CHAPTER 9
• Structure and
Molecular
Bonding
1
Introduction
• Bonds - Attractive forces that hold
atoms together in compounds
• Valence Electrons - The electrons
involved in bonding are in the
outermost (valence) shell.
2
Valence
Valence Electrons
Electrons
Electrons are divided between core and
valence electrons
B 1s2 2s2 2p1
Core = [He] , valence = 2s2 2p1
Br [Ar
[Ar]] 3d10 4s2 4p5
Core = [Ar
[Ar]] 3d10 , valence = 4s2 4p5
3
Rules of the Game
No. of valence electrons of a
main group atom = Group
number
•For Groups 1A1A-4A, no. of bond pairs =
group number.
• For Groups 5A -7A, BP’
BP’s = 8 - Grp.
Grp. No.
4
Rules of the Game
•No. of valence electrons of an atom = Group
number
•For Groups 1A1A-4A, no. of bond pairs = group
number
• For Groups 5A -7A, BP’
BP’s = 8 - Grp.
Grp. No.
•Except for H (and sometimes atoms of
3rd and higher periods),
BP’
BP’s + LP’
LP’s = 4
This observation is called the
OCTET RULE
5
Lewis Dot Formulas
of Atoms
.
H
.
Li
..
Be
..
He
..
..
..
..
..
..
. .. O . .. F . .. N e ..
B. .C. . N
.
.
..
..
6
Ionic Bonding
Formation of Ionic Compounds
•
An ion is an atom or a group of atoms
possessing a net electrical charge.
1. positive (+) ions or cations
•
These atoms have lost 1 or more electrons.
2. negative (-) ions or anions
•
These atoms have gained 1 or more electrons.
7
Formation of
Ionic Compounds
• Monatomic ions consist of one atom.
• Examples:
– Na+, Ca2+, Al3+ - cations
– Cl-, O2-, N3- -anions
• Polyatomic ions contain more than one
atom.
– NH4+ - cation
– NO2-,CO32-, SO42- - anions
8
Formation of
Ionic Compounds
• Reaction of Group IA Metals with
Group VIIA Nonmetals
G - 1 metal G - 17 nometal
2 Li (s) + F2(g) → 2 LiF(s)
silver
yellow
white solid
solid
gas
with an 842 o C
melting point
9
Formation of Ionic Compounds
1s
Li ↑↓
F ↑↓
2s
↑
↑↓
2p
↑↓↑↓↑
These atoms form ions with these configurations.
Li+ ↑↓
F- ↑↓
Li .
[He]
[Ne]
same configuration as
↑↓ ↑↓ ↑↓ ↑↓
+
same configuration as
..
.. .
..F
+
Li
..
[.. F.. ..]
10
Formation of Ionic Compounds
• General trend
• Cations become isoelectronic with the
preceding noble gas.
• Anions become isoelectronic with the
following noble gas.
11
Formation of
Ionic Compounds
• In general for the reaction of 1 metals and 17
nonmetals, the reaction equation is:
2 M(s) + X2 → 2 M+ X-(s)
– where M is the metals Li to Cs
– and X is the nonmetals F to I.
M
X
Electronically this is occurring.
ns
np
ns
np
↑
→ M+
↑↓ ↑↓ ↑↓ ↑
→ X- ↑↓ ↑↓ ↑↓ ↑↓
12
Formation of
Ionic Compounds
• Next we examine the reaction of 2
metals with 17 nonmetals.
• One example is the reaction of Be and
F2.
Be(s) + F2(g) →BeF2(g)
13
Formation of
Ionic Compounds
• The valence electrons in these two
elements are reacting in this fashion.
2s
2p
2s
2p
2+
Be [He] ↑↓
→ Be
F [He] ↑↓ ↑↓ ↑↓ ↑ → F- ↑↓ ↑↓ ↑↓ ↑↓
14
Formation of
Ionic Compounds
Be ..
..
.F
..
..
. F
..
..
2+
Be
..
..
2 .. F ..
..
• The remainder of the 2 metals and 17nonmetals
react similarly.
• Symbolically this can be represented as:
M(s) + X2 → M2+ X2M can be any of the metals Be to Ba.
X can be any of the nonmetals F to Cl.
15
Formation of
Ionic Compounds
• For the reaction of 1 metals with 16
nonmetals, a good example is the
reaction of lithium with oxygen.
• The reaction equation is:
4 Li(s) + O 2(g) → 2 Li 2+ O 2(s-)
16
Formation of
Ionic Compounds
• Draw the electronic configurations for Li, O,
and their appropriate ions.
You do it!
2s
2p
2s
2p
Li [He] ↑
→ Li1+
O [He] ↑↓ ↑↓ ↑ ↑
→ O2- ↑↓ ↑↓ ↑↓ ↑↓
Draw the Lewis dot formula representation of
this reaction.
17
Formation of
Ionic Compounds
Simple Binary Ionic Compounds Table
• Reacting Groups Compound General Formula
Example
1
2
3
1
2
3
+
+
+
+
+
+
17
17
17
16
16
16
MX
MX2
MX3
M2X
MX
M2X3
NaF
BaCl2
AlF3
Na2O
BaO
Al2S3
18
Formation of
Ionic Compounds
• Reacting Groups Compound General Formula
Example
1 + 15
M3X
Na3N
2 + 15
M3X2
Mg3P2
3 + 15
MX
AlN
H, a nonmetal, forms ionic compounds with 1
and 2 metals for example, LiH, KH, CaH2, and
BaH2.
Other hydrogen compounds are covalent.
19
Formation of
Ionic Compounds
• Coulomb’s Law – inverse square law
(q )(q )
+
F∝
−
2
d
where
F = force of attraction between ions
q = magnitude of charge on ions
d = distance between center of ions
20
Formation of
Ionic Compounds
• Force - Small ions with high ionic charges
>> large ions with small ionic charges
Al32+ O32- > Ca 2+ O 2- > K1+ Cl-
21
Covalent Bonding
• Atoms share electrons.
• If the atoms share 2 electrons a single
covalent bond is formed.
• 4 electrons - a double bond.
• 6 electrons - a triple bond.
• The atoms have a lower potential energy when bound.
22
Formation of Covalent Bonds
• This figure shows the potential energy of
an H2 molecule as a function of the
distance between the two H atoms.
23
Writing Lewis Formulas:
• 1. Sum the number of valence electrons for atoms
present.
• 2. Add or subtract electrons for the charge.
• 3. Identify the central atom (one that requires more e- to
complete octet – less e-neg if in same group) and draw a
skeletal structure.
• 4. Place a bond between each atom (2 e- per)
• 5. Fill in octet of outer atoms.
• 6. Complete octet of central atom – if deficient make
multiple bonds.
24
Formation of Covalent Bonds
• Use Lewis dot formulas
• 1. H molecule formation representation.
H.
2.
+
H .. H or H2
H.
HCl molecule formation
..
..
.
..
..
or HCl
H . Cl
H . + . Cl
..
..
25
Writing Lewis Formulas:
The Octet Rule
• Lewis octet rule - representative
elements usually attain stable noble gas
electron configurations in most of their
compounds.
• Distinguish between bonding (or shared)
electrons and nonbonding (or unshared or
lone pairs) of electrons.
26
Lewis Formulas for Molecules and
Polyatomic Ions
• Homonuclear diatomic molecules.
– 1. Two atoms of the same element.
1. Hydrogen molecule, H2.
H .. H
2.
or
Fluorine, F2.
.. ..
..
.
.
F . F . or
.. ..
3. Nitrogen, N2.
·· N ·· ·· ·· N ··
H H
..
..
..F
or
..
.
F.
..
·· N N ··
27
Lewis Formulas for Molecules and
Polyatomic Ions
• Next, look at heteronuclear diatomic
molecules.
–
1. hydrogen fluoride, HF
··
H .. F ··
··
2.
··
H F ··
··
hydrogen chloride, HCl
. ··
H . Cl ··
··
3.
or
or
··
H Cl··
··
hydrogen bromide, HBr
. ··
H . Br··
··
or
··
H Br··
··
28
Lewis Formulas for Molecules and
Polyatomic Ions
• Water, H2O
•Ammonia molecule , NH3
29
Lewis Formulas for Molecules and
Polyatomic Ions
•Water, H2O
··
H ·· O ··
··
H
• Ammonia molecule , NH3
··
H ·· N ·· H
··
H
30
Lewis Formulas for Molecules and
Polyatomic Ions
• Polyatomic ions.
• One example is the ammonium ion , NH4+.
H +
··
H ·· N ·· H
··
H
•Notice that the atoms other than H in these
molecules have eight electrons around them.
31
Writing Lewis Formulas:
The Octet Rule
• Example:
Write Lewis dot and dash
formulas for the sulfite ion, SO32-.
32
Writing Lewis Formulas:
The Octet Rule
• Sulfite ion, SO32-.
·· ··
·· O · S
·
·· ··
·· O
··
··
·· O ·· 2··
·
·
or
··
·· O
··
··
S
·· O ··
··
·· 2O ··
··
33
Double and
even triple
bonds are
commonly
observed for C,
N, P, O, and S
H2CO
SO3
C2F4
•
•
O
••
C
•
•
O
••
34
Lewis Structures
• Example:
Write Lewis dot and dash
formulas for sulfur trioxide, SO3.
35
Lewis Structures
• Example:
Write Lewis dot and dash
formulas for sulfur trioxide, SO3.
··
·· O · S ·· ·· O ··
·
·· ··
··
· O ··
·
··
or
··
·· O
··
S O ··
··
·· O ··
··
36
Resonance
• There are three possible structures for SO3.
–.
·O
·
··
S
·· O ·
·· ·
··
·· O
··
··
O ··
··
··
O ··
··
S
·O·
· ·
··
·O
·
··
S O ··
··
·· O ··
··
oTwo or more Lewis formulas are necessary to show the
bonding in a molecule, we must use equivalent resonance
structures to show the molecule’s structure.
oDouble-headed arrows are used to indicate resonance formulas.
37
Resonance
• Resonance is a flawed method of
representing molecules.
– There are no single or double bonds in
SO3.
S
O
O
O
38
Sulfur
Sulfur Dioxide,
Dioxide, SO
SO22
1. Central atom = S
2. Valence electrons = 18 or 9
pairs
O S O
••
••
••
•
•
•
•
••
••
3. Form double bond so that S has an octet —
but note that there are two ways of doing this.
OR bring in
right pair
bring in
left pair
••
•
•
O
••
••
S
••
•
•
O
••
39
Sulfur
Sulfur Dioxide,
Dioxide, SO
SO22
OR bring in
right pair
bring in
left pair
••
••
•
•
O
••
•
•
S
••
O
••
This leads to the following resonance
structures.
••
•
•
••
•
•
••
••
••
O S O
O S O
•
•
••
••
•
•
••
40
Writing Lewis Formulas:
Limitations of the Octet Rule
•
There are some molecules that violate the octet
rule.
–
1.
2.
3.
4.
For these molecules the N - A = S rule does not apply:
- Be.
- Group - 13.
-Odd number of total electrons.
-Central element must have a share of more than 8
valence electrons to accommodate all of the
substituents. (I.e. some S, P)
41
Writing Lewis Formulas:
Limitations of the Octet Rule
• Example: Write Lewis formula for BBr3.
··
.B
··
·· Br · B · ··
·
· Br ··
··
··
··
·· Br ·
·
··
··
·· Br .
··
or
··
·· Br
··
B
·· Br ··
··
··
Br ··
··
42
Sulfur Tetrafluoride
Tetrafluoride,, SF4
• Central atom =
• Valence electrons = ___ or ___
pairs.
• Form sigma bonds and distribute
electron pairs.
•
•
••
••
••
F
S
F
••
••
•
•
F
••
••
••
F
••
•
•
•
•
55 pairs
pairs around
around the
the SS
atom.
atom. A
A common
common
occurrence
occurrence outside
outside the
the
2nd
2nd period.
period.
43
Writing Lewis Formulas:
Limitations of the Octet Rule
• Example: Write dot and dash formulas
for AsF5.
44
Formal Atom Charges
• Atoms in molecules often bear a charge (+ or -).
• The predominant resonance structure of a
molecule is the one with charges as close to 0 as
possible.
• Formal charge
= Group number
– 1/2 (no. of bonding electrons)
- (no. of LP electrons)
45
Calculated Partial Charges in
CO22
Yellow = negative & red = positive
Relative size = relative charge
46
Thiocyanate Ion, SCN-
6 - (1/2)(2) - 6 = -1
5 - (1/2)(6) - 2 = 0
••
•
•
S
C
N
•
•
••
4 - (1/2)(8) - 0 = 0
47
Thiocyanate Ion, SCN-
••
••
•
•
S
C
N
•
•
•
•
••
S
C
N
•
•
••
••
•
•
S
C
N
•
•
••
Which is the most stable resonance form?
48
Calculated Partial Charges
in SCN-
All atoms negative, but
most on the S
••
•
•
S
C
N
•
•
••
49
Dipole Moments
• Asymmetric charge distribution
• The dipole moment has the symbol µ.
• µ is the product of the distance,d, separating
charges of equal magnitude and opposite
sign, and the magnitude of the charge, q.
50
Dipole Moments
• For example, HF and HI:
a
+
δ H - Fδ
a
-
1.91 Debye units
+
δ H -Iδ 0.38 Debye units
51
Dipole Moments
• There are some nonpolar molecules that have
polar bonds.
• There are two conditions that must be true for a
molecule to be polar.
1. There must be at least one polar bond present
or one lone pair of electrons.
2. The polar bonds, if there are more than one,
and lone pairs must be arranged so that their
dipole moments do not cancel one another.
3. Examples (water, CF4, CO2, NH3, NH4+)
52
Polar Molecules: The Influence of
Molecular Geometry
• Molecular geometry affects molecular
polarity.
– Due to the effect of the bond dipoles and
how they either cancel or reinforce each
other.
A
B A
linear molecule
nonpolar
A
B
A
angular molecule
polar
53
Polar Molecules: The Influence of
Molecular Geometry
•
Polar Molecules must meet two
requirements:
1. One polar bond or one lone pair of
electrons on central atom.
2. Neither bonds nor lone pairs can be
symmetrically arranged that their
polarities cancel.
54
Polar and Nonpolar Covalent Bonds
• Covalent bonds in which the electrons are
shared equally are designated as nonpolar
covalent bonds.
– Nonpolar covalent bonds have a symmetrical
charge distribution.
·· N ·· ·· ·· N ··
or
H .. H
·· N N ··
or
H H
55
Polar and Nonpolar Covalent Bonds
• Polar covalent bonds - electrons are not
shared equally -different
electronegativities.
H
F
Electronegativities 2.1
4.0
1424
3
1.9
Difference = 1.9 very polar bond
56
Polar and Nonpolar Covalent Bonds
• Compare HF to HI.
H
Electronegativities
I
2.1
2.5
1
424
3
0.4
Difference = 0.4 slightly polar bond
57
Two Simple Theories of Covalent
Bonding
• Valence Shell Electron Pair Repulsion Theory
– Commonly designated as VSEPR
– Principal originator
• R. J. Gillespie in the 1950’s
• Valence Bond Theory (Chapter 10)
– Involves the use of hybridized atomic orbitals
– Principal originator
• L. Pauling in the 1930’s & 40’s
58
VSEPR Theory
• VSEPR - electron density around the
central atom are arranged as far apart
as possible to minimize repulsions.
• Five basic molecular shapes
59
VSEPR Theory
1 Two regions of high electron density
around the central atom.
60
VSEPR Theory
2 Three regions of high electron density around the
central atom.
61
VSEPR Theory
3 Four regions of high electron density around the
central atom.
62
VSEPR Theory
4 Five regions of high electron density
around the central atom.
63
VSEPR Theory
5 Six regions of high electron density around the
central atom.
64
VSEPR Theory
1. Electronic geometry - locations of regions
of electron density around the central atom(s).
2. Molecular geometry - arrangement of
atoms around the central atom(s).
Electron pairs are not used in the
molecular geometry determination just
the positions of the atoms in the
molecule are used.
65
VSEPR Theory
• An example of a molecule that has the
same electronic and molecular
geometries is methane - CH4.
• Electronic and molecular geometries are
tetrahedral.
H
H C
H
H
66
VSEPR Theory
• An example of a molecule that has
different electronic and molecular
geometries is water - H2O.
• Electronic geometry is tetrahedral.
• Molecular geometry is bent or angular.
H
H C
H
H
67
VSEPR Theory
• Lone pairs of electrons (unshared pairs)
require more volume than shared pairs.
– Consequently, there is an ordering of repulsions
of electrons around central atom.
• Criteria for the ordering of the repulsions:
68
VSEPR Theory
1 Lone pair to lone pair is the strongest repulsion.
2 Lone pair to bonding pair is intermediate
repulsion.
3 Bonding pair to bonding pair is weakest
repulsion.
• Mnemonic for repulsion strengths
lp/lp > lp/bp > bp/bp
• Lone pair to lone pair repulsion is why bond
angles in water are less than 109.5o.
69
Molecular Shapes and Bonding
• In the next sections we will use the
following terminology:
A = central atom
B = bonding pairs around central atom
U = lone pairs around central atom
• For example:
AB3U designates that there are 3 bonding pairs
and 1 lone pair around the central atom.
70
Linear Electronic Geometry:AB2
Species (No Lone Pairs of Electrons
on A)
• Some examples of molecules with this
geometry are:
BeCl2, BeBr2, BeI2, HgCl2, CdCl2
• All of these examples are linear,
nonpolar molecules.
• Important exceptions occur when the
two substituents are not the same!
BeClBr or BeIBr will be linear and polar!
71
Trigonal Planar Electronic Geometry:
AB3 Species (No Lone Pairs of
Electrons on A)
• Some examples of molecules with this
geometry are:
BF3, BCl3
• All of these examples are trigonal planar,
nonpolar molecules.
• Important exceptions occur when the three
substituents are not the same!
BF2Cl or BCI2Br will be trigonal planar and polar!
72
Tetrahedral Electronic Geometry: AB4
Species (No Lone Pairs of Electrons
on A)
• Some examples of molecules with this
geometry are:
CH4, CF4, CCl4, SiH4, SiF4
• All of these examples are tetrahedral,
nonpolar molecules.
• Important exceptions occur when the four
substituents are not the same!
CF3Cl or CH2CI2 will be tetrahedral and polar!
73
Tetrahedral Electronic Geometry: AB4
Species (No Lone Pairs of Electrons on
A)
74
Tetrahedral Electronic Geometry:
AB3U Species (One Lone Pair of
Electrons on A)
• Some examples of molecules with this
geometry are:
NH3, NF3, PH3, PCl3, AsH3
• These molecules are our first examples of
central atoms with lone pairs of electrons.
Thus, the electronic and molecular geometries are
different.
All three substituents are the same but molecule is
polar.
• NH3 and NF3 are trigonal pyramidal, polar
molecules.
75
Tetrahedral Electronic Geometry:
AB2U2 Species (Two Lone Pairs of
Electrons on A)
• Some examples of molecules with this geometry
are:
H2O, OF2, H2S
• These molecules are our first examples of
central atoms with two lone pairs of electrons.
Thus, the electronic and molecular geometries are
different.
Both substituents are the same but molecule is polar.
• Molecules are angular, bent, or V-shaped and
polar.
76
Trigonal Bipyramidal Electronic
Geometry: AB5, AB4U, AB3U2, and
AB2U3
Some examples of molecules with this
geometry are:
PF5, AsF5, PCl5, etc.
• These molecules are examples of central
atoms with five bonding pairs of electrons.
The electronic and molecular geometries are the
same.
• Molecules are trigonal bipyramidal and
nonpolar when all five substituents are the
same.
If the five substituents are not the same polar
molecules can result, AsF4Cl is an example.
77
Trigonal Bipyramidal Electronic
Geometry: AB5, AB4U, AB3U2, and
AB2U3
•
If lone pairs are incorporated into the trigonal
bipyramidal structure, there are three possible
new shapes.
1. One lone pair - Seesaw shape
2. Two lone pairs - T-shape
3. Three lone pairs – linear
•
The lone pairs occupy equatorial positions
because they are 120o from two bonding pairs
and 90o from the other two bonding pairs.
–
Results in decreased repulsions compared to lone pair
in axial position.
78
•
Trigonal Bipyramidal Electronic
Geometry: AB5, AB4U, AB3U2, and
AB2U3
AB4U molecules have:
1. trigonal bipyramid electronic geometry
2. seesaw shaped molecular geometry
3. and are polar
•
One example of an AB4U molecule is
SF4
79
Trigonal Bipyramidal Electronic
Geometry: AB5, AB4U, AB3U2, and
AB2U3
Molecular Geometry
H
H C
H
H
80
Trigonal Bipyramidal Electronic
Geometry: AB5, AB4U, AB3U2, and
AB2U3
• AB3U2 molecules have:
1. trigonal bipyramid electronic geometry
2. T-shaped molecular geometry
3. and are polar
• One example of an AB3U2 molecule is
IF3
81
Trigonal Bipyramidal Electronic
Geometry: AB5, AB4U, AB3U2, and
AB2U3
Molecular Geometry
H
H C
H
H
82
Trigonal Bipyramidal Electronic
Geometry: AB5, AB4U, AB3U2, and
AB2U3
• AB2U3 molecules have:
1.trigonal bipyramid electronic geometry
2.linear molecular geometry
3.and are nonpolar
• One example of an AB3U2 molecule is
83
Trigonal Bipyramidal Electronic
Geometry: AB5, AB4U, AB3U2, and
AB2U3
Molecular Geometry
H
H C
H
H
84
Octahedral Electronic Geometry: AB6,
AB5U, and AB4U2
• Some examples of molecules with this
geometry are:
SF6, SeF6, SCl6, etc.
• These molecules are examples of central
atoms with six bonding pairs of electrons.
• Molecules are octahedral and nonpolar
when all six substituents are the same.
If the six substituents are not the same polar
molecules can result, SF5Cl is an example.
85
Octahedral Electronic Geometry: AB6,
AB5U, and AB4U2
•
If lone pairs are incorporated into the octahedral
structure, there are two possible new shapes.
1. One lone pair - square pyramidal
2. Two lone pairs - square planar
•
The lone pairs occupy axial positions because
they are 90o from four bonding pairs.
– Results in decreased repulsions compared to lone pairs
in equatorial positions.
86
Octahedral Electronic Geometry: AB6,
AB5U, and AB4U2
• AB5U molecules have:
1.octahedral electronic geometry
2.Square pyramidal molecular geometry
3.and are polar.
• One example of an AB4U molecule is
IF5
87
Octahedral Electronic Geometry: AB6,
AB5U, and AB4U2
Molecular Geometry
H
H C
H
H
88
Octahedral Electronic Geometry: AB6,
AB5U, and AB4U2
• AB4U2 molecules have:
1.octahedral electronic geometry
2.square planar molecular geometry
3.and are nonpolar.
• One example of an AB4U2 molecule is
XeF4
89
Octahedral Electronic Geometry: AB6,
AB5U, and AB4U2
Molecular Geometry
Polarity
H
H C
H
H
90
Compounds Containing
Double Bonds
• Ethene or ethylene, C2H4, is the
simplest organic compound containing a
double bond.
• Compound must have a double bond to
obey octet rule.
91
Compounds Containing
Double Bonds
Lewis Dot Formula
H ·
·
C ··
H ··
· H
·· C ·
·· H
H
H
C
or
C
H
H
92
Bond Properties
• What is the effect of bonding and
structure on molecular properties?
Free rotation
around C–
C–C single
bond
No rotation
around C=C
double bond
93
Bond Order
# of bonds between a pair of atoms
Double bond
Single bond
Acrylonitrile
Triple
bond
94
Bond Order
Fractional bond orders occur in molecules
with resonance structures.
••
••
N
N
Consider NO2-
•• • •••
••
••
O
O
O
O
••
•• • •••
••
The N—
—O bond order = 1.5
N—O
N
Bond order =
Total # of e - pairs used for a type of bond
Total # of bonds of that type
Bond order =
3 e - pairs in N — O bonds
2 N — O bonds
95
Bond
Bond Order
Order
Bond order is proportional to two important
bond properties:
(a) bond strength
(b) bond length
414 kJ
110 pm
123 pm
745 kJ
96
Bond Length
• Bond length is
the distance
between the
nuclei of two
bonded
atoms.
97
Bond Length
Bond length
depends on size
of bonded
atoms.
H—F
H—Cl
Bond
Bond distances
distances measured
measured
in
in Angstrom
Angstrom units
units where
where 11
A
A == 10
10--22 pm.
pm.
H—I
98
Bond Length
Bond length
depends on
bond order.
Bond
Bond distances
distancesmeasured
measured
in
in Angstrom
Angstrom units
unitswhere
where11
-2
AA== 10
10-2pm.
pm.
99
Bond Strength
• —measured by the energy req’d to break a bond. See
Table 9.10.
•
BOND
STRENGTH (kJ/mol)
H—H
436
C—C
346
C=C
602
C≡C
835
N≡N
945
The GREATER the number of bonds (bond order) the
HIGHER the bond strength and the SHORTER the
bond.
100
101
Bond Strength
• Measured as the energy req’d to break a
bond. See Table 9.10
102
Bond Strength
• Measured as the energy req’d to break a
bond. See Table 9.10.
•
BOND
STRENGTH (kJ/mol)
H—H
C—C
C=C
C≡C
N≡N
436
346
602
835
945
The GREATER the number of bonds (bond order) the
HIGHER the bond strength and the SHORTER the
bond.
103
Molecular Polarity
Water
Boiling point
= 100 ˚C
Methane
Boiling point
= -161 ˚C
Why do water and
methane differ so
much in their
boiling points?
Why do ionic compounds dissolve in
water?
104
Bond
Bond Polarity
Polarity
• Three molecules with
polar, covalent bonds.
• Each bond has one
atom with a slight
negative charge ((-δ)
and and another with
a slight positive
charge (+ δ)
105
Electronegativity, χ
χ is a measure of the ability of an
atom in a molecule to attract
electrons to itself.
Concept proposed by Linus Pauling 1901-1994
1901
1901-1994
106
Electronegativity
Figure 9.14
107
Molecular Polarity
Molecules will be polar if
a) bonds are polar
AND
b) the molecule is NOT “symmetric”
symmetric”
All above are NOT polar
108
Polar or Nonpolar?
Nonpolar?
Compare CO2 and H2O. Which one is
polar?
109
Polar or Nonpolar?
Nonpolar?
• Consider AB3 molecules: BF3, Cl2CO, and
NH3.
110
Molecular Polarity, BF33
F
B
F
F
B atom is
positive and
F atoms are
negative.
B
—F bonds in BF33 are polar.
B—F
But molecule is symmetrical
and NOT polar
111
Molecular Polarity, HBF22
H
B
F
F
B atom is
positive but
H & F atoms
are negative.
B
—F and B—
—H bonds in HBF22
B
B—F
B—H
are polar. But molecule is NOT
symmetrical and is polar.
112
Is Methane, CH44, Polar?
Methane is symmetrical and is NOT
polar.
113
Is CH33F Polar?
C
—F bond
C—F
bond is
is very
very polar.
polar.
Molecule
Molecule is
is not
not symmetrical
symmetrical
and
and so
so is
is polar.
polar.
114
CH4 … CCl4
Polar or Not?
•
Only CH4 and CCl4 are NOT polar. These are the only two molecules that are “symmetrical.”
115