Chemistry 4362

1
Chapter 9 Chemical Bonding I Lewis Theory
Read/Study Chapter 9 MGC Homework Due April
17, 2008 at 1150 p.m. MGC Quiz Due by April
20, 2008 at 1150 p.m. Chapters 7, 8, and 9
2
Chapters 9 and 10

• Chemical Bond Types
• Ionic
• Pure Covalent
• Polar Covalent
• Coordinate Covalent
• Metallic

• Properties of Chemical Bonds
• Bond Length
• Bond Strength
• Bond Angles
• Bond Order

3
Chapters 9 and 10

• Chemical Bond Types
• Ionic
• Pure Covalent
• Polar Covalent
• Coordinate Covalent
• Metallic

4

• Your Basic Chemical Bonding Tool Kit
• Lewis Symbols
• The Octet Rule
• Lewis Structures
• Resonance Structures
• Formal Charges

Chapter 9

• Your Advanced Chemical Bonding Tool Kit
• Valence Shell Electron Pair Repulsion
• Valence Bond Theory (VB)
• Molecular Orbital Theory (MO)

Chapter 10
5

• Key Questions
• Why do atoms combine to form
• compounds and molecules?
• Why do they combine in definite
• proportions by mass?
• Why do molecules have characteristic
• shapes?

• Important Terms
• Valence - The combining capacity of an
• element.
• Chemical Bond - A strong link among
• atoms in a molecule or crystal.

6
Overview of Bonding Types
Ionic Bonding - An ionic bond is a chemical bond
that results from an electrostatic
attraction among oppositely charged ions in a
compound. They form when electrons are
transferred from one atom to another to form ions.
Na Ne 3s1 Cl Ne 3s2 3p5
Na Ne Cl Ne-
Cl-
Chapter 9
Na
7
Overview of Bonding Types
Covalent Bonding - A covalent bond is a chemical
bond that results from a sharing of electrons
among the atoms in a compound.
e-
e-



Energy
H H
Chapter 9
e-
H2
e-


8
Overview of Bonding Types
Polar Covalent Bonding - A covalent bond
that occurs when the atoms unequally share one
or more pairs of electrons. This happens when
the atoms have different electronegativities.
e- e-
e- e-
F

H
e-
e-
Energy
e- e-
e- e-
Chapter 9
e- e-
e- e-
d
d-
H
F
e- e-
9
Overview of Bonding Types
Coordinate Covalent Bonding (Dative) - A
dative bond is a covalent bond that occurs when
the two shared electrons are donated to the bond
by the same atom. The donating atom is the
donor or Lewis Base and the accepting atom is
the acceptor or Lewis Acid.
..
..
F
F
F BF F
..
..
..
..
..
..
..
F BF

..
..
..
..
..
F
Donor
..
..
Acceptor
Tetrafluoroborate ion
10
Overview of Bonding Types
Metallic Bonding - The force of attraction
that holds metal atoms together in a metallic
lattice. It results from the fact that the
valence electrons (outer shell electrons) are NOT
bound to a particular atom but are free to be
shared by all of the atoms - they are delocalized.

11
Summary of Bonding Types
Ionic Bonding
Covalent Bonding
Polar Covalent Bonding
Electropositive Electronegative
Electronegative Electronegative
Metallic Bonding
Electropositive Electropositive
12
Assignment
State whether each of the following has ionic,
covalent, polar covalent, or metallic bonding
H2O NaI S4 HCl
Polar
Polar
Ionic
Covalent
K NH4Cl Rb C (diamond)
Metallic
Dative Polar and ionic
Metallic
Covalent Network
13

• Your Basic Chemical Bonding Tool Kit
• Lewis Symbols
• The Octet Rule
• Lewis Structures
• Resonance Structures
• Formal Charges

Chapter 9

• Your Advanced Chemical Bonding Tool Kit
• Valence Shell Electron Pair Repulsion
• Valence Bond Theory (VB)
• Molecular Orbital Theory (MO)

Chapter 10
14
Lewis Electron Dot Symbols
A symbol for an atom or ion consisting of
the chemical symbol for the element surrounded by
a number of dots equal to the number of valence
electrons in the atom.
Alkaline Earth Metals!!
H 1s1 H Li He 2s1 Li Na Ne 3s1 Na K
Ar 4s1 K Rb Kr 5s1 Rb Cs Xe 6s1 Cs
Be He 2s2 Be Mg Ne 3s2 Mg Ca Ar
4s2 Ca Sr Kr 5s2 Sr Ba Xe 6s2 Ba
15
Lewis Electron Dot Symbols
..
Assignment
F He 2s22p5 F Cl Ne 3s23p5 Cl Br Ar
4s24p5 Br I Kr 5s25p5 I
.
Draw the Lewis Symbols for the following -
..
.
..
C P Tl Ge As Po
.
..
.
Used Primarily with Main Group Elements
16
Ionic Bonding
Ionic Bonding - An ionic bond is a chemical bond
that results from an electrostatic
attraction among oppositely charged ions in a
compound. They form when electrons are
transferred from one atom to another to form ions.
Na Ne 4s1 Cl Ne 3s2 3p5
NaNe Cl Ne-
Cl-
Na
17
Ionic Bonding - Whats yours is mine.
Whats mine you can have
1. Metal Ion Formation - Metals in Groups 1 and
2 tend to give up one and two electrons,
respectively, to form 1 and 2 ions that are
isoelectronic with the preceding noble gas.
Positive ions are called cations.
Li
Li e-
He 2s1
He
Sr2 2 e-
Sr
2
Kr 5s2
Kr
18
2. Negative Ion Formation - The atoms of
non- metals in Groups 15, 16, and 17 tend to
gain electrons as to fill their valence shell s
and p orbitals. Thus, they become isoelectronic
with the noble gas at the end of their respective
period.
.. Br e-
.. Br
-
.
..
2-
.. Se
.. Se 2 e-
..
. P 3 e-
.. P
3-
..
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3. The Octet Rule - A rule expressing
the tendency of some main group elements to
obtain a total of 8 electrons in their valence
shell. There are MANY exceptions!
Sn2 2e- Kr 4d10 5s2
Sn Kr 4d10 5s2 5p2
Tl e- Xe 4f14 5d106s2
Tl Xe 4f14 5d106s2 6p1
20

• 4. Formation of Ionic Bonds
• Transfer of electrons to form ions.
• Electrostatic attraction among the ions to
• form an ionic crystal lattice.

.. Ca O
.. Ca2 O
2-
..
-

-
-
66 Coordination Octahedral Arrangement
-

-
-
-
21
Electrolytes
22
Covalent Bonding - Share and share alike!
Covalent bonding involves the sharing of
electrons.
1. Lewis Structures of Molecules H
H HH H O
H HOH
Octet
..
..
..
..
Shared Lone Pairs Pairs
The total number of electrons in a Lewis
structure of a molecule is the sum of the valence
electrons in the individual atoms.

23
H
..
H


HOH CC HCCH
..


H
H
Single Bonds
Double Bond
Triple Bond
H
H
H-O-H CC H-C?C-H
H
H

• Find the number of electrons in the Lewis
• structure by adding up the valence electrons
• of all the atoms in the molecule or ion add one
• extra electron for each negative charge subtract
• one electron for each positive charge.

24
1 C _at_ 4 valence electrons 4 V.E. 2 O _at_ 6
valence electrons 12 V.E. Total Valence
Electrons 16 V.E.
CO2

• Draw a skeletal structure of the molecule or
• ion by arranging atoms and putting one single
• bond between atoms that are bonded to each
• other.

O-C-O or OCO

• Distribute the remaining electrons so as to
• satisfy the octet rule as closely as possible.

OCO
..
..
25
..
But what about.
OCO ???
..
Both structures obey the octet rule. Which
one is RIGHT???
2. Formal Charges - A positive, negative, or
zero value assigned formally to the atoms in a
Lewis structure it is calculated for each atom
using the following formula Formal Charge V.E.
- (unshared electrons) 1/2(shared electrons)
Purpose - To provide a method for predicting
the best Lewis structure for molecules with
more than one bonding possibility.
26
..
..
..
..
..
..
or OS-O
O-SO
..
..
FCOL 6 - 6 1/2(2) 6 - 6 1 6 - 7 -1
FCS 6 - 2 1/2(6) 6 - 2 3 6 - 5 1
FCOR 6 - 4 1/2(4) 6 - 4 2 6 - 6 0
..
..
..
..
..
..
O-SO
OS-O
..
..
-1 1
1 -1
Resonance Structures
27

• Rules for Using Formal Charges
• The most stable structures have the least
• formal charge.
• Structures in which adjacent atoms have
• formal charges of the same sign tend to be
• unstable.
• Structures in which positive charges are on
• more electronegative atoms are not as stable.

Since both of the SO2 structures are
equivalent (not identical), they both
contribute equally and should be shown as
resonance forms.
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3. Resonance Structures - Two or more
equivalent ways to depict the bonding in a
molecule or ion two or more equivalent and
legitimate Lewis structures for the same molecule
or ion.
Assignment
Draw all appropriate resonance structures for
the nitrate ion, NO3-, the carbonate ion,
CO3-, and the sulfur trioxide molecule, SO3.
Class Exercise Construct appropriate Lewis
structures for CO2, calculate formal charges
for the atoms, and determine if resonance
structures are important.
29

• 4. Exceptions to the Octet Rule
• Atoms with more than 8 electrons - this is
• possible for elements in row 3 and beyond in
• the periodic table due to the fact that
• d-orbitals become available to handle addi-
• tional electrons.

..
..
Cl Cl P Cl Cl Cl
..
F F F S F
F F
..
..
..
..
..
..
..
..
..
..
..
..
..
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• Atoms with fewer than 8 electrons - This
• occurs in electron deficient compounds.

..
F B F F
1 - 1
F B F F
..
..
..
..
..
..
..
..
Cl Be Cl Cl Be Cl
..
..
1 - 2 1
31

• Free Radicals - An atom, molecule, or ion
• that contains an odd number of
• unpaired electron.

..
..
..
O - Cl - O
..
..
.
32

• Polar Covalent Bonding - All are equal but
• some are more equal than others
• Bond Polarity - The result of an uneven charge
• distribution between two atomic nuclei that
• bonded to each other. It is due to the fact that
• different elements have different levels of
• attraction - electronegativity - for the electron
• pairs being shared.
• Non-Polar Bond - A pure covalent bond
• wherein the two atoms sharing the electrons
• have identical attraction for them - they have
• the same electronegativity.

33
.. ..
HH ClCl NN
.. ..
Non-Polar Bonds (Pure Covalent Bonds)

• Electronegativity - The general tendency of an
• atom or group of atoms to attract SHARED
• electrons.
• 1. Electronegativity Scales - Relative values
• assigned to the elements in the Periodic
  Table
• to represent their attraction for electrons in
• a CHEMICAL BOND. (This is NOT the same
• as ELECTRON AFFINITY.)

34
2. Electronegativity Differences - The greater
the difference in electronegativity between
two atoms that are bonded together, the more
polar the bond will be.
E.N. for Cl 3.2 E.N. for H 2.2 D E.N.
for HCl 1.0
?
? -
H - Cl
Pauling Scale
3. Polar Molecules - A molecule is polar IF it
has one or more polar bonds AND is
unsymmetrical in charge distribution.
35
Assignment
State whether or not each of the
following molecules is polar N2 H2O SF6
CCl4 SO2
Non-Polar Polar Non-Polar Non-Polar Polar
36
Coordinate Covalent Bonding (Dative Bonding) - I
have plenty so you may share
This type of covalent bonding involves a
sharing of two electrons that have BOTH been
provided by only one of the atoms.
H O H H3O H
..
H3NBF3
H3N BF3
37
Chapter 10 Chemical Bonding II Molecular Shapes,
Valence Bond Theory, and Molecular Orbital Theory
Read/Study Chapter 10 MGC Homework Due April
23, 2008 at 1150 p.m. MGC Quiz Due April 25,
2008 at 1150 p.m.
38

• Your Basic Chemical Bonding Tool Kit
• Lewis Symbols
• The Octet Rule
• Lewis Structures
• Resonance Structures
• Formal Charges

Chapter 9
Chapter 10

• Your Advanced Chemical Bonding Tool Kit
• Valence Shell Electron Pair Repulsion
• Valence Bond Theory (VB)
• Molecular Orbital Theory (MO)

39
Valence Shell Electron Pair Repulsion A
model of chemical bonding that allows the shapes
of molecules to be predicted by making the
logical assumption that electron pairs in
molecules tend to stay as far apart as possible.
VSEPR is a sophisticated use of Lewis
structures to determine the geometry of
polyatomic ions and molecules.
40

• Central Atom - An atom in a molecule or
• ion that is bonded to two or more other
• atoms.
• Molecules with one Central Atom
• 1. Write the Lewis structure.

..
..
O C O

• 2. Count the VSEPR electron pairs on the
• Central Atom.

A. Count each bond - single, double, or triple -
as ONE VSEPR pair.
41
B. Count each lone pair as one VSEPR pair.
..
..
Total of 2 VSEPR pairs on Carbon!
O C O
C. Place VSEPR pairs around the Central Atom so
that they are as far apart as possible.
C 180o
42
D. Using ATOMS, not electron pairs, determine
the geometry.
O - C - O
Geometry is Linear!
E. In determining geometry, take
into consideration the electron pair
re- pulsions (1) Bond Pair - Bond Pair
Repulsion (2) Bond Pair - Lone Pair
Repulsion (3) Lone Pair - Lone Pair Repulsion
BP-BP lt BP-LP lt LP-LP
Greater Repulsion
43
Geometry Types
VSEPR Pairs Geometry 2
Linear 3 Trigonal Planar
4 Tetrahedral 5
Trigonal Bipyramidal 6 Octahedral
44
Assignment
Determine the geometry of the following
species Ammonia - NH3 Sulfur Hexafluoride -
SF6 Nitrate Ion - NO3- Water - H2O Ammonium
Ion - NH4 Sulfur Dioxide - SO2 Iodine
Heptafluoride - IF7 Chlorite Ion - ClO2-
45

• Your Basic Chemical Bonding Tool Kit
• Lewis Symbols
• The Octet Rule
• Lewis Structures
• Resonance Structures
• Formal Charges

Chapter 9
Chapter 10

• Your Advanced Chemical Bonding Tool Kit
• Valence Shell Electron Pair Repulsion
• Valence Bond Theory (VB)
• Molecular Orbital Theory (MO)

46
Valence Bond Theory

• 1. Limitations of Lewis Structures and VSEPR
• Gives only information about geometry.
• Is based on the octet rule which has many
• exceptions.
• Cannot adequately explain bonding in
• species such as Li2 and H2.
• Does not reflect the QUANTUM nature of
• electrons.

47

• 2. Quantum Mechanical Theories of Bonding
• Valence Bond Theory - Involves the over-
• lapping of atomic orbitals from THE SAME
• atom.
• Molecular Orbital Theory - Involves the
• formation of MOLECULAR orbitals around
• two or MORE nuclei in a molecule. The
• molecular orbitals are formed by the over-
• lapping of atomic orbitals from DIFFERENT
• atoms.

48
3. Valence Bond Theory
A. The Simplest View - Bonds are formed by the
simple overlap of atomic orbitals from two
different atoms.
Chapter 10
H
H
H 1s1 H-H s1s2 H 1s1
I 5 p1x
I 5 p1x
I-I s5p2
49
A single bond consists of 2 electrons of
opposite spin. The electrons are in a Sigma Bond
(s).
Sigma Bond - A bond resulting from the overlap of
two atomic orbitals from DIFFERENT
atoms, resulting in the build-up of electron
density along the interatomic axis.
500
Repulsion
H H
Energy
0
ATTRACTION
H H
74 pm - 436 kJ/mol
50
B. Orbital Hybridization
Cl Be Cl
He 2s2
Ne 3s2 3p5
Ne 3s2 3p5
How do we explain the bonding in BeCl2
using Valence Bond Theory? We must invoke.
51
Orbital Hybridization!!!
An Imaginary mixing process in which the
orbitals of an atom rearrange to form new atomic
orbitals called Hybrid Orbitals.
2p
2p
Energy
sp hybrid
s p
sp Hybridized Be Atom
2s
52
To explain what is known about the bonding
in BeCl2 using VB theory, we must find a way
to un-pair the paired electrons in Be to make
it suitable for bonding with Cl.
1. Hybridize the same number of atomic
orbitals as there are bonds to explain. 2.
Promote the appropriate number of electrons into
the hybrid orbitals. 3. Form bonds with other
atoms by overlapping the hybrid orbitals with
either simple atomic orbitals or hybrid orbitals
on the other atoms.
53
BF3
sp2 hybridization
Unhybridized p
2p
Energy
sp2 hybrids
s 2 p
sp2 Hybridized B Atom
2s
54
sp3 hybridization
NH3
Lone Pair
2p
Energy
sp3 hybrids
s 3 p
sp3 Hybridized N Atom
2s
55
sp3d hybridization
PF5
Unhybridized d orbitals
3d
3p
Energy
sp3d hybrids
s 3 p d
sp3d Hybridized P Atom
3s
56
sp3d2 hybridization
SF6
Unhybridized d orbitals
3d
3p
Energy
sp3d2 hybrids
s 3 p 2 d
sp3d2 Hybridized S Atom
3s
57

• Summary of Key Hybridizations
• sp - linear
• sp2 - trigonal planar
• sp3 - tetrahedral
• sp3d - trigonal bipyramidal
• sp3d2 - octahedral

58
C. Multiple Bonding (1) Review - Single bonds
are ALWAYS s - Bonds. The electron density
is predominantly between the nuclei on the
interatomic axis.
s-s s-bond
s1
s1
s-p s-bond
s1
p1
59
p1
p1
p-p s-bond
sp sp s-sp s-bonds
s1
s1
(2) Multiple Bonds - Consist of one s-bond and
one or more p-bonds bonds that form by
side-wise or lateral overlap of parallel
p-orbitals.
60
Double Bond Formation
H2CCH2
Unhybridized p orbital
2p
Energy
sp2 hybrids
s 2 p
2s
The sp2 hybrid orbitals form a trigonal
plane perpendicular to the unhybridized
p-orbital. They form s-bonds while two parallel
p-orbitals form p-bonds.
61
Assignment
Describe the structure and bonding in O2
using Lewis structures and valence bond theory.
Show all steps. Do these structures show any
unpaired electrons in dioxygen?
Use valence bond theory to explain why the
bond angles in ammonia are NOT 90o.
H N - H H
Why Not??
62

• Your Basic Chemical Bonding Tool Kit
• Lewis Symbols
• The Octet Rule
• Lewis Structures
• Resonance Structures
• Formal Charges

Chapter 9
Chapter 10

• Your Advanced Chemical Bonding Tool Kit
• Valence Shell Electron Pair Repulsion
• Valence Bond Theory (VB)
• Molecular Orbital Theory (MO)

63
4. Molecular Orbital Theory A.
Assumptions (1) During bonding, atomic
orbitals from DIFFERENT atoms are
transformed into new orbitals with
different shapes, energies, and electron
density distri- butions. (2) This is
brought about by the overlap- ping of
atomic orbitals among different atoms.
Chapter 10
64
(3) Molecular Orbitals are the allowed states
for an electron moving in the electric field
generated by two or more nuclei. The Aufbau
principle, the Pauli Exclusion principle, and
Hunds Rule of Maximum Multiplicity are all
used to fill Molecular Orbitals.
B. Rules (1) The total number of molecular
orbitals is the same as the number of atomic
orbitals combined.
65
(2) Bonding molecular orbitals have lower
energy than the parent atomic orbitals while
antibonding orbitals have higher energy than
the parent atomic orbitals.
s1s
1s
1s
Atomic Orbitals
Atomic Orbitals
s1s
Molecular Orbitals
66
s1s
H
H
1s
1s
s1s
(3) A molecule is stable with respect to its
atoms whenever the number of bonding electrons
is greater than the number of antibonding
electrons.
67
s1s
H2 (s1s)2
1s
1s
s1s
H2 (s1s)1
H2- (s1s)2 (s1s)1
68
He2 (s1s)2 (s1s)1
1s
1s
He2 (s1s)2 (s1s)2
Non-Bonding
69
C. Bond Order - Bond Order ½(bonding
electrons - antibonding electrons)
H2 (s1s)2 B.O. 1/2(2 - 0) 1
H2 (s1s)1 B.O. 1/2(1 - 0) 1/2
H2- (s1s)2(s1s)1 B.O. 1/2(2 - 1)
1/2
He2 (s1s)2(s1s)1 B.O. 1/2(2 -
1) 1/2
He2 (s1s)2(s1s)2 B.O. 1/2(2 -
2) 0
70
D. Homonuclear Diatomic Molelcules of n 2
The n 2 shell has n2 or 4 atomic
orbitals. Therefore, two identical atoms can form
8 molecular orbitals.
2p
2p
Li2 - N2 pattern
2s
2s
1s
1s
71
s2p
ppy, ppz
ppy, ppz
O2 - Ne2 pattern
s2px
s2s
s2s
O2 is paramagnetic!!
s1s
s1s
(s1s)2 (s1s) 2 (s2s) 2 (s2s) 2 (s2p) 2 (p2p) 2
(p2p) 2
72
The molecular orbital diagrams for
hetero- nuclear diatomic molecules similar to
those of homonuclear diatomic molecules.
However, the atomic orbital energy levels are
different, thus causing the molecular orbital
diagrams to be unsymmetrical.
2s
2s
O atom
C atom
CO Molecule
73
Some Final Thoughts. 1. So which is the most
accurate picture of the chemical bond? Is
it Lewis Structures with its resonance
forms, formal charges, and the often
violated octet rule? 2. Maybe VSEPR with its
predictive powers is best! 3. But, of course,
the Valence Bond theory uses quantum
mechanical concepts and gives us more
information about what happens to atomic
orbitals during bond formation.
74
4. Yes, but the Molecular Orbital Theory is
more sophisticated and intellectually
satisfying. 5. Perhaps we should use them all
in different contexts for different
purposes. After all, they each are a window
on the world of molecular reality. We need
to understand them ALL their limitations and
their strengths.
Which one should you use? The simplest one that
will answer the question you are asking!!
75
Any Questions on Chapters 9 and 10?