Modern Chemistry
Chapter 6
Chemical Bonding
Sections 1-5
Introduction to Chemical Bonding
Covalent Bonding & Molecular Compounds
Ionic Bonding & Ionic Compounds
Metallic Bonding
Molecular Geometry
Chapter 6 Section 1 Intro to Chem
Bonding pages 175-177
1
Section 1
Introduction to
Chemical Bonding
Chapter 6 Section 1 Intro to Chem
Bonding pages 175-177
2
Chapter Vocabulary
Chemical bond
Ionic bonding
Covalent bonding
Nonpolar covalent bond
Polar
Polar covalent bond
Chapter 6 Section 1 Intro to Chem
Bonding pages 175-177
3
Types of Chemical Bonding
• Chemical Bond – a
mutual electrical
attraction between
the nuclei and
valence electrons
of different atoms
that binds the
atoms together.
Chapter 6 Section 1 Intro to Chem
Bonding pages 175-177
4
Types of Chemical Bonding
p. 176
• Ionic Bonding – a bond that results
from the electrical attraction
between cations and anions
Chapter 6 Section 1 Intro to Chem
Bonding pages 175-177
5
Ionic Bonding
p. xx
Chapter 6 Section 1 Intro to Chem
Bonding pages 175-177
6
Types of Chemical Bonding
p. 176
• Covalent Bonding – results from the
sharing of electron pairs between
atoms
Chapter 6 Section 1 Intro to Chem
Bonding pages 175-177
7
Covalent Bonding
p. xx
Chapter 6 Section 1 Intro to Chem
Bonding pages 175-177
8
Ionic
or
Covalent?
The type of bond
can be estimated
by calculating the
difference in the
elements’
electronegativity.
Chapter 6 Section 1 Intro to Chem
Bonding pages 175-177
9
p. 161
Electronegativities on Page 161
Chapter 6 Section 1 Intro to Chem
Bonding pages 175-177
10
Type of Bond
HCl
H = 2.1 Cl = 3.0
3.0 – 2.1 = 0.9
Polar Covalent
H2
H = 2.1
2.1 - 2.1 = 0
Nonpolar Covalent
Chapter 6 Section 1 Intro to Chem
Bonding pages 175-177
11
Electronegative Differences
p. xx
Chapter 6 Section 1 Intro to Chem
Bonding pages 175-177
12
Practice Problems
Find the type of bond of the following…
• H2S
• Cs2S
• SCl6
• NCl3
p.177
1. PC 2. I 3. PC 4. NPC
Chapter 6 Section 1 Intro to Chem
Bonding pages 175-177
13
Polar and Nonpolar
• Nonpolar Covalent
–Electrons are shared equally by
the bonded atoms
–Balanced distribution of charge
Chapter 6 Section 1 Intro to Chem
Bonding pages 175-177
14
Polar and Nonpolar
• Polar Covalent
–Electrons are shared equally by
the bonded atoms because of the
atom unequal attraction
–Uneven distribution of charge
Chapter 6 Section 1 Intro to Chem
Bonding pages 175-177
15
What is it?
p. xx
Chapter 6 Section 1 Intro to Chem
Bonding pages 175-177
16
Electron Dot Diagrams
• Shows the valence electrons as dots
around the symbol
Chapter 6 Section 1 Intro to Chem
Bonding pages 175-177
17
Electron Dot Diagram
p. xx
Chapter 6 Section 1 Intro to Chem
Bonding pages 175-177
18
Section 1 Homework
Section Review Page 177 #1-6
Chapter 6 Section 1 Intro to Chem
Bonding pages 175-177
19
Section 2
Covalent Bonding and
Molecular Compounds
Chapter 6 Section 2 Covalent Bonding and
Molecular Compounds pages 178-189
20
Chapter Vocabulary
Molecule
Molecular compound
Chemical Formula
Molecular Formula
Bond Energy
Octet Rule
Electron-Dot Notation
Lewis Structures
Structural Formula
Single Bond
Multiple Bonds
Resonance
Chapter 6 Section 2 Covalent Bonding and
Molecular Compounds pages 178-189
21
Molecules-a neutral group of atoms
held together by covalent bonds.
Single molecules can exist on their
own.
Molecular Compound-a chemical
compound
whose
simplest units
are molecules
Chapter 6 Section 2 Covalent Bonding and
Molecular Compounds pages 178-189
22
Chemical Formula-indicate the
relative numbers of atoms of each
kind in a chemical compound using
symbols and subscripts.
Molecular Formula-shows the type
and number of atoms combined in a
single molecule of a molecular
formula.
Chapter 6 Section 2 Covalent Bonding and
Molecular Compounds pages 178-189
23
Sucrose and Water molecules
represent molecular formulas.
Chapter 6 Section 2 Covalent Bonding and
Molecular Compounds pages 178-189
24
PE changes during the
formation of a
covalent bond.
Separated atoms do
not affect each other.
PE decreases as the
atoms are drawn
together., then at a
minimum when
attractive forces are
balanced by repulsive forces. PE increases
when repulsive between like charges
outweighs attraction between opposite
charges.
Chapter 6 Section 2 Covalent Bonding and
Molecular Compounds pages 178-189
25
Characteristics of the
Covalent Bond
Bond energy—the energy required to
to BREAK a chemical bond
As bond length decreases the
strength of the bond (bond energy)
increases.
Chapter 6 Section 2 Covalent Bonding and
Molecular Compounds pages 178-189
26
The Octet Rule
Chapter 6 Section 2 Covalent Bonding and
Molecular Compounds pages 178-189
27
Exceptions to the
Octet Rule
Hydrogen and Helium – 2 valence eBoron– 6 valence eExpanded valence—more than 8 veFluorine
Oxygen
Chlorine
Chapter 6 Section 2 Covalent Bonding and
Molecular Compounds pages 178-189
28
Lewis Structures
Formulas in which atomic symbols
represent nuclei and inner-shell
electrons, dot-pairs, or dashes
between two atomic symbols
represent electron pairs in covalent
bonds, and dots adjacent to only one
atomic symbol represent unshared
electrons.
Chapter 6 Section 2 Covalent Bonding and
Molecular Compounds pages 178-189
29
Chapter 6 Section 2 Covalent Bonding and
Molecular Compounds pages 178-189
30
Practice
Draw Lewis
Structures for
CH3I
SiH4
NH3
PF3
H2S
Chapter 6 Section 5 Covalent Bonding and
Molecular Compounds pages 178-189
31
Multiple Covalent Bonds
Double and Triple bonds are referred
to as multiple bonds.
Double bonds share two pairs of
electrons.
Triple bonds share three pairs of
electrons.
Chapter 6 Section 5 Covalent Bonding and
Molecular Compounds pages 178-189
32
Multiple Covalent Bonds
Drawing multiple bonds differs only in the last
step:
a.Count the electron in the Lewis structure to be
sure that the number of valence electrons used
equals the number available.
b.b. If too many electron have been used,
subtract one or more lone pairs until the total
number of valence electron is correct. Then
move one or more lone electron pairs to existing
bonds between non-hydrogen atoms until the
outer-shells of all atoms are completely filled.
Chapter 6 Section 5 Covalent Bonding and
Molecular Compounds pages 178-189
33
Practice
Draw Lewis Structures for CO2 and
HCN
Chapter 6 Section 5 Covalent Bonding and
Molecular Compounds pages 178-189
34
Resonance Structures
Bonding in a molecules or ions that
cannot be correctly represented by a
single Lewis structure.
Chapter 6 Section 5 Covalent Bonding and
Molecular Compounds pages 178-189
35
Section 2 Review
Section Review p. 189,#1-5
Chapter 6 Section 5 Covalent Bonding and
Molecular Compounds pages 178-189
36
Section 3
Ionic Bonding and
Ionic Compounds
Chapter 6 Section 3 Molecular
Geometry pages 190-194
37
Section 4
Metallic Bonding
Chapter 6 Section 4 Metallic
Bonding pages 195-196
38
Section 5
Molecular Geometry
Chapter 6 Section 5 Molecular
Geometry pages 197-207
39
VSEPR Theory
• Valence-Shell
Electron-Pair
Repulsion
• Repulsions between the set of
valence-level electrons surrounding
an atom causes these sets to be
oriented as far apart as possible.
Chapter 6 Section 5 Molecular
Geometry pages 197-207
40
VSPRE & Molecular Geometry
p. xx
Chapter 6 Section 5 Molecular
Geometry pages 197-207
41
Geometry & Lone Pairs
Chapter 6 Section 5 Molecular
Geometry pages 197-207
42
Molecular Geometry
LINEAR
:
:
:
:
:F
- BeF:-
Example formula: BeF2
Type of molecule: AB2
Bond angle: 180°
Shared pairs on the central atom: 2
Unshared pairs on the central atom: 0
Chapter 6 Section 5 Molecular
Geometry pages 197-207
43
Molecular Geometry
TRIGONAL PLANAR
:
:
F:
:
B
:
:F
:
:F:
Example formula: BF3
Type of molecule: AB3
Bond angle: 120°
Shared pairs on the central atom: 3
Unshared pairs on the central atom: 0
Chapter 6 Section 5 Molecular
Geometry pages 197-207
44
Molecular Geometry
TETRAHEDRAL
H
C
H
H
H
Example formula: CH4
Type of molecule: AB4
Bond angle: 109.5°
Shared pairs on the central atom: 4
Unshared pairs on the central atom: 0
Chapter 6 Section 5 Molecular
Geometry pages 197-207
45
Molecular Geometry
ANGULAR
O
H
H
Example formula: H2O
Type of molecule: AB2E2
Bond angle: 105°
Shared pairs on the central atom: 2
Unshared pairs on the central atom: 2
Chapter 6 Section 5 Molecular
Geometry pages 197-207
46
Molecular Geometry
TRIGONAL PYRAMIDAL
:
N
H
H
H
Example formula: NH3
Type of molecule: AB3E
Bond angle: 107°
Shared pairs on the central atom: 3
Unshared pairs on the central atom: 1
Chapter 6 Section 5 Molecular
Geometry pages 197-207
47
Molecular Geometry
• Unshared pairs occupies more space
around the central atom than shared
pairs
• Unshared pairs repel other electrons
more strongly than shared pairs
• Multiple bonds are treated the same
as single bonds
• Polyatomic ions are treated like
molecules.
Chapter 6 Section 5 Molecular
Geometry pages 197-207
48
Molecular Geometry
• Try
– CO2
– ClO3
1-
• Practice Problems page 201
• Try
– CF4
– NO3 1-
Chapter 6 Section 5 Molecular
Geometry pages 197-207
49
Hybridization
• The mixing of two or more atomic
orbitals of similar energies on the
same atom to produce new hybrid
atomic orbitals of equal energy
• Example CH4
C =   _ _ __
1s 2s 2p
 _ _ _ _
1s sp3
Chapter 6 Section 5 Molecular
Geometry pages 197-207
50
Hybridization
• s and p orbitals have different
shapes
• The 2s & 2p hybridize to make four
identical orbitals
– named sp3
– The 3 is from the three p orbitals used
– But the 1 is not written for the s
Chapter 6 Section 5 Molecular
Geometry pages 197-207
51
Hybridization
• All sp3 orbitals have the same
energy
– Higher than 2s but
– Lower than 2p
• Hybrid orbitals – orbitals of equal
energy produced by the combination
of two or more orbitals.
Chapter 6 Section 5 Molecular
Geometry pages 197-207
52
Hybridization
N = 
1s

1s
O = 
1s

1s

2s

sp3

2s

sp3
_ _ _
2p
_ _ _
 _ _
2p
 _ _
Chapter 6 Section 5 Molecular
Geometry pages 197-207
53
Hybridization
Be =  
1s 2s
 _ _ __
1s sp
B =   _ __ __
1s 2s 2p
 _ _ _ __
1s sp2
Chapter 6 Section 5 Molecular
Geometry pages 197-207
Uses one
p orbital
Uses two
p orbitals
54
Hybridization
p. xx
Chapter 6 Section 5 Molecular
Geometry pages 197-207
55
Hybrid Orbital Animation
p. xx
Chapter 6 Section 5 Molecular
Geometry pages 197-207
56
p. xx
Comparing Molecular & Ionic Compounds
Chapter 6 Section 5 Molecular
Geometry pages 197-207
57
Molecule Polarity
δ+
δ-
2.1 H
- Cl
3.0
Lower EN
Higher EN
polar bond = dipole
• Dipole: created by equal but
opposite charges that are separated
by a short distance
Chapter 6 Section 5 Molecular
Geometry pages 197-207
58
Molecule Polarity
• Molecule polarity for compounds
with more than one bond depends
on …
bond polarity
and
molecule geometry.
Chapter 6 Section 5 Molecular
Geometry pages 197-207
59
Molecule Polarity
1. Draw the Lewis Structure true to
shape. Example NH3
:
N
H
H
H
Chapter 6 Section 5 Molecular
Geometry pages 197-207
60
Molecule Polarity
:
2. Find all the partial positive and
negatives for each atom in the
molecule
δ- 3.0
N
δ
H
+
H
H
2.1
δ+
δ+
+
Look
at each
High EN
= δ- bond.
Low
EN
=
δ
Chapter 6 Section 5 Molecular
Geometry pages 197-207
61
Molecule Polarity
3. Look at around the “outside” of the
molecule.
:
δ-
N
δ+
H
H
H
δ+
δ+
All the same δ Chapter
= NP;
Different δ = P
6 Section 5 Molecular
Geometry pages 197-207
62
Molecule Polarity
1. Draw the Lewis Structure true to
shape. Example CH4
H
C
H
H
H
Chapter 6 Section 5 Molecular
Geometry pages 197-207
63
Molecule Polarity
2. Find all the partial positive and
negatives for each atom in the
molecule
δ+
2.1
H
C
δ+
H
2.1
H
2.5
δ2.1
δ+
H
2.1
δ+
+
Look
at each
High EN
= δ- bond.
Low
EN
=
δ
Chapter 6 Section 5 Molecular
Geometry pages 197-207
64
Molecule Polarity
3. Look at around the “outside” of the
molecule.
H
C
δ+H
H
δ+
δ-
H
δ+
δ+
All the same
Carbon
is notδon
= the
NP;
“outside”.
Different δ = P
Chapter
6 Section
5 Molecular
Geometry pages 197-207
65
Intermolecular Forces
• The force of attraction between
molecules to make (solids or) liquids
• Boiling point is a good measure of
the strength of intermolecular forces
• Weaker than covalent bonds, ionic
bonds and metallic bonds
Chapter 6 Section 5 Molecular
Geometry pages 197-207
66
Molecule Polarity
δ+
H - Cl
δ-
δ+
H - Cl
δ-
Dipole-dipole force: the force of
attraction between polar molecules
Chapter 6 Section 5 Molecular
Geometry pages 197-207
67
Dipole Dipole Animation
p. xx
Chapter 6 Section 5 Molecular
Geometry pages 197-207
68
p. xx
Comparing Dipole Dipole Forces
Chapter 6 Section 5 Molecular
Geometry pages 197-207
69
Hydrogen Bonding
• H-F, H-O or H-N bonds have a large
electronegativity difference
• These bonds are very polar.
• Molecules with these bonds have
very strong dipole-dipole forces
Chapter 6 Section 5 Molecular
Geometry pages 197-207
70
Hydrogen Bonding
p. xx
Chapter 6 Section 5 Molecular
Geometry pages 197-207
71
Hydrogen Bonding
• The intermolecular force in which
a Hydrogen atom that is bonded to
Nitrogen or Oxygen or Fluorine
is attracted to
an unshared pair of electrons
of the N, O or F of another molecule
Chapter 6 Section 5 Molecular
Geometry pages 197-207
72
Hydrogen Bonding
• Compare
PH3 & NH3
H2O & H2S
Page 204
Chapter 6 Section 5 Molecular
Geometry pages 197-207
73
Dipole Induced Dipole
p. xx
Chapter 6 Section 5 Molecular
Geometry pages 197-207
74
Induced Dipole
• Polar molecules cause a dipole in a
nonpolar molecule
O
O
:
H
δ+
:
O
δ-
:
δ-
δ+
:
H
δ+
Chapter 6 Section 5 Molecular
Geometry pages 197-207
75
London Dispersion Forces
• Nonpolar molecules don’t have
dipoles
• However at any instance the
electron distribution may be uneven.
• An instantaneous dipole can occur
and induce dipoles in other
molecules
Chapter 6 Section 5 Molecular
Geometry pages 197-207
76
London Dispersion Force
p. xx
Chapter 6 Section 5 Molecular
Geometry pages 197-207
77
London Dispersion Forces
• London dispersion forces – the
intermolecular attraction resulting
from the constant motion of
electrons and the creation of
instantaneous dipoles
• Very weak intermolecular forces
• London forces increase with
increasing atomic or molar mass.
Chapter 6 Section 5 Molecular
Geometry pages 197-207
78
Lewis Structures
Practice
• C2H4
• BeF2
• AsH3
• IBr
• CHCl3
• CN 1• N2O2
Chapter 6 Section 5 Molecular
Geometry pages 197-207
79
Lewis Structures
Practice
• C2Cl4
• SCl2
• AsF5
• CI2Cl2
• BF3
• NO 1• CH2O
• IO3 1Chapter 6 Section 5 Molecular
Geometry pages 197-207
80
Section 5 Homework
Chapter 6 Section 5 Worksheet
Chapter 6 Section 5 Molecular
Geometry pages 197-207
81