Chapter 4
Covalent
Compounds
Prepared by
Andrea D. Leonard
University of Louisiana at Lafayette
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Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Covalent Compounds
Introduction to Covalent Bonding
Covalent bonds result from the sharing of electrons
between two atoms.
•A covalent bond is a two-electron bond in which
the bonding atoms share the electrons.
•A molecule is a discrete group of atoms held
together by covalent bonds.
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Covalent Compounds
Introduction to Covalent Bonding
Unshared electron pairs are called nonbonded
electron pairs or lone pairs.
Atoms share electrons to attain the electronic
configuration of the noble gas closest to them
in the periodic table.
•H shares 2 e−.
•Other main group elements are stable when they
possess an octet of e− in their outer shell.
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Covalent Compounds
Covalent Bonding and the Periodic Table
Lewis structures are electron-dot structures for
molecules. They show the location of all valence e−.
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Covalent Compounds
Covalent Bonding and the Periodic Table
Covalent bonds are formed when two nonmetals
combine, or when a metalloid bonds to a nonmetal.
How many covalent bonds will a particular atom form?
•Atoms with one, two, or three valence e−
generally form one, two or three bonds,
respectively.
•Atoms with four or more valence electrons form
enough bonds to give an octet.
predicted
number of bonds
=
8 – number of valence e−
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Covalent Compounds
Covalent Bonding and the Periodic Table
Number of bonds
+ Number of lone pairs
= 4
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Sample Problem 4.1
Question
How many covalent bonds are predicted for each atom:
(a) B
(b) N
Solution
(a) B has 3 valence electrons  three covalent bonds
(octet rule)
(b) N has 5 valence electrons  three covalent bonds
(8-5 = 3)
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Lewis Structures
A molecular formula shows the number and identity
of all of the atoms in a compound, but not which
atoms are bonded to each other.
A Lewis structure shows the connectivity between
atoms, as well as the location of all bonding and
nonbonding valence electrons.
General rules for drawing Lewis structures:
•Draw only valence electrons.
•Give every main group element (except H) an
octet of e−.
•Give each hydrogen two e−.
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Lewis Structures
HOW TO Draw a Lewis Structure
Step [1]
Arrange the atoms next to each other that
you think are bonded together.
•Place H and halogens on the periphery, since they
can only form one bond.
H
For CH4:
H C
H
H
not
H C
H
H
H
This H cannot form
two bonds.
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Lewis Structures
HOW TO Draw a Lewis Structure
•Use the common bonding patterns from Figure 4.1
to arrange the atoms
H
For CH5N:
H C
H
Place four atoms
around C, since C
generally forms
four bonds.
H
N H
H
not
H C
H
N H
H
Place three atoms
around N, since N
generally forms
three bonds.
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Lewis Structures
HOW TO Draw a Lewis Structure
Step [2] Count the valence electrons.
•For main group elements, the number of valence
e− is equal to the group number.
•The sum gives the total number of e− that must
be used in the Lewis structure.
For CH3Cl:
1 C x 4e− = 4e−
3 H x 1e− = 3e−
1 Cl x 7e− = 7e−
14 total valence e−
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Lewis Structures
HOW TO Draw a Lewis Structure
Step [3]
Arrange the electrons around the atoms.
•Place one bond (two e−) between every two atoms.
•For main group elements, give no more than 8 e−.
•For H, give no more than 2 e−.
•Use all remaining electrons to fill octets with lone
pairs, beginning with atoms on the periphery.
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Lewis Structures
HOW TO Draw a Lewis Structure
H
For CH3Cl:
H C
e−
2 on
each H
H
4 bonds x 2e− = 8 e−
Cl
e−
8
on Cl
+ 3 lone pairs x 2e− = 6 e−
14 e−
All valence e− have
been used.
•If all valence electrons are used and an atom still
does not have an octet, proceed to Step [4].
Step [4]
Use multiple bonds to fill octets when
needed.
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Lewis Structures
HOW TO Draw a Lewis Structure—Multiple Bonds
•Convert one lone pair to one bonding pair of
electrons for each two electrons needed to
complete an octet.
•A double bond contains four electrons in two
two-electron bonds.
O
O
•A triple bond contains six electrons in three
two-electron bonds.
N
N
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Lewis Structures
Multiple Bonds
Example Draw the Lewis Structure for C2H4.
Step [1]
Step [2]
Arrange the atoms.
H C
C
H
H
H
Count the valence e−.
2 C x 4 e− = 8 e−
4 H x 1 e− = 4 e−
12 e− total
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Lewis Structures
Multiple Bonds
Step [3]
Add the bonds and lone pairs.
H C
C
H
H
C still does not
have an octet.
H
5 bonds x 2 e− = 10 e−
+ 1 lone pair x 2 e− = 2 e−
12 e−
All valence e− have
been used.
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Lewis Structures
Multiple Bonds
Step [4]
To give both C’s an octet, change one
lone pair into one bonding pair of
electrons between the two C atoms,
forming a double bond.
H–C–C–H
H H
H C C
H
H H
Answer
Each C now has an octet.
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Lewis structure of C2H4 (ethylene)
(1) Arrange the atoms. Each C get 2 H’s.
(2) Count the electrons
(3) Add the bonds.
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Example: Lewis structure of HCO3(1) Arrange the atoms.
(2) Count the electrons
(3) Add the bonds.
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There are two structures, A and B, for bicarbonate anion.
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Exceptions to the Octet Rule
•Most of the common elements generally follow
the octet rule.
•H is a notable exception, because it needs only
2 e− in bonding.
•Elements in group 3A do not have enough
valence e− to form an octet in a neutral molecule.
F
F
B F
only 6 e− on B
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Exceptions to the Octet Rule
•Elements in the third row have empty d orbitals
available to accept electrons.
•Thus, elements such as P and S may have
more than 8 e− around them.
O
HO
P OH
OH
10 e− on P
O
HO
S OH
O
12 e− on S
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Exceptions to the Octet Rule
•Elements in the third row have empty d orbitals
available to accept electrons.
•Thus, elements such as P and S may have
more than 8 e− around them.
O
HO
P OH
OH
10 e− on P
O
HO
S OH
O
12 e− on S
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Resonance
When drawing Lewis structures for polyatomic ions:
•Add one e− for each negative charge.
•Subtract one e− for each positive charge.
Answer
For –CN:
C
N
1 C x 4 e− = 4 e−
1 N x 5 e− = 5 e−
C
N
C
N
−
All valence e−
Each atom
are used, but
has an octet.
C lacks an octet.
–1 charge = 1 e−
10 e− total
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Resonance
Drawing Resonance Structures
•Resonance structures are two Lewis structures
having the same arrangement of atoms but a
different arrangement of electrons.
•Two resonance structures of HCO3−:
•Neither Lewis structure is an accurate
representation for HCO3−.
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Resonance
Drawing Resonance Structures
•The true structure is a hybrid of the two resonance
structures.
•Resonance stabilizes a molecule by spreading out
lone pairs and electron pairs in multiple bonds
over a larger region of space.
•A molecule or ion that has two or more resonance
structures is resonance-stabilized.
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Naming Covalent Compounds
HOW TO Name a Covalent Molecule
Example
Name each covalent molecule:
(a) NO2
Step [1]
(b) N2O4
Name the first nonmetal by its element
name and the second using the suffix
“-ide.”
(a) NO2
nitrogen oxide
(b) N2O4
nitrogen oxide
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Naming Covalent Compounds
HOW TO Name a Covalent Molecule
Step [2]
Add prefixes to show the number of
atoms of each element.
•Use a prefix from Table 4.1 for each element.
•The prefix “mono-” is usually omitted.
Exception: CO is named carbon monoxide
•If the combination would place two vowels next
to each other, omit the first vowel.
mono + oxide = monoxide
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Naming Covalent Compounds
HOW TO Name a Covalent
Molecule
(a) NO2
nitrogen dioxide
(b) N2O4
dinitrogen tetroxide
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Molecular Shape
Valence Shell Electron Pair Repulsion
(VSEPR) Theory
•To determine the shape around a given atom,
first determine how many groups surround the
atom.
•A group is either an atom or a lone pair of
electrons.
•Use the VSEPR theory to determine the shape.
•The most stable arrangement keeps the groups
as far away from each other as possible.
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Molecular Shape
Two Groups Around an Atom
•Any atom surrounded by only two groups is
linear and has a bond angle of 180o.
•An example is CO2:
•Ignore multiple bonds in predicting geometry.
Count only atoms and lone pairs.
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Molecular Shape
Three Groups Around an Atom
•Any atom surrounded by three groups is
trigonal planar and has bond angles of 120o.
•An example is H2CO:
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Molecular Shape
Four Groups Around an Atom
•Any atom surrounded by four groups is
tetrahedral and has bond angles of 109.5o.
•An example is CH4:
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Molecular Shape
Four Groups Around an Atom
•If the four groups around the atom include one
lone pair, the geometry is a trigonal pyramid
with bond angles of ~109.5o.
•An example is NH3:
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Molecular Shape
Four Groups Around an Atom
•If the four groups around the atom include two
lone pairs, the geometry is bent and the bond
angle is 105o (i.e., close to 109.5o).
•An example is H2O:
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Sample problem 4.7
Question: Determine the shape of the following
molecules:
Solution: See Table 4.2 page 36
linear
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Electronegativity and Bond Polarity
•Electronegativity is a measure of an atom’s
attraction for e− in a bond.
•It tells how much a particular atom “wants” e−.
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Electronegativity and Bond Polarity
•If the electronegativities of two bonded atoms
are equal or similar, the bond is nonpolar.
•The electrons in the bond are being shared
equally between the two atoms.
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Electronegativity and Bond Polarity
•Bonding between atoms with different electronegativities yields a polar covalent bond or dipole.
•The electrons in the bond are unequally shared
between the C and the O.
•e− are pulled toward O, the more electronegative
element; this is indicated by the symbol δ−.
•e− are pulled away from C, the less electronegative
element; this is indicated by the symbol δ+.
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Electronegativity and Bond Polarity
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Polarity of Molecules
The classification of a molecule as polar or nonpolar
depends on:
•the polarity of the individual bonds
•the overall shape of the molecule
Nonpolar molecules generally have:
•no polar bonds
•individual bond dipoles that cancel
Polar molecules generally have:
•only one polar bond
•individual bond dipoles that do not cancel
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Polarity of Molecules
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Sample Problem 4.8
Question: Use electronegativity values to classify
each bond as nonpolar, polar covalent or ionic:
(a) Cl2 (b) HCl (c) NaCl
Solution
Pages 38 and 41
< 0.5 nonpolar; 0.5-1.9 polar covalent; >1.9 ionic
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Sample problem 4.9: Determine whether each molecule is
polar or nonpolar: (a) H2O (b) CO2
(a) H2O is a polar
molecule with net
dipole.
(b) CO2 is a nonpolar
molecule with no net
dipole.
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Sample problem 4.10: (a) Place lone pairs in the
structure of glycolic acid (b) Give the shape around
each atom in red (c) Label all polar bonds
Solution
All C-O and O-H bonds are polar (red bonds).
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