RESONANCE STRUCTURES
For certain molecules or molecular ions, two or more
EQUIVALENT Lewis structures can be drawn.
The equivalent structures have the same formal charge and
are equal in energy.
Example: CO32- , carbonate ion
-2
2
O
O
1
C
O
3
Three equivalent structures for the carbonate ion – all have
the same formal charge.
2
2
2
-1
0
O
-1
O
1
C
0
-1
O
-1
0
O
O
3
1
C
O
0
-1
-1
O
O
3
1
C
0
0
O
3
These equivalent structures are called RESONANCE
STRUCTURES.
The double bond electron pair is DELOCALIZED over all
three C-O bonds and hence all three C-O bonds are between
pure-single and pure-double bonds.
Note: This differs from the two structures for N3- which had
different bonding arrangements for the atoms, and hence not
equivalent.
The carbonate ion is NOT a mixture of the three structures,
nor does any one of the three exist independently.
The real structure is an average of the three structures, and
properties of these bonds are such that they are intermediate
between single and double bonds.
All three C-O bonds identical, the bond lengths are between a
single bond between C and O and a double bond between C
and O
Other examples of molecules that have more than one
equivalent resonance structures are NO3-, SO3
EXCEPTIONS TO THE OCTET RULE
Although the octet rule is very useful, it does not apply in all
cases. There are three classes of exceptions
1) Compounds with odd number of electrons
2) Compounds with a shortage of electrons, less
than the octet.
3) Compounds with an abundance of electrons,
exceeding the octet
1) Molecules with ODD number of valence electrons
An example is nitrogen monoxide, NO.
Molecules with an odd number of electrons cannot satisfy the
octet rule for all its atoms.
In this case the octet rule must be given up for one of the
atoms by leaving an unpaired lone electron.
ALL BONDING ELECTRONS MUST BE PAIRED.
NO is stable molecule with 11 valence electrons
Two Lewis structures can be drawn for NO
N=O
N=O
(i)
(ii)
In (i) N has an octet, but not O; in (ii) O has an octet but not N
Calculate the formal charge of each atom for (i) and (ii) to
determine which is the more favored structure.
-1
N=O
(i)
+1
0
N=O
0
(ii)
Structure (ii) is favored since the formal charges on both
atoms is zero; in (i) N is -1 and O is +1, with the more
electronegative atom, O, possessing a +1 formal charge
2) Electron Deficient Molecules
Example BF3
Lewis model predicts the following structure for BF3
F
F
-1
B
F
+1
However, experimental evidence suggests that the structure
of BF3 is:
F
F
0
B
F
0
F
F
F
B
-1
(i)
F
+1
F
B
F
0
0
(ii)
(i) assigns a +1 formal charge to a very electronegative
atom which is undesirable
(ii) is the observed structure, even though B does not have
an octet, the formal charges on all atoms are zero.
3) Valence-Shell Expansion
The elements in Groups IIIA to VII A in the third period and
beyond show a tendency to surround themselves with
more than 8 electrons.
It is possible for these elements through “valence shell
expansion” for the central atom to have more than 8
electrons
In cases like this, after accounting for bonding electrons,
assign lone pairs to the outer atoms to give them octets.
If any electrons still remain, assign them to the central
atoms as lone pairs.
Example: SF6
Number of valence electrons = 6 + (6x7) = 48
Total number of electrons needed for each atom to satisfy an
octet = 7x8 = 56
Number of shared electrons = 56-48 = 8
8 electrons, 4 bonds, is not enough bonds for SF6
Hence, assign one bond pair to each bond, which for SF6 is 6.
Then assign lone pairs to the outer atoms to give them an
octet. If any electrons remain, assign them to the central
atom.
F
F
F
S
F
F
F
F
F
F
S
F
Count all electrons
(6x6) + (6x2) = 48 electrons
Formal charge on all atoms: 0
F
F
POCl3 - phosphoryl chloride
O
0
Cl
-1
+1
P
O
Cl
0
0
0
Cl
0
0
P
0
Cl
Cl
(i)
(ii)
i) non zero formal charges
ii) valence shell expansion for P
Cl
0
Shapes of Molecules
Molecules have definite shape i.e. definite geometries.
For example, H2O is bent and CO2 is linear.
H
O
104.5o
H
O= C= O
180o
Lewis stuctures do not provide information on a molecule’s
geometry, i.e. shape.
VALENCE SHELL ELECTRON-PAIR REPULSION (VSEPR)
THEORY: predicts molecular shapes based on Lewis
stuctures
VSEPR derives from the realization that bonding and nonbonding electron pairs will locate themselves in space so as
to minimize repulsive interactions.
The geometry of the electron pairs, and the covalent bonds
that form, can be predicted using VSEPR
Molecular geometry depends on the total number of electron
pairs
To use the VSEPR theory to predict molecular geometry,
need to know how electron pairs can be arranged around a
central atom so as to minimize electron-electron repulsion.
An electron pairs arranges itself in three-dimensional space
such that it is equally far apart from all the other electron
pairs.
The resulting arrangement of the electron pairs in a molecule
depends on the NUMBER of electron pairs.
Note: For VSEPR, the electron pairs we refer to are the
valence electrons
For a central atom surrounded by TWO electron pairs, the
optimum geometry which minimizes repulsion between the
electrons pairs is LINEAR
180o
LINEAR
For a central atom surrounded by THREE electron pairs, the
optimum geometry which minimizes repulsion between the
electrons pairs is TRIGONAL PLANAR
120o
TRIGONAL PLANAR
For a central atom surrounded by FOUR electron pairs, the
optimum geometry which minimizes repulsion between the
electrons pairs is TETRAHEDRAL
109.47o
TETRAHEDRAL
For a central atom surrounded by FIVE electron pairs, the
optimum geometry which minimizes repulsion between the
electrons pairs TRIGONAL BIPYRAMIDAL
90o
120o
TRIGONAL BIPYRAMIDAL
For a central atom surrounded by SIX electron pairs, the
optimum geometry which minimizes repulsion between the
electrons pairs OCTAHEDRAL
90o
90o
OCTAHEDRAL
We have been talking about the location of the electron pairs
about the central atom.
This maybe different from molecular geometries i.e. the
structure or shape of a molecule
Water has a “BENT” structure with the H-O-H angle
measured to be 104.5o.
H
O
104.5o
H
If O is bonded to 2 H atoms, why is the structure of water not
linear?
MOLECULAR GEOMETRY
Molecular geometry is the arrangement of ATOMS about the
central atom.
If there are no non-bonding electrons, then the electron
geometry and the molecular geometry is the same.
If there are non-bonding electrons on the central atom then
the molecular geometry differs from the electron geometry
but is obtained from knowledge of the electron geometry
about the central atom.
To determine the molecular geometry about the central atom
of a molecule:
1) Write the best Lewis structure for the molecule or ion
2) Count the total number of ATOMS attached to the central
atom plus the number of NONBONDING ELECTRON PAIRS
on the central atom .
3) Using the VSEPR geometries establish the ELECTRON
geometry about the central atom.
4) Finally examine the placement of atoms and identify the
molecular geometry.
Step 1: Write the best Lewis structure for the molecule
Total number of valence electrons = 1+1+6 = 8
Number of electrons needed for each to complete an octet
(remember for H need two electrons) = 2 + 2 + 8 = 12
Shared electrons = 12- 8=4 => two bonds
Number of unpaired electrons = 4 => 2 lone pairs on O
Lewis structure of H2O is:
H
O H
2) Count the total number of atoms attached to the central
atom plus the number of non-bonding electron pairs on
the central atom
H
O H
For water,
Total number of electron pairs =
2 pairs of bonding electrons (bp) + 2 pairs of non-bonding
electrons (np) = 4 pairs
3) Then using the VSEPR geometries establish the electron
geometry about the central atom.
According to VSEPR theory: For a central atom surrounded
by four electron pairs, the optimum geometry which
minimizes repulsion between the electrons pairs is
tetrahedral
Arrange the two nonbonding electron pairs and the two H
atoms at the apices of the tetrahedron, with O at the center
O
H
H
4) Finally examine the placement of atoms and identify the
molecular geometry.
To determine the molecular geometry of water, look at the
arrangement of just the atoms (in this case, H,O,H)
O
H
H
Looking at the arrangement of the atoms in H2O, we see that
the molecule is bent.
O
H
104.5o
H
For a tetrahedral geometry, expect that the H-O-H angle is
109.5o
The measured H-O-H angle is 104.5o, slightly smaller than
predicted.
Non-bonding electron pairs are more diffuse than bonding
electron pairs, taking up more space.
Model
Draw the Lewis structure to determine the number of bonding
electron pairs and number of nonbonding pairs on the central
atom (a double or triple bond counts as a single electron pair
when using VSEPR theory).
Arrange the bonding and nonbonding electron pairs around
the central atom in a geometry predicted by VSEPR.
Position the atoms bonded to the central atom where the
bonding electrons are positioned.
Determine the molecular geometry based on the arrangement
of the atoms
Two electron pairs
Example: CO2
Lewis structure:
O=C=O
Number of electron pairs around C = 2 (count the double
bond as a single electron pair).
Both are bonding pairs
VSEPR predicts a linear molecular geometry for CO2
Model
Three electron pairs
F
BF3 – Lewis structure
F
B
F
0
0
The three electron pairs around the central B atom are all
bonding pairs.
According to VSEPR the three electron pairs around the
central atom are positioned at the vertices of an equilateral
triangle
In the case of BF3, since the three electron pairs are bonding
pairs, the three F atoms are positioned at the vertices of the
equilateral triangle.
Model
F
B
F
120o
F
trigonal planar
NO2- Lewis structure:
O=N-O
For VSEPR count two bonding electron pairs and one
nonbonding electron pair.
Hence the three electron pairs are situated at the vertices
of an equilateral triangle.
The two O atoms are situated at two of the vertices and the
nonbonding electron pair at the third.
The molecular geometry if NO2- is not trigonal planar, but
bent.
The O-N-O angle is less than 120o due to the presence of the
nonbonding pair on N
-1
N
O
<120o
bent
>120o
O
Four electron pairs
For molecules with four electron pairs, the electron pairs
surround the central atom in a TETRAHEDRAL geometry.
The molecular geometry depends on the number of
bonding and nonbonding electron pairs on the central
atom.
Examples:
H
H
C
H
H
methane, CH4
tetrahedral
H
N
H
H
ammonia, NH3
pyramidal
O
H
H
water, H2O
bent
Model
Five Electron Pairs
Electronic geometry is trigonal bipyramidal.
However, unlike the 2, 3 and 4 electron geometries all vertices
of a trigonal bipyramid are NOT equivalent.
There are two AXIAL positions and three EQUATORIAL
positions.
axial
equatorial
90o
120o
equatorial
equatorial
axial
90o
120o
Angle between
axial-central atom-axial = 180o
axial-central atom-equatorial = 90o
equatorial -central atom- equatorial = 120o
The positions of the nonbonding electron pairs relative to
each other and the bonding pairs have to account for the
non-equivalence of the axial and equatorial positions.
To position the nonbonding and bonding electrons:
nb-nb repulsion > nb-b > b-b
Hence, position nonbonding pairs first, then bonding pairs
accounting for the repulsion between the electron pairs.
PCl5
Cl
Cl
Cl P
Cl Cl
Cl
Cl
Cl
P
Cl
Cl
trigonal
bipyramidal
Model
SF4
F
F
S
F F
Position the nonbonding electron pair on S at the
equatorial position since this minimizes nb-b repulsions
F
F
S
F
F
“see-saw”
Angles in SF4 are different from the predicted 180o, 120o
and 90o due to the non-bonding electron pair
F
F
187o
S
102o
F
87o
F
Other examples
ClF3 - distorted T
XeF2 - linear
Model
Six electron pairs
Octahedral geometry
All vertices are the equivalent, hence bonding and
nonbonding electron pairs can be positioned at any vertex
OCTAHEDRAL
Molecular geometry determined by number and position at
atoms and nonbonding electron pairs
Model
Examples
F
F
F
S
F
F
F
F
Br
F
F
F
F
Xe
F
F
F
SF6
F
BrF5
XeF4
octahedral
square pyramidal
square planar
Molecular Properties
Bond Lengths
The bond distance is the distance between the two nuclei of
the two atoms bonded together.
Bond lengths depend on the elements and whether the bond
between the two elements is single, double or triple
Typically:
single bond length > double bond length > triple bond length
Bond Energies
If two atoms in a diatomic molecule are pulled far enough
apart the bond between the atoms breaks.
The energy needed to break the bond is called the bond
energy or bond dissociation energy.
Bond energies depend on the nature of the atoms bonded
together and whether the bond is single, double or triple.
Typically single bond energies < double bond energy < triple
bond energy
Dipole Moments
The bonding electron pair in HCl is not shared equally
between the H and Cl atoms.
This is because Cl is more electronegative than H, and tends
to attract the electron pair closer towards it.
Hence the electron distribution for the bonding pair of
electrons is not evenly distributed over both the H and Cl
atoms, but is more concentrated over the Cl atom.
The more electronegative atom develops a slight negative
charge and the more electropositive atom a slight positive
charge
This charge separation results in a DIPOLE MOMENT, and a
POLAR HCl bond.
The dipole moment is defined as the product of the
magnitude of the separated charges and the distance
between the centers of the charges.
The dipole moment is a vector
H
Cl
The larger the electronegative difference between the bonded
atoms, larger is the magnitude of the dipole moment and
hence more polar is the bond.
All heteronuclear diatomic molecules (A-B. where A≠B) are
polar molecules with the degree of polarity depending on the
electronegativity differences between the two atoms.
Heteronuclear diatomics have non-zero dipole moments
All homonuclear diatomic molecules (A2) are non-polar, with
zero dipole moments
Polyatomic molecules
Bonds between atoms of different elements will each have a
dipole moment associated with them.
The magnitude of the dipole moment of a bond depends on
the electronegativity differences between the bonded atoms.
However, whether the molecule as a whole has a dipole
moment depends on the geometry of the molecule.
The dipole moment is a vector and hence the dipole moment
of the molecule is the vector sum of the dipole moments of
each bond in the molecule.
Methane, CH4
H
H
C
H
H
The net dipole moment for CH4 is zero
Water, H2O
H
H
O
O
H
H
Model