that, in this structure, the boron atom has only six valence shell electrons, but the octet rule is obeyedby the fluorine atoms.

We might conclude from this one example that boron atoms obey a sextet rule. However, boron will form a stableion with hydrogen, $B{H}_4^{-}$ , in which the boron atom does have a complete octet. In addition, $B{F}_3$ will react with ammonia $N{H}_3$ for form a stable compound, $N{H}_{3}B{F}_3$ , for which a Lewis structure can be drawn in which boron has acomplete octet, shown here .

Compounds of aluminum follow similar trends. Aluminum trichloride, $\mathrm{Al}{\mathrm{Cl}}_3$ , aluminum hydride, $\mathrm{Al}{H}_3$ , and aluminum hydroxide, $\mathrm{Al}{\left(OH\right)}_3$ , all indicate a valence of 3 for aluminum, with six valenceelectrons in the bonded molecule. However, the stability of aluminum hydride ions, $\mathrm{Al}{H}_4^{-}$ , indicates that Al can also support an octet of valence shellelectrons as well.

We conclude that, although the octet rule can still be of some utility in understanding the chemistry of Boronand Aluminum, the compounds of these elements are less predictable from the octet rule. This should not be disconcerting, however. Theoctet rule was developed in on the basis of the observation that, for elements in Groups IV through VIII, the number of valenceelectrons plus the most common valence is equal to eight. Elements in Groups I, II, and III do not follow this observation mostcommonly.

Resonance structures

Another interesting challenge for the Lewis model we have developed is the set of molecules for which it ispossible to draw more than one structure in agreement with the octet rule. A notable example is the nitric acid molecule, $HN{O}_3$ , where all three oxygens are bonded to the nitrogen. Two structurescan be drawn for nitric acid with nitrogen and all three oxygens obeying the octet rule.

In each structure, of the oxygens not bonded to hydrogen, one shares a single bond with nitrogen while the othershares a double bond with nitrogen. These two structures are not identical, unlike the two freon structures in , because the atoms are bonded differently in the two structures.

Review and discussion questions