7
C hapter
Covalent Bonding
Chemistry 4th Edition
McMurry/Fay
Dr. Paul Charlesworth
Michigan Technological University
The Covalent Bond
•
Covalent bonds are formed by sharing at least
one pair of electrons.
Prentice Hall ©2004
Chapter 07
The Covalent Bond
•
01
Slide 2
02
Every covalent bond
has a characteristic
length that leads to
maximum stability.
•
This is the bond
length.
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Chapter 07
Slide 3
1
The Covalent Bond
•
03
Energy required to break a covalent bond in an isolated
gaseous molecule is called the bond dissociation energy.
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Chapter 07
Polar Covalent Bonds
•
Bond polarity is due
to electronegativity
differences between
atoms.
•
Pauling
Electronegativity : is
expressed on a scale
where F = 4.0
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Chapter 07
Polar Covalent Bonds
Slide 4
01
Slide 5
02
Pauling Electronegativities
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Chapter 07
Slide 6
2
Polar Covalent Bonds
•
03
% Ionic Character: As a general rule for two atoms
in a bond, we can calculate an electronegativity
difference ( ?EN ): ?EN = EN(Y) – EN(X) for X–Y
bond.
If ?EN < 0.5 the bond is covalent.
If ?EN < 2.0 the bond is polar covalent.
If ?EN > 2.0 the bond is ionic.
Prentice Hall ©2004
Chapter 07
Slide 7
Polar Covalent Bonds
•
04
Using electronegativity values, predict whether the
following compounds are nonpolar covalent, polar
covalent, or ionic:
SiCl4
CsBr
FeBr 3
CH4
HCl
CCl4
NH3
H2O
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Chapter 07
Electron-Dot Structures
•
Slide 8
01
Using electron-dot (Lewis) structures, the valence
electrons in an element are represented by dots.
•
Valence electrons are those electrons with the
highest principal quantum number ( n).
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Chapter 07
Slide 9
3
Electron-Dot Structures
02
•
The electron-dot structures provide a simple, but
useful, way of representing chemical reactions.
•
Ionic:
•
Covalent:
Prentice Hall ©2004
Chapter 07
Slide 10
Electron-Dot Structures
•
Single Bonds:
C
H
H
H
H
03
H
C
H
H
H
H
H
•
Double Bonds:
C
C
C
C
H
H
•
Triple Bonds:
Prentice Hall ©2004
C
C
H
C
C
Chapter 07
Drawing Lewis-Dot Structures
H
Slide 11
01
•
Rule 1: Count the total valence electrons.
•
Rule 2: Draw the structure using single bonds.
•
Rule 3: Distribute the remaining electron pairs
around the peripheral atoms.
•
Rule 4: Put remaining pairs on central atom.
•
Rule 5: Share lone pairs between bonded atoms to
create multiple bonds.
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Chapter 07
Slide 12
4
Drawing Lewis-Dot Structures
•
02
NH2 F Amino Fluoride: In this molecule, nitrogen
is the central atom.
•
Rule 1: Number of electrons = 5 + (2 x 1) + 7 = 14 = 7 pairs
H
N H
H N
H
H
F
F
Rule 2
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Rule 3
N H
F
Rule 4
Chapter 07
Drawing Lewis-Dot Structures
Prentice Hall ©2004
Chapter 07
Drawing Lewis-Dot Structures
Slide 13
03
Slide 14
04
•
Polyatomic molecules with central atoms below
the second row ten:
•
In this compound there are 10 valence electrons on
bromine; this is called an expanded octet. The
extra pairs go into unfilled d orbitals.
Prentice Hall ©2004
Chapter 07
Slide 15
5
Drawing Lewis-Dot Structures
•
•
05
Draw electron- dot structures for:
C3 H8
H2 O2
CO2
N2 H4
CH5 N
C2 H4
C2 H2
Cl2 CO
Draw electron- dot structures for:
SF4
SF6
XeOF4
XeF4
H3 S+
HCO3 –
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XeF5+
Chapter 07
Slide 16
Resonance Structures
•
How is the double bond formed in O3?
Move lone pair from
this oxygen?
O
O
O
O
O
or
O
Or from this
oxygen?
•
01
O
O
O
The correct answer is that both are correct,
but neither is correct by itself.
Prentice Hall ©2004
Chapter 07
Slide 17
Resonance Structures
•
02
When multiple structures can be drawn, the actual
structure is an average of all possibilities.
•
The average is called a resonance hybrid. A
straight double-headed arrow indicates resonance.
O
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O
O
O
Chapter 07
O
O
Slide 18
6
Resonance Structures
•
03
The nitrate ion, NO3 –, has three equivalent oxygen
atoms, and its electronic structure is a resonance
hybrid of three electron-dot structures. Draw them.
•
Draw as many electron-dot resonance structures
as possible for: SO2, CO3 2– , HCO2 – , SO4 2– , PO 43–.
Prentice Hall ©2004
Chapter 07
Slide 19
Formal Charge
•
Formal Charge:
•
Determines the best
01
resonance structure.
•
We determine formal
charge and estimate
the more accurate
representation.
Prentice Hall ©2004
Chapter 07
Slide 20
Formal Charge
02



 # of bonding -e 

Formal Charge= # of Valence e-  − 
−  # of nonbonding -e

 
 

2


•
Calculate the formal charge and determine the
most favorable of the following electron dot
structures:
SO2
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NO3–
NCO–
Chapter 07
N2O
O3
CO32–
Slide 21
7
Molecular Shapes: VSEPR
•
01
The approximate
shape of molecules
is given by ValenceShell Electron- Pair
Repulsion (VSEPR).
Prentice Hall ©2004
Chapter 07
Molecular Shapes: VSEPR
Slide 22
02
•
Step 01: Count the total electron groups.
•
Step 02: Arrange electron groups to maximize
separation.
•
Groups are collections of bond pairs between two
atoms or a lone pair .
•
Groups do not compete equally for space:
Lone Pair > Triple Bond > Double Bond > Single Bond
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Chapter 07
Molecular Shapes: VSEPR
•
Slide 23
02
Two Electron Groups: Electron groups point in
opposite directions.
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Chapter 07
Slide 24
8
Molecular Shapes: VSEPR
•
Three Electron Groups: Electron groups lie in the
same plane and point to the corners of an
equilateral triangle.
Prentice Hall ©2004
Chapter 07
Molecular Shapes: VSEPR
•
03
Slide 25
06
Four Electron
Groups:
•
Electron groups
point to the
corners of a
regular
tetrahedron.
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Chapter 07
Molecular Shapes: VSEPR
Slide 26
09
Five Electron Groups: Electron groups point to the
corners of a trigonal bipyramid.
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Chapter 07
Slide 27
9
Molecular Shapes: VSEPR
•
11
Six Electron Groups:
Electron groups point to
the corners of a regular
octahedron.
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Chapter 07
Slide 28
Molecular Shapes: VSEPR
Electron Groups
Lone Pairs Bonds
12
Geometry
Examples
2
0
2
Linear
BeCl2
3
0
3
Trigonal planar
B F3
3
1
2
Bent
S O2
4
0
4
Tetrahedral
CH 4
4
1
3
Trigonal pyramidal
NH 3
4
2
2
Bent
H 2O
5
0
5
5
1
4
See- saw
S F4
5
2
3
T-Shaped
ClF3
5
3
2
linear
I3 -
6
0
6
Octahedral
S F6
6
1
5
Square pyramidal
SbCl5 2-
6
2
4
Square planar
XeF4
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Trigonal bipyramidal
PCl5
Chapter 07
Slide 29
Molecular Shapes: VSEPR
•
13
Draw the Lewis electron-dot structure and predict
the shapes of the following molecules or ions:
O3
H3 O+
XeF2
PF 6 –
XeOF4
AlH4 –
BF4 –
SiCl4
ICl4–
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Chapter 07
AlCl 3
Slide 30
10
Valence Bond Theory
01
•
1. Covalent bonds are formed by overlapping of
atomic orbitals , each of which contains one
electron of opposite spin.
•
2. Each of the bonded atoms maintains its own
atomic orbitals , but the electron pair in the
overlapping orbitals is shared by both atoms.
•
3. The greater the amount of orbital overlap, the
stronger the bond.
Prentice Hall ©2004
Chapter 07
Valence Bond Theory
•
Linus Pauling:
•
Wave functions from
Slide 31
02
s orbitals & p orbitals
could be combined to
form hybrid atomic
orbitals.
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Chapter 07
Valence Bond Theory
•
Slide 32
03
sp hybrid:
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Chapter 07
Slide 33
11
Valence Bond Theory
•
sp2 hybrid:
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Chapter 07
Valence Bond Theory
•
Slide 34
05
sp2 hybrid (p bond):
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Chapter 07
Valence Bond Theory
•
04
Slide 35
06
sp3 hybrid:
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Chapter 07
Slide 36
12
Valence Bond Theory
•
sp3d hybrid:
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Chapter 07
Valence Bond Theory
•
Slide 37
08
sp3d 2 hybrid:
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Chapter 07
Molecular Orbital Theory
•
07
Slide 38
01
The molecular orbital (MO) model provides a
better explanation of chemical and physical
properties than the valence bond (VB) model .
•
Atomic Orbital: Probability of finding the electron
within a given region of space in an atom.
•
Molecular Orbital: Probability of finding the
electron within a given region of space in a
molecule.
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Chapter 07
Slide 39
13
Molecular Orbital Theory
•
Additive combination of orbitals ( σ) is lower in
energy than two isolated 1s orbitals and is called a
bonding molecular orbital .
Prentice Hall ©2004
Chapter 07
Molecular Orbital Theory
•
Slide 40
03
Subtractive combination of orbitals ( σ∗) is higher
in energy than two isolated 1s orbitals and is called
an antibonding molecular orbital .
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Chapter 07
Molecular Orbital Theory
•
02
Slide 41
04
Molecular Orbital Diagram for H2 :
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Chapter 07
Slide 42
14
Molecular Orbital Theory
•
Molecular Orbital Diagrams for H2– and He 2 :
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Chapter 07
Molecular Orbital Theory
•
05
Slide 43
06
Additive and subtractive combination of p orbitals
leads to the formation of both sigma and pi orbitals.
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Chapter 07
Molecular Orbital Theory
•
Slide 44
07
Second- Row MO Energy Level Diagrams:
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Chapter 07
Slide 45
15
Molecular Orbital Theory
•
MO Diagrams Can Predict Magnetic Properties:
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Chapter 07
Molecular Orbital Theory
•
08
Slide 46
09
Bond Order is the number of electron pairs shared
between atoms.
•
Bond Order is obtained by subtracting the number
of antibonding electrons from the number of
bonding electrons and dividing by 2.
Prentice Hall ©2004
Chapter 07
Molecular Orbital Theory
Slide 47
10
•
Construct an MO diagram for He2+ ion. Determine
the bond order and whether the ion is likely to be
stable?
•
The B2 and C2 molecules have MO diagrams
similar to N2. What MOs are occupied in B2 and C2,
and what is the bond order in each? Would any of
these be paramagnetic?
Prentice Hall ©2004
Chapter 07
Slide 48
16