Chapter 8 – Covalent Bonding
The
unspoken
hero:
“Covalent
Bond”
Review of Chapter 7
• In Chapter 7, we learned about electrons being
transferred (“given up” or “stolen away”)
• This type of “tug of war” between
a METAL and NONMETAL
is called an IONIC BOND,
which results in a SALT being formed
Chapter 8.1 – Molecular Compounds
• Covalent Bonds - atoms held
together by SHARING electrons
between NONMETALS
Salt versus Molecules
•Molecule - A group of
atoms joined together by a covalent bond
•Compound - a group of two or more elements
bonded together (Ionic or Covalent).
Monatomic vs. Diatomic Molecules
• Most molecules can be monatomic or diatomic
•Diatomic Molecule - molecule
consisting of two atoms
•There are 7 diatomic molecules (SUPER 7) –
•H2, O2, N2, Cl2, Br2, I2, F2
Properties of Molecular Compounds
• Liquids or gases at room temperature
• Lower Melting Points than
Ionic Compounds (means weaker bonds than ionic)
Molecular Formulas
• Molecular Formula –
formula of a molecular compound
• Shows how many atoms of each
element a molecule contains
•Example
H2O
contains 3 atoms (2 atoms of H, 1 atom of O)
C2H6 contains 8 atoms (2 atoms of C, 6 atoms of H)
Practice
How many atoms total and of each do the
following molecular compounds contain?
2
1. H2
2
2. CO
3
3. CO2
4
4. NH3
9
5. C2H6O
Practice: True or False
1. All molecular compounds are composed of
2.
3.
4.
5.
atoms of two or more elements.
All compounds are molecules.
Molecular compounds are composed of
two or more nonmetals.
Atoms in molecular compounds exchange
Share
electrons.
Molecular compounds have higher melting
and boiling points than ionic compounds.
Lower
Ionic versus Covalent
IONIC
COVALENT
Bonded Name
Salt
Molecule
Bonding Type
Transfer e-
Share e-
Types of Elements
Metal & Nonmetal
Nonmetals
Physical State
Solid
Solid, Liquid, or Gas
Melting Point
High (above 300ºC)
Low (below 300 ºC)
Solubility
Dissolves in Water
Varies
Conductivity
Good
Poor
Chapter 8.2 – Covalent Bonding
• Remember that ionic compounds
transfer electrons in order
to attain a noble gas electron configuration
• Covalent compounds form by
sharing electrons to attain
a noble gas electron configuration
• Regardless of the type of bond, the
Octet Rule still must
be obeyed (8 valence electrons)
•All elements need 8 except Hydrogen, which needs 2.
Single Covalent Bond
• A Single Covalent Bond consists of two atoms
held together by sharing 1 pair
of electrons (2 e-)
Electron Dot Structure
Shared versus Unshared Electrons
• A Shared Pair is a pair of
valence electrons that is shared
between atoms
• An Unshared Pair is a pair
of valence electrons that is not shared
between atoms
Practice Lewis Dot Structures
Chemical
Formula
F2
H2 O
# of
Valence
Electrons
14
8
NH3
8
CH4
8
Single Line
Bond
Structure
# of
Remaining
Electrons
F-F
12
Lewis Dot
Structure
Octet Check
All Atoms=8
Hydrogen=2
F, F = 8
4
O=8
H, H = 2
2
N=8
3H=2
H-O-H
0
C=8
4H=2
Double Covalent Bonds
• Sometimes atoms need to share 2 or 3 pairs of
electrons
• Hydrogen will NEVER form
double or triple bonds.
• Double Bond - bond that involves
2 shared pairs of electrons (4 e-)
Triple Covalent Bond
• Triple Bond - bond that
involves 3 shared pairs of electrons (6 e-)
Covalent Bonds
Practice Lewis Dot Structure
Chemical
Formula
O2
CO2
N2
HCN
# of
Valence
Electrons
12
16
10
10
Single Line
Bond
Structure
# of
Remaining
Electrons
Lewis Dot
Structure
Octet Check
All Atoms=8
Hydrogen=2
O-O
10
2O=8
O-C-O
12
C=8
2O=8
N-N
H-C-N
8
6
2N=8
C=8
N=8
H=2
Chapter 8: Basic Concepts of Chemical Bonding
Drawing Lewis Structures of Molecules
If the compound contains more than 2 atoms:
• how are the atoms bonded and,
• if there are nonbonding electron, where
are they?
Chemical
Bonding
Chapter 8: Basic Concepts of Chemical Bonding
Molecules with a central atom :
NH3, PCl3, CHCl3
Central atom is generally the
first in the molecular formula and the
most electronegative one.
H
H
N
H
H
Cl
P
Cl
Cl
Cl
C
Cl
Cl
Chemical
Bonding
Chapter 8: Basic Concepts of Chemical Bonding
…unless the first element is Hydrogen :
O
H2O
HCN
H
H
C
H
N
(same order as in formula)
Chemical
Bonding
How to find the # of bonds in a
lewis structure
**Doesn’t work for molecule over an octet
1. Find the total # of valence electrons.
2. Use the formula to find the number of bonds.
# of val e- needed (all have 8 or 2 e-)
- # of val e- available
= ____/2 to find the # of bonds
(# val e- –
# val e- available)
2
=
# of bonds
1. Find the total # of valence electrons.
2. Use the formula to find the number of bonds.
# of val e- needed (all have 8 or 2 e-)
- # of val e- available
= ____/2 to find the # of bonds
Ex: Find the number of bonds for each molecule or
compound and write the lewis dot structure
# of Bonds
LDS
a.) CO
b.) C2F4
c.) C2H6
3
6
7
**Doesn’t work for molecule over an octet
Chapter 8: Basic Concepts of Chemical Bonding
Rules for Drawing Lewis Structures
(1) Sum valence electrons from all atoms:
8
these are the ones that need to be distributed
NH3
H
(2) Connect atoms by covalent bonds:
count electrons left
2
(3) Complete "octets" of atoms around
central atom
n/a
H
H
(4) Place any leftover electrons on the
central atom.
2 leftovers
Check that central atom has octet
(5) If there are not enough electrons to give
the central atom an octet, try multiple bonds
N
n/a
H
N
H
H
Chemical
Bonding
Chapter 8: Basic Concepts of Chemical Bonding
(1) Sum valence electrons from all atoms:
these are the ones that need to be distributed
(2) Connect atoms by covalent bonds:
count electrons left
(3) Complete "octets" of atoms around
central atom
(4) Place any leftover electrons on the
central atom.
2
Check that central atom has octet
CO
10
C-O
8
(treat C as central)
left
(5) If there are not enough electrons to give
the central atom an octet, try multiple bonds
C-O
C-O
Chemical
Bonding
Chapter 8: Basic Concepts of Chemical Bonding
(1) Sum valence electrons from all atoms:
these are the ones that need to be distributed
(2) Connect atoms by covalent bonds:
count electrons left
SF2
20
F-S-F
16
(3) Complete "octets" of atoms around
central atom
(4) Place any leftover electrons on the
central atom.
4
Check that central atom has octet
left
(5) If there are not enough electrons to give
the central atom an octet, try multiple bonds
Chemical
Bonding
Lewis Structure for Ions
• If a molecule has a positive charge,
subtract that many electrons from the
total valence electrons available.
•
𝑁𝐻4+
has 8 electrons available.
• If a molecule has a negative charge,
add that many electrons to the
total valence electrons available.
• 𝑁𝑂3− has 24 electrons available.
• 𝑆𝑂4−2 has 32 electrons available.
Chemical
Bonding
For ions, the charge is generally
indicated by
square brackets and the sign
Chemical
Bonding
Chapter 8: Basic Concepts of Chemical Bonding
Ions
(1) Sum valence electrons from all atoms:
these are the ones that need to be distributed
(2) Connect atoms by covalent bonds:
count electrons left
(3) Complete "octets" of atoms around
central atom
(4) Place any leftover electrons on the
central atom.
Check that central atom has octet
NH4+
8
0 left
n/a
n/a
(5) If there are not enough electrons to give
the central atom an octet, try multiple bonds
Chemical
Bonding
Chapter 8: Basic Concepts of Chemical Bonding
Ions
(1) Sum valence electrons from all atoms:
these are the ones that need to be distributed
(2) Connect atoms by covalent bonds:
count electrons left
ClO220
O-Cl-O
16 left
(3) Complete "octets" of atoms around
central atom
(4) Place any leftover electrons on the
central atom.
Check that central atom has octet
4 left
(5) If there are not enough electrons to give
the central atom an octet, try multiple bonds
Chemical
Bonding
Chapter 8: Basic Concepts of Chemical Bonding
Exceptions to the Octet rule
On occasion, an atom in a molecule does not have an
octet of valence electrons:
• If the molecule has an
odd number of valence electrons
• an atom may have
less than an octet [mainly Be, B]
• an atom may have
more than an octet [periods 3-7]
Chemical
Bonding
Chapter 8: Basic Concepts of Chemical Bonding
Exceptions to the Octet rule: odd number of electrons
(1) Sum valence electrons from all atoms:
these are the ones that need to be distributed
(2) Connect atoms by covalent bonds:
count electrons left
17
O-N-O
NO2
17
(3) Complete "octets" of atoms around
central atom
(4) Place any leftover electrons on the
central atom.
Check that central atom has octet
1 left
(5) If there are not enough electrons to give
the central atom an octet, try multiple bonds
Chemical
Bonding
Chapter 8: Basic Concepts of Chemical Bonding
Exceptions to the Octet rule: less than an octet (B or Be)
(1) Sum valence electrons from all atoms:
these are the ones that need to be distributed
(2) Connect atoms by covalent bonds:
count electrons left
24
BF3
18 left
(3) Complete "octets" of atoms around
central atom
(4) Place any leftover electrons on the
central atom.
Check that central atom has octet
(5) If there are not enough electrons to give
the central atom an octet, try multiple bonds
0 left
Chemical
Bonding
More Than Eight Electrons
• The only way PCl5 can
exist is if phosphorus
has 10 electrons
around it.
• Periods 3-7
can expand its orbitals.
Chemical
Bonding
More Than Eight Electrons
Even though we can draw a Lewis structure for the
phosphate ion that has only 8 electrons around the
central phosphorus, the better structure puts a
double bond between the phosphorus and one of
the oxygens.
Chemical
Bonding
More Than Eight Electrons
• This eliminates the charge on the phosphorus
and the charge on one of the oxygens.
• The lesson is: When the central atom is on the
3rd row or below and expanding its octet
eliminates some formal charges, do so.
Chemical
Bonding
Chapter 8: Basic Concepts of Chemical Bonding
Exceptions to the Octet rule: more than an octet
(1) Sum valence electrons from all atoms:
these are the ones that need to be distributed
(2) Connect atoms by covalent bonds:
count electrons left
(3) Complete "octets" of atoms around
central atom
42
BrF5
32
2 left
(4) Place any leftover electrons on the
central atom.
Check that central atom has octet
(5) If there are not enough electrons to give
the central atom an octet, try multiple bonds
Chemical
Bonding
Bond Dissociation Energy
• Bond Dissociation Energy - energy required to
break a bond
between two atoms


A large bond dissociation energy
corresponds to a strong bond
which makes it unreactive
Carbon has strong bonds, which makes carbon
compounds stable and unreactive
Chemical
Bonding
Chapter 8.3 - Bonding Theories
• Determining shape
through bonds
VSEPR Theory
• VSEPR Theory predicts
the 3D shape of molecules

According to VSEPR, the shape of the
molecule adjusts so that the
electrons are far apart
A Few VSEPR Shapes
Nine possible molecular shapes
VSEPR Theory


Unshared pairs of electrons are
important in predicting the shapes of molecules

Each bond (single, double, or triple) and
unshared pair is considered
a steric number

Use the steric number to predict the
molecular geometry
VSEPR can only be used with the central atom
Molecule Lewis Dot
Structure
H 2O
Steric #
4
Shape
Bent
CO2
2
XeF4
6
Linear
Square
Planar
Hybrid Orbitals
• VSEPR is good at describing the molecular
shapes, but not the types of bonds formed
• Orbital hybridization
provides information about both molecular
bonding and molecular shape

In hybridization, several atomic orbitals
mix to form hybrid orbitals
Bond Hybridization
• Hybridization Involving Single Bonds – sp3 orbital
 Ethane (C2H6)
• Hybridization Involving Double Bonds – sp2 orbital

Ethene (C2H4)
• Hybridization Involving Triple Bonds – sp orbital

Ethyne (C2H2)
Gets another orbital added for each atom
and lone pair around atom of interest.
Start with sp…
…then sp2…
(what is the orbital
hybridization of C?)
…then sp3
Chapter 8.4 – Polar Bonds and Molecules
• There are two types of covalent bonds
Nonpolar Bonds (share equally)
 Polar Bonds (share unequally)

Polar Covalent
• Polar Bond - unequal sharing of
electrons between two atoms (ex: HCl)
• In a polar bond, one atom typically has a negative
charge, and the other atom has a positive charge
Nonpolar Covalent Bond
•
Nonpolar Bond - equal sharing
of electrons between two atoms (Cl2, N2, O2)
Classification of Bonds
You can determine the type of bond between two
atoms by calculating the difference in
electronegativity values
between the elements
Type of Bond
Nonpolar Covalent
Electronegativity
Difference
0 0.4
Polar Covalent
0.5
1.9
Ionic
2.0
4.0
Practice
What type of bond is HCl? (H = 2.1, Cl = 3.1)
Difference = 3.1 – 2.1 = 1.0
Therefore it is polar covalent bond.
Your Turn To Practice

N(3.0) and H(2.1)
Mg(1.2)

H(2.1) and H(2.1)


Ca(1.0) and Cl(3.0)
and O(3.5)
H(2.1) and F(4.0)
Polar molecules –
• All molecules
with lone pairs
(unless it’s linear)
• All molecules
surrounded by
different atoms.
Nonpolar molecule –
• Linear molecules
with the same atoms
• All molecules
surrounded by the
same atoms.
Dipole
When
there is unequal sharing of electrons, a
dipole exists
Dipole
- a molecule with
two poles with opposite
charges
Represented
by an
arrow pointing towards
the more negative end
Practice Drawing Dipoles
P- Br
P = 2.1
Br = 2.8
P –Br
+
Practice
H(2.1) – Cl(3.0)
 C(2.5) - F(4.0)
 H(2.1)– F(4.0)

-
Attractions Between Molecules
• There are also attractions between
molecules
• Intermolecular attractions
are weaker than ionic, covalent, and
metallic bonds
• There are 2 main types of attractions between
molecules: Van der Waals and Hydrogen
Van der Waals Forces
• Van der Waals
forces consists of the two weak attractions
between molecules
1. dipole
interactions –
polar molecules
attracted to one
another
2. dispersion
forces –
caused by
the motion
of electrons
(weakest of
all forces)
Hydrogen Bond

Hydrogen Bonds - forces
where a hydrogen atom is weakly
attracted to an unshared electron
pair of another atom

Strongest of all intermolecular forces