# C has 4 valence electrons, O has six electrons. The total number of electrons is 4 + 2(6) = 16.

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1 129 Lewis Structures G. N. Lewis hypothesized that electron pair bonds between unlike elements in the second (and sometimes the third) row occurred in a way that electrons were shared such that each element has an octet of electrons. Of course, hydrogen cannot have an octet. Look at HF, H 2 O, NH 3, CH 4 : In all of these systems, the heavy, central atom has eight electrons surrounding it. Let s also look at multiple bonding, as in CO 2 : C has 4 valence electrons, O has six electrons. The total number of electrons is 4 + 2(6) = 16. If each of the three atoms has a separate octet, 3(8) = 24 electrons would be required. Therefore = 8 electrons are shared in two bonds. Note that oxygen has two lone (unshared) pairs of electrons, carbon has none A more complicated system: O 3 : number of valence electrons is 6 per atom. 3(6) = 18 electrons If each atoms has a separate octet, then 24 electrons are required = 6 electrons to be shared. We can t share the electrons equally in the bonds, so we put two electrons in one bond, four in another. But, the two bonds are identical, so we have two structures like: The arrow means that the two structures are resonance structures

2 130 Exceptions to the Octet Rule Electron deficient compounds: BF 3 The boron only has a sextet of electrons around it. Octet expansion: I 3 - Atoms that have accessible d-electrons can expand their octets. The central iodine in I 3 - has ten electrons around it. Odd-electron compounds: NO N has 5 valence electrons, O has six. So, 11 electrons are available, five available for sharing. Four go into the double bond, the remaining electron is unpaired on the nitrogen atom. Note that for NO 2, there are two resonance structures. The odd electron is on the nitrogen, and there is no lone pair on N.

3 131 Electron Pair Repulsions and Molecular Structure By considering how many electron pairs surround the central atom in a compound, we can predict the geometry by considering how to minimize the repulsions among all electron pairs. First, we determine the number of electron pairs around the central atom. The electron pairs will arranger themselves to minimize repulsions. We distinguish between bonding electron pairs and lone (nonbonding) pairs. Three kinds of electron pair interactions occur: Bonding pair binding pair Lone pair lone pair Lone pair-bonding pair The prediction of molecular geometries based on these electron pair interactions is called the Valence Shell Electron Pair Repulsion (VSEPR) model. Sometimes also called the Gillispie-Nyholm model This diagram summarizes the geometries arising from various numbers of electron pairs. Sometimes called steric number

4 132 Examples We begin by counting electron pairs around the central atom in a molecule. Let s consider single bonds only. First, draw the Lewis structure. For counting purposes, multiple electron pairs in a multiple bond count as a single pair. The electron pairs will arrange themselves in a way that minimizes repulsion. Example 1: BeCl 2 electron. Be has a (2s) 2 configuration. Each Cl atom provides one bonding With two pairs of electrons, the structure is linear. Example 2: BF 3 As noted above, there are three electron pairs. The electrons arrange themselves around the equator of the sphere. The trigonal planar structure results. Example 3: This is one of the most common examples. H 2 O, NH 3, CH 4 are examples With four electron pairs, the charges arrange themselves at the apexes of a tetrahedron. For H 2 O, two hydrogen occupy two positions at the tetrahedron, two lone pairs also occupy two positions.

5 133 For NH 3, one lone pair occupies a position on the tetrahedron: The lone pairs provide distortions from perfect tetrahedral symmetry The single lone pair provides a distortion of about 2.5, while two lone pairs provide a distortion of about 5 Example 4: Five electron pairs. PCl 5 P has an s 2 p 3 configuration, so 5 electrons from each of five Cl atoms yields 10 electrons. Trigonal bipyramid give minimum energy structure. Now let s consider I = 10 electrons Of the structures given, we must now minimize the lone pair interactions. These are the most repulsive. The last structure (linear) is the most stable.

6 134 Finally, XeF 4 is a case for six electron pairs. Xe contributes 4 pairs. The four F atoms contribute four more electrons. The base structure is the octahedron. Having lone pairs at 180 produces the lowest energy structure. Multiple Bonds When we work out VSEPR structures for systems with multiple bonds, we only count a single electron pair for assigning the appropriate geometry. Consider CO 2 : For a VSEPR analysis, we have only two electron pairs on carbon, and the structure is linear. For O 3, ignoring the multiple bonds, note that there are three electron pairs involved in the bonding. So the molecule is bent.

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