Chapter 4
Molecular Structure and Orbitals
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Table of Contents
• (4.1)
Molecular structure: The VSEPR model
• (4.2)
Bond polarity and dipole moments
• (4.3)
Hybridization and the localized electron model
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Question to Consider
• The spicy flavor of chili peppers is attributed to a complex
molecule with multiple hybridizations
• What is the name of this molecule?
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Section 4.1
Molecular Structure: The VSEPR Model
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The VSEPR Model
• Molecular structure is the three-dimensional arrangement of
atoms in a molecule
• Valence shell electron-pair repulsion (VSEPR) model
• The structure around a given atom is determined principally by
minimizing electron pair repulsions
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Electron Structures (1 of 3)
• Linear structure can be observed in BeCl2
• Each electron pair on Be is shared with a Cl atom
• BF3 shows a trigonal planar structure
• Each electron pair is shared with a fluorine atom
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Electron Structures (2 of 3)
• When there are four pairs of electrons around an atom, they
take up a tetrahedral structure
• The bond angle for such a structure is 109.5 degrees
• In the presence of a lone pair, the molecular structure is a
trigonal pyramid
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Electron Structures (3 of 3)
• Consider the structure of NH3, which has one lone pair
• The arrangement of electron pairs is tetrahedral, but the arrangement
of atoms is not
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Bonding Pairs and Lone Pairs
• Bonding pairs are shared between two nuclei
• Electrons can be close to either nucleus
• They are relatively confined between the two nuclei
• Lone pairs center around just one nucleus, and both electrons
choose that nucleus
• Lone pairs need more space than bonding pairs
• They compress the angles between bonding pairs
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Effect of Lone Pairs
• The bond angle between bonding pairs decreases as the
number of lone pairs increases on the central atom
Number of Lone Pairs
Bond Angle
CH4
NH3
H2O
0
1
2
109.5°
107°
104.5°
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Concept Check (1 of 6)
• Arrange the following molecules in the increasing order of bond
angle:
• H2O, CH4, SF6, BF3, NH3 , BeF2
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Electron Structures
• When there are five electron pairs, the structure that produces
minimal repulsion is a trigonal bipyramid
• It consists of two trigonal-based pyramids that share a common base
• The best arrangement for six pairs of electrons around a given
atom is the octahedral structure
• This structure has 90-degree bond angles
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Table 4.2 - Structures of Molecules That Have
Four Electron Pairs Around the Central Atom
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Table 4.3 - Structures of Molecules with Five
Electron Pairs Around the Central Atom (1 of 2)
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Table 4.3 - Structures of Molecules with Five
Electron Pairs Around the Central Atom (2 of 2)
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Concept Check (2 of 6)
• Determine the shape and bond angles for each of the following
molecules:
• HCN
• PH3
• SF4
• O3
• KrF4
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Problem Solving Strategy - Steps to Apply the
VSEPR Model
• Draw the Lewis structure for the molecule
• Count the electron pairs and arrange them in the way that
minimizes repulsion
• Put the pairs as far apart as possible
• Determine the positions of the atoms from the way the electron
pairs are shared
• Determine the name of the molecular structure from the
positions of the atoms
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Interactive Example 4.2 - Prediction of Molecular
Structure II (1 of 4)
• When phosphorus reacts with excess chlorine gas, the
compound phosphorus pentachloride (PCl5) is formed. In the
gaseous and liquid states, this substance consists of PCl5
molecules, but in the solid state it consists of a 1:1 mixture of
PCl4+ and PCl6 − ions. Predict the geometric structures of PCl5,
PCl4+ , and PCl6− .
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Interactive Example 4.2 - Prediction of Molecular
Structure II (2 of 4)
• Solution
• The Lewis structure for PCl5 is shown
• Five pairs of electrons around the phosphorus atom require a trigonal
bipyramidal arrangement
• When the chlorine atoms are included, a trigonal bipyramidal molecule
results:
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Interactive Example 4.2 - Prediction of Molecular
Structure II (3 of 4)
• The Lewis structure for the PCl4+ ion [5 + 4(7) − 1 = 32 valence
electrons] is shown below
• There are four pairs of electrons surrounding the phosphorus atom in
the PCl4+ ion, which requires a tetrahedral arrangement of the pairs
• Since each pair is shared with a chlorine atom, a tetrahedral PCl4+
cation results
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Interactive Example 4.2 - Prediction of Molecular
Structure II (4 of 4)
• The Lewis structure for PCl6− [5 + 6(7) + 1 = 48 valence electrons] is
shown below
• Since phosphorus is surrounded by six pairs of electrons, an
octahedral arrangement is required to minimize repulsions, as shown
below in the center
• Since each electron pair is shared with a chlorine atom, an octahedral
PCl6− anion is predicted
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The VSEPR Model and Multiple Bonds
• While using the VSEPR model, a double bond must be
considered as one effective pair
• The two pairs involved in the double bond are not independent pairs
• The double bond acts as one center of electron density that repels
other electron pairs
• With molecules that exhibit resonance, any one of the
resonance structures can be used to predict its molecular
structure using the VSEPR model
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Interactive Example 4.4 - Structures of Molecules
with Multiple Bonds (1 of 3)
• Predict the molecular structure of the sulfur dioxide molecule.
Is this molecule expected to have a dipole moment?
• Solution
• First, determine the Lewis structure for the SO2 molecule, which has
18 valence electrons
• The expected resonance structures are:
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Interactive Example 4.4 - Structures of Molecules
with Multiple Bonds (2 of 3)
• To determine the molecular structure, count the electron pairs around
the sulfur atom
• In each resonance structure the sulfur has one lone pair, one pair in a single
bond, and one double bond
• Counting the double bond as one pair yields three effective pairs around the
sulfur
• A trigonal planar arrangement is required, which yields a V-shaped
molecule
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Interactive Example 4.4 - Structures of Molecules
with Multiple Bonds (3 of 3)
• Thus the structure of the SO2 molecule is expected to be V-shaped,
with a 120-degree bond angle
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Molecules Containing No Single Central Atom
• The VSEPR model can accurately determine the structure of
complicated molecules such as methanol
• Lewis structure:
• There are four pairs of electrons around the C and O atoms,
which give rise to a tetrahedral arrangement
• Space requirements of the lone pairs distort the arrangement
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Figure 4.8 - The Molecular Structure of Methanol
(a)The arrangement of electron pairs and atoms around the
carbon
(b)The arrangement of bonding and lone pairs around oxygen
(c) The molecular structure
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Accuracy of the VSEPR Model
• It aptly predicts the molecular structures of most molecules
formed from non-metallic elements
• It can be used to predict the structures of molecules with
hundreds of atoms
• It fails to determine the molecular structure in certain instances
• Phosphine (PH3) and ammonia (NH3) have similar Lewis structures
but different bond angles—94 degrees and 107 degrees, respectively
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Section 4.2
Bond Polarity and Dipole Moments
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Dipole Moment (1 of 2)
• A molecule that has a center of positive charge and a center of
negative charge is said to be dipolar or to possess dipole
moment
• It is represented by an arrow pointing to the negative charge center
• The tail indicates the positive charge center
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Dipole Moment (2 of 2)
• Electrostatic potential diagrams can also be used to represent
dipole moment
• The colors of visible light are used to show variation in distribution of
charge
• Red - Most electron-rich region
• Blue - Most electron-poor region
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Bond Polarity Trends
• Any diatomic molecule with polar bonds will exhibit dipole
moments
• This behaviour can also be exhibited by polyatomic molecules
• Few molecules possess polar bonds but lack dipole moment
• Occurs when the individual bond polarities are arranged in a manner
that they cancel each other out
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Table 4.4 - Types of Molecules with Polar Bonds
but No Resulting Dipole Moment
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Example 4.5 - Bond Polarity and
Dipole Moment (1 of 7)
• For each of the following molecules, show the direction of the
bond polarities and indicate which ones have a dipole moment:
• HCl
• Cl2
• SO3 (planar molecule with the oxygen atoms spaced evenly around
the central sulfur atom)
• CH4 (tetrahedral with the carbon atom at the center)
• H2S (V-shaped with the sulphur atom at the point)
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Example 4.5 - Bond Polarity and
Dipole Moment (2 of 7)
• Solution
• The HCl molecule:
• The electronegativity of chlorine is greater than that of hydrogen
• Thus the chlorine will be partially negative, and the hydrogen will be partially
positive
• The HCl molecule has a dipole moment:
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Example 4.5 - Bond Polarity and
Dipole Moment (3 of 7)
• The Cl2 molecule:
• The two chlorine atoms share the electrons equally
• No bond polarity occurs and the Cl2 molecule has no dipole moment
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Example 4.5 - Bond Polarity and
Dipole Moment (4 of 7)
• The SO3 molecule:
• The electronegativity of oxygen is greater than that of sulfur
• This means that each oxygen will have a partial negative charge, and the sulfur
will have a partial positive charge
• The bond polarities arranged symmetrically as shown cancel, and the molecule
has no dipole moment
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Example 4.5 - Bond Polarity and
Dipole Moment (5 of 7)
• The CH4 molecule:
• Carbon has a slightly higher electronegativity than does hydrogen
• This leads to small partial positive charges on the hydrogen atoms and a small
partial negative charge on the carbon:
• The bond polarities cancel, and the molecule has no dipole moment
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Example 4.5 - Bond Polarity and
Dipole Moment (6 of 7)
• The H2S molecule:
• Since the electronegativity of sulfur is slightly greater than that of hydrogen, the
sulfur will have a partial negative charge, and the hydrogen atoms will have a
partial positive charge, which can be represented as:
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Example 4.5 - Bond Polarity and
Dipole Moment (7 of 7)
• This case is analogous to the water molecule, and the polar bonds result in a
dipole moment oriented as shown:
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Section 4.3
Hybridization and the Localized Electron
Model
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Hybridization
• It refers to the mixing of the native atomic orbitals to form
special orbitals for bonding
• Atoms may adopt a different set of atomic orbitals or hybrid
orbitals from those in the free state
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sp3 Hybridization
• It can be observed upon combination of one 2s and three 2p
orbitals
• Whenever an atom requires a set of equivalent tetrahedral
atomic orbitals, this model assumes that the atom adopts a set
of sp3 orbitals
• The atom becomes sp3 hybridized
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Figure 4.15 - The Formation of sp3 Hybrid Orbitals
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Concept Check (3 of 6)
• What is the valence electron configuration of a carbon atom?
• Why can’t the bonding orbitals for methane be formed by an
overlap of atomic orbitals?
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Concept Check (4 of 6)
• Why can’t sp3 hybridization account for the ethylene molecule?
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sp2 Hybridization
• Gives a trigonal planar arrangement of atomic orbitals with
bond angles of 120 degrees
• It occurs on the combination of one 2s and two 2p orbitals
• One p orbital is not used
• It is oriented perpendicular to the plane of the sp2 orbitals
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Types of sp2 Hybridized Bonds
• Sigma () bond
• Electron pair is shared in an area centered on a line running between
the atoms
• Pi () bond
• Forms double and triple bonds by sharing electron pair(s) in the space
above and below the σ bond using the unhybridized p orbitals
• A double bond always consists of one  bond and one  bond
• If an atom is surrounded by three effective pairs, a set of sp2
hybrid orbitals is required
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Figure 4.24 - A Carbon–Carbon Double Bond
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sp Hybridization
• It occurs upon combination of one s and one p orbital
• Two effective pairs around an atom always require sp hybridization of
that atom
• It follows a linear arrangement of atomic orbitals
• p orbitals that remain unchanged upon hybridization are used
in the formation of  bonds
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Figure 4.31 - Bonding in CO2 Part (a)
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Concept Check (5 of 6)
• Draw the Lewis structure for HCN
• Which of the hybrid orbitals are used?
• Draw HCN and:
• Show all the bonds between the atoms
• Label each  or  bond
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dsp3 Hybridization
• It is a combination of one d, one s, and three p orbitals
• It results in a trigonal bipyramidal arrangement of five
equivalent hybrid orbitals
• The image illustrates hybrid orbitals in a phosphorus atom
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d2sp3 Hybridization
• An atom is d2sp3 hybridized when there is a combination of two
d, one s, and three p orbitals
• It results in an octahedral arrangement of six equivalent hybrid
orbitals
• The image illustrates the orbitals in a sulfur atom
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Interactive Example 4.9 - The Localized Electron
Model IV (1 of 2)
• How is the xenon atom in XeFe4 hybridized?
• Solution
• XeFe4 has six pairs of electrons around xenon that are arranged
octahedrally to minimize repulsions
• An octahedral set of six atomic orbitals is required to hold these
electrons, and the xenon atom is d2sp3 hybridized
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Interactive Example 4.9 - The Localized Electron
Model IV (2 of 2)
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Concept Check (6 of 6)
• For each of the following molecules, determine:
a)Bond angle
b)Expected hybridization of the central atom
NH3
SO2
KrF2
CO2
ICl5
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Figure 4.36 - The Relationship of the Number of Effective Pairs, Their
Spatial Arrangement, and the Hybrid Orbital Set Required (1 of 2)
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Figure 4.36 - The Relationship of the Number of Effective Pairs, Their
Spatial Arrangement, and the Hybrid Orbital Set Required (2 of 2)
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Problem Solving Strategy: Using the Localized
Electron Model
• Draw the Lewis structure(s)
• Determine the arrangement of electron pairs, using the VSEPR
model
• Specify the hybrid orbitals needed to accommodate the
electron pairs
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Interactive Example 4.10 - The Localized Electron
Model V (1 of 6)
• For each of the following molecules or ions, predict the
hybridization of each atom, and describe the molecular
structure
a. CO
b. BF4−
c. XeF2
• Solution
a. CO
• The CO molecule has 10 valence electrons, and its Lewis structure is : C  O :
• Each atom has two effective pairs, which means that both are sp hybridized
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Interactive Example 4.10 - The Localized Electron
Model V (2 of 6)
• The triple bond consists of a σ bond produced by the overlap of an sp orbital from
each atom and two  bonds produced by the overlap of 2p orbitals from each atom
• The lone pairs are in the sp orbitals
• Since the CO molecule has only two atoms, it must be linear
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Interactive Example 4.10 - The Localized Electron
Model V (3 of 6)
• b. BF4−
• The BF4− ion has 32 valence electrons
• The Lewis structure shows four pairs of electrons around the boron atom, which
means a tetrahedral arrangement:
• This requires sp3 hybridization of the boron atom
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Interactive Example 4.10 - The Localized Electron
Model V (4 of 6)
• Each fluorine atom also has four electron pairs and can be assumed to be sp3
hybridized
• The BF4− ion’s molecular structure is tetrahedral
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Interactive Example 4.10 - The Localized Electron
Model V (5 of 6)
c. XeF2
• The XeF2 molecule has 22 valence electrons
• The Lewis structure shows five electron pairs on the xenon atom, which requires
trigonal bipyramidal arrangement:
• Note that the lone pairs are placed in the plane where they are 120 degrees apart
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Interactive Example 4.10 - The Localized Electron
Model V (6 of 6)
• To accommodate five pairs at the vertices of a trigonal bipyramid requires that the
xenon atom adopt a set of five dsp3 orbitals
• Each fluorine atom has four electron pairs and can be assumed to be sp3
hybridized
• The XeF2 molecule has a linear arrangement of atoms
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