Chemistry 11
Resource: Chang’s Chemistry Chapter 9
Objectives
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Describe the covalent bond as the electrostatic
attraction between a pair of electrons and
positively charged nuclei.
Describe how the covalent bond is formed as a
result of electron sharing.
Deduce the Lewis structures of molecules and
ions for up to four electron pairs on each atom.
State and explain the relationship between the
number of bonds, bond length and bond strength.
Predict whether a compound of two elements
would be covalent from the position of the
elements in the periodic table or from their
electronegativity values.
Objectives
Predict the relative polarity of bonds from
electronegativity values.
 Predict the shape and bond angles for
species with four, three, and two negative
charge centers on the central atom using
the valence shell electron pair repulsion
theory (VSEPR).
 Predict whether or not a molecule is polar
from its molecular shape and bond
polarities.
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The covalent bond
A bond in which two electrons are
shared by two atoms.
 Covalent compounds are compounds
that contain only covalent bonds.
 Most of the molecules found in nature
contain covalent bonds.
 Model
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The covalent bond
Ionic vs. covalent bonding
Quality
Mechanism
Ionic
“giving up”
electrons
Covalent
“sharing” two
electrons
Force of
attraction
between
oppositely
charged ions
between a pair
of electrons and
positively
charged nuclei
Why do atoms form covalent bonds?
The covalent bond
Covalent bonds are formed to achieve
stability—the octet rule.
 An atom other than hydrogen tends to
form bonds until it is surrounded by eight
valence electrons (or the configuration
of a noble gas).
 Exceptions to the octet rule
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The covalent bond
Ways of representing the covalent bond
Cl - Cl
The covalent bond
Draw the Lewis structures for:
 H2
 CH4
The covalent bond
H–H
Hi
H–C–H
Hi
The covalent bond
Draw the Lewis structure for ammonia
(NH3)
The covalent bond
Ammonia (NH3)
The covalent bond
Draw the Lewis structures for:
 H 2O
 PCl3
The covalent bond
lone pair
Multiple covalent bonds
Number of electrons shared
Number of covalent
bonds
1
Number of
electrons shared
1 pair
2
3
2 pairs
3 pairs
Multiple covalent bonds
Representations of a double covalent bond
O=O
Multiple covalent bonds
How do multiple covalent bonds affect the
atoms that are bonded together?
Is a single bond stronger or weaker than a
double bond?
Multiple covalent bonds
Bond energies of some diatomic
molecules
Bond
Bond energy
(kJ/mol)
C–C
C=C
C≡C
C–N
347
620
812
276
C=N
615
Multiple covalent bonds
Average bond lengths of common covalent
bonds
Bond
Bond length
(pm)
C–C
C=C
C≡C
C–N
154
133
120
143
C=N
138
Multiple covalent bonds
The greater the number of covalent bonds,
the stronger and shorter the bond
between two atoms.
Electronegativity
Recall the definition of electronegativity.
 In covalent bonds, we expect the
electrons to be equally shared.
 Consider the covalently bonded
compound HF (p 357 Figure 9.4)
 Why do the electrons “spend more time”
around the fluorine nucleus?
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Electronegativity
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We can predict the relative
electronegativity of an element using the
periodic table based on the following
factors:
 shielding
 nuclear charge
Electronegativity
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This difference in electronegativity results
in a polar covalent bond.
Electronegativity
So what’s the difference between a polar covalent
bond and an ionic bond?
Electronegativity
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The 2.0 rule of electronegativity:
A difference in electronegativity greater
than or equal to 2.0 will most likely
indicate an ionic bond.
How else can we predict whether or not a
bond between two elements is ionic or
covalent?
Electronegativity
Classify the following bonds as ionic, polar
covalent, or covalent:
(a) the bond in hydrochloric acid (HCl)
(b) the bond in potassium fluoride (KF)
(c) the CC bond in ethane (H3CCH3)
Electronegativity
(a)
(b)
(c)
the bond in HCl – polar covalent
the bond in KF - ionic
the CC bond in ethane (H3CCH3) –
purely covalent
Typically, only atoms of the same element
can be joined by a pure covalent bond.
Practice exercises
Covalent bonding:
p 379, #9.30-9.32
Electronegativity and bond type
p 379 #9.36, 38, 40
Lewis structures
p 379 #9.44