Chapter 5
Compounds and Their Bonds
Covalent Bonds
Naming and Writing Formulas of
Covalent Compounds
Bond Polarity
1
Ionic Bond
H2, A Covalent Molecule


In hydrogen, two hydrogen atoms share their
electrons to form a covalent bond.
Each hydrogen atom acquires a stable outer
shell of two (2) electrons like helium (He).
H + H
H : H = HH = H2
hydrogen molecule
2
A covalent bond between two hydrogen atoms is
shown in this picture.
Fig 5.1 A covalent bond is the result of
attractive and repulsive forces between atoms.
3
Spherical 1S orbital of two individual hydrogen
atoms bends together and overlap to give an egg
shaped region in the hydrogen molecule. The
shared pair of electrons in a covalent bond is
often represented as a line between atoms.

4
Bond length: The optimum distance
between nuclei involved in a covalent bond.
If the atoms are too far apart, the attractive
forces are small and no bond exists. If the
atoms are too close, the repulsive interaction
between the nuclei is so strong that it pushes
the atoms apart, Fig 5.2.

5
When two chlorine atoms approach each
other, the unpaired 3p electrons are shared by
both atoms in a covalent bond. Each chlorine
atom in the Cl2 molecule now have 6 electrons
in its own valence shell and sharing two giving
each valence shell octet.

6
Diatomic Elements

As elements,
the following
share electrons
to form
diatomic,
covalent
molecules.
7
In addition to H2 and Cl2, five other elements
always exist as diatomic molecule.

8
Learning Check
What is the name of each of the following diatomic
molecules?
H2
hydrogen
N2
nitrogen
Cl2
_______________
O2
_______________
I2
_______________
9
Solution
What are the names of each of the following diatomic
molecules?
H2
hydrogen
N2
nitrogen
Cl2
chlorine
O2
oxygen
I2
iodine
10
Covalent Bonds in NH3

The compound NH3 consists of a N atom and three
H atoms.


N  and 3 H 


By sharing electrons to form NH3, the electron dot
structure is written as
H
Bonding pairs

H:N:H

Lone pair of electrons
11
Number of Covalent Bonds

Often, the number of covalent bonds formed
by a nonmetal is equal to the number of
electrons needed to complete the octet.
12
Dot Structures and Models of
Some Covalent Compounds
13
Multiple Bonds
 Sharing one pair of electrons is a single bond.
X:X
or
X–X
 In multiple bonds, two pairs of electrons are
shared to form a double bond or three pairs
of electrons are shared in a triple bond.
X:
:X
X:::X
or
X =X
or
X≡X
14
Multiple Bonds in N2



In nitrogen, octets are achieved by sharing
three pairs of electrons.
When three pairs of electrons are shared, the
multiple bond is called a triple bond.
octets

N

+ 

N




N:::N
triple bond
15
5.4 Coordinate Covalent Bonds
Coordinate Covalent Bond: The covalent bond
that forms when both electrons are donated by the
same atom.

16
Fig 5.7 Electronegativities and the periodic table

17
5.6 Drawing Lewis Structure
1. Draw skeletal structure of compound showing
what atoms are bonded to each other. Put the
unique element ( or least electronegative atom)
in the center.
2. Count total number of valence e-. Add 1 for
each negative charge. Subtract 1 for each
positive charge.
3. Complete an octet for all atoms except
hydrogen
4. If structure contains too many electrons, form
double and triple bonds on central atom as
needed.
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18
Write the Lewis structure of nitrogen trifluoride (NF3).
Step 1 – N is less electronegative than F, put N in center
Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5)
5 + (3 x 7) = 26 valence electrons
Step 3 – Draw single bonds between N and F atoms and complete
octets on N and F atoms.
Step 4 - Check, are # of e- in structure equal to number of valence e- ?
3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons
F
N
F
F
19
Write the Lewis structure of the carbonate ion (CO32-).
Step 1 – C is less electronegative than O, put C in center
Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4)
-2 charge – 2e4 + (3 x 6) + 2 = 24 valence electrons
Step 3 – Draw single bonds between C and O atoms and complete
octet on C and O atoms.
Step 4 - Check, are # of e- in structure equal to number of valence e- ?
3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons
Step 5 - Too many electrons, form double bond and re-check # of e-
O
C
O
2 single bonds (2x2) = 4
1 double bond = 4
8 lone pairs (8x2) = 16
Total = 24
O
20
Two possible skeletal structures of formaldehyde (CH2O)
H
C
O
H
H
C
H
O
An atom’s formal charge is the difference between the
number of valence electrons in an isolated atom and the
number of electrons assigned to that atom in a Lewis
structure.
formal charge
on an atom in
a Lewis
structure
=
total number
total number
of valence
of nonbonding
electrons in electrons
the free atom
-
1
2
(
total number
of bonding
electrons
)
The sum of the formal charges of the atoms in a molecule
or ion must equal the charge on the molecule or ion.
21
H
-1
+1
C
O
formal charge
on an atom in
a Lewis
structure
H
=
C : 4 eO : 6 e2H :2x1 e12 e-
2 single bonds (2x2) = 4
1 double bond = 4
2 lone pairs (2x2) = 4
Total = 12
total number
total number
of valence
of nonbonding
electrons in electrons
the free atom
-
1
2
(
total number
of bonding
electrons
)
formal charge
= 4 -2 -½ x 6 = -1
on C
formal charge
= 6 -2 -½ x 6 = +1
on O
22
H
H
0
C
formal charge
on an atom in
a Lewis
structure
0
O
=
C – 4 eO – 6 e2H – 2x1 e12 e-
2 single bonds (2x2) = 4
1 double bond = 4
2 lone pairs (2x2) = 4
Total = 12
total number
total number
of valence
of nonbonding
electrons in electrons
the free atom
-
1
2
(
total number
of bonding
electrons
)
formal charge
= 4 - 0 -½ x 8 = 0
on C
formal charge
= 6 -4 -½ x 4 = 0
on O
23
Formal Charge and Lewis Structures
1. For neutral molecules, a Lewis structure in which there
are no formal charges is preferable to one in which
formal charges are present.
2. Lewis structures with large formal charges are less
plausible than those with small formal charges.
3. Among Lewis structures having similar distributions of
formal charges, the most plausible structure is the one in
which negative formal charges are placed on the more
electronegative atoms.
Which is the most likely Lewis structure for CH2O?
H
-1
+1
C
O
H
H
H
0
C
0
O
24
5.7 Shape of Molecules
VSEPR


The shape of a molecule is predicted from the
geometry of the electron pairs around the central
atom.
In the valence-shell electron-pair repulsion
theory (VSEPR), the electron pairs are arranged
as far apart as possible to give the least amount of
repulsion of the negatively charged electrons.
25
Two Electron Pairs



In a molecule of BeCl2, there are two bonding
pairs around the central atom Be. (Be is an
exception to the octet rule.)
The arrangement of two electron pairs to
minimize their repulsion is 180° or opposite
each other.
The shape of the molecule is linear.
26
Two Electron Pairs with
Double Bonds



The electron-dot structure for CO2 consists of two
double bonds to the central atom C.
Because the electrons in a double bond are held
together, a double bond is counted as a single unit.
Repulsion is minimized when the double bonds are
placed opposite each other at 180° to give a linear
shape.
27
Three Electron Pairs



In BF3, there are 3 electron pairs around the central
atom B. (B is an exception to the octet rule.)
Repulsion is minimized by placing three electron
pairs in a plane at angles of 120°, which is a trigonal
planar arrangement.
The shape with three bonded atoms is trigonal
planar.
28
Two Bonding Pairs and A
Nonbonding Pair




In SO2, there are 3 electron units around the
central atom S.
Two electron units are bonded to atoms and
one electron pair is a nonbonding pair.
Repulsion is minimized by placing three
electron pairs in a plane at angles of 120°,
which is trigonal planar.
The shape with two bonded atoms is bent.
29
Four Electron Pairs



In CH4, there are 4 electron pairs around the
central atom C.
Repulsion is minimized by placing four
electron pairs at angles of 109°, which is a
tetrahedral arrangement.
The shape with four bonded atoms is called
tetrahedral.
30
Three Bonding Atoms and One
Nonbonding Pair




In NH3, there are 4 electron pairs around the N.
Three pairs are bonded to atoms and one is a
nonbonding pair.
Repulsion is minimized by placing four electron
pairs at angles of 109°, which is a tetrahedral
arrangement.
The shape with three bonded atoms is pyramidal.
31
Two Bonding Atoms and Two
Lone Pairs




In H2O, there are 4 electron pairs around O.
Two pairs are bonded to atoms and two are
nonbonding pairs.
Repulsion is minimized by placing four
electron pairs at angles of 109° called a
tetrahedral arrangement.
The shape with two bonded atoms is called
bent.
32
Some Steps Using VSEPR to
Predict Shape




Draw the electron dot structure.
Count the charged clouds around the central
atom.
Arrange the charged clouds to minimize
repulsion.
Determine the shape using the number of
bonded atoms in the electron arrangement.
33
Summary of Electron
Arrangements and Shapes
Number of atoms bonded
to the central atom
34
The shape depends on the number of charged clouds
surrounding the atom as summarized in Table 5.1

35
Learning Check
Use VSEPR theory to determine the shape
of the following molecules or ions.
1) tetrahedral
2) pyramidal 3) bent
A. PF3
B. H2S
C. CCl4
D. PO43-
36
Solution
Use VSEPR theory to determine the shape of
the following molecules or ions.
1) tetrahedral 2) pyramidal
3) bent
A. PF3
2) pyramidal
B. H2S
3) bent
C. CCl4
1) tetrahedral
D. PO431) tetrahedral
37
Comparing Nonpolar and Polar
Covalent Bonds
38
Electronegativity
 Electronegativity is the attraction of an atom
for shared electrons.
 The nonmetals have high electronegativity
values with fluorine as the highest.
 The metals have low electronegativity values.
39
Fig 5.7 Electronegativities and the periodic table

40
41
Nonpolar Covalent Bonds
 The atoms in a nonpolar covalent bond have
electronegativity differences of 0.4 or less.
 Examples:
Atoms Electronegativity Type of
Difference
Bond
N-N
3.0 - 3.0 = 0.0
Nonpolar covalent
Cl-Br
3.0 - 2.8 = 0.2
Nonpolar covalent
H-Si
2.1 - 1.8 = 0.3
Nonpolar covalent
42
Polar Covalent Bonds
 The atoms in a polar covalent bond have
electronegativity differences of 0.5 to 1.9.
 Examples:
Atoms Electronegativity Type of
Difference
Bond
O-Cl
3.5 - 3.0 = 0.5
Polar covalent
Cl-C
3.0 - 2.5 = 0.5
Polar covalent
O-S
3.5 - 2.5= 1.0
Polar covalent
43
Comparing Nonpolar and Polar
Covalent Bonds
44
Polar, Nonpolar and Ionic Bond
45
Ionic Bonds
 The atoms in an ionic bond have
electronegativity differences of 2.0 or more.
 Examples:
Atoms Electronegativity
Type of
Difference
Bond
Cl-K
3.0 – 0.8
= 2.2
Ionic
N-Na
3.0 – 0.9
= 2.1
Ionic
46
Predicting Bond Type
47
Learning Check
Identify the type of bond between the
following as
1) nonpolar covalent
2) polar covalent
3) ionic
A. K-N
B. N-O
C. Cl-Cl
48
Solution
A. K-N
3) ionic
B. N-O
2) polar covalent
C. Cl-Cl
1) nonpolar covalent
49
5.9 Polar Molecules


Entire molecule can be polar if electrons
are attracted more strongly to one part of
the molecule than to another.
Molecule’s polarity is due to the sum of
all individual bond polarities and lonepair contribution in the molecule.
50

Molecular polarity is represented by an arrow
pointing at the negative end and is crossed at
the positive end to resemble a positive sign.
51

Molecular polarity depends on the shape of
the molecule as well as the presence of polar
covalent bonds and lone-pairs.
52
Would a linear water molecule be
Polar?

H

..

O H
..
Why is water not linear?
53
54
Learning Check
Identify each of the following molecules as
1) polar or 2) nonpolar. Explain.
A. PBr3
B. HBr
C. Br2
D. SiBr4
55
Solution
Identify each of the following molecules as
1) polar or 2) nonpolar. Explain.
A. PBr3 1) polar; pyramidal
B. HBr 1) polar; polar bond
C. Br2
2) nonpolar, nonpolar bond
D. SiBr4 2) nonpolar; dipoles cancel
56
Naming Covalent Compounds
In the name of a
covalent compound,
the first nonmetal is
named followed by
the name of the
second nonmetal
ending in –ide.
 Prefixes indicate the
number of atoms of
each element.

57
Learning Check
Complete the name of each covalent compound:
CO
carbon ______oxide
CO2
carbon _______________
PCl3
phosphorus ___________
CCl4
carbon _______________
N2O
______________________
58
Solution
Complete the name of each covalent compound:
CO
carbon monoxide
CO2
carbon dioxide
PCl3
phosphorus trichloride
CCl4
carbon tetrachloride
N2O
dinitrogen monoxide
59
Formulas and Names of Some
Covalent Compounds
60
Learning Check
Select the correct name for each compound.
A. SiCl4 1) silicon chloride
2) tetrasilicon chloride
3) silicon tetrachloride
B. P2O5 1) phosphorus oxide
2) phosphorus pentoxide
3) diphosphorus pentoxide
C. Cl2O7 1) dichlorine heptoxide
2) dichlorine oxide
3) chlorine heptoxide
61
Solution
Select the correct name for each compound.
A. SiCl4 3) silicon tetrachloride
B. P2O5
3) diphosphorus pentoxide
C. Cl2O7 1) dichlorine heptoxide
62
Chapter Summary





Covalent bond: Bond formed by sharing of electrons
between the atoms.
Molecule: A group of atoms held together by
covalent bonds.
Coordinate covalent bond: Bond formed when a
filled orbital containing lone pair of electrons on one
atom overlaps a vacant orbital on another atom.
Molecular formula: Formula that shows the numbers
and kinds of atoms in a molecule.
Lewis structure: shows how atoms are connected in a
molecule.
63
Chapter Summary Contd.




Molecules have specific shapes that depend on the
number of electron charge clouds surrounding the
various atoms (VSEPR model)
Bonds between atoms are polar covalent if the
bonding electrons are not shared equally between
the atoms.
The ability of an atom to attract electrons in a
covalent bond is the atom’s electronegativity.
Molecular compounds have lower melting points
and boiling points than ionic compounds.
64