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CHM 123
Chapter 7
7.9 Molecular shapes and VSEPR theory
VSEPR theory proposes that the geometric arrangement of terminal atoms, or groups of atoms about a central
atom in a covalent compound, or charged ion, is determined solely by the repulsions between electron pairs
present in the valence shell of the central atom
In the valence-shell electron-pair repulsion theory (VSEPR), the electron groups around a central atom
• are arranged as far apart from each other as possible, WHY?
• have the least amount of repulsion of the negatively charged electrons.
• have a geometry around the central atom that determines molecular shape.
Electrons in bonds and in lone pairs can be thought of as
“charge clouds” that repel one another and stay as far
apart as possible, this causing molecules to assume
specific shapes.
Working from an electron-dot structure, count the
number of “charge clouds,” and then determine the
molecular shape.
The following are the “parent” electronic structures upon which VSEPR is based. These structures show how to
minimize the energy of the structure by placing 2, 3, 4, 5 or 6 electron groups (charge clouds) as far apart around a
central atom as possible in three-dimensional space.
Electrons in bonds and in lone pairs can be thought of as “charge clouds” that repel one another and stay as far
apart as possible, this causing molecules to assume specific shapes.
Working from an electron-dot structure, count the number of “charge clouds,” and then determine the molecular
shape.
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The Effect of Lone Pairs
 lone pair groups “occupy more space” on the central atom
o because their electron density is exclusively on the central atom rather than shared like bonding
electron groups
 relative sizes of repulsive force interactions is:
 Lone Pair – Lone Pair > Lone Pair – Bonding Pair > Bonding Pair – Bonding Pair
 this effects the bond angles, making them smaller than expected
In each of these examples, the electron
pairs are arranged tetrahedrally, and two or
more atoms are bonded in these
tetrahedral directions to give the different
geometries.
Lone pairs are absolutely critical part of
the electronic structure (charge clouds)
that contributes to the shape of the
molecule, but only the attached atoms are
included in deriving the shape name.
How many electron groups (charge clouds) are around the central atom in the following?
SO2
PCl5
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Total # of egroups on
central atom
“Parent” electronic
geometry
2
3
3
4
4
4
5
5
5
5
6
6
6
Linear
Trigonal Planar
Trigonal Planar
Tetrahedral
Tetrahedral
Tetrahedral
Trigonal Bipyramidal
Trigonal Bipyramidal
Trigonal Bipyramidal
Trigonal Bipyramidal
Octahedral
Octahedral
Octahedral
#
Bonde
d
atoms
2
3
2
4
3
2
5
4
3
2
6
5
4
# Lone
pairs
Idealized molecular
shape
Idealized
bond angles
0
0
1
0
1
2
0
1
2
3
0
1
2
Linear
Trigonal Planar
Bent
Tetrahedral
Trigonal Pyramidal
Bent
Trigonal Bipyramidal
Seesaw
T-shaped
Linear
Octahedral
Square Pyramidal
Square Planar
180o
120 o
120 o
109.5 o
109.5 o
109.5 o
90 o, 120 o, 180 o
90 o, 120 o, 180 o
90 o, 180 o
180 o
90 o, 180 o
90 o, 180 o
90 o, 180 o
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Real bond angles vs. Idealized bond angles
VSEPR predicts the idealized bond angle(s) by assuming that all electron groups take up the same amount of space.
Since lone pairs are attracted to only one nucleus, they expand into space further than bonding pairs, which are
attracted to two nuclei. As a result, real molecules that has lone pairs on the central atom often have bond angles
that are slightly different than the idealized prediction.
Central atom without lone pairs has the
same real bond angle as the idealized angle.
The exceptions to this are square planar
shapes and linear (derived from trigonal
bipyramidal electronic structure) shapes where the
lone pairs offset one another, thus causing no
deviation from ideality.
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Examples: Name the shape and give the idealized bond angles for the following
Lewis structure
Shape
Idealized bond angle
Real bond angle
Molecules with no central atom
Many molecules don’t have a “central” atom but many
“central” atoms. These molecules don’t fit into the
shape names that we’ve learned. However, we can
give an approximated shape and bond angle to each
“central” atom at a time.
Example: Give the approximate shape at the
numbered “central” atoms
Drawing 3-D structures
In order to draw 3-D structure, chemists use dark wedges to indicate a
bond projecting forward (out of the page) and dashed to indicate a bond
going away (going back into the page) and a normal line to indicate a bond
in the plane of the page
Examples: Draw 3-D structure for the following molecules
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6
7.10 – Valence Bond Theory
When a covalent bond is formed, there is shared electron density between the nuclei of the bonded atom
The simultaneous attraction of the shared electron density for both nuclei holds the atoms together, forming a
covalent bond
Valence Bond Theory: A quantum mechanical model which shows how electron pairs are shared in a covalent
bond.
Bond forms between two atoms when the following conditions are met:
•
Covalent bonds are formed by overlap of atomic orbitals, each of which contains one electron of opposite
spin.
•
Each of the bonded atoms maintains its own atomic orbitals, but the electron pair in the overlapping
orbitals is shared by both atoms.
•
The greater the amount of overlap, the stronger the bond.
In some cases, atoms use “simple” atomic orbital (e.g., 1s, 2s, 2p, etc.) to form bonds. In other case, they use a
“mixture” of simple atomic orbitals known as “hybrid” atomic orbitals.
Two special names for covalent bonds of Organic molecules
Sigma (σ) bonds
Created when “head on” overlap
occurs of orbital
Pi (π) bonds
Created when “side on” overlaps occurs of orbital (p
orbitals)
Pi bonds are usually weaker than sigma bon. From the perspective of quantum mechanics,
this bond's weakness is explained by significantly less overlap between the component porbitals due to their parallel orientation. This is contrasted by sigma bond which form
bonding orbitals directly between the nucleus of the bonding atoms, resulting in greater
overlap and a strong sigma bond
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Hybrid orbital – orbitals are used to describe bonding that is obtained by taking combinations of atomic orbitals of
the isolated atoms.
1. When forming hybrid orbitals, the number of hybrid orbitals formed equals the number of orbitals
mathematically combined or “mixed”. For example, if an s orbital is combined with a p orbital the result
is two “sp” hybrid orbitals.
2. Hybrid orbitals have orientations around the central atom that correspond to the electron-domain
geometry predicted by the VSEPR Theory
How can the bonding in CH4 be explained?
sp3 Hybridization of Carbon
- Mixing 1s and all 3p atomic orbital
- Four sp3 hybridized orbitals equal in size, energy and shape
- Responsible for sigma bond ( single bond)
- Tetrahedral shape
Consider methane (CH4)
3-D representation of methane (CH4)
The sp2 Hybridization of Carbon
- Mixing 1s and 2p atomic orbitals
- Three sp2 hybridized orbitals equal in size, energy and shape
- One 2p unhybridized orbital
- Responsible for σ bond ( single bond)
- One π bond (double bond)
- Trigonal planar shape
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Consider ethane (C2H4)
3-D representation of ethane (C2H4)
The sp Hybridization of Carbon
- Mixing 1s and 1p atomic orbitals
- Two sp hybridized orbitals equal in size, energy and shape
- Two 2p unhybridized orbital
- one σ bond ( single bond)
- Two π bond (triple bond)
- Linear
Consider ethyne (C2H2)
3-D representation of ethyne (C2H2)
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Five types of hybrid are shown below
# of e- groups around central atom
Hybrid orbitals used
Orientation of Hybrid Orbitals
2
sp
3
sp2
4
sp3
5
sp3d
6
sp3d2
Bonding to O and N
Like Carbon, O and N can participate in single bond and multiple bonds compose of σ and π
*Note: the lone pair or non bonding e- pair occupies space just as bonded atom
Examples:
1. Predict the hybridization, geometry, and bond angle for the carbon, oxygen and nitrogen atoms in the
following molecules
2. Assign hybridization for all carbon atoms and identify the angle at C-C-C.
3. What hybrid orbitals would be expected for the central atom in each of the following?
a) SF2
b)
BrF3
c)
ClF4+
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