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Chapter 9
Molecular Geometry:
Shape Determines Function
Chapter Outline
• 9.1 Molecular Shape
• 9.2 Valence-Shell Electron-Pair Repulsion
Theory (VSEPR)
• 9.3 Polar Bonds and Polar Molecules
• 9.4 Valence Bond Theory
• 9.5 Shape and Interactions with Large
Molecules
• 9.6 Chirality and Molecular Recognition
• 9.7 Molecular Orbital Theory
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Molecular Shape
• Chemical/physical properties are related
to molecular shape.
• Lewis structures
• Show atoms and bonds, but not spatial
orientations (3D).
• Molecular models
• Show orientations and bond angles; help us
understand physicochemical properties.
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Lewis Structures vs. Models
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Molecular Shape (cont.)
• Bond angle:
• Angle (in degrees) defined by lines joining
the centers of two atoms to the center of a
third atom to which they are covalently
bonded
• Not always predictable from Lewis
structures
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Chapter Outline
• 9.1 Molecular Shape
• 9.2 Valence-Shell Electron-Pair Repulsion
Theory (VSEPR)
• Central Atoms with No Lone Pairs
• Central Atoms with Lone Pairs
• 9.3 Polar Bonds and Polar Molecules
• 9.4 Valence Bond Theory
• 9.5 Shape and Interactions with Large
Molecules
• 9.6 Chirality and Molecular Recognition
• 9.7 Molecular Orbital Theory
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Valence-Shell Electron-Pair
Repulsion Theory (VSEPR)
• VSEPR Theory:
• A model predicting that the arrangement of valence
electron pairs around a central atom minimizes
repulsion to produce the lowest-energy orientation.
• Electron-pair geometry:
• Three-dimensional arrangement of bonding e– pairs
and lone pairs electrons about a central atom
• Molecular geometry:
• 3-dimensional arrangement of atoms in a molecule.
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VSEPR: Electron-Pair
Geometry
• To determine electron-pair geometry:
• Draw Lewis structure (see Chapter 8).
• From Lewis structure, determine steric
number (SN):
•
• Determine optimal spatial arrangement of
electron pairs (bonding + nonbonding) to
minimize repulsion.
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Molecular Geometry: Central
Atom with No Lone Pairs
• Molecular geometry = electron-pair
geometry
• Determine steric number (SN):
• SN = 2 (two atoms bonded to central atom)
• geometry  linear
• SN = 3 (three atoms bonded to central atom)
• geometry  trigonal planar
• SN = 4  tetrahedral
• SN = 5  trigonal bipyramidal
• SN = 6  octahedral
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Central Atoms with No Lone
Pairs
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Central Atoms with No Lone
Pairs (cont.)
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Geometric Forms
• Examples:
• CO2
BF3
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CCl4
PF5
SF6
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Practice: Molecular
Geometry (No Lone Pairs)
• Determine the molecular geometry of:
a) H2CO (C is central atom)
•
b) CH4
- Collect and Organize: We are given
molecular formulas and asked to predict
their molecular geometry.
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Practice: Molecular
Geometry (No Lone Pairs)
• Determine the molecular geometry of:
a) H2CO (C is central atom)
•
b) CH4
- Analyze: We can use the periodic table to determine the
number of valence electrons for each atom. From the
molecular formula and valence electrons, we can draw
Lewis structures. From the Lewis structures we can
determine the SN. From the SN we can predict the electronpair geometry. Since there are no lone pairs, the electronpair geometry is the same as the molecular geometry.
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Practice: Molecular
Geometry (No Lone Pairs)
• Determine the molecular geometry of:
a) H2CO (C is central atom)
•
b) CH4
-Solve: H2CO  C is the central atom, with single
bonds to each H atom and a double bond to O, SN
= 3. Molecular geometry = trigonal planar.
CH4  C is central atom, with single bonds to
each H atom, SN = 4. Molecular geometry =
tetrahedral.
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Practice: Molecular
Geometry (No Lone Pairs)
• Determine the molecular geometry of:
a) H2CO (C is central atom)
•
b) CH4
- Think About It: The molecular geometries are
consistent with VSEPR theory for 3 and 4
electron clouds. It is worth noting that the C
atom always has 4 bonds, but a double bond
counts as only one electron cloud, resulting in a
trigonal planar geometry.
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Central Atoms with Lone
Pairs
• Molecular geometry  electron-pair
geometry
• Replace bonding pair(s) with lone pair(s).
• Example: SO2 (SN = 3)
• Three electron pairs (2 bonding + 1 lone pair)
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Central Atoms w/ Lone Pairs
(cont.)
• Bond angles less than predicted
• Electron pair repulsion!
• Lone pair–lone pair = greatest repulsion.
• Lone pair–bonding pair
• Bonding pair–bonding pair = least repulsion.
• Multiple bonds > single bonds
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Molecular Geometry: SN =
4
Note: bond angles decrease as # of
lone pairs increases.
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Molecular Geometry: SN = 4
(cont.)
Two lone pairs = greater repulsion,
decreased bond angle.
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Molecular Geometry: SN =
5
Note: lone pairs occupy equatorial positions.
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Molecular Geometry: SN = 6
Note:
bond angles = 90
(geometries w/ more
than 2 lone pairs are
possible.)
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Practice: Molecular Geometry
•
What are the molecular geometries of
the ions: SCN– and NO2– ?
- Collect and Organize: We are given the
molecular formulas for two polyatomic
ions and asked to predict the molecular
geometries.
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Practice: Molecular Geometry
•
What are the molecular geometries of
the ions: SCN– and NO2– ?
-Analyze: We can use the periodic table to determine
the number of valence electrons for each atom. From
the molecular formula and valence electrons, we can
draw Lewis structures. From the Lewis structures, we
can determine the SN. From the SN, we can predict
the electron-pair geometry. Making note of the
number of bonding pairs and lone pairs, we can
identify the molecular geometry.
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Practice: Molecular Geometry
•
What are the molecular geometries of
the ions: SCN– and NO2– ?
Solve:
SCN–  As the least electronegative element and the one with
the greatest bonding capacity, C is the central atom. Although
there are several possible resonance structures, they all have C
with SN = 2 and no lone pairs. Molecular geometry = linear.
NO2–  With N as the central atom, the Lewis structure has N
with SN = 3 (two bonding pairs and one lone pair). Again,
although there are two possible resonance structures, they both
have the same SN value. Molecular geometry = bent, with bond
angle <120 due to the extra repulsive energy of the lone pair
on N.
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Practice: Molecular Geometry
•
What are the molecular geometries of
the ions: SCN– and NO2– ?
-Think About It: It is worth noting that both
molecular structures have two bonding
electron clouds, but different molecular
geometries due to the differences in steric
number.
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Chapter Outline
• 9.1 Molecular Shape
• 9.2 Valence-Shell Electron-Pair Repulsion
Theory (VSEPR)
• 9.3 Polar Bonds and Polar Molecules
• What Makes a Molecule Polar?
• Dipole Moments
• 9.4 Valence Bond Theory
• 9.5 Shape and Interactions with Large
Molecules
• 9.6 Chirality and Molecular Recognition
• 9.7 Molecular Orbital Theory
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Polar Bonds and Polar
Molecules
• Requirements for polar molecule:
• 1. Polar bonds (i.e. covalent bond between
atoms with ΔEN).
• 2. Nonuniform distribution of polar bonds.
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Polar Molecules (cont.)
• Bond dipole:
• Separation of
electrical charge
created when
atoms with different
EN form a covalent
bond.
• Polar molecule:
• Vectors of bond
dipoles whose sum
> zero.
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Measuring Polarity
• Dipole moment () – a quantitative
expression of the polarity of a molecule.
• Units = debyes (D); 1 D = 3.34 × 10–30 C·m )
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Chapter Outline
• 9.1 Molecular Shape
• 9.2 Valence-Shell Electron-Pair Repulsion
Theory (VSEPR)
• 9.3 Polar Bonds and Polar Molecules
• 9.4 Valence Bond Theory
• Bonds from Orbital Overlap
• Hybridization (sp3, sp2, sp, sp3d, sp3d2)
• 9.5 Shape and Interactions with Large
Molecules
• 9.6 Chirality and Molecular Recognition
• 9.7 Molecular Orbital Theory
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Atomic Orbitals and Bonds
•
•
Valence Bond Theory (Linus Pauling)
•
Quantum mechanics-based model
•
Covalent bond = overlap of half-filled
orbitals
Sigma () bond:
•
Covalent bond in which the highest
electron density lies between the two
atoms along the bond axis.
Overlap of 1s orbitals
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Hybridization: sp3 Orbitals
• Hybridization – the mixing of atomic
orbitals to generate new sets of orbitals
that form covalent bonds with other
atoms.
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Tetrahedral Sigma Bonds
• Tetrahedral orientation of sp3 hybridized
orbitals = tetrahedral molecular geometry
Overlap of 1s with
sp3 orbitals 
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Other sp3 Hybrid Examples
Note: lone
pairs (nonbonding)
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Trigonal Planar: sp2
Hybridization
Unhybridized p orbitals form double bonds.
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sp2 Hybridization (cont.)
• pi () bond: a covalent bond in which
electron density is greatest around—not
along—the bonding axis.
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Linear: sp Hybridization
Form triple
bond (one 
and two π
bonds).
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Trigonal Bipyramidal: sp3d
Hybridization
Formed by mixing one s, one d, and three p orbitals.
Example: PF5 – five sigma bonds
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Octahedral: sp3d2 Hybridization
Formed by mixing one s, two d, and three p orbitals.
Example: SF6 – six sigma bonds
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Practice: Hybrid Orbitals
What are the hybridizations of the central
atoms of the ions: SCN– and NO2– ?
- Collect and Organize: Note that these are the
same molecules for which we determined
molecular geometry back in section 9.2. Using
Lewis structures and VSEPR, we determined
that the electronic geometry around the central
carbon in SCN– was linear (SN = 2), and the
electronic geometry around the central N atom
in NO2– was trigonal planar (SN = 3).
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Practice: Hybrid Orbitals
What are the hybridizations of the central
atoms of the ions: SCN– and NO2– ?
- Analyze: From the steric number of the
central atoms and valence bond theory, we
can determine the hybridization around the
central atom based on electron-pair
geometry.
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Practice: Hybrid Orbitals
What are the hybridizations of the central
atoms of the ions: SCN– and NO2– ?
Solve:
SCN– is linear (SN = 2), so the
hybridization must be sp.
NO2– is trigonal planar (SN=3), so the
hybridization must be sp2.
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Chapter Outline
• 9.1 Molecular Shape
• 9.2 Valence-Shell Electron-Pair Repulsion
Theory (VSEPR)
• 9.3 Polar Bonds and Polar Molecules
• 9.4 Valence Bond Theory
• 9.5 Shape and Interactions with Large
Molecules
• Shape and Molecular Recognition
• Delocalized Electrons
• 9.6 Chirality and Molecular Recognition
• 9.7 Molecular Orbital Theory
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Molecular Recognition
• Molecular recognition:
• The process by which molecules interact with
other molecules in living tissues to produce a
biological effect.
• Example: ethylene (ripening agent)
All atoms in
same plane.
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Delocalization of Electrons
• Delocalization:
spreading of
electrons in
alternating single
and double
bonds over three
or more atoms in
a molecule
a)
b)
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Aromatic Compounds
• Aromatic compound:
• A cyclic, planar compound with delocalized 
electrons above and below the plane of the molecule.
• Example: polycyclic aromatic hydrocarbons (PAH)
• Planar shape may allow intercalation in DNA
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Chapter Outline
• 9.1 Molecular Shape
• 9.2 Valence-Shell Electron-Pair Repulsion
Theory (VSEPR)
• 9.3 Polar Bonds and Polar Molecules
• 9.4 Valence Bond Theory
• 9.5 Shape and Interactions with Large
Molecules
• 9.6 Chirality and Molecular Recognition
• Optical Isomerism
• 9.7 Molecular Orbital Theory
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Isomerism
• Isomers – compounds with same
molecular formula but with atoms
arranged differently in three-dimensional
space.
Optical isomers
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Optical Isomers = Chirality
• Chirality – property of a molecule that is
not superimposable on its mirror image.
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Chapter Outline
• 9.1 Molecular Shape
• 9.2 Valence-Shell Electron-Pair Repulsion Theory
(VSEPR)
• 9.3 Polar Bonds and Polar Molecules
• 9.4 Valence Bond Theory
• 9.5 Shape and Interactions with Large Molecules
• 9.6 Chirality and Molecular Recognition
• 9.7 Molecular Orbital Theory
•
•
•
•
Molecular Orbitals of Hydrogen and Helium
MOs of Homonuclear Diatomic Molecules
MOs of Heteronuclear Diatomic Molecules
MOs of N2+ and Spectra of Auroras
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Molecular Orbital (MO) Theory
• Based on mixing of atomic orbitals of
similar shapes and energies to form
molecular orbitals (MOs) that belong to
the molecule as a whole
• The number of MOs formed is equal to the
number of atomic orbitals combined.
• MOs represent discrete energy states;
orbitals spread out over entire molecule.
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Types of Molecular Orbitals
• Bonding orbitals:
• Region of increased electron density between
nuclear centers that hold atoms together
• Are lower in energy (more stable) than atomic
orbitals from which they are formed
• Antibonding orbitals:
• Region of electron density that destabilize
the molecule because they do not increase
electron density between nuclear centers
• Less stable than atomic orbitals from which
they are formed
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Molecular Orbital Diagrams
• MO diagram:
• Energy-level diagram for molecular showing
the relative energies and electron occupancy
of the MOs for a molecule.
• Sigma (σ) bond:
• Covalent bond with the highest electron
density along the bond axis.
• Pi (π) bond:
• Bond formed by mixing of atomic orbitals not
oriented along the bonding axis in a
molecule.
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Molecular Orbital Diagram: H2
The two 1s orbitals mix to yield two sigma
MOs (1 bonding/1 antibonding).
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Bond Order and Stability
Bond order = 1/2 (# bonding e– – # antibonding e–)
Bond order in H2 – = ½
(Stable)
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Bond order in He2 = 0
(Not stable)
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MO Guidelines
1. The total # of MOs = the # of AOs orbitals
mixed.
2. Orbitals with similar energy/shape mix more
effectively than do those of different
energy/shape.
3. Orbitals of different n (different sizes/energies)
result in less effective mixing.
4. A MO can accommodate two electrons.
5. Electrons fill MO diagrams according to
Hund’s rule.
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MO Scheme for O2
• Electron configuration
for O2:
(σ2s)2(σ2s*)2(σ2p)2(π2p)4(π2p
*)2
• Bond order = ½ (8 – 4) = 2
• O2 has two bonds
• O2 has two unpaired
electrons in π2p*
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MO Scheme for N2
• Electron configuration for
N2: (σ2s)2(σ2s*)2(σ2p)2(π2p)4
• Bond order = ½ (8 – 2)
=3
 N2 has three bonds.
 N2 has no unpaired
electrons.
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Paramagnetic vs. Diamagnetic
• Paramagnetism:
• Atoms or molecules having unpaired
electrons are attracted to magnetic fields.
• Example: O2
• Diamagnetism:
• Atoms or molecules having all paired
electrons are repelled by magnetic fields.
• Example: N2
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MO for NO
• Zeff alters MO
diagram; atomic
orbitals for O are
lower in energy.
• Odd electron in π*2p ,
closer in energy to the
2p atomic orbitals of N
atom.
• Bond order = ½ (8 – 3) =
2.5
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MO for N2, N2+: Emission
Spectra
Crimson red
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Blue-violet
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