Name
Unit Title: Covalent Bonding and Nomenclature
Text Reference: Pages 189-193
Date
Intermolecular Forces of Attraction
Intramolecular vs. Intermolecular
So far in our discussion of covalent bonding, we have been discussing intramolecular
forces. Intramolecular forces are forces within a molecule that hold atoms together, that
is, covalent bonds. Today we are going to look at intermolecular forces. Intermolecular
forces are forces between molecules that hold molecules to each other. These
intermolecular forces are collectively referred to as van der Waals forces. Van der Waals
forces are much weaker than covalent bonds.
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.
The strength of the intermolecular forces of attraction can also be used to compare
boiling and melting points. The more strongly the molecules are attracted to each other,
the higher the boiling and melting points.
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.
• London dispersion forces arise from the motion of valence electrons.
o 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.
o 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.
o In general the larger the electron cloud, the more likely the molecule is to form
temporary dipoles.
• London forces are the weakest type of intermolecular forces of attraction.
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Dipole-Dipole Forces
•
Dipole-dipole forces of attraction exist between polar molecules.
o
Polar molecules contain uneven distributions of charge.
o
The negative dipole of one molecule is attracted to the positive dipole of another
molecule.
•
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+
H−Cl
↓
↑
↓
↑
Cl –H δ-
Dipole-dipole forces
↓
↑
Cl−H
δ- δ+
δ- δ+
•
Dipole-dipole forces of attraction are stronger than London dispersion 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.
•
Examples of hydrogen bonding.
o Hydrogen bonding occurs in pure substances. The hydrogen bonding is
represented by a dotted line.
H
O
H
O
H
o
H
H
H
H N
H N
H
H
water
ammonia
Hydrogen bonding can also occur in mixtures.
H
F
H
H N
H
H
•
H F
H F
hydrogen fluoride
O
H
H N
H
H
hydrogen fluoride and ammonia
water and ammonia
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.
You Try It
1.
List the intermolecular forces of attraction in order of increasing strength.
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
b.
CO2
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 is a polar
molecule. NH3 exhibits hydrogen bonding because it is 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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c.
HI
d.
BeH2
Comparing Boiling Points
Two factors that affect boiling point are the mass of the substance (molar mass) and the
strength of the intermolecular forces of attraction. The stronger the intermolecular
forces of attraction the higher the boiling point.
Examine the table below.
Boiling Points of Halogens
Name
Formula
fluorine
chlorine
bromine
iodine
F2
Cl2
Br2
I2
Physical State at
Room Temperature
gas
gas
liquid
solid
Molar Mass
(g/mol)
38.0
70.9
159.8
253.8
Boiling Point
(K, at 1 atm)
85.0
239.1
331.9
457.4
1.
What relationship exists between the mass of the halogens and the boiling point?
2.
Arrange the halogens in order of increasing intermolecular strength of attraction.
Justify your answer.
4
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.
1.
Which group of elements has the lowest boiling points for each period? Why do
they have the lowest boiling points for each period?
2.
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?
3.
With the exception of H2O, NH3, and HF, why do the boiling points generally
increase within a group?
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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
b.
a liquid with a high boiling point
c.
a gas at room temperature
d.
a liquid with a low-boiling point
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?
3.
Arrange the following according to increasing boiling point: H2O, H2S, CO2. Justify
your ranking.
4.
Arrange the following according to increasing boiling point: CH4, CI4, CF4. Justify
your ranking.
5.
HF, NH3, and H2O all exhibit hydrogen bonding. H2O, however, has stronger
intermolecular forces of attraction and therefore a higher boiling point. Suggest a
possible explanation as to why H2O has stronger hydrogen bonding.
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