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A reaction is shown with three Lewis diagrams. The left diagram shows a boron atom single bonded to three fluorine atoms, each with three lone pairs of electrons. There is a plus sign. The next structure shows a nitrogen atom with one lone pair of electrons single bonded to three hydrogen atoms. A right-facing arrow leads to the final Lewis structure that shows a boron atom single bonded to a nitrogen atom and single bonded to three fluorine atoms, each with three lone pairs of electrons. The nitrogen atom is also single bonded to three hydrogen atoms. The bond between the boron atom and the nitrogen atom is colored red.

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.

Two Lewis structures are shown. The left shows a phosphorus atom single bonded to five chlorine atoms, each with three lone pairs of electrons. The right shows a sulfur atom single bonded to six fluorine atoms, each with three lone pairs of electrons.
In PCl 5 , the central atom phosphorus shares five pairs of electrons. In SF 6 , sulfur shares six pairs of electrons.

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:

Two Lewis structures are shown. The left shows an iodine atom with one lone pair single bonded to five fluorine atoms, each with three lone pairs of electrons. The right diagram shows a xenon atom with two lone pairs of electrons single bonded to four fluorine atoms, each with three lone pairs of electrons.

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:
    Two Lewis diagrams are shown. The left depicts a xenon atom single bonded to two fluorine atoms. The right shows a xenon atom single bonded to six fluorine atoms.
  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:
    A Lewis diagram shows a xenon atom with three lone pairs of electrons single bonded to two fluorine atoms, each with three lone pairs of electrons.
    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:
    This structure shows a xenon atom single bonded to six fluorine atoms. Each fluorine atom has three lone pairs of electrons.

Check your learning

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 ICl 4 .

Answer:

Two Lewis structures are shown. The left depicts a bromine atom with two lone pairs of electrons single bonded to three chlorine atoms, each with three lone pairs of electrons. The right shows an iodine atom, with two lone pairs of electrons, single boned to four chlorine atoms, each with three lone pairs of electrons. This structure is surrounded by brackets and has a superscripted negative sign.
Got questions? Get instant answers now!

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.

Questions & Answers

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s.
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Damian
That is a great question Damian. best way to answer that question is to Google it. there are hundreds of applications for buck minister fullerenes, from medical to aerospace. you can also find plenty of research papers that will give you great detail on the potential applications of fullerenes.
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Mostly, they use nano carbon for electronics and for materials to be strengthened.
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s. Reply
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Practice Key Terms 9

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Source:  OpenStax, Ut austin - principles of chemistry. OpenStax CNX. Mar 31, 2016 Download for free at http://legacy.cnx.org/content/col11830/1.13
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