Download Chapter 10 Molecular Shapes and Valence Bond Theory Lewis dot

Chapter 10
Molecular Shapes and Valence Bond Theory
10.1 Artificial Sweeteners: Fooled by Molecular Shape
(Suggested Reading)
10.2 VSEPR Theory: The Five Basic Shapes
10.3 VSEPR Theory: The Effect of Lone Pairs
10.4 VSEPR Theory: Predicting Molecular Geometries
10.5 Molecular Shape and Polarity
10.6 Valence Bond Theory: Orbital Overlap as a Chemical Bond
10.7 Valence Bond Theory: Hybridization of Atomic Orbitals
(sp, sp2 and sp3 only)
1
Lewis dot structure
Molecular Shape
• Lewis dot structures only gives us an idea of the
electron distribution in the species. There is NO IDEA
about the molecular geometry, which depends on the
relative position of terminal atoms around the central
atom.
• We will connect the electron distribution in a Lewis dot
structure → molecular geometry by using the ValenceShell Electron-Pair Repulsion (VSEPR) theory.
VSEPR = Groups of electrons repel each other,
ending up as far from each other as physically
possible.
2
1
VSEPR Model
• Although the Theory states “repulsion of
ELECTRON PAIRS”
………..it is actually repulsion of ELECTRON
GROUPS because lone pair, single, double
and triple bond pairs are treated as ONE PAIR
of electrons in VSEPR theory
That is:
A lone pair is ONE GROUP of electrons
A single bond is ONE GROUP of electrons
A double bond is ONE GROUP of electrons
A triple bond is ONE GROUP of electrons
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Electron distribution vs. geometry
Electron
distribution
Molecular geometry
Looks at “shape” of
electron group distribution
Looks at “shape” of
nuclear positions around
the central atom
INCLUDES lone pairs
No terminal nuclei on lone
pairs means we IGNORE
(“can’t see”) all lone pairs.
Bonding
pair of eLone pair
of e-
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2
Electron distribution vs. Geometry
If the central atom has no lone pairs on it, then the electron group
distribution and the molecular geometry are the same !
5
Molecular Geometries
Examples of geometries of molecules with
no lone pairs around central atom and bond angles
Lewis
Structure
For CH4
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3
Molecular Geometries
Its molecular
shape
Lewis structure for PCl5
Lewis structure for SF6
7
The Effect of Lone Pairs on Shape
• Lone pair groups “occupy more space” on the
central atom.
– because their electron density is exclusively on the
central atom rather than shared like bonding electron
groups
• relative sizes of repulsive force interactions are:
• Lone Pair – Lone Pair > Lone Pair – Bonding
Pair > Bonding Pair – Bonding Pair
• This affects the bond angles, making them
smaller than expected.
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4
The Effect of Lone Pairs on Shape
IF NO LONE PAIRS
-The molecule’s shape will be one of the basic
molecular geometries if all the electron groups
are bonds (see slide 5)
LONE PAIRS AROUND CENTRAL
-Molecules with lone pairs will have distorted
bond angles but the shape will be a derivative
of one of the basic shapes (See Table 10.1(lone
pair column) for examples)
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Electron and Molecular Geometries
10
5
Electron and Molecular Geometries
Electron
groups
11
Electron and Molecular Geometries
Electron
groups
12
6
The Effect of Lone Pairs on Shape
Detailed examples of 4 molecules
(1) When there are three electron groups around the central and one of
them is a lone pair, the resulting molecular shape is called a bent
shape.
The bond angle is <120°.
SO2
Example of 2 bonding and one lone pair
<1200
Bent shape
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The Effect of Lone Pairs on Shape
(2) When there are four electron groups around the central
and one of them is a lone pair
Lewis structure of
NH3
<1090
Example of 3 bonding
and 1 lone pair
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7
The Effect of Lone Pairs on Shape
(3) When there are five electron groups around the central
and one of them is a lone pair
A “seesaw” shape
Example of 4 bonding
And one lone pair
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The Effect of Lone Pairs on Shape
(4) When there are six electron groups around the central
and one of them is a lone pair
All <900
Lewis structure of
BrF5
Example of 5 bonding
And one lone pair
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8
Bond Angles
You also need to know the bond angles of all the shapes for:
1) The basic molecular geometries AND
2) the geometries of molecules having lone pairs (Table 10.1)
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Figuring out Molecular Shapes
1.
Draw the Lewis dot structure
2.
Determine the number of electron groups on the
central atom to get electron geometry (see Slide 5).
If no lone pairs around central atom then you have the
molecule’s molecular geometry.
3.
If central atom has lone pairs, use the number of
bonded and lone pairs and the arrangement (Table
10.1) to determine resulting molecular geometry.
4.
Draw the 3 D structure as best you can.
3.
Determine bond angles.
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9
Drawing 3D shapes
• By convention, the central atom is put in the plane of the paper.
• Put as many other atoms as possible in the same plane and indicate
with a straight line.
• For atoms in front of the plane, use a solid wedge.
• For atoms behind the plane, use a hashed wedge
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Practice:
Draw the 1) Lewis structures 2) the molecular shapes ( 3D
structure) around their central atom and
3) the bond angles for the following molecules?
(a)
(b)
(c)
(d)
(e)
H2O
PF5
SeCl4
KrCl2
IF5
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10
Dipole Moment and Molecular Shape
• If there are polar covalent bonds in a molecule, the
molecule MAY OR MAY NOT have a permanent dipole
moment.
• A permanent dipole moment means that there is a
partially negative and a partially positive site that is
permanent.
• To determine if a molecule has a permanent dipole
moment, we add together the vectors of all the polar
covalent bonds (and their dipole moments.)
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Molecular Dipole Moment
• A simple permanent
dipole is HCl
••
H - Cl :
••
δ+ δ−
Dipole moment = 3.34 D
••
It has a polar covalent bond.
Since
H - Cl
: there is
••
only one bond, this one vector of
charge
describes the permanent dipole.
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11
Molecular Dipole Moment
• Water has two polar covalent bonds and two dipoles
• The permanent dipole moment in water can be seen
by adding together the charge separation vectors of
the two polar covalent O-H bonds.
Dipole moment = 1.94 D
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Molecular Dipole Moment
But a molecule with more than one polar bond MIGHT
NOT have a permanent dipole moment .
If there is symmetry of the (identical) polar bonds the
resultant vector sums may add up to zero.
An example is carbon dioxide CO2
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12
Predicting dipole moments from Geometry
Note:
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Predicting dipole moments from Geometry
Practice—Decide whether the following
molecules are polar, given the EN values.
EN values
O = 3.5
N = 3.0
Cl = 3.0
S = 2.5
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13
Molecular Properties from molecular dipole
moments
• As we will see in next Chapter (section 11.3)
molecules with partial positive and negative
charges will attract the opposite regions on other
molecules of the same type. Such intermolecular
forces affect the molecular properties of the
compound, i.e., It is a liquid or a solid.
It will also affect the compound’s boiling point.
Example: two isomers of C2H2Cl2
cis-1,2-dichloroethane
trans-1,2-dichloroethane
boiling point
60oC
48oC
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Explain why the b.p. are different by drawing the dipole moments of each.
10.6 Valence Bond Theory (Orbital overlap)
Covalent bonds form between atoms when:
1. Orbitals in the atoms overlap to create molecular
bonding orbitals.
2.
Each molecular bonding orbital has no more than 2
electrons in it.
ALSO…. bond formation occurs between two atomic orbitals
containing one electron and
…..Covalent bonds are strongest when there is
maximum orbital overlap between atomic orbitals.
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14
Valence Bond Theory (Orbital overlap)
Here are some favourable atomic orbital
overlaps for H2 and HCl
H
1s
↑
H
1s
↑
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10.7 Hybridization of atomic orbitals
• But there are problems that arise from this
simple theory. The number of partially filled or
empty atomic orbitals did not always predict the
number of bonds or orientation of bonds.
For Carbon whose valence atomic orbital is =
2s22px12py12pz0 would predict two or three
bonds that are 90° apart, rather than four
bonds that are 109.5° apart in CH4.
• To adjust for these inconsistencies, it was
postulated that the valence atomic orbitals could
hybridize (mix) before bonding took place.
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15
Unhybridized C orbitals predict the
wrong bonding and geometry
Hybridization of C is to mix all the
2s and 2p orbitals to get four equal
orbitals that point to the corners of a
tetrahedron.
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Hybridization
• Many atoms hybridize their orbitals to
maximize bonding.
– Hybridizing is mixing different types of orbitals
to make a new set of degenerate orbitals.
– sp, sp2, sp3, sp3d, sp3d2
– more bonds = more full orbitals = more stability
• Same types of atom can have different
hybridizations depending on the compound.
C = sp, sp2, sp3
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16
Carbon Hybridizations
Unhybridized
↑↓
↑
↑
2p
2s
sp hybridized
↑
↑
↑
2sp
↑
2p
sp2 hybridized
↑
↑ ↑
2sp2
sp3 hybridized
↑
↑
↑
↑
2p
↑
2sp3
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Hybridization
• The number of standard atomic orbitals
combined equals the number of hybrid orbitals
formed.
– H cannot hybridize!
.
• The number and type of standard atomic
orbitals combined determine the shape of the
hybrid orbitals.
• The particular kind of hybridization that occurs
is the one that yields the lowest overall energy
for the molecule.
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17
Hybridization
Orbital Diagram of the sp3 Hybridization of C
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Formation of sp3
Hybrid Orbitals
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18
Hybridization
sp3 Hybridized Atoms: Orbital Diagrams
• Place electrons into hybrid and unhybridized valence orbitals
as if all the orbitals have equal energy.
• Lone pairs generally occupy hybrid orbitals.
Unhybridized atom
2s
↑↓
2s
↑ ↑
2p
C
↑
↑ ↑ ↑
2p
N
↑
↑ ↑ ↑
2sp3
↑
↑↓
sp3 hybridized atom
↑ ↑ ↑
2sp3
37
Hybridization
Methane Formation
with sp3 C
Ammonia Formation
with sp3 N
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19
Other Types of Hybrid Orbitals
• A total of n atomic orbitals combine to give n
hybrid orbitals of a given kind
39
Resulting Shapes of Hybrid Orbitals
40
20
Determining hybrid orbital diagrams
1.
2.
3.
4.
Draw the Lewis dot structure
Use VSEPR theory to predict electron group
arrangement
Use Table on slide 38 to determine what hybrid orbitals
have the same arrangement
Create the hybrid orbital diagram based on changing
the ground state diagram of the central atom
Practice:
Describe the bonding of BF3 in terms of
hybrid orbital theory.
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Types of Bonds
• A sigma (σ) bond results when the
interacting atomic orbitals point along the axis
connecting the two bonding nuclei.
-between s-to-s, hybrid-to-hybrid, s-to-hybrid, axis p
• A pi (π) bond results when the bonding
atomic orbitals are parallel to each other and
perpendicular to the axis connecting the two
bonding nuclei.
– between unhybridized parallel p orbitals
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21
Types of Bonds
43
Orbital Diagrams with Hybridization
• “Overlap” between a hybrid orbital on one
atom and a hybrid or nonhybridized orbital
on another atom results in a σ bond.
• “Overlap” between unhybridized p orbitals
on bonded atoms results in a π bond.
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22
An example of a Double Bond from hybridization
H2C=CH2
Sp2 orbitals
Double bond
using 1 sp2 and the 2p
2p orbitals
In ethene (C2H4) the hybrid
orbital is an sp2 (not sp3)
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An example of a Triple Bond
C2H2 (ethyne) HC≡CH
using sp hybridized
Using 2p atomic orbitals
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23
Example of sp2 Hybridized C and O in CH2O
Formaldehyde
pC
↑
π
σ
sp2 C ↑ ↑ ↑
↑
↑↓ ↑↓
sp2 O
σ
↑
σ
pO
↑
↑
1s H 1s H
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Practice—Draw the orbital diagram for
hybridization of B and O atom. How many σ
the
and π bonds would you expect each to form?
sp2
Unhybridized atom
↑↓
2s
B
↑
2p
O
↑↓
2s
↑↓ ↑ ↑
2p
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24
HCN Orbital Diagram-sp hybridized
↑
pC
sp C
↑
2π
↑
↑
σ
↑
↑
pN
↑
↑↓
sp N
s
↑
1s H
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sp3d Hybridized -Orbital Diagrams
Unhybridized atom
↑↓
3s
↑↓
3s
↑ ↑ ↑
3p
↑↓ ↑ ↑
3p
sp3d hybridized atom
P
3d
↑
↑ ↑ ↑
3sp3d
S ↑↓ ↑
3d
↑ ↑
↑
↑
3sp3d
(nonhybridizing d orbitals not shown)
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SOF4 Orbital Diagram
↑
dS
↑
↑
σ
σ
σ
↑
↑
σ
↑
↑
sp3d S ↑
↑
π
σ
pO
↑
↑
↑↓ ↑↓
sp2 O
↑
2p F 2p F 2p F 2p F
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