Covalent Bonding
Chapter 8
Ch. 8 Vocabulary
2
 Covalent bond
 Exothermic Reaction
 Molecule
 Structural Formula
 Lewis Structure
(Ch. 8.3)
 Polar Covalent Bond
 Sigma bond
 Pi bond
 Bond Dissociation
Energy
 Endothermic Reaction
OBJECTIVES
3
 Draw Lewis Structures of Covalent compounds
 Name covalent compounds
 Write formulas for covalent compounds
 Describe characteristics of covalent molecules
 Review electronegativity and compare they types of
bonds of different molecules using the
electronegativity scale
Covalent Bonding
 Remember the Octet Rule:
 Similar arrangement of valence electrons
 Electron arrangements determines chemical properties
 Presents a model of chemical stability
 RULE: Atoms become stable by having 8 electrons
in their outer energy level (2 for smaller atoms)
 When they have gotten 8 electrons they have achieved
NOBLE GAS CONFIGURATION (NGC)
 One way to get NGC is by transferring electrons as
in an ionic bond.
4
Covalent Bonding
 The other way to achieve Noble Gas Configuration:
 Sharing Electrons
 Take the case of Water (H2O)
 Hydrogen can’t transfer its electron, otherwise it would
be just a proton and not a noble gas configuration.
 It also can’t gain one. Hydrogen and oxygen can’t both
gain electrons.
5
Sharing of Electrons
 In the case of NaCl, chlorine has a much stronger
affinity for electrons and sodium holds its valence
electron very weakly.
 In the case of H2O, both hydrogen and oxygen have
similar affinities.
 In other words, the attraction for electrons is not
strong enough.
 So what do they do?
6
Molecular Elements
 Molecules can vary greatly in size
 Can be just two atoms (CO) to thousands or millions of
atoms (DNA)
 :C O:
 Two or more atoms of the same element can form a
covalent bond – this is called a molecular element.
7
Molecular Elements
H-H
N N O=O
F-F
Cl-Cl
Br-Br
 They are:
I -I
 Hydrogen, Nitrogen, Oxygen, Fluorine, Chlorine, Bromine, and Iodine
 Their Formulas are:
H2, N2, O2, F2, Cl2, Br2, and I2
 All are gases except Br2 (liquid) and I2 (Solid)

 Diatomic Molecules –
 Seven non metal elements are found naturally as
molecular elements of two identical atoms
8
Molecular Elements
 Allotropes
 Examples:
9
Molecular Elements
 Allotropes of Phosphorus
Red
White
Black
10
Electron Dot Structures
(Lewis Symbols)
• Remember (Chapter 5) electron dot structures (Lewis
symbols).
• The number of electrons available for bonding are
indicated by unpaired dots.
• Place the electrons one four sides of a square around the
·
element symbol, e.g. ·C·
·
• Lewis Structures – quick method – Pair up the single
electrons in each atom to form a bond.
11
Electron Dot Structures
(Lewis Symbols)
12
Chlorine
forms
a
covalent
bond
with
itself
Cl2
13
Cl
Cl
How
will
two
chlorine
atoms
react?
14
Cl Cl
octet
15
Cl Cl
The octet is achieved by
each atom sharing the
electron pair in the middle
16
Lone pair
Cl Cl
It is called a SINGLE BOND
also called a sigma (σ) bond
Lone Pairs are non-bonding electrons or unshared
electrons
The Cl2 molecule has how many lone pairs? 6
17
Single bonds are abbreviated
with a dash
Cl Cl
This is the chlorine molecule,
Cl2
Contrast this with ionic bonding - next slide
18
NaCl
This is the formation of an ionic bond.
Na
+
Cl
-
electron transfer
and the formation of ions
Cl2
This is the formation of a covalent bond.
Cl
Cl
sharing of a pair of electrons
and the formation of molecules
19
Covalent Bonding
20
Molecular Lewis Structures – Other
Single Covalent Bonds
Bonding pairs ?
Lone pairs?
21
O2
Oxygen is also one of the diatomic molecules
22
O
O
Each atom has two unpaired electrons
23
O
O
Oxygen atoms are highly electronegative.
So both atoms want to gain two electrons.
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O O
Both electron pairs are shared.
25
O O
6 valence electrons
plus 2 shared electrons
= full octet
26
O O
two bonding pairs,
making a double bond
(a sigma (σ) and a pi (π) bond)
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O O O =O
For convenience, the double bond
can be shown as two dashes.
How many bonding pairs in O2?
How many lone pairs in O2?
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O =O
This is the oxygen molecule,
O2
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σ and π Bonds
 Single bond (Cl2) has one σ bond
 Bond directly between the atoms
 Double bond (O2 and C2H4) have two bonds
–one σ and one π
 σ bond goes directly between the atoms
 π bond goes above the axis
30
σ and π Bonds in a Molecule
31
Strength of Covalent Bonds
 The energy required to break a covalent bond is called
the bond dissociation energy.
 As number of bonds increase, length gets shorter and
bond energy increases.
 So, H-H bonds are fairly easy to break and makes
hydrogen gas reactive.
 Nitrogen has a triple bond, which takes a lot of energy
to break and is pretty much inert.
:N  N:
32
Strength of Covalent Bonds
 Exothermic reactions release energy. They occur
when the bonds of the products are more stable than
the bonds of the reactants. In other words, energy is
released forming the new bonds in the products.
Example: CH4 + O2  H2O + CO2 + energy
• Endothermic reactions absorb energy. The occur
when bonds of the reactants are more stable than the
bonds of the products. In other words, energy has to
be entered to make the products.
Example: 2H2O(l) + energy  2H2(g) + O2(g)
33
Chapter 8.2 - Formulas and Names
 Rules for Naming Binary Inorganic Compounds
 Write out the name of the first nonmetal (Left
most or bottom most first)
 Follow it by the name of the second nonmetal
and end in –ide
 Add a prefix to the name of each element to
denote how many are present.

Example: SO2 is sulfur dioxide because sulfur is below
oxygen in the same group.
34
Prefixes for Molecular Compounds
Number of Atoms
1
2
3
4
5
6
7
8
9
10
Prefix
MonoDiTriTetraPentaHexaHeptaOctaNonaDeca-
35
Formulas and Names
 Rules for Naming Binary Inorganic Compounds
 Omit mono- if there is only one atom of the first
element
 If o-o or a-o vowels appear next to each other, the first of
the pair is omitted for easier pronunciation.

Example: NO is Nitrogen Monoxide rather than
mononitrogen monooxide
36
Practice Formulas and Names
Formula
Name
 NO
 NO2
 N2O
 N2O5
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More Practice
Formula or Name
Name or Formula
 CCl4
 CO
 Diarsenic Trioxide
 Sulfur hexafluoride
 P2O5
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Common Names
 Some molecules are so common they have common
names.
 What is the common name of dihydrogen monoxide?
 Another common compound is NH3. What is this
commonly called?
39
Common Names of Compounds
 Acids
 HCl – Hydrochloric Acid
 H2SO4 – Sulfuric Acid
 H3PO4 – Phosphoric Acid
 HNO3 – Nitric Acid
 HC2H3O2 – Acetic Acid (vinegar)
40
Common Names of Compounds
 Bases
 NaOH – Sodium Hydroxide
 KOH – Potassium Hydroxide
 NH3 - Ammonia
41
Organic Compounds
 Organic Compounds
 Made up of carbon
 Carbon is able to bond with other carbon atoms to form
long chains, rings, sheets, and larger networks
 Carbon forms 4 covalent bonds

The simplest with 4 hydrogens to form methane (CH4)
42
Organic Compounds
 Hydrocarbons – compounds containing just hydrogen
and carbon
 The first was methane (CH4)
 The names of others change as more carbons are added.
43
Organic Compounds
Formula
CH4
C 2 H6
C 3 H8
C4H10
C5H12
C6H14
C7H16
C8H18
C9H20
C10H22
Name
Methane
Ethane
Propane
Butane
Pentane
Hexane
Heptane
Octane
Nonane
Decane
#of Carbons
1
2
3
4
5
6
7
8
9
10
44
Ch. 8.3 - Molecular Structures
 Structural Formulas use letter symbols and bond
symbols to show relative positions of atoms.
 You have already done some like H-H, O=O, N≡N,
HCl, and CCl4
 What happens when there are more than two types of
atoms and/or lone pairs are involved?
45
Apply to CO2
• Central atom: C
• Total number valence electrons: 16
(4 from C and 6 each from O)
• Draw skeleton structure: O-C-O
• Place electrons on Os – form octet
• All atoms have octet?
• Draw final structure
46
Try SO2 (neutral molecule)






Central atom S
Total number valence electrons 18
Draw skeleton structure O – S - O
Place electrons on Os – form octet
Place remaining electrons on S
All atoms have octet?
 No? need to form multiple bonds
 Draw final structure
47
Try ClO4- (anion)
 Central atom
 Total valence electrons
 Skeleton structure
48
ClO4- (anion)
 Place remaining electrons on Os
 Place remaining electrons on Cl.
 All atoms have an octet?
49
Ch. 8.4 - Molecular Shapes
 The shape of a molecule determines many of its
physical and chemical properties.
• Molecular geometry (shape) can be determined
with the Valence Shell Electron Pair Repulsion
model, or VSEPR model which minimizes the
repulsion of shared and unshared atoms around
the central atom.
50
Molecular Shapes
 Electron pairs repel each other and cause molecules to
be in fixed positions relative to each other.
 Unshared electron pairs also determine the shape of a
molecule.
 Electron pairs are located in a molecule as far apart as
they can be.
51
52
53
Practice
Other examples from PhET.
Draw the Lewis Structures and note the shapes of the
following:
 NF3
 CS2
 BH3
 NH4+
 SO42-
54
Exceptions to the Octet
Rule
• There are three classes of exceptions to the octet rule:
•
•
•
Molecules with an odd number of electrons;
Molecules in which one atom has less than an octet;
Molecules in which one atom has more than an octet.
Odd Number of Electrons
• Few examples. Generally molecules such as ClO2, NO,
and NO2 have an odd number of electrons.
N O
N O
55
Exceptions to the Octet
Rule
Less than an Octet
• Relatively rare.
• Molecules with less than an octet are typical for
compounds of Groups 1, 2, and 13 (3A).
• Most typical example is BF3.
56
Exceptions to the Octet
Rule
•
•
•
•
More than an Octet
This is the largest class of exceptions.
Atoms from the 3rd period on down can accommodate
more than an octet.
From the third period, the d-orbitals are low enough in
energy to participate in bonding and accept the extra
electron density.
Example: SF6
57
Ch. 8.5 – Electronegativity and
Polarity
 Atoms form bonds to increase their stability.
 They acquire an octet of electrons and get noble gas
configuration.
 The types of bonds we looked at so for are...
Ionic and Covalent
58
Electronegativity and Polarity
 Bonding involves a sharing of electrons.
 Some are shared equally, slightly or none at all.
 Ionic bonds – sharing is so unequal that it is
considered a transfer of electrons
 Covalent – Pure covalent bonds are shared equally,
such as with O2 or Cl2
 Most bonds fall in between.
59
Electronegativity
 Electronegativity – a measure of the ability of an
atom in a bond to attract electrons.
 You can make a comparison between two atoms.
 The difference in electronegativity (ΔEN, pronounced
‘delta E N’) tells you the type of bond.
60
Electronegativity
 The higher the number, the more electronegative the
atom is, i.e. the more it wants the electrons.
 Electronegativity increases across a period. Why?
 Protons increase in the nucleus, the same energy level
being filled and the pull of the nucleus increases the
attraction for the valence electrons. (Higher Z* across a
period)
61
Electronegativity
Use your Electronegativity table to find ΔEN for the
following ionic compounds.

 ΔEN = 3.0
 NaCl
 ΔEN = 2.1
 CsF
 ΔEN = 3.3
 KBr
 ΔEN = 2.0
Which one is “most” ionic?
Which one is “least” ionic?
 LiF
62
Electronegativity
 What happens when the electronegativity numbers are
similar?
 Then the difference between them is very small.
 The bond is then more described as sharing between
the atoms and the bond is covalent.
63
Electronegativity
 What happens when the electronegativity is the same?
 ΔEN = 0 and the bond is described as pure covalent
bond.
 Ex: The F-F bond. EN for F = 4.0 and 4.0 for the other
F. So, ΔEN = 0.
 All electrons are shared equally for the other diatomic
molecules.
64
Electronegativity
 What happens when the ΔEN is greater than 0 and
≤0.5?
 The bonds are still considered covalent. The
compounds they make tend to be gases or low-boiling
liquids at room temperature.
 Example: Find ΔEN for C-H
C EN =2.5; H EN =2.1 ; ΔEN = 0.4
The molecules CH4, C2H6, C3H8 are gases at RT.
65
Electronegativity
 So, ionic bonds are on our scale are between ΔEN≥ 2.0
and ΔEN=3.3.
 (Mostly) Covalent bonds are between ΔEN=0.0 and
ΔEN≤0.5.
 (Pure Covalent bonds are ΔEN=0.0)
 What happens between 0.5 and 2.0?
66
Electronegativity
 Between ΔEN > 0.5 and ΔEN < 2.0, there is a partial
transfer of the shared electrons to the more
electronegative element.
 The bond that forms is called polar covalent.
 A polar covalent bond has some degree of ionic
character.
67
Electronegativity
 So what happens in a polar covalent bond?
 The electron spends time around each element, but
more time around the more electronegative element.
 ΔEN= EN (Cl) – EN(H) = 3.0 – 2.1 = 0.9
δ-
δ+
H
Cl
68
Bond Polarity
HCl is POLAR
because it
has a positive end
and a negative end.
(difference in
electronegativity)
+d -d
••
H Cl
••
••
Cl has a greater share
in bonding electrons
than does H.
Cl has slight negative charge (-d) and H
has slight positive charge (+ d)
69
Practice – Calculate the ∆EN and
classify the type of bond
Pair of Atoms
∆EN and Bond Type
 Ca-S
 Ba-O
 C-Br
 Ca-F
 H-Br
70
Electronegativity
 What is the ΔEN for O-H bonds in water?
 O EN =
 H EN =
 Δ EN =
 How does this affect the properties of water?
71
Electronegativity
ΔEN=1.4
ΔEN=0.0
ΔEN=0.3
ΔEN=0.4
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