Chapter 5
Chemical Bonding:
The Covalent Bond
Model
Chapter 5
Chapter Outline
5.1
5.2
5.3
5.4
5.5
5.6
5.7
5.8
5.9
5.10
5.11
5.12
The covalent bond model
Lewis structures for molecular compounds
Single, double, and triple covalent bonds
Valence electrons and number of covalent bonds
formed
Coordinate covalent bonds
Systematic procedures for drawing Lewis structures
Bonding in compounds with polyatomic ions present
Molecular geometry
Electronegativity
Bond polarity
Molecular polarity
Recognizing and naming binary molecular compounds
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Section 5.1
Covalent Bond Model
Key Differences Between Ionic and Covalent Bonding
1.  Ionic bonds form between a metal and a
nonmetal
–  Covalent bonds usually form between nonmetals
2.  Ionic bonds involve electron transfer
–  Covalent bonds involve electron sharing
3.  Ionic compounds do not contain discrete
molecules
–  Basic structural unit of covalent bonds is a molecule
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Section 5.1
Covalent Bond Model
Key Differences Between Ionic and Covalent Bonding
4.  All ionic compounds are solids at room
temperature
–  Covalent compounds can be solids, liquids, or gases
5.  Soluble ionic solids form aqueous solutions that
conduct electricity
–  Soluble covalent compounds usually produce a
nonconducting aqueous solution
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Section 5.1
Covalent Bond Model
Covalent Bond
•  A chemical bond resulting from two nuclei
attracting the same shared electrons
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Section 5.1
Covalent Bond Model
Hydrogen Molecule
•  Electron sharing can occur only when electron
orbitals from two different atoms overlap
–  Produces increased stability
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Section 5.1
Covalent Bond Model
Lewis Notation for the Hydrogen Atom
•  Two shared electrons help each hydrogen atom
achieve a helium noble-gas configuration
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Section 5.1
Covalent Bond Model
A covalent bond occurs when:
a. there is a transfer of electrons only between similar or
identical atoms.
b. there is a transfer of electrons only between similar
atoms.
c. there is a sharing of valence electrons between similar or
identical atoms.
d. there is a sharing of electrons between a metal and
nonmetal.
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Section 5.1
Covalent Bond Model
A covalent bond occurs when:
a. there is a transfer of electrons only between similar or
identical atoms.
b. there is a transfer of electrons only between similar
atoms.
c. there is a sharing of valence electrons between similar or
identical atoms.
d. there is a sharing of electrons between a metal and
nonmetal.
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Section 5.2
Lewis Structures for Molecular Compounds
Bonding Electrons
•  Pairs of valence electrons that are shared
between atoms in a covalent bond
•  Shared electron pairs are represented with
dashes
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Section 5.2
Lewis Structures for Molecular Compounds
Nonbonding Electrons
•  Pairs of valence electrons on an atom that are
not involved in electron sharing
•  Referred to as unshared electron pairs or lone
electron pairs
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Section 5.2
Lewis Structures for Molecular Compounds
What is the difference between bonding electrons
and lone electron pairs?
a. Bonding electrons are valence electrons that are shared between
atoms and lone electron pairs are nonbonding electrons.
b. Bonding electrons are shared between two identical atoms and lone
electron pairs are shared between two nonidentical atoms.
c. Bonding electrons and lone electron pairs are interchangeable terms
used to identify electrons in a covalent bond.
d. Bonding electrons are involved in the formation of covalent bonds
and lone electron pairs are involved in formation of ionic bonds.
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Section 5.2
Lewis Structures for Molecular Compounds
What is the difference between bonding electrons
and lone electron pairs?
a. Bonding electrons are valence electrons that are shared between
atoms and lone electron pairs are nonbonding electrons.
b. Bonding electrons are shared between two identical atoms and lone
electron pairs are shared between two nonidentical atoms.
c. Bonding electrons and lone electron pairs are interchangeable terms
used to identify electrons in a covalent bond.
d. Bonding electrons are involved in the formation of covalent bonds
and lone electron pairs are involved in formation of ionic bonds.
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Section 5.3
Single, Double, and Triple Covalent Bonds
Single Covalent Bond
•  Covalent bond in which two atoms share one
pair of electrons
H–H
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Section 5.3
Single, Double, and Triple Covalent Bonds
Double Covalent Bond
•  Covalent bond in which two atoms share two
pairs of electrons
O=C=O
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Section 5.3
Single, Double, and Triple Covalent Bonds
Triple Covalent Bond
•  Covalent bond in which two atoms share three
pairs of electrons
N≡N
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Section 5.3
Single, Double, and Triple Covalent Bonds
How many electrons are shared in a double
covalent bond?
a. 2
b. 4
c. 6
d. 8
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Section 5.3
Single, Double, and Triple Covalent Bonds
How many electrons are shared in a double
covalent bond?
a. 2
b. 4
c. 6
d. 8
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Section 5.4
Valence Electrons and Number of Covalent Bonds Formed
•  Atoms of nonmetallic elements have a strong
tendency to form a specific number of covalent
bonds
•  Number of bonds formed is equal to the number
of electrons the nonmetallic atom must share to
obtain an octet of electrons
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Section 5.4
Valence Electrons and Number of Covalent Bonds Formed
Oxygen (6 Valence Electrons, 2 Octet Vacancies)
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Section 5.4
Valence Electrons and Number of Covalent Bonds Formed
Nitrogen (5 Valence Electrons, 3 Octet Vacancies)
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Section 5.4
Valence Electrons and Number of Covalent Bonds Formed
Carbon (4 Valence Electrons, 4 Octet Vacancies)
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Section 5.4
Valence Electrons and Number of Covalent Bonds Formed
How many covalent bonds can oxygen form?
a. 1
b. 2
c. 4
d. 8
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Section 5.4
Valence Electrons and Number of Covalent Bonds Formed
How many covalent bonds can oxygen form?
a. 1
b. 2
c. 4
d. 8
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Section 5.5
Coordinate Covalent Bonds
•  Covalent bond in which both electrons in a
shared pair come from one of the two atoms
involved in the bond
•  Enable an atom that has two vacancies in its
valence electron shell to share a pair of
nonbonding electrons on another atom
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Section 5.5
Coordinate Covalent Bonds
The HOClO (HClO2) Molecule
•  In HOCl, all the bonds are ordinary covalent
bonds
•  In HClO2, the new chlorine–oxygen bond is a
coordinate covalent bond
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Section 5.5
Coordinate Covalent Bonds
Figure 5.3 - Formation of a Regular Covalent Bond Vs a
Coordinate Covalent Bond
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Section 5.5
Coordinate Covalent Bonds
•  Once a coordinate covalent bond forms, it is
indistinguishable from other covalent bonds in a
molecule
•  Atoms participating in coordinate covalent bonds
deviate from their regular bonding patterns
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Section 5.5
Coordinate Covalent Bonds
When does a coordinate covalent bond form?
a. When there is a sharing of four electrons between atoms
b. When both electrons of a shared pair come from one of
the two atoms in the bond
c. When each electron of a shared pair comes from each
atom involved in the bond
d. Both (b) and (c)
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Section 5.5
Coordinate Covalent Bonds
When does a coordinate covalent bond form?
a. When there is a sharing of four electrons between atoms
b. When both electrons of a shared pair come from one of
the two atoms in the bond
c. When each electron of a shared pair comes from each
atom involved in the bond
d. Both (b) and (c)
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Section 5.6
Systematic Procedures for Drawing Lewis Structures
Steps for Writing Lewis Structures
1.  Calculate the total number of valence electrons
available in the molecule by adding together the
valence electron counts for all atoms in the
molecule
Example - SO2
Sulfur contains 6 valence electrons
Oxygen contains 6 valence electrons
Total number = 6 + 2(6) = 18
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Section 5.6
Systematic Procedures for Drawing Lewis Structures
Steps for Writing Lewis Structures
2.  Write the chemical symbols of the atoms in the
molecule in the order in which they are bonded
to one another
–  Place a single covalent bond, involving two electrons,
between each pair of bonded atoms
–  Determine the central atom which appears only once
in the formula
Example - SO2
O:S:O
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Section 5.6
Systematic Procedures for Drawing Lewis Structures
Steps for Writing Lewis Structures
3.  Add nonbonding electron pairs to the structure
such that each atom bonded to the central atom
has an octet of electrons
–  Remember that for hydrogen, an octet is only two
electrons
Example - SO2
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Section 5.6
Systematic Procedures for Drawing Lewis Structures
Steps for Writing Lewis Structures
4.  Place any remaining electrons on the central
atom of the structure
Example - SO2
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Section 5.6
Systematic Procedures for Drawing Lewis Structures
Steps for Writing Lewis Structures
5.  If there are not enough electrons to give the
central atom an octet:
–  Use one or more pairs of nonbonding electrons on the
atoms bonded to the central atom to form double or
triple bonds
Example - SO2
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Section 5.6
Systematic Procedures for Drawing Lewis Structures
Steps for Writing Lewis Structures
6.  Count the total number of electrons in the
completed Lewis structure to make sure it is
equal to the total number of valence electrons
available for bonding, as calculated in Step 1
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Section 5.6
Systematic Procedures for Drawing Lewis Structures
Concept Check
•  Draw a Lewis structure for each of the following
molecules
H2
F2
HF
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Section 5.6
Systematic Procedures for Drawing Lewis Structures
Concept Check
•  Draw a Lewis structure for each of the following
molecules
H2
H H
F2
F F
HF
H F
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Section 5.6
Systematic Procedures for Drawing Lewis Structures
Concept Check
•  Draw a Lewis structure for each of the following
molecules
NH3
CO2
CCl4
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Section 5.6
Systematic Procedures for Drawing Lewis Structures
Concept Check
•  Draw a Lewis structure for each of the following
molecules
H
NH3
H N H
CO2
O C O
CCl4
Cl
Cl C Cl
Cl
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40
Section 5.6
Systematic Procedures for Drawing Lewis Structures
What is the total number of dots in the Lewis
structure for SO2?
a. 6
b. 12
c. 16
d. 18
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Section 5.6
Systematic Procedures for Drawing Lewis Structures
What is the total number of dots in the Lewis
structure for SO2?
a. 6
b. 12
c. 16
d. 18
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Section 5.7
Bonding in Compounds with Polyatomic Ions Present
Ionic Compounds Containing Polyatomic Ions
•  Covalent bonding exists within the polyatomic
ion and ionic bonding exists between it and ions
of opposite charge
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Section 5.7
Bonding in Compounds with Polyatomic Ions Present
Lewis Structure of Potassium Sulfate, K2SO4
•  Polyatomic ion charge is not localized on a
particular atom but rather is associated with the
ion as a whole
•  Ionic charge is shown outside the brackets
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Section 5.7
Bonding in Compounds with Polyatomic Ions Present
How many valence electrons are present in the
polyatomic ion SO42-?
a. 28
b. 30
c. 32
d. 38
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Section 5.7
Bonding in Compounds with Polyatomic Ions Present
How many valence electrons are present in the
polyatomic ion SO42-?
a. 28
b. 30
c. 32
d. 38
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Section 5.8
Molecular Geometry
•  Description of the three-dimensional
arrangement of atoms within a molecule
•  Key factor in determining the physical and
chemical properties of a substance
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Section 5.8
Molecular Geometry
VSEPR Theory
•  VSEPR - Valence shell electron-pair repulsion
•  Set of procedures for predicting the molecular
geometry of a molecule using the information
contained in the molecule’s Lewis structure
•  Structure around a given atom is determined
principally by minimizing electron pair repulsions
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Section 5.8
Molecular Geometry
VSEPR Electron Group
•  Collection of valence electrons present in a
localized region about the central atom in a
molecule
•  Four electrons in a double bond or six electrons
in a triple bond are localized in the region
between two bonded atoms
–  Similar to the localization of two electrons on a single
bond
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Section 5.8
Molecular Geometry
Steps to Apply the VSEPR Model
1.  Draw a Lewis structure for the molecule and
identify the specific atom for which geometrical
information is desired
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Section 5.8
Molecular Geometry
Steps to Apply the VSEPR Model
2.  Determine the number of VSEPR electron
groups present about the central atom
–  No distinction is made between bonding and
nonbonding electron groups
–  Single, double, and triple bonds are all counted
equally as one electron group
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Section 5.8
Molecular Geometry
Steps to Apply the VSEPR Model
3.  Predict the VSEPR electron group arrangement
about the atom by assuming that the electron
groups orient themselves in a manner that
minimizes repulsions
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Section 5.8
Molecular Geometry
VSEPR: Two Electron Groups
•  Carbon dioxide (CO2) and hydrogen cyanide
(HCN)
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Section 5.8
Molecular Geometry
VSEPR: Three Electron Groups
•  Formaldehyde (H2CO) and sulfur dioxide (SO2)
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Section 5.8
Molecular Geometry
VSEPR: Four Electron Groups
•  Methane (CH4), ammonia (NH3), and water
(H2O)
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Section 5.8
Molecular Geometry
Molecules with More Than One Central Atom
•  Molecular shape of molecules that contain more
than one central atom can be obtained by:
–  Considering each central atom separately and then
combining the results
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Section 5.8
Molecular Geometry
Molecules with More Than One Central Atom
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Section 5.8
Molecular Geometry
Concept Check
•  Determine the shape for each of the following
molecules, and mention their bond angles
HCN
NH3
O3
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Section 5.8
Molecular Geometry
Concept Check
•  Determine the shape for each of the following
molecules, and mention their bond angles
HCN - Linear, 180°
NH3 - Trigonal pyramid, 107°
O3 - Bent, 120°
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Section 5.8
Molecular Geometry
What theory is used to predict the molecular
geometry of a molecule?
a. Valence shell electron pair repulsion theory
b. Valence electron pair repulsion theory
c. Electron pair repulsion theory
d. Valence electron pair shell repulsion theory
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Section 5.8
Molecular Geometry
What theory is used to predict the molecular
geometry of a molecule?
a. Valence shell electron pair repulsion theory
b. Valence electron pair repulsion theory
c. Electron pair repulsion theory
d. Valence electron pair shell repulsion theory
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Section 5.9
Electronegativity
•  Ability of an atom in a molecule to attract shared
electrons to itself
•  Measure of the relative attraction that an atom
has for the shared electrons in a bond
•  Increases from left to right across periods, and
from bottom to top within groups of the periodic
table
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Section 5.9
Electronegativity
Figure 5.11 - Pauling Electronegativity Values
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Section 5.9
Electronegativity
Concept Check
•  What is the general trend for electronegativity
across rows and columns on the periodic table?
•  Explain the trend
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Section 5.9
Electronegativity
Concept Check
•  If lithium and fluorine react, which element will
have more attraction for an electron? Why?
•  In a bond between fluorine and iodine, which has
more attraction for an electron? Why?
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Section 5.9
Electronegativity
Concept Check
•  If lithium and fluorine react, which element will
have more attraction for an electron? Why?
•  In a bond between fluorine and iodine, which has
more attraction for an electron? Why?
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Section 5.9
Electronegativity
Which of the following elements has the highest
electronegativity?
a. H
b. He
c. O
d. Cl
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Section 5.9
Electronegativity
Which of the following elements has the highest
electronegativity?
a. H
b. He
c. O
d. Cl
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Section 5.10
Bond Polarity
•  Measure of the degree of inequality in the
sharing of electrons between two atoms in a
chemical bond
•  Greater the electronegativity difference between
the two bonded atoms, the greater the polarity of
the bond
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Section 5.10
Bond Polarity
Nonpolar Covalent Bond
•  Covalent bond in which there is equal sharing of
electrons between two atoms
Example - H2
H–H
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Section 5.10
Bond Polarity
Polar Covalent Bond
•  Covalent bond in which there is unequal sharing
of electrons between two atoms
Example - HCl
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Section 5.10
Bond Polarity
Polar Covalent Bond and Fractional Charges
•  Creates fractional positive and negative charges
on atoms
•  Head of the arrow is positioned above the more
electronegative element
–  Tail is positioned above the less electronegative
element
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Section 5.10
Bond Polarity
Figure 5.13 - Chemical Bond Type Classification Based on
Electronegativity Difference
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Section 5.10
Bond Polarity
Exercise
•  Arrange the following bonds in the descending
order of polarity
a) N–F
O–F
C–F
b) C–F
N–O
Si–F
c) Cl–Cl
B–Cl
S–Cl
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Section 5.10
Bond Polarity
Exercise
•  Arrange the following bonds in the descending
order of polarity
a) N–F
C–F
b) C–F
Si–F
c) Cl–Cl
B–Cl
O–F
N–F
N–O
C–F
B–Cl
S–Cl
C–F
O–F
Si–F
N–O
S–Cl
Cl–Cl
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Section 5.10
Bond Polarity
Concept Check
•  Which of the following bonds would be the least
polar yet still be considered polar covalent?
Mg–O
C–O
O–O
Si–O
N–O
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Section 5.10
Bond Polarity
Concept Check
•  Which of the following bonds would be the least
polar yet still be considered polar covalent?
Mg–O
C–O
O–O
Si–O
N–O
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Section 5.10
Bond Polarity
Concept Check
•  Which of the following bonds would be the most
polar without being considered ionic?
Mg–O
C–O
O–O
Si–O
N–O
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Section 5.10
Bond Polarity
Concept Check
•  Which of the following bonds would be the most
polar without being considered ionic?
Mg–O
C–O
O–O
Si–O
N–O
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Section 5.10
Bond Polarity
Which statement best describes a polar covalent
bond?
a. In a polar covalent bond, there is equal sharing of
electrons between two atoms.
b. In a polar covalent bond, there is unequal sharing of
electrons between two atoms.
c. In a polar covalent bond, there is equal sharing of
electrons between identical atoms.
d. In a polar covalent bond, there is unequal sharing of
electrons between identical atoms.
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Section 5.10
Bond Polarity
Which statement best describes a polar covalent
bond?
a. In a polar covalent bond, there is equal sharing of
electrons between two atoms.
b. In a polar covalent bond, there is unequal sharing of
electrons between two atoms.
c. In a polar covalent bond, there is equal sharing of
electrons between identical atoms.
d. In a polar covalent bond, there is unequal sharing of
electrons between identical atoms.
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Section 5.11
Molecular Polarity
•  Measure of the degree of inequality in the
attraction of bonding electrons to various
locations within a molecule
•  Polar molecule: Molecule in which there is an
unsymmetrical distribution of electron charge
•  Nonpolar molecule: Molecule in which there is
a symmetrical distribution of electron charge
•  Depends on bond polarities and molecular
geometry
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Section 5.11
Molecular Polarity
Nonpolar Molecule: CO2
•  Effects of the two polar bonds are canceled as a
result of the oxygen atoms being arranged
symmetrically around the carbon atom
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Section 5.11
Molecular Polarity
Polar Molecules: H2O and HCN
•  In H2O, bond polarities do not cancel one
another because of the nonlinearity of the
molecule
•  In HCN, nitrogen is more electronegative than
carbon and hydrogen
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Section 5.11
Molecular Polarity
Concept Check
True or false
•  A molecule that has polar bonds will always be
polar
–  If true, explain why
–  If false, provide a counter-example
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Section 5.11
Molecular Polarity
Concept Check
True or false
•  A molecule that has polar bonds will always be
polar
–  If true, explain why
–  If false, provide a counter-example
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Section 5.11
Molecular Polarity
Let’s Think About It
•  Draw the Lewis structure for SiO2
•  Does SiO2 contain polar bonds?
•  Is the molecule polar or nonpolar overall? Why?
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Section 5.11
Molecular Polarity
Concept Check
•  Which of the following molecules are polar?
F2
HF
NH3
SO2
CCl4
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Section 5.11
Molecular Polarity
Concept Check
•  Which of the following molecules are polar?
F2
HF
NH3
SO2
CCl4
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Section 5.11
Molecular Polarity
Which of the following is a polar compound?
a. CO2
b. BeCl2
c. H2O
d. CH4
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90
Section 5.11
Molecular Polarity
Which of the following is a polar compound?
a. CO2
b. BeCl2
c. H2O
d. CH4
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Section 5.12
Recognizing and Naming Binary Molecular Compounds
Binary Molecular Compound
•  Molecular compound in which only two
nonmetallic elements are present
•  Full name of the nonmetal of lower
electronegativity is given first followed by:
–  A separate word containing the stem of the name of
the more electronegative nonmetal
–  The suffix –ide
•  Numerical prefixes, giving numbers of atoms,
precede the names of both nonmetals
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Section 5.12
Recognizing and Naming Binary Molecular Compounds
Binary Covalent Compounds
•  Examples CO2
Carbon dioxide
SF6
Sulfur hexafluoride
N2O4
Dinitrogen tetroxide
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93
Section 5.12
Recognizing and Naming Binary Molecular Compounds
Table 5.1 - Numerical Prefixes for Numbers 1 Through 10
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94
Section 5.12
Recognizing and Naming Binary Molecular Compounds
Table 5.2 - Selected Binary Molecular Compounds that have
Common Names
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95
Section 5.12
Recognizing and Naming Binary Molecular Compounds
Exercise
•  Which of the following compounds is named
incorrectly?
a) 
b) 
c) 
d) 
e) 
NO2
P2O5
PCl3
SO3
ICl
nitrogen dioxide
phosphorus pentoxide
phosphorus trichloride
sulfur trioxide
iodine monochloride
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96
Section 5.12
Recognizing and Naming Binary Molecular Compounds
Exercise
•  Which of the following compounds is named
incorrectly?
a) 
b) 
c) 
d) 
e) 
NO2
P2O5
PCl3
SO3
ICl
nitrogen dioxide
phosphorus pentoxide
phosphorus trichloride
sulfur trioxide
iodine monochloride
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97
Section 5.12
Recognizing and Naming Binary Molecular Compounds
What does the prefix penta- mean?
a. 2
b. 3
c. 4
d. 5
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Section 5.12
Recognizing and Naming Binary Molecular Compounds
What does the prefix penta- mean?
a. 2
b. 3
c. 4
d. 5
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99
Chapter 5
Concept Question 1
An unknown is thought to be an aqueous salt
solution; however, it was determined not to be a
conductor of electrical current. Through analysis, it
is found that the unknown contains carbon,
hydrogen, and chlorine in a ratio of 1:1:3. Which
one of the following represents the Lewis structure,
molecular geometry, and classification of the
unknown?
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100
Chapter 5
Concept Question 1
; triangular pyramidal; covalent compound
••
••
H
••
a.  Cl C
Cl
••
Cl
; tetrahedral; ionic compound
••
••
H
••
b.  Cl C Cl
••
Cl
••
H
•• ••
C Cl
d.  Cl
•• •• ••
Cl
; tetrahedral; covalent compound
••
••
••
••
••
••
••
••
••
••
••
••
••
••
••
••
; tetrahedral; covalent compound
••
H
•• •• ••
c.  Cl C Cl
•• •• ••
Cl
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101
Chapter 5
Concept Question 1
; triangular pyramidal; covalent compound
••
••
H
••
a.  Cl C
Cl
••
Cl
; tetrahedral; ionic compound
••
••
H
••
b.  Cl C Cl
••
Cl
••
H
•• ••
C Cl
d.  Cl
•• •• ••
Cl
; tetrahedral; covalent compound
••
••
••
••
••
••
••
••
••
••
••
••
••
••
••
••
; tetrahedral; covalent compound
••
H
•• •• ••
c.  Cl C Cl
•• •• ••
Cl
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102
Chapter 5
Concept Question 2
A liquid molecular compound is given to you and you are told to
handle it with caution. It can cause damage or death if it is
consumed in large quantities, if it enters your lungs, if you are hit
with a frozen chunk of this material, or if a hot solution of it is
spilled on you; however, this liquid is necessary to sustain
life. The liquid is a pure molecular compound, has no taste, and
is known to contain hydrogen and oxygen. Identify the liquid, its
molecular geometry, and whether it is polar or nonpolar.
a. Water; linear; polar
b. Water; linear; nonpolar
c. Water; angular; nonpolar
d. Water; angular; polar
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103
Chapter 5
Concept Question 2
A liquid molecular compound is given to you and you are told to
handle it with caution. It can cause damage or death if it is
consumed in large quantities, if it enters your lungs, if you are hit
with a frozen chunk of this material, or if a hot solution of it is
spilled on you; however, this liquid is necessary to sustain
life. The liquid is a pure molecular compound, has no taste, and
is known to contain hydrogen and oxygen. Identify the liquid, its
molecular geometry, and whether it is polar or nonpolar.
a. Water; linear; polar
b. Water; linear; nonpolar
c. Water; angular; nonpolar
d. Water; angular; polar
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104