Chemistry Matter and Change Chapter 8
COVALENT
BONDING
CHAPTER 8 MAIN IDEA
Covalent bonds form
when atoms share
electrons.
8.1 Main Idea
Atoms gain
stability when
they share
electrons and
form covalent
bonds.
REVIEW VOCABULARY &
CONCEPTS
Chemical bond
Valence electrons
Electronegativity
Lewis structure
NEW VOCABULARY &
CONCEPTS
Covalent bond
Molecule
Sigma (σ)bond
Pi (π)bond
Endothermic reaction
Exothermic reaction
Single bond
Double bond
Triple bond
WHY DO ATOMS BOND?
Sharing electrons takes less
energy than being “alone”
Octet is usually the most stable
electron configuration
WHAT IS A COVALENT BOND?
Shared pairs of electrons
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SINGLE COVALENT BONDS
Two atoms share one pair of
electrons
Sigma bond (σ)
_
Either : or for a Lewis Structure
· · H  H·· H
H+
or
H-H
MULTIPLE COVALENT BONDS
Two atoms share more than one
pair of electrons
One pair is a sigma bond (σ)the
others are pi (π)bonds
Double bond shares 2 pair of
electrons
Triple bond shares 3 pair of
electrons
COMPARING BONDS
Sigma (σ) Bonds
Pi (π) Bonds
 Single
 Paired
 Centered
 Parallel
s
p
orbital
electrons
 One pair of
electrons is
shared
orbital
electrons
 Multiple pairs of
electrons are
shared
MOLECULE
A neutral group of atoms joined
together by covalent bonds.
DIATOMIC MOLECULES
Molecules made up of two
atoms.
There are 7 diatomic molecules.
H2,
N2, O2,
F2, Cl2, Br2, I2
DIATOMIC MOLECULES
Br I N Cl H O F
I couldn't
exist
without
you!
Oh, Ha Ha!
HYDROGEN
H· + ·H 
Hydrogen
Atom
Hydrogen
Atom
H:H
Hydrogen
molecule
Shared
electron
pair
The hydrogen molecule has a single
covalent bond.
The electronegativity of each
hydrogenatom is the same, so the
electrons are shared.
MOLECULAR COMPOUNDS
Compounds comprised of
molecules.
Ionic Compounds
Molecular
Compounds
Crystal Lattice
Molecule
Metal with non-metal Non-metal with
or polyatomic ions
non-metal
Solid
Solid, liquid or gas
Types of
Elements
Physical
State
Melting Point High
> 300 C
Solubility in Generally high
water
Electrical
Good conductor
conductivity
of solution
Low
<300 C
Generally low
Poor to none
VIDEO
http://www.glencoe.com/sites/co
mmon_assets/science/cmc/cim/an
imations/ch8_1.swf
LEWIS STRUCTURES AND OCTET
Practice by drawing
H2
O2
+
N2
H2O
CO2
+
LEWIS STRUCTURES AND OCTET
Practice by drawing
H2
+
-
+
LEWIS STRUCTURES AND OCTET
Practice by drawing O2
σ
- -
-
+
- -
- -
-
-
-
+
-
π
-
+
- -
-
+
-
STRENGTH OF COVALENT BONDS
Strength depends on distance of
the atoms from each other
With more bonds comes
stronger bonds
O2 is stronger than H2
single
bond
<
Double
bond
<
Triple
bond
CHEMICAL BONDS
22
CHEMICAL BONDS
Electronegativity is an atom’s affinity for
electrons.
Differences in electronegativity dictate
how electrons are distributed in covalent
bonds.
- nonpolar covalent bonds = equal sharing
of electrons
- polar covalent bonds = unequal sharing
of electrons
23
BONDS AND ENERGY
Endothermic reactions require
additional energy for bonds to
break
Exothermic reactions release
energy when the bonds break
(spontaneous)
Section 8.2 Chemistry Matter & Change
NAMING
MOLECULES
8.2 Main Idea
Specific rules are
used when naming
binary molecular
compounds, binary
acids and oxyacids
OBJECTIVES
Translate molecular formulas
into binary molecular names.
Name acidic solutions.
REVIEW VOCABULARY &
CONCEPTS
Ionic bond
Covalent bond
Formula unit
Oxyanion
Naming ionic substances
Molecule
NEW VOCABULARY &
CONCEPTS
Oxyacid
Binary acid
NAMING BINARY MOLECULAR
COMPOUNDS
1. Name the first element using
the entire name of the
element.
2. The second element in the
formula is named using the
root and suffix “-ide.”
3. Prefixes are used to indicate
the numbers of each element.
COMMON PREFIXES
Number of
atoms
1
2
3
4
5
Prefix
MonoDiTriTetraPenta-
Number of
atoms
6
7
8
9
10
Prefix
HexaHeptaOctaNonaDeca-
NAMING BINARY MOLECULAR
COMPOUNDS
Exceptions to the rules:
The first element never uses
“mono-”
There
is an understood 1 if
nothing is specified.
Awkward vowels can be
dropped
NAMING BINARY MOLECULAR
COMPOUNDS EXCEPTION
 The first element never uses
“mono-”
There
is an understood 1 if
nothing is specified.
 CO2 is carbon dioxide not
monocarbon dioxide
 CO is carbon monoxide, not
monocarbon monoxide
PRACTICE NAMING BINARY
COMPOUNDS
FORMULA
SYSTEM NAME
N2O
dinitrogen monoxide
NO
nitrogen monoxide
N2O3
dinitrogen trioxide
NO2
nitrogen dioxide
N2O4
dinitrogen tetroxide
N2O5
dinitrogen pentoxide
NO3
nitrogen trioxide
NAMING BINARY ACIDS
The first word has the prefix
“hydro-” followed by the root of
the second element followed by
“-ic”
The second word is “acid”
HCl-
hydrochloric acid
PRACTICE NAMING BINARY
ACIDS
Hydrobromic acid
Hydrophosphic acid
Hydrosulfuric acid
HF
H2Se
HI
NAMING OXYACIDS
First word consists of the root of the
oxyanion (with prefixes if needed)
followed by a suffix as specified
If oxyanion ends in
Oxyanion ending
New suffix for acid
-ate
-ic
-ite
-ous
The second word is always “acid.”
NAMING OXYACIDS
Relationshi General
p
name
Oxyanion
name
Example
name
Example
formula
One more
oxygen than
(root)ic
Per(root)ic
acid
Perchlorate
Perchloric
acid
HClO4
Root(ic) acid
chlorate
Chloric acid
HClO3
Root(ous)
acid
Chlorite
Chlorous
acid
HClO2
One less
oxygen than
root(ic)
Two less
Hypo(root)ou
Hypochlorou
oxygens than
Hypochlorite
s acid
s acid
root(ic)
HClO
NAMING OXYACIDS
First word consists of the root of the
oxyanion (with prefixes if needed)
followed by a suffix as specified
The second word is always “acid.”
Compound
Oxyanion
Acid suffix
Acid name
HClO3
Chlorate
-ic
Chloric acid
HClO2
Chlorite
-ous
Chlorous acid
HNO3
Nitrate
-ic
Nitric Acid
HNO2
Nitrite
-ous
Nitrous Acid
Look at the
formula of the
molecule
Is it an acid?
Yes
Is there oxygen
present in the
compound?
No
Hydro(root) ic acid
Yes
No
Name the first element using a
prefix if necessary
Name the second element
indicating the number of atoms
and changing the suffix to –ide.
Root + -ic if the anion ends in –ate,
or
Root + -ous if the anion ends in –ite,
then
acid
PRACTICE NAMING OXYACIDS
Acid Formula
Acid Name
H2CO3
Carbonic Acid
HBrO3
Bromic Acid
H2C
Carbonic Acid
CH3COOH or CH3CO2H
Acetic acid
H3PO4
Phosphoric acid
H2SO4
Sulfuric acid
H2SO3
Sulfurous Acid
CAN YOU
Translate molecular formulas
into binary molecular names.
Name acidic solutions.
Section 8.3 Chemistry Matter & Change
MOLECULAR
STRUCTURE
8.3 Main Idea
Structural
formulas show
the relative
positions of
atoms within a
molecule.
OBJECTIVES
List the basic steps used to
draw Lewis structures.
Explain why resonance occurs
and identify resonance
structures.
Identify three exceptions to the
octet rule and name molecules in
which these exceptions occur.
REVIEW VOCABULARY &
CONCEPTS
Ionic bond
Covalent bond
Lewis-dot structure
Octet
NEW VOCABULARY &
CONCEPTS
Structural formula
Resonance
Coordinate covalent bond
Space filling model
Ball and stick model
MOLECULAR STRUCTURES
Many different ways to depict
the same thing
Molecular
formula
Structural formula
Lewis Structure
Space-filling model
Ball-and-stick model
MOLECULAR FORMULA
Indicates number of each
element in a molecule
H2O
C6H12O6
NH3
STRUCTURAL FORMULAS
Molecular model that shows the
relative positions of the atoms
O
H
H
LEWIS STRUCTURE
Shows shared electrons and
lone pairs
Shared
pairs are usually depicted
as lines
Lone pairs may be lines or dots
SPACE-FILLING MODEL
Atoms are shown in relative
size and position accounting for
lone pairs
BALL AND STICK MODELS
Show relative position of atoms
and bonds
Easier
to see double and triple
bonds and bond angles
HINTS FOR LEWIS STRUCTURES
1. Carbon is usually in the middle
2. Group 1 and 17 elements are
always at ends
3. Atoms that are less numerous
are usually in the middle
(~Polyatomic Ions)
SIMPLE PRACTICE FOR LEWIS
STRUCTURES
H2O
HCl
CH3F
ORGANIC PRACTICE WITH
LEWIS STRUCTURES
Methane CH4
Ethane
C2H6
Propane C3H8
TRICKIER PRACTICE
CO2
HCN
BH3
NH4+
TRICKIER ORGANIC PRACTICE
C2H6
C2H4
C2H2
ethane
ethene
ethyne
RESONANCE STRUCTURES
More than one valid Lewis
structure is possible
Differ
in position of electrons, not
position of atoms
HINTS FOR FORMING LEWIS
STRUCTURES
Carbon is usually in the middle
Group 1 and 17 elements are always at ends
Atoms that are less numerous are usually in the
middle
Hydrogen always forms one single bond
Oxygen has two bonding electrons and two lone
pairs
Nitrogen has three bonding electron and one lone
pair
Group 13 elements have three bonding electrons
and zero lone pairs
Elements in groups 1, 2, 13 may break octet by
having fewer electrons
Phosphorus is a mess!
REALLY HARD
PCl5
SF6
ORGANIC CHEATS
Common groups are often listed
differently in organic molecules
Ex: we write NH2CH2COOH
instead of C2NO2H5
CAN YOU
List the basic steps used to
draw Lewis structures.
Explain why resonance occurs
and identify resonance
structures.
Identify three exceptions to the
octet rule and name molecules in
which these exceptions occur.
Section 8.4 Chemistry Matter and Change
MOLECULAR
SHAPES
Main Idea
The VSEPR model
is used to
determine
molecular shape
OBJECTIVES
Summarize VSEPR bonding
theory
Predict the shape of, and bond
angles in, a molecule
Define hybridization
REVIEW VOCABULARY &
CONCEPTS
Atomic orbital
s orbital
p orbital
NEW VOCABULARY &
CONCEPTS
VSEPR theory
Hybridization
Trigonal planar
Trigonal pyramidal
Tetrahedral
Trigonal bipyramidal
Linear
Bent
Octahedral
VSEPR MODEL
Valence Shell Electron Pair
Repulsion
Electrons repel each other
Electrons
stay as far away from
each other as possible
Lone pairs occupy more space
than shared electrons
HYBRIDIZATION
Occurs because hybrid orbitals
are more stable in some cases,
notably carbon
Remodel the “hotel” to make the
“rooms” equal
CARBON HYBRIDIZATION
C: 1s22s22p2  1s22sp3
GETTING STARTED
Begin with Lewis Structure
Start with a tetrahedral and
modify as needed
Remember to use 3-D space not
2-D space
VIDEO
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imations/ch8_2.swf
CH4: A TYPICAL TETRAHEDRAL
TETRAHEDRAL SHAPE
Four atoms are coming from the
center atom at equal angles
109.5° angle
TRIGONAL PYRAMIDAL
Three atoms coming from
central atom and one lone pair
Bond
Lone
angle 107°
pairs occupy more space than
shared pairs
BENT
Two atoms coming from central
atom and two lone pairs
104.5° bond angle
SUMMARY OF COMMON
MOLECULAR SHAPES
Molecul Tota Share
e
l
d
pair pairs
s
Lon
e
pair
s
Hybrid
orbital
Molecular shape
BeCl2
2
2
0
sp
Linear
AlCl3
3
3
0
sp2
Trigonal planar
CCl4
4
4
0
sp3
Tetrahedral
NH3
4
3
1
sp3
Trigonal
pyramidal
H2O
4
2
2
sp3
Bent
NbBr5
5
5
0
sp3d
Trigonal bipyramidal
SF
6
6
0
sp3d2
Octahedral
TETRAHEDRAL BASED SHAPES
Tetrahedral
Trigonal Pyramidal
Bent
NON-TETRAHEDRAL BASED SHAPES
Linear
Trigonal Planar
FEWER ELECTRONS THAN OCTET!
NON-TETRAHEDRAL BASED SHAPES
Trigonal Bipyramidal
Octahedral
MORE ELECTRONS THAN OCTET!
CAN YOU…
Summarize VSEPR bonding
theory
Predict the shape of, and bond
angles in, a molecule
Define hybridization
Chapter 8.5 Chemistry Matter and Change
ELECTRONEGATIVITY
AND POLARITY
Main Idea
A chemical bond’s
character is
related to each
atom’s attraction
for the electrons
in the bond.
OBJECTIVES
Describe how electronegativity
is used to determine bond type.
Compare and contrast polar
and nonpolar covalent bonds and
polar and nonpolar molecules
Generalize about the
characteristics of covalently
bonded compounds
REVIEW VOCABULARY &
CONCEPTS
Electronegativity
Electron affinity
Ionic bond
Covalent bond
NEW VOCABULARY &
CONCEPTS
Polar covalent bond
Electronegativity difference
Molecular polarity
Intermolecular forces
Solubility
Dipole
ELECTRON AFFINITY, ELECTRONEGATIVITY,
AND BOND CHARACTER
Electronegativity difference determines
the character of the bond between atoms
Most bonds are partly ionic and partly
covalent
Diatomics are 100% covalent
Electronegativity difference
Bond Character
> 1.7
Mostly ionic
0.4 - 1.7
Polar covalent
< 0.4
Mostly covalent
0
Nonpolar covalent
BOND CHARACTER
IONIC BONDS
Created when the
electronegativity difference is
greater than 1.7
Elements are far apart on
periodic table
Example:
Na and Cl
Na = 0.9
Cl = 3.0
Difference is 2.1  ionic bond
Electrons are transferred from Na to Cl

NON-POLAR COVALENT BONDS
Created when electronegativity
difference is 0
Usually
atoms of the same
element
Diatomics

COVALENT BONDS
Created when electronegativity
difference is < 0.4
Elements are usually close
together on the periodic table
IBr
POLAR-COVALENT BONDS
Created when electronegativity
difference is between 0.4 and 1.7
Most bonds fall into this
category
Electrons are shared unequally
H2O
POLAR COVALENT BONDS
Molecules have a partial
positive and a partial negative
side
DIPOLE INTERACTIONS
Reactions between oppositely
charged ends of a polar molecule
POLAR COVALENT BONDS
Polar bonds that are equal in
all directions in a molecule
create non-polar molecules
SOLUBILITY RULES
Like dissolves like
Polar
substances dissolve polar
substances
Non-polar substances dissolve
non-polar substances.
CAN YOU
Describe how electronegativity
is used to determine bond type.
Compare and contrast polar
and nonpolar covalent bonds and
polar and nonpolar molecules
Generalize about the
characteristics of covalently
bonded compounds