Chemical Bonding (Predicting Bond Types)
Lewis (Electron) Dot Diagrams
Binary Molecular Nomenclature
Exceptions to the Octet Rule
Coordinate Covalent Bonding
Resonance Structures
Molecular Shapes and Polarity
Intermolecular Forces of Attraction
What is a chemical bond?
A chemical bond is a strong
attractive force between
atoms or ions in a chemical
compound.
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Why do elements form
chemical bonds?
1.
2.
3.
Uncombined elements have relatively
high potential energy.
Atoms will gain, lose or share valence
electrons in order to chemically combine
with other atoms.
By combining with other atoms, atoms
decrease potential energy and create
more stable arrangements.
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What two factors determine
whether or not a chemical bond
will form?
1.
the electron configurations of
the atoms involved
2.
the attraction the atoms have
for electrons
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How is the type of chemical
bond formed between two
atoms determined?
The type of chemical bond formed
depends upon the degree to which
the valence electrons are shared
between the atoms.
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Covalent Bonding
In a covalent bond, valence
electrons are shared by the atoms.
Covalent bonds can be nonpolar or
polar.
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Nonpolar vs. Polar
In a nonpolar covalent bond, electrons
are shared equally. Bonding which
occurs between two atoms of the same
element is an example of nonpolar
covalent bonding.
Examples: H2, Br2, O2, N2, Cl2, I2, F2
In a polar covalent bond, electrons are
shared unequally.
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Ionic Bonding
In an ionic bond, valence electrons
are transferred between atoms.
One atom gains electrons to form a
negative ion (anion) and the other
atom loses electrons to form a
positive ion (cation).
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Ionic Bonding
Which category of elements tends to
gain electrons and form negative ions
(anions)?
nonmetals
Which category of elements tends to
lose electrons and form positive ions
(cations)?
metals
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Using differences in electronegativity to
determine bond type
Electronegativity is a measure of an atom’s ability
to attract electrons when chemically combining
with another element.
The higher the electronegativity value, the stronger
the attraction the atom has for another atom’s
electrons.
The degree to which bonding between atoms of
two elements is ionic or covalent can be estimated
by calculating the difference in the elements’
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electronegativities (ΔEN).
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Using differences in electronegativity to
determine bond type
Type of
Bond
Nonpolar
Covalent
Polar
Covalent
Ionic
∆EN (Difference in
Electronegativity)
≤0.2
0.2 to 1.7
≥1.7
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1
H
2.1
3
Li
1.0
11
Na
0.9
19
K
0.8
37
Rb
0.8
55
Cs
0.7
87
Fr
0.7
Periodic Table of Electronegativities
4
Be
1.5
12
Mg
1.2
20
Ca
1.0
38
Sr
1.0
56
Ba
0.9
88
Ra
0.9
21
Sc
1.3
39
Y
1.2
57
La
1.1
89
Ac
1.1
22 23
Ti
V
1.5 1.6
40 41
Zr Nb
1.4 1.6
72 73
Hf Ta
1.3 1.4
104 105
Rf Db
-
24
Cr
1.6
42
Mo
1.8
74
W
1.7
106
Sg
-
25
Mn
1.5
43
Tc
1.9
75
Re
1.9
107
Bh
-
26
Fe
1.8
44
Ru
2.2
76
Os
2.2
108
Hs
-
27
Co
1.9
45
Rh
2.2
77
Ir
2.2
109
Mt
-
28
Ni
1.8
46
Pd
2.2
78
Pt
2.2
29
Cu
1.9
47
Ag
1.9
79
Au
2.4
30
Zn
1.6
48
Cd
1.7
80
Hg
1.9
110
Uun
-
111
Uuu
-
112
Uub
-
5
B
2.0
13
Al
1.5
31
Ga
1.6
49
In
1.7
81
Tl
1.8
6
C
2.5
14
Si
1.8
32
Ge
1.8
50
Sn
1.8
82
Pb
1.8
7
N
3.0
15
P
2.1
33
As
2.0
51
Sb
1.9
83
Bi
1.9
8
O
3.5
16
S
2.5
34
Se
2.4
52
Te
2.1
84
Po
2.0
9
F
4.0
17
Cl
3.0
35
Br
2.8
53
I
2.5
85
At
2.2
113
114
Uuq
-
115
116
Uuh
-
117
2
He
10
Ne
18
Ar
36
Kr
3.0
54
Xe
2.6
86
Rn
2.4
118
Uuo
-
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Example 1: What type of bond would form
between an atom of nitrogen and an atom
of chlorine?
a.
b.
c.
d.
Nitrogen has an electronegativity value of
3.0.
Chlorine has an electronegativity value of
3.0.
The difference in the electronegativity values
for nitrogen and chlorine is
ΔEN = 3.0
- 3.0
= 0.0
Therefore the type of bond formed would be
nonpolar covalent. The electrons would be
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shared equally.
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Example 2: What type of bond would form
between an atom of hydrogen and an atom
of chlorine?
a.
b.
c.
d.
Hydrogen has an electronegativity value of
2.1.
Chlorine has an electronegativity value of
3.0.
The difference in the electronegativity values
for hydrogen and chlorine is
ΔEN = 3.0
- 2.1
= 0.9
Therefore the type of bond formed would be
polar covalent. The electrons would be
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shared unequally.
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Dipole
A bond formed between atoms which are not
shared equally is called a dipole.
a) In the bond formed between hydrogen and
chlorine, the chlorine would form the negative
dipole (symbolized by δ-) because it has the
higher electronegativity value.
b) The hydrogen would form the positive dipole
(symbolized by δ+) because it has the lower
electronegativity value.
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Example 3: What type of bond would form
between an atom of lithium and an atom of
chlorine?
a.
b.
c.
d.
Lithium has an electronegativity value of
1.0.
Chlorine has an electronegativity value of
3.0.
The difference in the electronegativity values
for lithium and chlorine is
ΔEN = 3.0
- 1.0
= 2.0
Therefore the type of bond formed would be
ionic. The electrons would be transferred
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between atoms
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Example 3: What type of bond would form
between an atom of lithium and an atom of
chlorine?
The lithium atom would lose electrons
and form a positive ion, also known as
a cation.
The chlorine atom would gain electrons
and form a negative ion, also known as
an anion.
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You Try It
1. Complete the following table.
Compound
KF
O2
ICl
Elements
Electronegativity
K
0.8
F
4.0
O
3.5
O
3.5
I
2.5
Cl
3.0
∆EN
Bond Type
3.2
Ionic
Nonpolar
0.0
Covalent
Polar
0.5
Covalent
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You Try It
2. For each of the bonds in question 1
that were polar covalent, identify the
negative dipole (δ-) and the positive
dipole (δ+).
ICl Iodine is the positive dipole and
chlorine is the negative dipole.
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You Try It
3. Elements that exist as two atoms
chemically bonded together are called
diatomic elements. The diatomic elements
are hydrogen, bromine, oxygen, nitrogen,
chlorine, iodine, and fluorine. (You need to
memorize the diatomic elements.) What
type of chemical bond exists between the
diatomic elements?
Nonpolar covalent
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You Try It
4. Using the three classifications of
bonds discussed, predict the type of
bond that is most likely to be present
in compounds made from elements of
groups 1 (1A) and 17 (7A).
Ionic
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You Try It
5. Using the three classifications of
bonds discussed, predict the type of
bond that is most likely to be present
in compounds made from elements of
groups 16 (6A) and 17 (7A).
Polar Covalent
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You Try It
6. Arrange the following chemical bonds
in order of least covalent to most
covalent: H-H, H-Cl, H-Br, Li-Cl
Li-Cl, H-Cl, H-Br, H-H
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Drawing Lewis Dot Diagrams for Atoms
The electrons that play the most
important role in determining whether or
not a chemical bond will form are the
valence electrons.
In a Lewis dot diagram, dots are placed
around the chemical symbol of an
element to illustrate the valence
electrons. The chemical symbol
represents the nucleus of the atom. Back to
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Drawing Lewis Dot Diagrams for Atoms
Examples
Group 1
Group 2
Group 13
Group 14
Group 15
Group 16
Group 17
H
Li
Group 18
He
Be
B
C
N
O
F
Ne
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Drawing Lewis Structures for Covalent Compounds
Types of Covalent Bonds
Single Covalent Bond – one pair
of valence electrons is shared.
Double Covalent Bond - two pairs
of valence electrons are shared.
Triple Covalent Bond - three pairs
of valence electrons are shared.
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Example 1: H2
H
H
H H
The two hydrogen atoms will
form a single, nonpolar covalent
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bond.
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Example 2: O2
O
O O
O
The two oxygen atoms will form a
double, nonpolar covalent bond. Back to
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Example 3: N2
N
N N
N
The two nitrogen atoms will form
a triple, nonpolar covalent bond. Back to
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Example 4: HCl
H
Cl
H Cl
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Example 5: NH3
H N H
H
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You Try It
HBr
CCl4
H 2O
C2H5Cl
C 2H 4
C 2H 2
H 2 O2
HCN
CO2
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Structural Formulas
Structural formulas can also be
used to show the arrangement of
atoms in molecules.
In a structural formula, dashes are
used to represent shared pairs of
electrons.
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Structural Formulas
Example: H2
HH
H–H
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Structural Formulas
Example: H2S
HS
H
H –S
H
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Structural Formulas
Example: N2
N N
N≡ N
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Binary Molecular Nomenclature
Compounds formed when atoms covalently
bond are called molecular compounds.
Binary molecular compounds are generally
composed of two nonmetallic elements.
When two nonmetallic elements combine,
they often do so in more than one way. For
example carbon can combine with oxygen to
form carbon dioxide, CO2 and carbon
monoxide, CO.
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Naming Binary Molecular Compounds
Prefixes are used to show how many atoms
of each element are present in each
molecule of the compound.
monoditritetrapenta-
1
2
3
4
5
hexaheptaoctanonadeca-
6
7
8
9
10
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Naming Binary Molecular Compounds
The names of molecular
compounds have this
form: (prefix + element
name) (prefix + element
root + ide)
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Naming Binary Molecular Compounds
The prefix mono is usually omitted if there is
just a single atom of the first element.
Example: CO2 is carbon dioxide not
monocarbon dioxide.
If the vowel combinations o-o or a-o appear
next to each other in the name, the first of
the pair is omitted to simplify the name.
Example: N2O is dinitrogen monoxide not
dinitrogen monooxide.
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You Try It
Name the following compounds.
a.
b.
c.
d.
e.
f.
CBr4
Cl2O7
N2O5
BCl3
PCl5
NO
a.
b.
c.
d.
Carbon tetrabromide
Dichlorine heptoxide
Dinitrogen pentoxide
Boron trichloride
e. Phosphorus pentachloride
f. nitrogen monoxide
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Writing Formulas for Binary Molecular Compounds
To write the formula for a binary molecular
compound you simply write down the
number of atoms of each element indicated
by the name.
Example: Carbon tetrachloride
CCl4
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You Try It
Write formulas for the following binary
molecular compounds.
a. dinitrogen tetrahydride
N2H4
b. carbon disulfide
c. iodine heptafluoride
CS2
IF7
d. sulfur dioxide
SO2
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Writing Formulas for Binary Molecular Compounds
A few molecular compounds have common
names that all scientists use in place of
formal names.
CH4 is methane
H2O is water
NH3 is ammonia
You need to memorize these.
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Exceptions to the Octet Rule
Some molecules are stable even
though the atoms do not all obtain
an octet.
There are three common
exceptions to the octet rule.
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Exception #1
In some molecules the central atom
has less than eight valence electrons.
This is called an incomplete octet.
Incomplete octets are common in
covalent compounds in which the
central atom is beryllium, boron or
aluminum.
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Exception #1
Example: BeH2
Beryllium has only four electrons
around it.
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Exception #2
Molecules almost always have an
even number of electrons, allowing
electrons to be paired, but there
are some exception in which there
are an odd number of electrons.
These exceptions usually involve
nitrogen.
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Exception #2
Example: NO
You will not be expected to draw
exceptions with odd numbers of
electrons in this course.
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Exception #3
In some molecules the central atom
has more than eight valence
electrons.
This is called an expanded octet.
Some common central elements that
have expanded octets are sulfur,
chlorine, bromine, iodine, xenon,
phosphorus, and arsenic.
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Exception #3
Example: SF6
Sulfur has twelve electrons around it.
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You Try It
BF3
AsH5
BeI2
PCl5
ClF5
XeF4
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Coordinate Covalent Bonding
Objectives
1. Define coordinate covalent
bonding and give an example.
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Coordinate Covalent Bonding
A coordinate covalent bond is
formed when one atom contributes
both electrons in a shared pair.
Example: CO
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Coordinate Covalent Bonding
and Polyatomic Ions
Polyatomic ions form coordinate
covalent bonds.
A polyatomic ion is covalently bonded
within itself, but is ionically bonded to
another atom or polyatomic ion to
form a neutral compound.
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Coordinate Covalent Bonding
and Polyatomic Ions
Example: NH4+
Example: SO42-
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You Try It
OH-
PO43-
ClO3-
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Resonance Structures
Objectives
1. Define resonance and draw
resonance structures for
molecules.
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Resonance Structures
Resonance occurs when more than
one valid Lewis structure can be
written for a particular molecule.
The different Lewis structures
possible for a molecule are referred
to as resonance structures.
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Resonance Structures
Let’s look at the Lewis structure for
the ozone, O3, molecule.
.. .. ..
O=O-O:
..
..
Another possible structure for the
ozone molecule is as follows:
.. .. ..
:O-O=O
..
..
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Resonance Structures
Notice that each structure indicates
that the ozone molecule has two
types of O-O bonds, one single
bond and one double bond.
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Resonance Structures
Based on what we just learned
about bond length, you would
expect the bond lengths between
the atoms to be different.
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Resonance Structures
Scientists, however, have
experimentally determined that the
bond lengths between the oxygen
atoms are identical.
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Resonance Structures
No one structure correctly describes
the ozone molecule. Scientists have
determined that the structure for
ozone is the average of the two
structures.
A double-headed arrow is used to
indicate resonance.
.. .. ..
O=O-O:
..
..
.. .. ..
:O-O=O
..
..
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You Try It
Resonance structures can often be written
for polyatomic ions. Draw the possible
resonance structures for NO2- .
.. .. ..
O=N-O:
..
..
.. .. ..
:O-N=O
..
..
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Molecular Shapes and Polarity
Objectives
1. Define VSEPR and given a chemical
formula of a simple molecule, identify
its geometric shape as linear, trigonal
planar, angular, tetrahedral, trigonal
pyramidal, trigonal bypyramidal, or
octahedral.
2. Using the shape of a molecule and
electronegativites of its atoms,
determine the polarity of the molecule.
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Molecular Shapes and Polarity
The valence shell electron pair repulsion
(VSEPR) theory can be used to predict the
three dimensional shapes of a molecule.
The main idea behind VSEPR theory is that
electron pairs (bonding and nonbonding) will
orient themselves so that repulsions
between electron pairs are minimized.
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Linear
Formula
HI
Lewis Drawing
Structure of Model
H
I
Bond
Angle
180°
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Linear
Formula
HCN
Lewis Drawing
Structure of Model
H
C
N
Bond
Angle
180°
HCN is a 3-atom linear molecule. Which atom is the central
atom in the HCN molecule? Carbon
How many pairs of nonbonding electrons on the central
Back to
atom of the HCN molecule? None
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Bent (also called angular)
Formula
H2O
Lewis Drawing
Structure of Model
Bond
Angle
104.5°
Which atom of the water molecule is the central atom?
Oxygen
How many pairs of nonbonding electrons on the central
atom of the H2O molecule?
Two
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How can you differentiate between a linear
molecule and a bent molecule in terms of
nonbonding electron pairs on the central
atom?
Linear molecules do not have nonbonding
electrons on the central atom. Bent
molecules have nonbonding electrons on
the central atom.
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Trigonal Planar
Formula
Lewis Drawing
Structure of Model
H2CO
Bond
Angle
120°
Which atom is the central atom? Carbon
How many atoms are bonded to the central atom? Three
How many nonbonded electrons are there on the central
atom?
Zero
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Trigonal Pyramidal
Formula
Lewis Drawing
Structure of Model
NI3
Bond
Angle
107°
Which atom is the central atom? Nitrogen
How many atoms are bonded to the central atom? Three
How many nonbonded electrons are there on the central
atom?
one
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How can you differentiate between a trigonal
planar molecule and a trigonal pyramidal
molecule in terms of nonbonding electron
pairs on the central atom?
Trigonal planar molecules do not have
nonbonding electrons on the central atom.
Trigonal pyramidal molecules have a pair of
nonbonding electrons on the central atom.
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Tetrahedral
Formula
Lewis Drawing
Structure of Model
CH4
Bond
Angle
109.5°
Which atom is the central atom? Carbon
How many atoms are bonded to the central atom? Four
How many nonbonded electrons are there on the central
atom?
zero
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Trigonal Bipyramidal
Formula
Lewis Drawing
Structure of Model
Bond
Angle
PH5
120°
90°
Which atom is the central atom? Phosphorus
How many atoms are bonded to the central atom? Five
How many nonbonded electrons are there on the central
atom?
zero
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Octahedral
Formula
Lewis Drawing
Structure of Model
SH6
Bond
Angle
120°
90°
Which atom is the central atom? Sulfur
How many atoms are bonded to the central atom? Six
How many nonbonded electrons are there on the central
atom?
zero
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Summary of Molecular Shapes
Name
Drawing of Shape
Number of Atoms
Bonded to
Central Atom
Number of Lone
Bond Angle
Pairs of Electrons
2-Atom Linear
not applicable
not applicable
180°
3-Atom Linear
2
2
0
2
180°
104.5°
Trigonal Planar
3
0
120°
Trigonal
Pyramidal
3
1
107°
Tetrahedral
4
0
109.5°
Bent
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Summary of Molecular Shapes
Name
Drawing of Shape
Number of Atoms
Bonded to
Central Atom
Number of Lone
Bond Angle
Pairs of Electrons
Trigonal
Bypryamidal
5
0
180°
Octahedral
6
0
180°
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Determining Molecular Polarity
The polarity of each bond, along with the
geometry of the molecule, determines the
polarity of the molecule.
A nonpolar molecule has an even
distribution of molecular charge.
A polar molecule has an uneven distribution
of molecular charge.
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Steps in Determining Molecular Polarity
First look at the geometric shape of the
molecule.
Molecules with nonbonding pairs of
electrons on the central atom are polar.
Which two shapes are always polar?
bent and trigonal pyramidal
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Steps in Determining Molecular Polarity
If the molecule does contain nonbonding pairs
of electrons on the central atom, the polarity is
determined by the atoms surrounding the
central atom.
If all of the atoms surrounding the central atom
are the same, the molecule is nonpolar. This
is because the bond dipoles will cancel out.
If all of the atoms surrounding the central atom
are not alike, the molecule is polar. The bond
dipoles will not cancel out.
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Steps in Determining Molecular Polarity
H-C≡N:
This molecule is polar.
H-Be-H
This molecule is nonpolar.
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You Try It
Determine the polarity of each of the
following molecules.
a.
b.
c.
d.
e.
HI
H2O
H2CO
NI3
CH4
polar
polar
polar
polar
nonpolar
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Intermolecular Forces of Attraction
Objectives
13. Define van der Waals forces, dipoledipole forces, hydrogen bonds, and
London forces.
14. Given a molecule, identify the
dominant type of intermolecular force
of attraction.
15. Given chemical formulas for two
substances, identify which type of
intermolecular forces they exhibit
and compare their boiling and
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freezing points.
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Intramolecular vs. Intermolecular
Intramolecular forces – forces within a molecule that
hold atoms together, that is, covalent bonds.
Intermolecular forces – forces between molecules that
hold molecules to each other.
These intermolecular forces are collectively referred to
as Van der Waals Forces.
They are much weaker than covalent bonds.
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Importance of Intermolecular Forces
The strength of the intermolecular forces can be used to
determine whether a covalent compound exists as a solid,
liquid, or gas under standard conditions.
Solids have the strongest intermolecular forces of attraction
between their particles.
The intermolecular forces of attraction between the
molecules of liquids are not as strong as those found
between the particles of a solid.
Gases have the weakest intermolecular forces of attraction
between their particles.
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Importance of Intermolecular Forces
The strength of the intermolecular forces
can also be used to compare melting and
boiling points.
The more strongly the molecules are
attracted to each other, the higher the boiling
and melting points.
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Types of Intermolecular Forces
London Dispersion Forces
London dispersion forces exist in all
covalent molecules, however; they are the
most noticeable between nonpolar
molecules and the nonbonding atoms of
noble gases.
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Types of Intermolecular Forces
London Dispersion Forces
London dispersion forces arise from the
motion of valence electrons.
From the probability distributions of orbitals,
it is concluded that the electrons are evenly
distributed around the nucleus. However, at
any one instant, the electron cloud may
become distorted as the electrons shift to an
unequal distribution.
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Types of Intermolecular Forces
London Dispersion Forces
It is during this instant that a molecule develops
a temporary dipole.
This temporary dipole introduces a similar
response in neighboring molecules, thus
producing a short-lived attraction between
molecules.
In general the larger the electron cloud, the
more likely the molecule is to form temporary
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dipoles.
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Types of Intermolecular Forces
London Dispersion Forces
London forces are the weakest type of
intermolecular forces of attraction.
Examples: CO2, H2, Ar
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Types of Intermolecular Forces
Dipole-Dipole Forces
Dipole-dipole forces of attraction exist
between polar molecules.
Polar molecules contain uneven distributions
of charge.
The negative dipole of one molecule is
attracted to the positive dipole of another
molecule.
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Example of Dipole-Dipole Forces HCl
HCl is a polar molecule. The hydrogen end of
the molecule forms the positive dipole because
it has the lower electronegativity. The chloride
end of the molecule forms the negative dipole
because it has the higher electronegativity. The
chloride end of the molecule is attracted to the
hydrogen end of a neighboring molecule.
δ+
δ-
H−Cl
δ+
δ-
H−Cl
↓
↓
↓
↓
↑
↑
↑
↑
Cl−H
δ-
δ+
Dipole-dipole forces
Cl−H
δ-
δ+
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Types of Intermolecular Forces
Dipole-Dipole Forces
Dipole-dipole forces of attraction are
stronger than London dispersion forces.
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Types of Intermolecular Forces
Hydrogen Bonding
Hydrogen Bonding is a special type of dipole-dipole
force. Since no electrons are shared or transferred,
hydrogen bonding is not a chemical bond.
Hydrogen bonding exists between where the very
electronegative elements of nitrogen, oxygen and
fluorine are covalently bonded to hydrogen.
Hydrogen bonding occurs between hydrogen and
the unbonded electron pairs of nearby N, O, or F
molecules
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Examples of Hydrogen Bonding
Hydrogen bonding occurs in pure
substances. The hydrogen bonding is
represented by a dotted line.
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Examples of Hydrogen Bonding
Hydrogen bonding can also occur in
mixtures.
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Examples of Hydrogen Bonding
Hydrogen bonding occurs in pure
substances. The hydrogen bonding is
represented by a dotted line.
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Types of Intermolecular Forces
Hydrogen Bonding
Hydrogen bonding is about ten times
stronger than ordinary dipole-dipole forces.
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Identifying the Types of Intermolecular Forces of Attractions
The chart below can help you identify the types of
intermolecular forces of attraction exhibited by a
substance. Reminder: London Dispersion Forces
are exhibited by all covalent molecules.
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You Try It
1. List the intermolecular forces of attraction
in order of increasing strength.
London dispersion forces, dipole-dipole
forces, hydrogen bonding
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You Try It
2. What type of intermolecular forces of attraction would be exhibited by each
of the following substances? Justify your answer. The first one has been
done for you. (Hint: Draw the Lewis Structure for the molecule in order to
help you determine the polarity of the molecule.)
a. NH3
London dispersion forces, dipole-dipole, hydrogen
bonding. NH3 exhibits London dispersion forces
because all covalent molecules exhibit London
dispersion forces. NH3 exhibits dipole-dipole forces
because it’s a polar molecule. NH3 exhibits
hydrogen bonding because it’s a polar molecule in
which hydrogen is bonded to a nitrogen, oxygen, or
fluorine atom. In this case, hydrogen is bonded to
nitrogen.
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You Try It
2. What type of intermolecular forces of attraction would be exhibited by each
of the following substances? Justify your answer. The first one has been
done for you. (Hint: Draw the Lewis Structure for the molecule in order to
help you determine the polarity of the molecule.)
b. CO2
London dispersion forces only
CO2 is a nonpolar molecule. Nonpolar molecules
only exhibit London dispersion forces.
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You Try It
3. What type of intermolecular forces of attraction would be exhibited by each
of the following substances? Justify your answer. The first one has been
done for you. (Hint: Draw the Lewis Structure for the molecule in order to
help you determine the polarity of the molecule.)
c. HI
London dispersion forces and dipole-dipole forces
HI exhibits London dispersion forces because all
covalent molecules exhibit London dispersion forces.
HI also exhibits dipole-dipole forces because it’s a
polar molecule. It does not exhibit hydrogen bonding
because since H is not bonded to O, N or F.
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You Try It
2. What type of intermolecular forces of attraction would be exhibited by each
of the following substances? Justify your answer. The first one has been
done for you. (Hint: Draw the Lewis Structure for the molecule in order to
help you determine the polarity of the molecule.)
d. NH3
London dispersion forces only
BeH2 is a nonpolar molecule. Nonpolar molecules
only exhibit London dispersion forces.
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Comparing Boiling Points
Two factors that affect boiling point are the
mass of the compound (molar mass) and
the strength of the intermolecular forces of
attraction. The stronger the intermolecular
forces of attraction the higher the boiling
point.
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Comparing Boiling Points
Examine the table below.
Boiling Points of Halogens
Physical State at
Molar Mass Boiling Point
Name Formula
Room
(g/mol)
(K, at 1 atm)
Temperature
gas
fluorine
F2
38.0
85.0
gas
chlorine
Cl2
70.9
239.1
liquid
bromine
Br2
159.8
331.9
solid
iodine
I2
253.8
457.4
1. What relationship exists between the mass of
the halogens and the boiling point?
The larger the molar mass, the higher the
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boiling point.
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Comparing Boiling Points
Examine the table below.
Boiling Points of Halogens
Physical State at
Molar Mass Boiling Point
Name Formula
Room
(g/mol)
(K, at 1 atm)
Temperature
gas
fluorine
F2
38.0
85.0
gas
chlorine
Cl2
70.9
239.1
liquid
bromine
Br2
159.8
331.9
solid
iodine
I2
253.8
457.4
2. Arrange the halogens in order of increasing
intermolecular strength of attraction. Justify your answer.
F2, Cl2, Br2, I2
The stronger the intermolecular forces of
attraction, the greater the boiling points.
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Comparing Boiling Points
3. The graph below is a plot of the boiling points of the hydrogen
compounds in the groups headed by fluorine (HF, HCl, HBr, and
HI), oxygen (H2O, H2S, H2Se, H2Te), nitrogen (NH3, PH3, AsH3,
SbH3), and carbon (CH4, SiH4, GeH4, SnH4). Use the graph below
to answer the following questions.
a. Which group of elements has the
lowest boiling points for each
period? Why do they have the
lowest boiling points for each
period?
The group headed by carbon has
the lowest boiling points for each
period. They are all nonpolar
molecules. Nonpolar molecules
exhibit weaker London dispersion
forces.
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Comparing Boiling Points
3. The graph below is a plot of the boiling points of the hydrogen
compounds in the groups headed by fluorine (HF, HCl, HBr, and
HI), oxygen (H2O, H2S, H2Se, H2Te), nitrogen (NH3, PH3, AsH3,
SbH3), and carbon (CH4, SiH4, GeH4, SnH4). Use the graph below
to answer the following questions.
b. Notice in each of the other three groups
that the first compound (H2O, NH3, and
HF) in each group has a significantly
higher boiling point than the other
elements in their groups. What
accounts for this phenomenon?
H2O, NH3, and HF all exhibit hydrogen
bonding. The other substances in the
groups exhibit dipole-dipole forces of
attraction which are not as strong as
hydrogen bonding. Since H2O, NH3, and
HF all exhibit hydrogen bonding they have
higher than expected boiling points.
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Comparing Boiling Points
3. The graph below is a plot of the boiling points of the hydrogen
compounds in the groups headed by fluorine (HF, HCl, HBr, and
HI), oxygen (H2O, H2S, H2Se, H2Te), nitrogen (NH3, PH3, AsH3,
SbH3), and carbon (CH4, SiH4, GeH4, SnH4). Use the graph below
to answer the following questions.
c. With the exception of H2O,
NH3, and HF, why do the
boiling points generally
increase within a group?
The boiling points increase
because the molar mass of the
compounds increases.
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You Try It
1.Determine whether each of the following
would more likely be formed by polar or
nonpolar molecules.
a. a solid at room temperature polar
b. a liquid with a high boiling point polar
c. a gas at room temperature nonpolar
d. a liquid with a low-boiling point nonpolar
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You Try It
2. Considering what you have learned about
forces between atoms and molecules,
why do you think all of the elements in
group 18 exist as gases at room
temperature?
The noble gases exhibit London dispersion
forces. London dispersion forces are the
weakest of the intermolecular forces of
attraction. Substances with weak
intermolecular forces of attraction tend to have
lower boiling points.
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You Try It
3. Arrange the following according to
increasing boiling point: H2O, H2S, CO2.
Justify your ranking.
CO2 < H2S < H2O
CO2 has only London dispersion forces.
H2S has dipole-dipole forces.
H2O has hydrogen bonding.
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You Try It
4. Arrange the following according to
increasing boiling point: CH4, CI4, CF4.
Justify your ranking.
CH4 < CF4 < CI4
All three molecules are nonpolar and thus only
have London dispersion forces between them.
The bigger the molecule, the more electrons
and thus the larger the temporary dipole. The
larger the temporary dipole, the stronger the
intermolecular force and thus the higher the
melting point.
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You Try It
5. NH3 is a gas at room temperature and H2O is
a liquid at room temperature. However, they
both exhibit hydrogen bonding. What does
that tell you about the strength of the
hydrogen bonding in H2O as compared to
NH3?
The hydrogen bonding in H2O is stronger than the hydrogen
bonding which occurs in NH3. H2O has two H atoms that can
potentially form four hydrogen bonds with surround water
molecules. There are exactly the right number of hydrogens and
lone so pairs that every one of them can be involved in hydrogen
bonding. In the case of ammonia, the amount of hydrogen
bonding is limited by the fact that each nitrogen has only one lone
pair.
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