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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, H2O, NH3, CH4: In all of these systems, the heavy, central atom has eight
electrons surrounding it.
Let’s also look at multiple bonding, as in CO2:
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 24 – 16 = 8 electrons are shared in two bonds. Note that oxygen has two lone
(unshared) pairs of electrons, carbon has none
A more complicated system:
O3: number of valence electrons is 6 per atom. 3(6) = 18 electrons
If each atoms has a separate octet, then 24 electrons are required. 24 – 18 = 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
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Exceptions to the Octet Rule
Electron deficient compounds: BF3
The boron only has a sextet of electrons around it.
Octet expansion: I3-
Atoms that have accessible d-electrons can expand their octets. The central iodine in I3has 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 NO2, there are two resonance structures. The odd electron is on the
nitrogen, and there is no lone pair on N.
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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”
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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: BeCl2
electron.
Be has a (2s)2 configuration. Each Cl atom provides one bonding
With two pairs of electrons, the structure is linear.
Example 2: BF3 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. H2O, NH3, CH4 are examples
With four electron pairs, the charges arrange themselves at the apexes of a tetrahedron.
For H2O, two hydrogen occupy two positions at the
tetrahedron, two lone pairs also occupy two positions.
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For NH3, 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.
PCl5 P has an s2p3 configuration, so 5 electrons from each of five Cl atoms
yields 10 electrons. Trigonal bipyramid give minimum energy structure.
Now let’s consider I3-. 7 + 2 + 1 = 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.
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Finally, XeF4 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 CO2: For a VSEPR analysis, we have only two electron pairs on carbon, and
the structure is linear.
For O3, ignoring the multiple bonds, note that there are three electron pairs involved in
the bonding. So the molecule is bent.