# 3.5 Lewis symbols and structures  (Page 5/15)

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## Hypervalent molecules

Elements in the second period of the periodic table ( n = 2) can accommodate only eight electrons in their valence shell orbitals because they have only four valence orbitals (one 2 s and three 2 p orbitals). Elements in the third and higher periods ( n ≥ 3) have more than four valence orbitals and can share more than four pairs of electrons with other atoms because they have empty d orbitals in the same shell. Molecules formed from these elements are sometimes called hypervalent molecules . [link] shows the Lewis structures for two hypervalent molecules, PCl 5 and SF 6.

In some hypervalent molecules, such as IF 5 and XeF 4 , some of the electrons in the outer shell of the central atom are lone pairs:

When we write the Lewis structures for these molecules, we find that we have electrons left over after filling the valence shells of the outer atoms with eight electrons. These additional electrons must be assigned to the central atom.

## Writing lewis structures: octet rule violations

Xenon is a noble gas, but it forms a number of stable compounds. We examined XeF 4 earlier. What are the Lewis structures of XeF 2 and XeF 6 ?

## Solution

We can draw the Lewis structure of any covalent molecule by following the six steps discussed earlier. In this case, we can condense the last few steps, since not all of them apply.

1. Calculate the number of valence electrons:
XeF 2 : 8 + (2 $×$ 7) = 22
XeF 6 : 8 + (6 $×$ 7) = 50
2. Draw a skeleton joining the atoms by single bonds. Xenon will be the central atom because fluorine cannot be a central atom:
3. Distribute the remaining electrons.
XeF 2 : We place three lone pairs of electrons around each F atom, accounting for 12 electrons and giving each F atom 8 electrons. Thus, six electrons (three lone pairs) remain. These lone pairs must be placed on the Xe atom. This is acceptable because Xe atoms have empty valence shell d orbitals and can accommodate more than eight electrons. The Lewis structure of XeF 2 shows two bonding pairs and three lone pairs of electrons around the Xe atom:

XeF 6 : We place three lone pairs of electrons around each F atom, accounting for 36 electrons. Two electrons remain, and this lone pair is placed on the Xe atom:

The halogens form a class of compounds called the interhalogens, in which halogen atoms covalently bond to each other. Write the Lewis structures for the interhalogens BrCl 3 and ${\text{ICl}}_{4}{}^{\text{−}}.$

## Key concepts and summary

Valence electronic structures can be visualized by drawing Lewis symbols (for atoms and monatomic ions) and Lewis structures (for molecules and polyatomic ions). Lone pairs, unpaired electrons, and single, double, or triple bonds are used to indicate where the valence electrons are located around each atom in a Lewis structure. Most structures—especially those containing second row elements—obey the octet rule, in which every atom (except H) is surrounded by eight electrons. Exceptions to the octet rule occur for odd-electron molecules (free radicals), electron-deficient molecules, and hypervalent molecules.

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