Molecular Geometry and
Bonding Theories
Chapter 9
In this chapter
• Molecular shapes
• Valence Shell Electron Pair Repulsion model.
• Molecular shapes and polarity.
• Valence bond theory.
• Hybridization.
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Molecular Shapes
• Lewis structures give atomic connectivity: they tell us
which atoms are physically connected to which.
• The shape of a molecule is determined by its bond
angles.
• Consider CCl4: experimentally we find all Cl-C-Cl
bond angles are 109.5°.
– Therefore, the molecule cannot be planar.
– All Cl atoms are located at the vertices of a tetrahedron
with the C at its center.
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Molecular Shapes
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Molecular Shapes
• The Valence Shell Electron Pair Repulsion (VSEPR)
theory proposed by Gillespie and Nyholm is a
reliable method for predicting shapes of covalent
molecules and polyatomic ions.
• The VSEPR theory is based on the idea that bond
and lone pair electrons in the valence shell of an
element repel each other and move as far apart as
possible.
• VSEPR is extremely successful in predicting the
structures of molecules and ions of the main group
elements, but is generally useless for transition
metal containing compounds.
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There are five fundamental geometries for molecular shape:
AB2
AB3
AB5
AB4
AB6
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Predicting the Shape
• Predict the shape of CH2Cl2
• CH2Cl2 has one central atom and 4 other atoms
around it. Draw the Lewis structure.
• Referring to the possible arrangements we can
have for such a molecule the shape can be
predicted.
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Molecular Shapes
• When considering the geometry about the central
atom, we consider all electrons (lone pairs and
bonding pairs).
• When naming the molecular geometry, we focus
only on the positions of the atoms.
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Molecular Shapes
• To determine the shape of a molecule, need to
distinguish between lone pairs (or non-bonding
pairs, those not in a bond) of electrons and bonding
pairs (those found between two atoms).
• The electron domain geometry is defined by the
positions in 3D space of ALL electron pairs (bonding
+ non-bonding). This becomes especially important
when the central atom has lone pairs.
• The electrons adopt an arrangement in space to
minimize e--e- repulsion.
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Lone pairs on Central Atom
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4
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Lone Pairs & Bond Angles
• The electron pair geometry is decided only by
looking at the positions of the electrons.
• The molecular geometry is decided only by
the positions of atoms, lone pairs are ignored
in the molecular geometry.
• Lone pairs are considered to be “fat”, while
bonding pairs are considered “skinny”.
• The increased volume of lone pairs makes
bond pairs squeeze together, reducing bond
angles.
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Lone Pairs & Bond Angles
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5
Lone Pairs & Bond Angles
• Experimentally, the H-X-H bond angle decreases on
moving from C to N to O:
H
H C H
H
109.5O
H N H
H
107O
O
H
H
104.5O
• Since the electrons in a bonding pair are attracted by
two nuclei they do not repel as much as lone pairs.
• The bond angle decreases as the number of lone pairs
on the central atom increases.
• Typically the relative strength of repulsion is:
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Multiple Bonds & Bond Angles
• Similarly, electrons in multiple bonds repel more
than electrons in single bonds.
• A multiple bond is counted as one electron domain.
• The Lewis structure of COCl2 (carbonyl chloride)
indicates the molecular and the electron pair
geometry to be trigonal-planar.
Cl
111.4o
Cl
C O
124.3o
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Molecules with Expanded Valence
Shells
• Atoms that have expanded octets have AB5 (trigonal
bipyramidal) or AB6 (octahedral) electron pair
geometries.
• For trigonal bipyramidal structures there is a plane
containing three electrons pairs. The fourth and fifth
electron pairs are located above and below this
plane.
• For octahedral structures, there is a plane containing
four electron pairs. Similarly, the fifth and sixth
electron pairs are located above and below this
plane.
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6
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Molecules with Expanded Valence
Shells
• To minimize e--e- repulsion, lone pairs are always
placed in equatorial positions.
• In an octahedron, all
positions are considered
to be equivalent and
there are no axial or
equatorial positions.
• If the molecule has any
lone pairs they are placed
above or below the plane.
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Shapes of Large Molecules
• In acetic acid, CH3COOH, there are three
central atoms.
• We assign the geometry about each central
atom separately.
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Determining Molecular Shapes
• Draw the Lewis structure.
• Determine the total number of bonding pairs
(count any multiple bonds as 1 pair) around
the central atom.
• Pick the appropriate electron pair geometry
and then choose the molecular shape that
matches the total # of single-bond pairs and
lone pairs.
• Predict the bond angles, remembering that the
lone pairs occupy more volume than bonding
pairs.
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Molecular Shape & Polarity
• Difference in electronegativity between two
atoms makes the bond between them polar.
• It is possible for a molecule to contain polar
bonds, but not be polar.
• For example, the bond dipoles in CO2 cancel
each other because CO2 is linear.
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8
Molecular Shape & Polarity
• In water, the molecule is not linear and the
bond dipoles do not cancel each other.
• Therefore, water is a polar molecule.
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Molecular Shape and Molecular
Polarity
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Questions not answered by VSEPR
• Why does a bond form?
• How do we account for molecular shape in
terms of quantum mechanics, meaning in
terms of orbital shapes?
• What are the orbitals that are involved in
bonding?
• Valence Bond Theory proposed by Linus
Pauling provides these answers.
• Provides a qualitative visual picture of
molecular structure and bonding.
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Valence Bond (VB) Theory
• Orbitals overlap to form a bond between
atoms.
• Two electrons of opposite spins can be
accommodated in the overlapping orbitals.
Usually one electron is supplied by each of the
2 bonded atoms.
• Overlapping orbitals result in a high probability
of the 2 electrons being located in the region
that is influenced by both nuclei. Both
electrons are simultaneously attracted by both
nuclei.
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Orbital Overlap
• Orbital overlap produces a single bond that is known
as a sigma (σ) bond.
• The electron density of the sigma bond is greatest
along its axis.
The 2s and 2p electrons that are not involved in the
bond are the lone pairs on the F atoms.
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Hybrid Orbitals
• Consider the CH4 molecule, which has a tetrahedral shape.
• C has the electron configuration 1s22s22p2 and H has 1s1.
Carbon has only 2 unpaired electrons.
• The 2p orbitals in C are at right angles to each other, the 1s
orbital in H is spherical.
• According to Pauling’s theory of orbital hybridization, a new
set of orbitals called hybrid orbitals can be created by
combining the appropriate # of s, p, d and f orbitals.
• This combination results in hybrid orbitals with a spatial
orientation that matches the geometry of the molecule.
• The number of hybrid orbitals is the same as the number of
atomic orbitals used in their construction.
construction.
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Hybridization in Methane
• By mixing the (one) 2s and (three) 2p valence
orbitals in the C atom a new set of hybrid
orbitals called the (four) sp3 orbitals is created.
• The four sp3 hybrid orbitals are tetrahedral in
shape.
• All four sp3 orbitals have identical shapes and
are equivalent in energy.
• Each H atom bonds to the C atom by orbital
overlap between the 1s orbital and the sp3
hybrid orbitals.
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sp2 Hybrid Orbitals
• sp2 hybrid orbitals are formed by the mixing of
one s and two p orbitals.
• The large lobes of sp2 hybrid orbitals lie in a
trigonal plane and the angles between them
are 120°.
• All molecules with trigonal planar electron pair
geometries have sp2 orbitals on the central
atom.
• One unhybridized p orbital is left as a result of
the formation of sp2 orbitals.
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sp Hybrid Orbitals
• Formed by the combination of one s and one
p orbital.
• The lobes of sp hybrid orbitals are 180º apart,
meaning the geometry is linear.
• All molecules with linear geometry will have sp
hybrid orbitals on the central atom.
• The formation of sp hybrid orbitals leaves 2
unhybridized p orbitals on the central atom.
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sp3d and sp3d2 Hybrid Orbitals
• Since there are only three p-orbitals, trigonal
bipyramidal and octahedral electron pair
geometries must involve d-orbitals.
• Trigonal bipyramidal electron pair geometries
require sp3d hybridization.
• Octahedral electron pair geometries require
sp3d2 hybridization.
• The electron pair geometry from VSEPR
theory always determines the hybridization.
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Deciding Hybridization
• Decide the hybridization of the central atom
for SF3+ and I3- ions.
+
F
S
I
I
F
I
F
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Multiple Bonds
• σ-bonds: electron density lies on
the axis between the nuclei. All
single bonds are σ-bonds.
• π-bonds: electron density lies
above and below the plane of the
nuclei.
• A double bond consists of one σbond and one π-bond.
• A triple bond has one σ-bond and
two π-bonds.
• Often, the p-orbitals involved in πbonding come from unhybridized
orbitals.
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Multiple Bonds
• Ethylene, C2H4, has:
– one σ- and one π-bond
– both C atoms are sp2 hybridized
– both C atoms have trigonal planar electron pair and
molecular geometries
– The formation of a π bond requires that the molecule be
planar.
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Multiple Bonds
Consider acetylene, C2H2
–The electron pair geometry of each C is linear & the C atoms
are sp hybridized;
–The sp hybrid orbitals form the C-C and C-H σ-bonds;
–There are two unhybridized p-orbitals which form the two πbonds.
–one π-bond is above and below the plane of the nuclei;
–one π-bond is in front and behind the plane of the nuclei.
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Delocalized π Bonding
• So far all the bonds we have encountered are
localized between two nuclei.
• In the case of benzene
– there are 6 C-C σ bonds, 6 C-H σ bonds,
– each C atom is sp2 hybridized,
– and there are 6 unhybridized p orbitals on each C atom.
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Delocalized π Bonding
• In benzene there are two options for the 3 π bonds
– localized between C atoms or
– delocalized over the entire ring (i.e. the π
electrons are shared by all 6 C atoms).
• Experimentally, all C-C bonds are the same length in
benzene.
• Therefore, all C-C bonds are of the same type (recall
single bonds are longer than double bonds).
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Multiple Bonds
• Every two atoms share at least 2 electrons.
• Two electrons between atoms on the same
axis as the nuclei are σ bonds.
• σ-bonds are always localized.
• If two atoms share more than one pair of
electrons, the second and third pair form πbonds.
• When resonance structures are possible,
delocalization is also possible.
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