Chapter 9Covalent Bonds
AgendaLab Review Quiz –
Review –Chapter 8 / 9
Test – Chapter 8/9
Section 1
 Why
Covalent Bonds
do atoms bond?
 To become noble or stable
 To achieve an octet (are exceptions)
Covalent Bonds
What is a covalent Bond?
 Elements
share electrons
 Majority form between nonmetallic elements
Result?
 A Molecule is formed
Lewis Dot
Review: In your notes draw the following dot structures
H
S
N
Cl
O
C
Ar
Groups and Bonds
Group 17 = 1 Bond
HCl

 Group
PH3
15 = 3 Bonds
Lewis structures
 Group
16 = 2 Bonds
H2S
 Group
 CCl4
14 = 4 Bonds
Sigma Bond
 The
single covalent bond is calls the…
“Sigma Bond”
 Shared
electrons between two atoms
Multiple Covalent Bonds
Why Multiple Bonds? Hint: Think Noble.
 To achieve an Octet!
Example: C2H4
 Draw the central atoms
 Attach the surrounding atoms
 Make sure each atom has an octet
Sigma and pi Bonds
Sigma Bonds

Two atoms share electrons
pi Bonds


Parallel orbitals over lap
Forms double bonds
Example: C2H4
Let’s take a closer look
Lets take a closer look
Strength and Energy
Bond Strength
 The shorter the bond length, the stronger the
bond, the greater the bond-dissociation energy
Bond Energy
 Endothermic
– more energy is needed to break
the bond than is released
 Exothermic
– more energy is released during bond
formation than is required to break it.
Naming Covalent
Section 2:
Naming Covalent Molecules
 Different
1.
2.
3.
than Ionic
First element = entire name
Second element = root + ide
Prefixes used to indicate the # of each type
present in compound
Prefixes
Prefixes for Covalent Molecules
Example
P205
Follow
your rules!
di Phosphorus
pent oxide
Naming Acids
2 types of acids
1. Binary Acids

2.
HCl, H2S, HBr, HCN
Oxyacids
H2SO4,
HClO3, HClO2
Binary Acids
Example: HCl
1. Hydro + root of second element

Hydrochlor…
Add –ic then acid
2.

Hydrochloric acid
Name the following:
HBr, HI, HF, HCN
Oxyacids
Example: H2SO4 and H2SO3
Root of oxyanion present
 Sulfur…
2. If oxyanion ends in …ate add -ic to the end
 H2SO4 = sulfuric acid
3. If oxyanion ends in …ite add –ous to the end
 H2SO3 = sulfurous acid
1.
Section 3
 Molecular
Structures
Molecular Structures
Section 3:
Structural
Uses
Formula
letter symbols and bonds to
show relative positions of atoms.
Lewis Structures
Determining
Lewis Structures
Step 1…
Predict the location of atoms
H is always terminal (end)
Central atom has the least
attraction for shared electrons
(closer to the left of the periodic
table)

Lewis Structures
Step 2…
Find
total # of valence electrons
Step 3…

Determine # of bonding pairs.
Divide # of valence electrons by 2
Step 4…
Place 1 pair (single bond) between
the central atom and terminal
atoms

Lewis Structures
Step 5…
Subtract
pairs used from total
possible pairs (step 3)
Place remaining pairs around
terminal and central atom (octet)
Step 6…

If central atom does not have
octet, use lone pairs as double
bonds.
Lewis Structures
:::
: :: : : : : : :
Example: Carbon dioxide
CO O
 Step 1- predict location
 Step 2 – Total Valence Electrons = 16
 Step 3 – Divide by 2 = 8 pairs
: :
 Step 4 – Central Atom bonds
 Step 5 – Place remaining pairs
 Step 6 – Check Octet rule
::
Lewis Structures
WHAT’S WRONG WITH THIS PICTURE?
• Carbon ~ octet?
• Move electron pairs on each
O to achieve octet around C
C O:
: :
: :
:O
Positive Charges…
Charge Molecules
 You must remove electrons from the total
electrons available for bonding according
to the charge.
 Example NH4+
Total Valence Electrons = 9
Subtract the charge (9-1 = 8)
Divide by 2 (8/2 = 4) ~ bonding pairs
H
H N H
H
+
Charged Molecules
Negative Charges…
 You must add electrons to the total
electrons available for bonding according
to the charge.
 Example PO43- = Total Valence Electrons = 29
ADD the charge (29 + 3 = 32)
Divide by 2 (32/2 = 16) ~ bonding pairs
3-
:
:O:
: :
: :
:O P O:
:O:
:
VSEPR Model
Valence Shell Electron Pair Repulsion

Electrons are located as far apart as they can be


Shared electron pairs repel one another
Lone pairs also repel (even more)
Hybrid Orbitals
 S and p orbitals change to form new equal orbits
 Each bond between atoms represents an s, p or d
orbit
Visualizing the Models
Example #1: BeCl2
 1st Draw the Lewis dot.

Determine the # of shared pairs and lone pairs around
the central atom.



Shared pairs = 2
Lone pairs = 0
2 Total hybrid bonds

S and p (sp)
Visualizing the Models
Example #1: AlCl3
(Exception to the Octet Rule)
 1st Draw the Lewis dot.

Determine the # of shared pairs
and lone pairs



Shared pairs = 3
Lone pairs = 0
3 Total hybrid bonds

s, p and another p (sp2)
Visualizing the Models
Example #1: CH4
 1st Draw the Lewis dot.
Determine the # of shared pairs and lone pairs
around the central atom!
Shared pairs = 4
Lone pairs = 0
4 Total hybrid bonds
s, p, p and another p (sp3)
Visualizing the Models
 Determine


2


H P H
H
the # of shared pairs and lone pairs
Shared pairs = 2
Lone pairs = 1
Total hybrid bonds
S and p (sp3)
Shape

:
Example #1: PH3
 1st Draw the Lewis dot.
Trigonal Pyramidal
Electronegativity and
Polarity
Even or uneven sharing of electrons.
Determined by the electronegativity
Identical atoms share evenly
Bonds between to different atoms.


one atom pulls the electrons closer
creates a relative negative and positive side of the
molecule
RULES: difference between electronegative #s
0.0 - 0.4 = Nonpolar covalent
0.4 - 1.7 = polar covalent
> 1.7 = ionic bond
Properties of Covalent
Compounds
Solubility


Polar Molecules soluble in polar substances
Non-polar in non-polar
Intermolecular force between molecules is called
the “van der Waals” force
3 types of intermolecular foces
1. Nonpolar (weak) = dispersion forces
2. Polar (weak)
= dipole-dipole force
3. Hydrogen Bonds


Very Strong
between H and (F, N, O)
Molecular Shapes
180o
120o
sp
sp2
109.5o
sp3
107.3o
104.5o
Molecular Shapes
90o /
120o
90
o
sp3d
sp3d2
Ionic or Covalent
Does the compound contain a metal?
NO
YES
Is the metal a
transition metal?
YES
Use I, II, III,
IV, V – to
indicate the
charge of the
metal
The compound is
covalent; use prefixes
NO
Don’t use
roman
numerals:
Don’t use
prefixes
Example:
FeO – Iron (II) oxide
Cu2S – Copper (I) Sulfide
Example:
N2O – dinitrogen monoxide
P2O5 – diphosphorus pentoxide
Example:
NaCl – sodium chloride
CaCl2 – calcium chloride