Simple Bonding
Theory
Review of Lewis Structures and VSEPR
Theory
Resonance
Some molecules may have more than one
valid Lewis structure. These structures differ in
the placement of multiple bonds.
In molecules with resonance, none of the
Lewis structures accurately represents the true
bonding in the molecule.
Resonance
The molecule SO2 has two resonance
structures:
: :
:
: :
: :
:
: :
[ O=S-O: ] ↔ [ :O-S=O]
The molecule has two equivalent bonds
between sulfur and oxygen.
Resonance
The sulfur-oxygen bonds are identicallonger than double bonds, and shorter than
single bonds.
:
:
:
:
: :
: :
[ O=S-O: ] ↔ [ :O-S=O]
The true structure of the molecule is in
between the two Lewis structures drawn.
Resonance
The Lewis structure and VSEPR theory
correctly predict the shape, polarity and bond
angles (≈120o) of the molecule.
However, this approach cannot accurately
predict the identical nature of the bonds. The
sulfur oxygen bonds are equivalent, and
somewhat shorter than single bonds, and slightly
longer than double bonds.
Formal Charges
Formal charge is a way to keep track of the
electrons in a covalent molecule. The formal
charges can also be used to determine if one
Lewis structure is more valid than another.
Formal Charges
The formal charge on an atom is a
comparison between the number of valence
electrons on each atom and the number of
electrons it has in the Lewis structure.
Formal Charges
Consider the ion SCN-1. There are three
valid Lewis structures for the ion.
-1
: :
: :
:
:
-1
-1
[:S=C=N:] ↔ [:S-C N:] ↔ [:S C-N:]
Formal charges can be used to determine the
major contributor(s) to the actual structure of
the ion.
Formal Charges
Divide the bonds in half and determine the
number of electrons on each atom.
-1
: :
: :
:
:
-1
-1
[:S=C=N:] ↔ [:S-C N:] ↔ [:S C-N:]
6e 4e 6e
7e 4e 5e
5e 4e 7e
Formal Charges
Compare the number of electrons in the
structure to the number of valence electrons.
-1
: :
: :
:
:
-1
-1
[:S=C=N:] ↔ [:S-C N:] ↔ [:S C-N:]
6e 4e 6e
6e 4e 5e
7e 4e 5e
6e 4e 5e
5e 4e 7e
6e 4e 5e
Formal Charges
The net charge is the formal charge on each
atom.
0
0
-1
0
-1
+1
0 -2
: :
: :
:
:
-1
-1 0
-1
[:S=C=N:] ↔ [:S-C N:] ↔ [:S C-N:]
6e 4e 6e
6e 4e 5e
7e 4e 5e
6e 4e 5e
5e 4e 7e
6e 4e 5e
Formal Charges
The net charge is the formal charge on each
atom.
0
0
-1
0
-1
+1
0 -2
: :
: :
:
:
-1
-1 0
-1
[:S=C=N:] ↔ [:S-C N:] ↔ [:S C-N:]
The sum of the formal charges must equal
the charge on the ion.
Formal Charges
There are two rules used to determine the
most likely Lewis structure(s).
1. Atoms try to achieve formal charges as
close to zero as possible.
2. Any negative formal charges should
reside on the most electronegative atoms.
Formal Charges
The third Lewis structure is unlikely, due to
the high formal charge on nitrogen.
0
0
-1
0
-1
+1
0 -2
: :
: :
:
:
-1
-1 0
-1
[:S=C=N:] ↔ [:S-C N:] ↔ [:S C-N:]
Since nitrogen is more electronegative than
sulfur, the first structure should be the major
contributor.
Formal Charges
The actual molecule will be somewhere in
between the first and second structures.
0
0
-1
0
-1
+1
0 -2
: :
: :
:
:
-1
-1 0
-1
[:S=C=N:] ↔ [:S-C N:] ↔ [:S C-N:]
The sulfur-carbon bond should be slightly
longer than a double bond, and the carbonnitrogen bond should be slightly shorter than a
double bond.
Formal Charges
Some Lewis structures violate the formal
charge guidelines (ex. CO and BF3).
In either example, if the “octet rule” is to be
satisfied, a formal charge of +1 must be placed
on the more electronegative element.
Formal Charges
In the case of BF3, experimental evidence suggests
there is some multiple bonding between the boron and
the fluorine atoms to create an extended π system.
VSEPR Theory/Molecular Shapes
Molecules with lone pairs of electrons on the
central atom are described as “distorted” from
ideal geometry.
NH3 has a bond angle of 106.6o
H2O has a bond angle of 104.5o
Hybridization of Distorted
Molecules
These molecules can be viewed as having
different hybridization in the bonds and lone
pairs.
Since lone pairs are more diffuse and
“bulky”, they will have greater “s” character.
Since bonds are more compact and
directional, they will have greater “p” character.
Hybridization of Distorted
Molecules
s = p-1
cos θ = s-1
p

Determine the hybridization of the bonding and
non-bonding orbitals on nitrogen in ammonia.
The bond angle is 106.6o.
“Expanded” Coordination
Elements below period 2 can use d orbitals
to make more than 4 bonds.
A trigonal bipyramid, with sp3d hybridization,
is the common structure for 5 atoms or lone
pairs. These molecules are not fully hybridized.
The trigonal plane is sp2 hybridized, and the
axial positions are a mixture of the pz and dz2
orbitals.
5 Coordinate Complexes
In molecules with trigonal planar geometry,
not all bond angles and positions are equivalent.
If the molecule contains lone pairs of
electrons, they will occupy the trigonal plane
rather than the axial sites. This will minimize
repulsion within the molecule.
5 Coordinate Complexes
The lone pair of electrons in the trigonal
plane will force the fluorine atoms closer
together.
F
F
101.6o
187o
S:
F
F
“Expanded” Coordination
Molecules with 6 atoms or lone pairs
typically take an octahedral shape, with sp3d2
hybridization. Some complexes with
coordination number 6 may be distorted prisms.
“Expanded” Coordination
Molecules with
coordination number
of 7 may have a
pentagonal
bipyramidal shape,
and those with a
coordination number
of 8 may be a square
antiprism.
Multiple Bonds and Distortion
Multiple bonds have an influence on
structure similar to that of lone pairs. The
diffuse nature of π bonds causes repulsion with
lone pairs and bonding pairs of electrons.
Electronegativity
Electronegativity is a measure of an atoms
ability to attract electrons from an atom to which
it is bonded.
Mulliken Electronegativites
Mulliken defined electronegativity as:
½ (electron affinity + ionization energy)
This approach yields very high
electronegativity values for He and Ne, even
though they do not form compounds.
Electronegativity Values
Electronegativity and Bond Angles
Electronegative atoms bonded to a less
electronegative central atom tend to draw electron
density away from the central atom, thus lowering
repulsion.
PF3
PCl3
PBr3
97.8o
100.3o
101o
Hybridization in Period 3 and Below
For central atoms in period 3 or below,
hybridization is less common.
AsH3
PH3
PF3
PCl3
PBr3
91.8o
93.8o
97.8o
100.3o
101o
Steric Number 5
In molecules such as PCl5, the five bonds are
not equal in length.
Steric Number 5
Any lone pairs preferentially occupy the
equatorial positions so as to minimize electron
pair repulsion.
Steric Number 5
In five-coordinate compounds, the more
electronegative (less bulky) atoms occupy the
axial positions.
Steric Number 5
Methyl groups, considered to be less
electronegative than fluorine atoms, occupy the
equatorial sites to minimize repulsion between
the bonding electron pairs.