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When the nonbonding electrons are used to satisfy the octet of the oxygen atom, the carbon atom has a total of only six valence electrons.

We therefore assume that the carbon and oxygen atoms share two pairs of electrons. There are now four bonds in the skeleton structure, which leaves only four nonbonding electrons.

This is enough, however, to satisfy the octets of the carbon and oxygen atoms. Every once in a while, we encounter a molecule for which it is impossible to write a satisfactory Lewis structure. There are three covalent bonds in the most reasonable skeleton structure for the molecule.

Because it takes six electrons to form the skeleton structure, there are 18 nonbonding valence electrons. Each fluorine atom needs six nonbonding electrons to satisfy its octet. Thus, all of the nonbonding electrons are consumed by the three fluorine atoms. As a result, we run out of electrons while the boron atom has only six valence electrons.

Because neither boron nor fluorine falls in this category, we have to stop with what appears to be an unsatisfactory Lewis structure. It is also possible to encounter a molecule that seems to have too many valence electrons. When that happens, we expand the valence shell of the central atom. Consider the Lewis structure for sulfur tetrafluoride SF 4 which contains 34 valence electrons.

There are four covalent bonds in the skeleton structure for SF 4. Because this requires using eight valence electrons to form the covalent bonds that hold the molecule together, there are 26 nonbonding valence electrons. Because there are four of these atoms, so we need 24 nonbonding electrons for this purpose.

But there are 26 nonbonding electrons in this molecule. We have already satisfied the octets for all five atoms, and we still have one more pair of valence electrons. We therefore expand the valence shell of the sulfur atom to hold more than eight electrons. This raises an interesting question: How does the sulfur atom in SF 4 hold 10 electrons in its valence shell? The electron configuration for a neutral sulfur atom seems to suggest that it takes eight electrons to fill the 3 s and 3 p orbitals in the valence shell of this atom.

But let's look, once again, at the selection rules for atomic orbitals. Because the 3 d orbitals on a neutral sulfur atom are all empty, one of these orbitals can be used to hold the extra pair of electrons on the sulfur atom in SF 4.

Click here to check your answer to Practice Problem 3. The only difference between these Lewis structures is the identity of the oxygen atom to which the double bond is formed. As a result, they must be equally satisfactory representations of the molecule. Interestingly enough, neither of these structures is correct. The two Lewis structures suggest that one of the sulfur-oxygen bonds is stronger than the other. There is no difference between the length of the two bonds in SO 2 , however, which suggests that the two sulfur-oxygen bonds are equally strong.

When we can write more than one satisfactory Lewis structure, the molecule is an average, or resonance hybrid , of these structures. The meaning of the term resonance can be best understood by an analogy. In music, the notes in a chord are often said to resonate they mix to give something that is more than the sum of its parts.

In a similar sense, the two Lewis structures for the SO 2 molecule are in resonance. They mix to give a hybrid that is more than the sum of its components. The fact that SO 2 is a resonance hybrid of two Lewis structures is indicated by writing a double-headed arrow between these Lewis structures, as shown in the figure above.

Click here to check your answer to Practice Problem 4. It is sometimes useful to calculate the formal charge on each atom in a Lewis structure.

The first step in this calculation involves dividing the electrons in each covalent bond between the atoms that form the bond.

The number of valence electrons formally assigned to each atom is then compared with the number of valence electrons on a neutral atom of the element. If the atom has more valence electrons than a neutral atom, it is assumed to carry a formal negative charge. If it has fewer valence electrons it is assigned a formal positive charge. Use the concept of formal charge to explain the meaning of the positive and negative signs in the following Lewis structure.

Click here to check your answer to Practice Problem 5.

The Easy Way to Teach About Valence Electrons: Activities for Kinesthetic Learners

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valance electron diagram homework help A summary of Electron Configuration and Valence Electrons in s Atomic Structure. the diagram above clearly shows that the 4s orbital is filled before the 3d Valence Electrons & Bohr Diagrams have a negative charge Neutrons are Neutral Valence Electrons Each electron shell can hold a certain number of electrons Draw a dot & cross diagram. Show transcribed image text The following Lewis diagram represents the valence electron configuration of a main-group element EX: If this element is in period 5, its valence electron configuration is. Expert Answer. Get this answer with Chegg Study View this answer. OR. Find your book.