CHAPTER 8
Bonding and Molecular Structure
Introduction
• Bonds: Attractive forces that hold atoms
together in compounds
• Valence Electrons: the outermost electrons
– -These e- are involved in bonding
2
Valence Electrons
Electrons are divided between core and
valence electrons
B 1s2 2s2 2p1
Core = [He] , valence = 2s2 2p1
Br [Ar] 3d10 4s2 4p5
Core = [Ar] 3d10 , valence = 4s2 4p5
3
Valence Electrons
-The number of valence electrons of a main
group atom is the Group number
-For Groups IA-IVA, number of bonding (unpaired)
electrons is equal to the group number
-For Groups VA -VIIA, number of bonding (unpaired)
electrons is equal to 8 - group number
4
Valence Electrons
-Except for H (and sometimes atoms of the 3rd
group and higher)
-The total number of valence electrons around
a given atom in a molecule will be eight:
OCTET RULE
- (with the exception of hydrogen) atoms in
molecules prefer to be surrounded by 8 electrons (or
have 4 bonds = 8 electrons)
5
Lewis Dot Formulas of Atoms
IA
IIA
IIIA
IVA
VA
VIA
VIIA
H
Li
VIIIA
He
Be
B C N O F
Ne
6
Ionic Bonding
An ion is an atom or a group of atoms
possessing a net electrical charge
•
-cations: positive (+) ions
•
These atoms have lost 1 or more electrons
1. -anions: negative (-) ions
•
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
Examples:
• NH4+ - cation
• NO2-,CO32-, SO42- - anions
8
Formation of Ionic Compounds
• General trend:
– metals become isoelectronic with the
preceding noble gas electron
configuration
– nonmetals become isoelectronic with the
following noble gas electron configuration
9
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
solid
yellow
gas
white solid
o
with an 842 C
melting point
10
Formation of Ionic Compounds
•
1s
2s
2p
Li   .
.
F   
These atoms form ions with these configurations.
Li+ 
F- 
Li .
.
   
+
..
..
.
F
..
[He]
same configuration as [Ne]
same configuration as
Li
+
..
.. ..
F
..
[ ]
11
Formation of Ionic Compounds
• In general:
the reaction of IA metals and VIIA nonmetals:
2 M(s) + X2  2 MX(s)
– where M is the metals Li to Cs
– and X is the nonmetals F to I
Electronically it looks like:
ns
np
ns
M

 M+ __
X
   
 X- 
np
__ __ __
  
12
Formation of Ionic Compounds
reaction of IIA metals with VIIA nonmetals:
Be(s) + F2(g) BeF2(g)
13
Formation of Ionic Compounds
The valence electrons in these two elements
react like:
2s
2p
2s
2p
Be [He]  __ __ __  Be2+ __ __ __ __
F [He]      F-    
Lewis dot structure representation:
14
Formation of Ionic Compounds
The remainder of the IIA metals and VIIA
nonmetals react similarly:
M(s) + X2  MX2
M can be any of the metals Be to Ba
X can be any of the nonmetals F to I
15
Formation of Ionic Compounds
For the reaction of IA metals with VIA
nonmetals:

2
4 Li (s)  O2(g)  2 Li O
2s 
Draw the valence electronic configurations for Li, O, and
their appropriate ions
16
Formation of Ionic Compounds
• Draw the electronic configurations for
Li, O, and their appropriate ions
2s
2p
Li [He] 
O [He]    
You do it!
2s
2p
 Li1+
 O2-    
Draw the Lewis dot formula representation of this
reaction
17
Formation of Ionic Compounds
Simple Binary Ionic Compounds Table
• Reacting Groups General Formula
IA + VIIA
IIA + VIIA
IIIA + VIIA
IA + VIA
IIA + VIA
IIIA + VIA
MX
MX2
MX3
M2X
MX
M2X3
Example
NaF
BaCl2
AlF3
Na2O
BaO
Al2S3
18
Formation of Ionic Compounds
• Reacting Groups General Formula
IA + VA
IIA + VA
IIIA + VA
M3X
M3X2
MX
Example
Na3N
Mg3P2
AlN
-H forms ionic compounds when bound to metals (IA and
IIA metals
For example: LiH, KH, CaH2, and BaH2
-When H is bound to nonmetals, the compounds are
covalent in nature
19
Formation of Covalent Bonds
• -potential energy of an H2 molecule as a
function of the distance between the two H
atoms
20
Covalent Bonding
• Atoms share electrons
•
•
•
•
If the atoms share:
2 electrons a single covalent bond is formed
4 electrons - a double covalent bond
6 electrons - a triple covalent bond
Atoms have a lower potential energy when
bound…this is a more favorable situation (why?)
21
Writing Lewis Formulas:
• 1. Add the number of valence electrons for all the atoms
that are present in the molecule
• 2. Add or subtract electrons based on the molecule’s (or
ion’s) charge
• 3. Identify the central atom and draw a skeletal
structure:
– -the one that requires the most e- to complete octet
– -the less electronegative
• 4. Place a bond between each atom (1 bond = 2 e-)
• 5. Fill in octet of outer atoms first
• 6. Finish by completing the octet of central atom
– – if you run out of e- then multiple bonds must be created
between the central atom and atoms bound to it
22
Writing Lewis Formulas
octet rule: representative elements usually
attain stable noble gas electron
configurations (8 valence e-) in most
compounds
You must distinguish the difference
between:
– -bonding electrons and nonbonding electrons
-shared (paired) and unshared (unpaired)
electrons
23
Formation of Covalent Bonds
• Lewis dot structures:
• 1. H2 molecule formation:
2. HCl molecule formation:
24
Lewis Structures
• Homonuclear diatomic molecules
– 1. Two atoms of the same element, H2:
H .. H
or
H H
2. Fluorine, F2:
3. Nitrogen, N2:
25
Lewis Structures
heteronuclear diatomic molecules
1. hydrogen fluoride, HF
2. hydrogen chloride, HCl
··
. ·· ·
H . Cl · or H Cl··
··
··
3. hydrogen bromide, HBr
26
Lewis Structures
• Water, H2O
•Ammonia molecule , NH3
27
Lewis Structures
• Polyatomic ions:
• ammonium ion NH4+
Notice that the N-atom in this molecule has eight
electrons around them (H does not)
28
Writing Lewis Formulas
• Sulfite ion, SO32-.
29
Double and even
triple bonds are
commonly
observed for C, N,
P, O, and S
H2CO
SO3
C2F4
30
Lewis Structures
• Example: Write Lewis dot and dash formulas
for sulfur trioxide, SO3
31
Resonance
• There are three possible structures for SO3:
·· O S
··
·· O ·
·· ·
·· ·
O·
··
··
·· O
··
S
·· O ··
··
O ··
··
··
·· O
S O ··
··
··
·· O ··
··
-Two or more Lewis formulas are necessary to show the
bonding in a molecule
-use equivalent resonance structures to show the molecule’s
structure
-Double-headed arrows are used to indicate resonance
formulas
32
Resonance
Resonance is a flawed method of
representing molecules
– -There are no single or double bonds in SO3
O
S
O
O
33
Sulfur Dioxide, SO2
1. Central atom =
2. Valence electrons = ___
or ___ pairs
3. Write the Lewis structure
4. Form double bond so that S has
an octet — but note that there are
two ways of doing this.
34
Limitations of the Octet Rule
•
There are some molecules that violate the octet
rule:
1. - Be
2. - Group IIIA
3. -Odd number of total electrons.
4. -Central element must have a share of more than
8 valence electrons to accommodate all of the
substituents. (i.e. S and P)
35
Limitations of the Octet Rule
• Example: Write Lewis formula for BBr3.
36
Sulfur Tetrafluoride, SF4
Central atom =
Valence electrons = ___ or ___
pairs.
Form sigma bonds and distribute
electron pairs.
5 pairs around the S
atom. A common
occurrence outside the
2nd period.
37
Limitations of the Octet Rule
• Example: Write dot structures for AsF5.
38
Formal Atomic Charges
• Atoms in molecules often bear a charge (+ or -)
• The predominant resonance structure of a
molecule is the one with charges on atoms as
close to 0 as possible
• Formal charge = Group number – 1/2 (# of
bonding electrons) - (# of Lone electrons)
•
•
= Group number – (# of bonds)
•
– (# of Lone electrons)
39
Formal Charge
CO2
..
..
..
..
40
Formal Charge
Thiocyanate Ion, SCN-
••
••
•
•
S
C
N
•
•
•
•
••
S
C
N
•
•
••
••
•
•
S
C
N
•
•
••
Which is the most stable resonance form?
41
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 9)
– Involves the use of hybridized atomic orbitals
– Principal originator
• L. Pauling in the 1930’s & 40’s
42
VSEPR Theory
electron densities around the central
atom are arranged as far apart as
possible to minimize repulsions (why?)
• Five basic molecular shapes:
• Linear, trigonal planar, tetrahedral, trigonal
bipyramidal, octahedral
43
VSEPR Theory
1. Two regions of high electron
density around the central atom.
44
VSEPR Theory
2. Three regions of high electron density
around the central atom.
45
VSEPR Theory
3. Four regions of high electron density
around the central atom.
46
VSEPR Theory
4. Five regions of high electron
density around the central atom.
47
VSEPR Theory
5. Six regions of high electron
density around the central atom.
48
VSEPR Theory
1. Electronic geometry(family):
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
49
VSEPR Theory
Lone pairs (unshared pairs) of electrons require more
volume than shared pairs
– -there is an ordering of repulsions of lone electrons
around central atom
Criteria for the ordering of the repulsions:
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.
50
Molecular Shapes and Bonding
• Symbolism:
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
51
Linear Electronic Geometry: AB2
Some examples of molecules with this
geometry:
BeCl2, BeBr2, BeI2, HgCl2, CdCl2
52
Trigonal Planar Electronic Geometry:
AB3
Some examples of molecules with this
geometry are:
BF3, BCl3
53
Tetrahedral Electronic Geometry: AB4
Some examples of molecules with this
geometry are:
CH4, CF4, CCl4, SiH4, SiF4
54
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
55
Tetrahedral Electronic Geometry: AB4
56
Tetrahedral Electronic Geometry: AB3U
Some examples of molecules with this
geometry are:
NH3, NF3, PH3, PCl3, AsH3
– -trigonal pyramidal
-electronic and molecular geometries are
different.
..
..
107.5°
57
Tetrahedral Electronic Geometry:
AB2U2
• Some examples of molecules with this
geometry are:
H2O, OF2, H2S
– -bent
-electronic and molecular geometries are different
104.5°
58
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
59
Trigonal Bipyramidal Electronic
Geometry: AB5, AB4U, AB3U2, and
AB2U3
 Some examples of molecules with this
geometry are:
PF5, AsF5, PCl5
axial
equatorial
axial
60
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.
2.
3.
One lone pair - Seesaw shape
Two lone pairs - T-shape
Three lone pairs – linear
The lone pairs occupy equatorial positions first:
-they are 120o from each other
-90o from the axial positions
– Results in decreased repulsions compared to lone pair in axial
position
axial
equatorial
61
axial
Trigonal Bipyramidal Electronic Geometry:
AB5, AB4U, AB3U2, and AB2U3
•
AB4U molecules have:
1. trigonal bipyramid electronic geometry
2. seesaw shaped molecular geometry
3. polar
•
One example of an AB4U molecule is
SF4
62
Trigonal Bipyramidal Electronic Geometry:
AB5, AB4U, AB3U2, and AB2U3
H
H C
H
H
63
Trigonal Bipyramidal Electronic Geometry:
AB5, AB4U, AB3U2, and AB2U3
•
AB3U2 molecules have:
1. 1. trigonal bipyramid electronic
geometry
2. T-shaped molecular geometry
3. polar
•
One example of an AB3U2 molecule is
IF3
64
Trigonal Bipyramidal Electronic Geometry:
AB5, AB4U, AB3U2, and AB2U3
H
H C
H
H
65
Trigonal Bipyramidal Electronic Geometry:
AB5, AB4U, AB3U2, and AB2U3
• AB2U3 molecules have:
1.trigonal bipyramid electronic geometry
2.linear molecular geometry
3.nonpolar
• One example of an AB3U2 molecule is
BrF2-
66
Trigonal Bipyramidal Electronic Geometry:
AB5, AB4U, AB3U2, and AB2U3
H
H C
H
H
67
Octahedral Electronic Geometry: AB6,
AB5U, and AB4U2
• Some examples of molecules with this
geometry are:
SF6, SeF6, SCl6, etc.
68
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 any position because they are all 90o
from all bonds positions:
–
–
-Additional lone pairs occupy the position 180º from the first set
of lone pairs
-This results in decreased repulsions compared to lone pairs in
the other positions
69
Octahedral Electronic Geometry: AB6,
AB5U, and AB4U2
• AB5U molecules have:
1.octahedral electronic geometry
2.Square pyramidal molecular geometry
3.polar.
• One example of an AB4U molecule is
IF5
70
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
71
Polarity and Electronegativity
Figure 8.11
72
Dipole Moments
• For example, HF and HI:

  H - F 1.91 Debye units

  H-I 0.38 Debye units
73
Dipole Moments
some “nonpolar molecules” that have polar
bonds
Two conditions to be polar:
1. 1. There must be at least one polar bond present or
one lone pair of electrons
2. 2. the molecule must be nonsymmetric
Examples: water, CF4, CO2, NH3, NH4+
74
Polar Molecules
• Molecular geometry affects molecular
polarity
– -they either cancel or reinforce each other
A B A
linear molecule
nonpolar
A
B
A
angular molecule
polar
75
Polar and Nonpolar Bonds
• Covalent bonds in which the electrons are
shared equally are designated as nonpolar
covalent bonds
– -Nonpolar covalent bonds have a symmetrical
charge distribution (electron distribution)
·· N ·· ·· ·· N ··
or
·· N N ··
H .. H
or
H H
76
Polar and Nonpolar Bonds
• Polar covalent bonds: electrons are not
shared equally
• -they have different electronegativities
Electronegativities:
H
F
2.1
4.0
Difference = 1.9 very polar bond
77
Polar and Nonpolar Bonds
• Compare HF to HI:
Electronegativities:
H
I
2.1
2.5
Difference = 0.4 slightly polar bond
more complicated geometries exist…
78
Bond Polarity
• Three molecules with polar
covalent bonds:
• -Each bond has one atom
with a slight negative
charge (-)
• -another with a slight
positive charge (+)
79
Polar or Nonpolar?
AB3 molecules: BF3, Cl2CO, and NH3
80
Polar or Nonpolar?
CO2 and H2O
Which one is polar?
81
CH4 … CCl4
Polar or Not?
• Only CH4 and CCl4 are NOT polar. These are the only
two molecules that are “symmetrical.”
82
Compounds Containing
Double Bonds
• Ethene or ethylene, C2H4, is the simplest
organic compound containing a double bond.
– -has a double bond to obey octet rule
Lewis Dot Formula
H·
H
·
·
·
C ·· ·· C
·· H
H ··
H
H
C
or
H
C
H
83
Double Bonds
• What is the effect of bonding and
structure on molecular properties?
s and p
Free rotation
around C–C single
bond
No rotation
around C=C
double bond
84
Bond Order
# of bonds between similar pairs of atoms
Double bond
Single bond
Acrylonitrile
Triple
bond
85
Bond Order
Consider NO2-:
••
N
••
N
•• • •••
••
••
O
O• • O
O
••
••
••
••
The N—O bond order = 1.5
Total # of bonds of one - type
Bond order =
Total # of atoms bound of that type
86
Bond Order
Bond order is proportional to two important
bond properties:
(a) bond strength
(b) bond length
414 kJ
110 pm
123 pm
745 kJ
87
Bond Length
the distance between
the nuclei of two
bonded atoms
88
Bond
Bond length
Lengthdepends
on size of bonded
atoms
H—F
H—Cl
Bond distances measured
in Angstrom units where 1
Å = 10-2 pm.
H—I
89
Bond length depends on bond
order
Bond distances measured
in Angstrom units where 1
Å = 10-2 pm.
90
Bond Strength
• Measure of the energy required to break
a bond
• See Table 9.10
•
BOND
H—H
C—C
C=C
CC
NN
STRENGTH (kJ/mol)
436 KJ
346 KJ
602 KJ
835 KJ
945 KJ
The GREATER the number of bonds (bond order) the HIGHER
91
the bond strength and the SHORTER the bond.