Molecular Geometry and
Chemical Bonding Theory
Page 232 # 2 – 5
Page 235 # 10, 12
Page 238 # 18, 20 - 24
What is the “shape” of the
molecule?
• How are valence electrons in a molecule
distributed among the orbitals?
• What are the shapes of these orbitals?
• What order are they occupied?
Molecular Geometry
VSEPR Theory
Valence Shell Electron Pair Repulsion Theory
• draw Lewis electron dot structure
• count the number of bonding electron pairs
about central atom (double and triple bonds
count as one pair for shape prediction)
• count the number of lone pairs of electrons
• match electron pair information to shapes
Shape: PCl3
Lewis Electron Dot Structure
.. .. ..
:Cl
.. : P
.. : Cl:
..
:Cl:
..
3 bond pairs
1 lone pair
=>
AB3E
trigonal pyramidal shape
Shape: IF5
Lewis Electron Dot Structure
.. ..
:F::F:
.. .. ..
: I :F:
.. .. ..
: ..F : :F.. :
5 bonds pairs
1 lone pair
=>
Square
pyramidal
shaped
AB
E
5
Shape: IF4-1
Lewis Electron Dot Structure
.. ..
-1
:F::F:
.. .. ..
: I :F:
.. .. ..
: ..F :
4 bond pairs
2 lone pairs
=> AB4E2
square planar shape
Shape: SO3
Lewis Electron Dot Structure
..
:O:S::O:
.. .. ..
: ..O :
3 bond pairs
0 lone pairs
=> AB3
trigonal planar shape
What would be expected to be the shape of
chloroform, CHCl3?
“see-saw”
square planar
tetrahedral
Central Themes of Valence Bond Theory
3) Hybridization of atomic orbitals. To explain the bonding in simple
diatomic molecules such as HF it is sufficient to propose the direct
overlap of the s and p orbitals of isolated ground state atoms. In cases
such as methane CH4 where 4 hydrogen atoms are bonded to a central
carbon atom it is impossible to obtain the correct bond angles.
Pauling proposed that the valence atomic orbitals in the molecule are
different from those in the isolated atoms.We call this Hybridization!
Hybrid Orbitals
http://www.colby.edu/chemistry/OChem/DEMOS/Orbitals.html
Hybrid Orbital Types - Periodic Groups
Hybrid Orbital Types
Groups in the Periodic Table Associated
SP
Group IIA
Alkaline Earth Elements
SP2
Group IIIA
Boron Family
SP3
Group IVA
Carbon Family**
SP3d
Group VA
Nitrogen Family
SP3d2
Group VIA
Oxygen Family
** The exception is carbon which can have: SP, SP 2, SP 3 hybrid orbitals
WHY DOESN’T THE ATOMIC ORBITAL APPROACH WORK ?
These orbitals are for the
atom - we can’t expect
that they are suitable for
the molecule.
2s
2px,2py,2pz
atomic
orbitals
During bonding ….
new orbitals form that
are more suitable for
making bonds.
After bonding (overlap)
we get a totally
new solution for the
new molecule.
sp,sp
2py,2pz
LCAO
hybrid atomic
orbitals
s, p, p ,n
overlap
molecular
orbitals
HYPOTHETICAL BONDING PROCESS
NOTE. Formally LCAO theory and Molecular Orbital theory are
two completely different approaches. You do not need to use
hybid orbitals to derive the molecular orbitals, combinations
of any type of function will do. Nevertheless, the abstraction
presented above is quite useful, as we will see quite soon.
HYBRID ORBITALS
PRECURSORS TO THE FINAL RESULT
We will concentrate first on hybrid orbitals.
ADVANTAGE
HYBRID ORBITALS allow us to correctly
predict the final shape of the molecule.
DISADVANTAGE
Hybrid orbitals are an abstraction and are
not the final MOLECULAR ORBITALS that
result after bonds form.
The types of hybrids common to organic
molecules follow ……
sp
LINEAR HYBRIDIZATION
FORMATION OF LINEAR HYBRID ORBITALS
2 pair
unused 2p
orbitals
unused
2p
y
2s
sp(1)
C
2 orbitals
filled
hybridization
x
sp(2)
2p
(1)
linear
(2)
sp hybrid orbitals
z
2p
SP HYBRID ORBITAL
sp
more density in
the bonding lobe
than in an sp2
orbital
smaller tail
than in an
sp2 orbital
Courtesy of
Professor George Gerhold
The sp Hybrid Orbitals in Gaseous BeCl2
Fig. 11.2 A&B
Fig. 11.2 C&D
2
sp
TRIGONAL PLANAR HYBRIDIZATION
3 pair in the valence shell
and either an incomplete
octet or a double bond
FORMATION OF TRIGONAL PLANAR HYBRID ORBITALS
3 pair
unused 2p
orbital
unused
sp2(2)
2p
2s
B
2p
3 orbitals
filled
hybridization
x
120o
sp2(1)
trigonal
planar
2p
sp2(3)
z
(1)
(2)
(3)
sp2 hybrid orbitals
SP2 HYBRID ORBITAL
cusp
sp2
more density in
the bonding lobe
than in an sp3
orbital
smaller tail
than in an
sp3 orbital
Courtesy of
Professor George Gerhold
Fig. 11.3
3
sp
TETRAHEDRAL HYBRIDIZATION
4 pair in the valence shell
(no double or triple bonds)
FORMATION OF TETRAHEDRAL HYBRID ORBITALS
New orbitals point
to the corners of a
tetrahedron.
4 pair
sp3(1)
2p
109o28’
2s
O
FILLED VALENCE
SHELL
hybridization
occurs when orbitals
are full and have
finished bonding
sp3(3)
sp3(4)
sp3(2)
tetrahedral geometry
(cartoon)
(1)
(2)
(3)
(4)
sp3 hybrid orbitals
FORMATION OF
SP3 HYBRID ORBITALS
2pz
(1)
(2)
2p
(3)
2s
2s
X
2py
These orbital shapes are
cartoons - actual shapes
are shown on the next
slide.
2px
sp3 hybridized
unhybridized
atom
[animation]
(4)
SP3 HYBRID ORBITAL
… and its cartoon
( cross section )
The hybrid orbital
has more density
in the bonding lobe
than a p orbital and
forms stronger bonds.
sp3
To avoid confusion the
back lobe is omitted
from the cartoons,
already shown, and the
front lobe is elongated
to show its direction.
The shape shown
is calculated from
quantum theory.
Courtesy of
Professor George Gerhold
omitted
ORIGIN OF THE SP3 DESIGNATION
add together, divide in four
hybridization
2s
2p
(1)
(2)
(3)
(4)
sp3 hybrid orbitals
each new orbital is
1/4 s + 3/4 p (25% s, 75% p)
S1P3 = SP3
( 1+3 ) = 4 parts total
ORIGIN OF THE SP3 ORBITAL SHAPE
2s orbital
sp3 hybrid
orbital
2p orbital
+
x
--
++
RECALL:
signs are mathematical
coordinates, not
electronic charge
HYBRIDIZATION
[animation]
Fig. 11.4
Bonding in Water
The sp3 Hybrid Orbitals in NH3 and H2O
Fig. 11.5
COMPARISON OF THE HYBRIDS
COMPARISON OF SPx HYBRID ORBITALS
bigger
“tail”
“cusp”
more “p” character
sp3
sp2
sp
more “s” character
Orbital plots courtesy of
Professor George Gerhold
more electron
density in the
bonding lobe
sp3
more
p-like
HYBRID ORBITALS
sp2
sp
COMPARISONS OF
BONDING DISTANCE
SIZE OF CUSP
SIZE OF TAIL
more
s-like
Orbital plots courtesy of
Professor George Gerhold
makes shorter, stronger bonds
The sp3d Hybrid Orbitals in PCl5
Fig. 11.6
The sp3d2 Hybrid Orbitals in SF6
Sulfur Hexafluoride -Fig. 11.7
SF6
Fig. 11.8
WHY DO HYBRIDS FORM ?
WHY DO HYBRIDS FORM?
1. Electron pair repulsions are minimized (= lower energy)
2. Stronger bonds (= lower energy) are formed
3. Hybrids have better directionality for forming bonds
4. Since promotion usually occurs, hybrids allow
more bonds to form (= lower energy)
CONSTRUCTION BLOCKS
THE HYBRIDS ARE “MOLECULAR LEGOS”
EACH IS USED IN A SPECIFIC BONDING SITUATION
HYBRID CONSTRUCTION BLOCKS
4 PAIR
TETRAHEDRAL
sp3
3 PAIR
TRIGONAL
PLANAR
sp2
X
sp
Z
120o
Y
X
Y
X
LINEAR
Y
109o28’
X
2 PAIR
Y
Y
180o
Z
Z
Z
Postulating the Hybrid Orbitals in a Molecule
Problem: Describe how mixing of atomic orbitals on the central atoms
leads to the hybrid orbitals in the following:
a) Methyl amine, CH3NH2
b) Xenon tetrafluoride, XeF4
Plan: From the Lewis structure and molecular shape, we know the
number and arrangement of electron groups around the central atoms,
from which we postulate the type of hybrid orbitals involved. Then we
write the partial orbital diagram for each central atom before and after
the orbitals are hybridized.
Postulating the Hybrid Orbitals in a Molecule
Problem: Describe how mixing of atomic orbitals on the central atoms
leads to the hybrid orbitals in the following:
a) Methyl amine, CH3NH2
b) Xenon tetrafluoride, XeF4
Plan: From the Lewis structure and molecular shape, we know the
number and arrangement of electron groups around the central atoms,
from which we postulate the type of hybrid orbitals involved. Then we
write the partial orbital diagram for each central atom before and after
the orbitals are hybridized.
Solution:
a) For CH3NH2: The shape is tetrahedral around the C and N atoms.
Therefore, each central atom is sp3 hybridized. The carbon atom has
four half-filled sp3 orbitals:
2s
2p
Isolated Carbon Atom
sp3
Hybridized Carbon Atom
The N atom has three half-filled sp3 orbitals and one filled with a
lone pair.
2s
2p
sp3
..
H
C
N
H
H
H
H
b) The Xenon atom has filled 5 s and 5 p orbitals with the 5 d orbitals
empty.
Isolated Xe atom
5s
5p
5d
Hybridized Xe atom:
sp3d2
5d
b) continued:For XeF4. for Xenon, normally it has a full octet of
electrons,which would mean an octahedral geometry, so to make the
compound, two pairs must be broken up, and bonds made to the four
fluorine atoms. If the two lone pairs are on the equatorial positions,
they will be at 900 to each other, whereas if the two polar positions are
chosen, the two electron groups will be 1800 from each other. Thereby
minimizing the repulsion between the two electron groups.
F
..
1800
Xe
F
..
F
F
F
Xe
F
F
F
Square planar
What would be the name of the hybrid
orbitals created by joining 1s type, 3 p type
and 2 d type orbitals?
spd
sp3d2
s3p2d
Water, H2O, and methane, CH4, have very similar molar
masses yet the boiling point for water is approximately
200oC higher than that for methane. The higher boiling
point is due to:
hydrogen bonding in water
polarity of water molecules
both of the above
none of the above
Carbon monoxide, CO and carbon dioxide,
CO2, are linear molecules. Which one is the
polar molecule:
CO
CO2
SIGMA BONDS
BOND LENGTHS AND STRENGTHS
Shorter = Stronger = (more) S-character
BOND LENGTHS AND OVERLAP
Overlap leads to the bond.
In general, the more overlap that
can be established, the stronger
the bond that is formed.
sp3
Shorter bonds
establish more
overlap and are
stronger bonds.
sp
Orbitals that
have more
s-character
place more (%)
electron density
in the bonding
lobe and form
stronger bonds.
BOND STRENGTHS - C-H SINGLE BONDS
C-H
bond energy molecule
bond
type
length
per mole
measured
Kcal ( KJ )
2p-1s
1.12 A
80
(335)
unhybridized
intermediate
C-H
sp3-1s
1.10 A
101
(422)
CH3CH3
C-H
sp2-1s
1.08 A
106
(444)
CH2=CH2
C-H
sp-1s
1.06 A
121
(506)
=
HC=CH
[CH2]
increasing
s-character
NOTE: more S = Shorter = Stronger
Hybridization makes stronger bonds.
BOND STRENGTHS - C-C SINGLE BONDS
C-C
bond
type
length
bond energy
per mole
molecule
measured
Kcal ( KJ )
-C-C-
sp3-sp3
1.54 A
83
(347)
CH3-CH3
-C-C=
sp3-sp2
1.51 A
109
(456)
CH3-CH=CH2
1.46 A
-
-
-C-C
sp3-sp
increasing
s-character
CH3-C CH
Typical sigma bonds have energies
in the range 80-120 kcal/mole
Sigma Bond s
Bond with the greatest electron density on a line
connecting the atomic nuclei
s-s type
p-p type
s-p type
s-sp3 type
s-sp2 type
s-sp type
p-sp3 type
p-sp2 type
p-sp type
sp3 - sp3 type
sp2-sp2 type
sp - sp type
etc., including all combinations of s, p, d, and
hybrid orbitals
Fig. 11.9
DIFFERENCES BETWEEN
SIGMA AND PI BONDS
PI (p) BONDS
In a multiple bond, the first bond is a sigma s bond
and the second and third bonds are pi p bonds.
p
C C
s
p p
C C
s
Pi bonds are formed differently than sigma bonds.
SYMMETRY DIFFERENCES IN SIGMA AND PI BONDS
SIGMA BONDS
symmetric to
rotation about
internuclear
axis
s 1s-2p
END-TO-END
OVERLAP
PI BONDS
SIDE-TO-SIDE
OVERLAP
s 2p-2p
not symmetric
p 2p-2p
ONLY ONE SIGMA BOND CAN FORM
You can’t form more than one sigma bond along the
internuclear axis.
When you try to form more than
one sigma bond, you get bent bonds.
X
X
Bent bonds are not
as strong as pi bonds.
Pi Bond p
Bond with the greatest electron density
above and below a line connecting the
atomic nuclei
p-p type
Bonding in Ethylene
Fig. 11.10
Restricted Rotation of
p-Bonded Molecules
A) Cis - 1,2 dichloroethylene
Fig. 11.12
B) trans - 1,2 dichloroethylene
Bonding in Acetylene
Fig. 11.11
TYPES OF ELECTRONS IN A MOLECULE
SYMBOLS FOR
BONDED ELECTRONS AND
NON-BONDED ELECTRONS
STANDARD SYMBOLS
Types of Bonds: s and p
..
s
H3C
O
n
:
p
CH3
Non-Bonded Pairs : n
PI BONDS
BOND LENGTHS
AND STRENGTHS
PI BONDS ARE WEAKER THAN SIGMA BONDS
BOND STRENGTHS - MULTIPLE BONDS
CC
bond
bond
type
bond
length
bond energy
per mole
Kcal
molecule
measured
(KJ)
C-C
sp3-sp3
1.54 A
83
(347)
CH3CH3
C=C
sp2-sp2
1.34 A
146
(611)
CH2=CH2
=
C=C
sp - sp
1.20 A
200 (837)
increasing
s-character
shorter
=
HC=CH
Typical pi bonds have a bond
energy of about 50-60 kcal/mole
The shorter sigma bond that
results is also stronger.
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy
( Aufbau principle )
2) An orbital has a maximum capacity of two electrons with
opposite spins ( Pauli exclusion principle )
3) Orbitals of equal energy are half filled, with spins parallel,
before any is filled ( Hund’s rule )
Fig. 11.15
Molecular Orbital Diagram
Molecular Orbital Diagram
Fig. 11.22
Fig. 11.23
Fig. 11.24
Molecular Orbital Diagram
Molecular Orbital Diagram
Dr. S. M. Condren