chapter 9 chemical bonding ii: molecular geometry and bonding · pdf filechapter 9 chemical...
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![Page 1: Chapter 9 Chemical Bonding II: Molecular Geometry and Bonding · PDF fileChapter 9 Chemical Bonding II: Molecular Geometry and Bonding Theories . Copyright McGraw-Hill 2009 2 9.1 Molecular](https://reader031.vdocuments.net/reader031/viewer/2022013113/5a9f761a7f8b9a84178cdb0d/html5/thumbnails/1.jpg)
Copyright McGraw-Hill 2009 1
Chapter 9
Chemical Bonding II:
Molecular Geometry
and Bonding
Theories
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9.1 Molecular Geometry
• Molecular geometry is the three-dimensional shape of a molecule.
• Geometry can be predicted using – Lewis dot structures– VSEPR model
CCl4
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• Molecules of the type ABx will be
considered
– A is the central atom
– B atoms surround the central atom
– x commonly has integer values from 1 to 6
– Examples:
All AB1 molecules are linear.
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AB2
AB3
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• VSEPR Model: Valence-Shell Electron-
Pair Repulsion Model
− Electron pairs move as far apart as possible to
minimize repulsions.
Lewis
structure
Electron-domain
geometry
Molecular
geometry
− Strategy to predict geometry:
− Electron domain is a lone pair or a bond
(the bond may be single, double, or triple).
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• Steps to determine Geometry
−Step #1: Draw the molecule’s Lewis structure.
−Step #2: Count the number of electron domains on
the central atom.
−Step #3: Determine the electron-domain geometry.
Examples
H – CN::O:
:O – N = O:
– :
::
:–
:F – Xe – F:::
: ::
• The electron-domain geometry is based on the
number of electron domains around the central
atom.
2 3 5
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• Electron domains and geometry
Number of
Electron DomainsElectron-Domain Geometry
2 Linear
3 Trigonal planar
4 Tetrahedral
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Number of
Electron DomainsElectron-Domain Geometry
5 Trigonal bipyramidal
6 Octahedral
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− Step #4: Determine the molecular geometry.
Step #2 4 e¯ domains
Example: Ammonia, NH3
The electron-domain geometry and the number of
bonded atoms determine the molecular geometry.
Step #1
:
H – N – H –
Helectron-domain
geometry
tetrahedral
Step #3
molecular geometry = trigonal pyramidal
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Note: The common molecular geometries are all
derived from these 5 electron-domain
geometries.
Linear
Bent
Trigonal planar
Bent
Trigonal pyramidal
Tetrahedral
T-shaped
Seesaw
Trigonal bipyramidal
Linear
T-shaped
Square planar
Square pyramidal
Octahedral
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For SF4, which geometry is correct?
or
• Axial and equatorial positions
ax
ax
eq
eqeq
ax = axial
eq = equatorial
The 5 electron
domains are not
all equivalent.
Why?
Fewest lone-pair –
bond-pair interactions
at angles of 90o
90º
120º
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Copyright McGraw-Hill 2009 12
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Exercise: For each of the following species,
i) draw the Lewis structure.
ii) determine the number of electron domains on
the central atom and its electron-domain
geometry.
iii) predict the molecular geometry.
a) NF3
b) CO32
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a) NF3
i)
ii) 4 electron domains on the central atom.
Electron-domain geometry: tetrahedral
iii) One lone pair on the central atom.
Molecular geometry: trigonal pyramidal
F N F
F
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a) CO32
i)
ii) 3 electron domains on the central atom.
Electron-domain geometry: trigonal planar
iii) No lone pairs on the central atom.
Molecular geometry: trigonal planar
O C O
O
2
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• Deviations from ideal bond angles
− All electron domains repel each other.
− The repulsion between domains depends on the
types of domains involved.
single
bond
triple
bond
double
bond
lone
pair< <<
Increasingly repulsion
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lone-pair - lone-pair repulsion
is greater than
lone-pair - bonding-pair repulsion
is greater than
bonding-pair - bonding-pair repulsion
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• Geometry with more than one central atom
CH3OH
C O
H
H
H
H
AX4AX4No lone pairs
2 lone pairs
tetrahedral
bent
C
H
H
O
H
H
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9.2 Molecular Geometry
and Polarity
+ -
H – F H – F
Bond dipoles are vectors and therefore are additive.
The HF bond is polar and HF has a dipole moment ().
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• Molecules with more than two atoms– Remember bond dipoles are additive since
they are vectors.
dipole moment > 0
dipole moment = 0
H2O
CO2
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cis-1,2-dichloroethene trans-1,2-dichloroethene
C C=Cl
H
Cl
H
C C=Cl
H
H
Cl
Example: Dichloroethene, C2H2Cl2, exists as three isomers.
polar
= 1.90 D
bp = 60.3ºC
nonpolar
= 0 D
bp = 47.5ºC
C C=Cl
Cl
H
H
1,1-dichloroethene
polar
= 1.34 D
bp = 31.7ºC
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9.3 Valence Bond Theory
• Electrons in molecules occupy atomic
orbitals.
• Covalent bonding results from the
overlap of atomic orbitals.
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Representation of singly-occupied and doubly-occupied
s and p atomic orbitals. Singly-occupied orbitals appear
light; doubly-occupied orbitals appear darker.
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Example: H(1s1) + H(1s1) H2
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Example: F(1s22s22p5) + F(1s22s22p5) F2
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Example: H(1s1) + F(1s22s22p5) HF
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9.4 Hybridization of
Atomic Orbitals• Valence bond theory cannot account for
many experimental observations.
Beryllium Chloride, BeCl2
linear
both bonds equivalent
VSEPRBe ClCl
• No unpaired electrons
• 2 types of overlap
with 2s and 2p
#1
#2
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• The atomic orbitals on an atom mix to
form hybrid orbitals.
– Orbital shapes (boundary surfaces) are
pictorial representations of wave functions.
– Wave functions are mathematical
functions.
– Mathematical functions can be combined.
• Hybridization of s and p orbitals
– sp hybridization
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• The two sp orbitals point in opposite directions
inline with one another.
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• Each Be sp orbital overlaps a Cl 3p orbital to yield
BeCl2.
2Cl + BeCl2Be
All bond angles 180o.
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#2 • 2 types of overlap with 2s and 2p
#3 • overlap with p-orbitals = 90o
Example: Boron trifluoride, BF3
. .F ....
B
F..
.... F
. .
. ...
VSEPR • trigonal planar
• all bonds equivalent
• only 1 unpaired electron
− sp2 hybridization
#1
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• The three sp2 orbitals point to the corners of an
equilateral triangle.
All bond angles 120o.
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• Each B sp2 orbital overlaps a F 2p orbital to yield BF3.
B3F + BF3
2s2 2p5
All bond angles 120o.
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#2 • 2 types of overlap with 2s and 2p
• overlap with p-orbitals = 90°
tetrahedral
all bonds equivalent
Example: Methane, CH4
VSEPRC
H
H
H
H
• only 2 unpaired electrons#1
− sp3 hybridization
#3
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• The sp3 hybrid orbitals point to the corners of a
tetrahedron.
All bond angles 109.5o.
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• Each C 2sp3 orbital overlaps a H 1s orbital to yield CH4.
C4H + CH41s1
All bond angles 109.5o.
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• Expanded octets
− Hybridization of s, p and d orbitals
Bond angles 120o and 90o.
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Lewis
structure
Number of
electron domainsType of
hybridization
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9.5 Hybridization in Molecules
Containing Multiple Bonds
• A sigma () bond is a bond in which
the shared electron density is
concentrated directly along the
internuclear axis between the two nuclei
involved in bonding.
– end-to-end overlap
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• Double and triple bonds consist of both
sigma () and pi () bonds.
• A pi () bond forms from the sideways
overlap of p orbitals resulting in regions
of electron density that are concentrated
above and below the plane of the
molecule.– sideways overlap
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Example: Ethylene, C2H4 C C=H
H
H
H
number of = 3
electron domains
Double bond = 1 bond + 1 bond
hybridization = sp2
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Example: Acetylene, C2H2 H– CC – H
number of = 2
electron domains
Triple bond = 1 bond + 2 bonds
hybridization = sp
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Exercise: How many pi bonds and sigma bonds are
in each of the following molecules? Describe the
hybridization of each C atom.
(a) 4 sigma bonds
(b) 5 sigma bonds, 1 pi bond
(c) 10 sigma bonds, 3 pi bonds
sp3 sp2
sp2 sp3 sp2
sp2
sp spC ClCl
H
H
C C
HH
H Cl
H3C C
H
C
H
C C H
(a) (b) (c)
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9.6 Molecular Orbital Theory
• Molecular Orbital Theory (MO)
– Atomic orbitals combine to form new molecular orbitals which are spread out over the entire molecule. Electrons are in orbitals that belong to the molecule as a whole.
– Molecular orbitals (wave functions) result from adding and/or subtracting atomic orbitals (wave functions).
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• Molecular orbital theory (MO)
– Atomic orbitals combine to form new,
molecular orbitals that are spread out over
the entire molecule. Electrons are now in
orbitals that belong to the molecule as a
whole.
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− Pictorial representation.
H + H H2
1s1 1s1 1s2
bonding orbital
antibonding orbital
*s1
s11s 1s
H atomic orbitals
H2 molecular orbitals
sum
difference
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• Lower energy and more stable than the atomic
orbitals that were added
• Types of molecular orbitals
− Bonding molecular orbital
• High electron density between the nuclei
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• Higher energy and less stable than the atomic
orbitals that were subtracted
− Antibonding molecular orbital
• Low electron density between the
nuclei
node
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bond order = 1/2 (2 – 0) = 1
single bondbond order = 1/2 (2 – 2) = 0
no bond
• Bond order
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bond order = 1/2 (2 – 0) = 1
single bond
bond order = 1/2 (2 – 2) = 0
no bond
Examples Using Antibonding Orbitals
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• Pi () molecular orbitals
– Wave functions representing p orbitals
combine in two different ways yielding
either orbitals or orbitals.
– End-to-end combination yields sigma ()
orbitals
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Atomic orbitals
+ +- -±
2pz 2pz
Molecular orbitals
+
- -+ +
- -
sum
difference
bonding orbital
antibonding orbital
*p2
p2
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Atomic orbitals
sum
difference
sum
difference
2px
±+
--
2px
+
+
-
+
+
-
-
+
-
+
+ --
Molecular orbitals
*p2
*p2
p2
p2
− Sideways combination yields pi () orbitals
2py
±
++
--
2py
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• Energy order of the 2p and 2p orbitals
changes across the period.
B2, C2, N2 O2, F2, Ne2
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Molecular orbital diagram of nitrogen, N2
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Molecular orbital diagram of oxygen, O2
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• Magnetism
– Diamagnetic substance
• A substance whose electrons are all
paired.
•
• Weakly repelled by magnetic fields.
– Paramagnetic substance
• A substance with one or more unpaired
electrons.
•
• Attracted by magnetic fields.
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9.7 Bonding Theories and
Descriptions of Molecules with
Delocalized Bonding
• In localized bonds the and bonding
electrons are associated with only two
atoms.
• Resonance requires delocalized
bonds when applying valence bond
theory.
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• In delocalized bonds the bonding
electrons are associated with more than
two atoms.
Examples:
NO3, CO3
2, C6H6 , O3
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Example: benzene, C6H6
Each C atom has 3
electron domains
Hybridization: sp2
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Key Points
• Molecular geometry
– VSEPR model
• Molecular geometry and polarity
• Valence bond theory
• Hybridization of atomic orbitals
– s and p
– s, p, and d
• Hybridization involving multiple bonds
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Key Points
• Molecular orbital theory
– Bonding and antibonding orbitals
– Sigma () molecular orbitals
– Pi () molecular orbitals
– MO diagrams