Chemistry 1A: Chapter 10
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Chapter 10: Chemical Bonding II: Molecular Shapes; VSEPR,
Valence Bond and Molecular Orbital Theories
Homework:
Read Chapter 10: Work out sample/practice exercises.
Suggested Chapter 10 Problems: 35, 41, 43, 45, 47, 51, 63, 65, 67, 71,
75, 77, 81, 83, 87, 89, 93
Check for the MasteringChemistry.com assignment and complete before due date
Molecular Shapes:
Properties of molecular substances depend on its 3D structure
Bonding neighbors, what is next to what (skeleton arrangement)
Type of bonding; polar, nonpolar, ionic
Shape and Polarity; overall do dipoles cancel or is there an overall dipole
moment
Limitations in Lewis Structures:
• Lewis theory predicts the number of electron regions (lone pair or any bond; single
double, triple), but does not determine actual bond angles.
• Lewis theory predicts trends in properties, but does not give good numerical
predictions of bond strength and bond length
• Lewis theory cannot write one correct structure for molecules where resonance is
important
• Lewis theory often does not correctly predict magnetic behavior of molecules.
Oxygen, O2, is paramagnetic, though the Lewis structure predicts it is diamagnetic
Valence Shell Electron Pair Repulsion (VSEPR) Theory: Three-dimensional
• Electron groups (all negatively charged) around the central atom are most stable
when they are as far apart as possible –valence shell electron pair repulsion theory.
Use all the information gained in the Lewis Dot Structure and convert it to a three
dimensional model to predict electronic and molecular shapes, angles, and polarity
of the molecule.
VSEPR Guidelines:
 Start with information from a Lewis Dot Structure
 Electronic and Molecular 3D shapes
 Bonds angles: When electron groups attach to different size atoms the
ideal bond angles are affected. Lone pairs (nonbonding) use more space.
 Polarity of whole substance (ionic, ion, nonpolar, polar molecule)
Chemistry 1A: Chapter 10
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Electronic and Molecular Geometry:
Count the electron regions. Electron regions will give an electronic shape
while the number of bonded versus nonbonded regions will give the
molecular shape.
# Electron
regions
2
3
4
5
6
Electronic
geometry
five basic
shapes
molecular
geometry
Linear
180˚
trigonal planar
120˚
tetrahedral
109.5˚
trig. bipyramidal
90˚, 120˚, 180˚
octahedral
90˚, 180˚
Linear
Trig planar,
bent
Tetrahedral,
Trig. pyramidal,
bent
trig.bipyramidal,
see saw,
T-shaped,
linear
octahedral,
square pyramidal,
square planar
Samples
Imperfect Geometry:
When electron groups attach to different size atoms the ideal bond
angles are affected
CH2O ideally should be trigonal planar with angles of 120° each.
In reality the angle between the smaller H atoms is smaller.
Lone pairs (nonbonding electrons)
use more space. Ideally four
regions should spread out to
angles of 109.5°. Notice how the
bond angles around the atoms are
forced closer together as the
unseen nonbonding electrons take
more space.
Website to try: ChemEdDL.org Click on molecules 360. This website shows the 3D
structure of many chemicals and allows you to rotate in three dimensions, showing
bonding, bond length, dipole arrows, dipole moment, etc.
Chemistry 1A: Chapter 10
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Writing 3D shapes on paper: May use lines and wedges.
Multiple Central Atoms:
Describe the shape around each central
atom separately.
atom
Polarity of the Molecule:
Polar: must have polar bonds (electronegativity difference between the
neighbor atoms with a measureable bond dipole moment) and an
unsymmetrical shape (lone pairs or varying atom neighbors)
Polarity affects properties: boiling points, solubilities (like dissolves like)
HCl and H2O are both polar
CO2 is nonpolar
Chemistry 1A: Chapter 10
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Valence Bond (VB) Theory: Three-dimensional
The Valence Bond theory is a quantum mechanical model that expands the previous
two theories to describe the electronic nature of covalent bonds.
• Valence bond theory applies principles of quantum mechanics to molecules
• A chemical bond between atoms occurs when atomic orbitals and hybridized
atomic orbitals interact with those in another atom to form a new molecular
orbital with two electrons.
 If orbitals align along the axis between the nuclei, sigma bonds which
directly overlap will form ( bonds). It is possible to rotate a sigma bond
 If orbitals align outside the axis, pi bonds form, which indirectly overlap
above and below ( bonds). Unable to rotate without breaking bonds. This
causes cis and trans structural isomers.
VB Guidelines:
 Use all the information from a Lewis Dot Structure
 Hybridizing some orbitals allow for more bonds and more stability
 Visualize orbital picture using atomic (s, p, d, f) and hybridized (sp, sp2,
sp3, sp3d, and sp3d2) orbitals
 Direct overlap orbitals, sigma () bonds
 Indirect overlap orbitals, pi () bonds
 All types of bonds have only one  bond. Double bonds have 1  and 1 
and triple bonds have 1  and 2  bonds
 Valence Bond (Bubble) Pictures draw the orbitals in balloon type pictures
 Delocalized  bonding occurs in substances with resonance
Chemistry 1A: Chapter 10
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Chemical bonds between atoms occur when atomic orbitals interact with those in
another atom to form a new molecular orbital with two electrons. Sigma bonds
(direct overlap) are stronger than pi bonds (indirect overlap).
Chemistry 1A: Chapter 10
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Double bond:
CH2O
C
3 sp2 hybridized orbitals and
1 p unhybridized orbital
H
1 s orbital on each
O
1 s unhybridized orbital and
3 p unhybridized orbitals
Triple bond:
C2H2
C
2 sp hybridized
orbitals and
2 p unhybridized
orbitals
H
1 s orbital on each
Limitations in Valence Bond Theory:
Valence Bond theory predicts bond strengths, bond lengths, and bond rigidity better
than Lewis theory.
• Other properties, such as the magnetic behavior of O2, of molecules are not
predicted well.
• VB theory views electrons as localized in overlapping atomic orbitals and it
doesn’t account for delocalization
Chemistry 1A: Chapter 10
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Molecular Orbital (MO) Theory:
The Molecular Orbital Theory is separate from the first three. This theory explains
the paramagnetic behavior found in O2 gas molecules.
• In MO theory, Schrödinger’s wave equation is applied to the molecule to
calculate a set of molecular orbitals
• Electrons and orbitals belong to the whole molecule – Delocalization
• A Bonding Molecular Orbital forms when wave functions combine
constructively, resulting in a molecular orbital with lower energy than the
original atomic orbitals. Most of the electron density is between the nuclei.
Lower energy-stabilizing
• The Antibonding* Molecular Orbital forms when wave functions combine
destructively, resulting in a molecular orbital with more energy than the
original atomic orbitals. Most of the electron density is outside the nuclei
creating nodes between nuclei. Higher energy-unstable
Sigma () 1s molecular orbitals (2s looks the same, but a bit bigger)
Sigma () px molecular orbitals
Chemistry 1A: Chapter 10
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Pi () py or pz molecular orbitals
MO Guidelines:
 Electrons belong to the molecule, not the individual atoms
 For this class, limit most of the discussion and examples to diatomic
species such as: H2, O2, CN-1, HF.
 Occasionally this gives a more accurate electronic structure than VB
 Combination of two atomic orbitals makes a molecular orbital
 Bonding orbitals are sigma or pi orbitals. Sigma orbitals directly overlap
and pi orbitals indirectly overlap
 Antibonding* sigma or pi orbitals create a node between the atoms with no
overlap
 Two atomic s orbitals combine to form a lower energy  bonding and a
higher energy * antibonding* orbital
 six atomic p orbitals combine to form lower energy bonding orbitals,
and 2 degenerate  orbitals and higher energy antibonding*
orbitals,and 2 degenerate  orbitals
 Predicts paramagnetic or diamagnetic behavior
 Predicts bond order
 Compares bond lengths and bond strengths
 For diatomic molecules with fewer than 15 total electrons like N2, energy
increases as follows: s, 1s*, 2s, 2s*, 2p,2p, 2p, 2p*, 2p*, 2p*
 For diatomic molecules with 15 or more total electrons like O2, energy
increases as follows: s, 1s*, 2s, 2s*, 2p,2p,2p, 2p*, 2p*, 2p*
Magnetic behavior of O2
Diatomic oxygen is attracted to
a magnetic field, indicating
paramagnetic behavior, so it
has unpaired electron(s)
Chemistry 1A: Chapter 10
Heteronuclear Diatomic Elements and Ions:
• The more electronegative an atom is, the
lower in energy are its orbitals
• Lower energy atomic orbitals contribute
more to the bonding MOs
• Higher energy atomic orbitals contribute
more to the antibonding MOs
• Nonbonding MOs remain localized on the
atom donating its atomic orbitals
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Chemistry 1A: Chapter 10
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Polyatomic Molecular Orbitals:
• Atomic orbitals of all the atoms in a molecule, even those with 3 or more atoms,
combine to make a set of molecular orbitals, delocalized over the entire molecule
• Predictions made using molecular orbital theory, (especially resonance molecules
and predicting magnetic properties), match the real molecule properties better than
either Lewis or Valence bond theories.
Ozone, O3:
MO theory predicts equivalent bond lengths due to
the delocalized electrons.
Chemistry 1A: Chapter 10
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Molecular Shapes, Handedness and Drugs:
The shapes of molecules can dramatically change its characteristics. Mirror images
have different biological properties due to the specific shapes of receptor sites in the
body. For a molecule to exhibit handedness it needs four different groups attached to
a carbon.
Identify the electronic and molecular geometries, angles, and VB hybridization
a)
h)
b)
i)
c)
j)
d)
k)
e)
l)
f)
g)
m)
Chemistry 1A: Chapter 10
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Fill in the following tables: First page follows octet and duet rules, second page has extended octets.
#of electron number of
electronic
molecular
bond angles
rough
an example
regions and bonded
geometry
geometry
3-D
molecule or ion
VB hybrid
atoms
name
name
sketch
H2
CO
any
1
linear
linear
O−−O
(180)
HF
N2
CN-1
CO2
3
3
sp2
120
bent
or angular
4
109.5
trigonal
pyramidal
H2O
Chemistry 1A: Chapter 10
#of electron
regions and
VB hybrid
number of
bonded
atoms
5
electronic
geometry
name
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molecular
geometry
name
bond angles
rough
3-D
sketch
an example
molecule or ion
trigonal
bipyramidal
see-saw
180
3
90
(120)
5
sp3d
2
6
octahedral
BrF5
6
sp3d2
square
planar
Chapter 10: Chemical Bonding II: Molecular Shapes. Valence Bond and Molecular Orbital Theories
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Examples:
1. The valence bond hybrid atomic orbitals sp3 are used by both C in CH4 and O in H2O. Yet, the
bond angles between atoms in H2O are less than in CH4. Explain.
2. Describe completely the main features of each of the following and explain what useful information
we gain from each.
a)
b)
c)
d)
Lewis Structures
Valence Shell Electron Pair Repulsion (VSEPR) theory
Valence Bond (VB) theory
Molecular Orbital (MO) theory
3. a) Draw all possible resonance Lewis structures for NO3-1. Include formal charges and the correct
angles.
b)
Draw the "realistic" hybrid resonance structure with appropriate angles that takes and
average of the Lewis structures in part a. Include formal charges (fractions) and bond
orders (fractions). Include nonbonding electrons on central atom but not on terminal
atoms.
c)
Sketch the valence bond (bubble) probability picture of one of the NO3-1 resonances.
Identify and label the hybridized orbitals. Identify sigma and pi bonds.
4. Draw and identify the cis and trans isomers for 1,2-dichloroethene, C2H2Cl2
5.
For each of the following:
B2,
Ne2, O2
a)
Give the molecular orbital (MO) energy diagram for each.
b)
Write the MO configurations for O2
1s)2
c)
Give the bond order of each B2,
Ne2, O2
d)
List the species in decreasing order of bond energy and stability
e)
Identify each as diamagnetic or paramagnetic?
f)
Using the bond order information, which is least expected to exist. Explain why.
g)
Which would have the shortest bond length? Explain.
Chapter 10: Chemical Bonding II: Molecular Shapes. Valence Bond and Molecular Orbital Theories
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6. Complete the following table for the indicated substances.
Electronegativities: Na = 0.9, N = 3.0, O = 3.5, F = 4.0, Cl = 3.0, Br = 2.8, I = 2.5
substance
SO2
C2H4O2
ICl5
NaBrO3
a) Draw the best
Lewis
structure(s),
resonances, and
structural isomers
if any with octet
b) Include formal
charges if they
are not zero
c) Indicate polar
bonds with dipole
arrows toward
the more
electronegative
name electronic
geometry around
central atom
give hybrid
orbital for center
name molecular
geometry around
central atom
show 3-D sketch
with atoms &
bonds in it
give all bond
angles
how many sigma
bonds? how
many pi bonds?
is it an ionic
compound, polar
or nonpolar
molecule or an
ion?
Draw the VB
hybrid resonance
(bubble) picture
Chapter 10: Chemical Bonding II: Molecular Shapes. Valence Bond and Molecular Orbital Theories
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7. Complete the following table for the indicated substances.
substance
SCN-1
I3-1
SF6
a)Draw the best
Lewis structure(s),
resonances, and
structural isomers
if any with octet
b) Include formal
charges if they are
not zero
c) Indicate polar
bonds with dipole
arrows toward the
more
electronegative
name electronic
geometry around
central atom
give hybrid orbital
for center
name molecular
geometry around
central atom
show 3-D sketch
with atoms &
bonds in it
give all bond
angles
how many sigma
bonds? how many
pi bonds?
is it an ionic
compound, polar
or nonpolar
molecule or an
ion?
Draw the VB
hybrid resonance
(bubble) picture
K2SO3
Answer questions
below for SO3-2