CHAPTER 11:
INTERMOLECULAR FORCES
AP Chemistry
States of Matter
• The fundamental difference between states of matter is
the distance between the particles
States of Matter
• The state of a substance
is in at a particular
temperature and pressure
depends on two
antagonistic entities:
The kinetic energy of the
particles
The strength of the
attractions between the
particles
Intermolecular Forces
• The attractions between molecules are not nearly as
strong as the intramolecular attractions that hold
compounds together
• They are, however, strong enough to control properties
such as boiling and melting points, vapor pressures, and
viscosities
• These intramolecular forces are:
• London Dispersion Forces
• Dipole-Dipole Interactions
• Hydrogen Bonding
• ** Ion-Dipole Interactions
• The first 2 above are referred to as Van der Waals forces
London Dispersion Forces
• While the electrons in the 1s orbital of helium repel each
other (and, therefore, tend to stay far away from each
other), it does happen that they occasionally wind up on
the same side of the atom.
• At that instant, then, the helium atom is polar, with an
excess of electrons on the left side and a shortage on the
right side
London Dispersion Forces
• Another helium nearby, then, would have a dipole induced
in it, as the electrons on the left side of helium atom 2
repel the electrons in the cloud on helium atom 1.
• London dispersion forces, or dispersion forces, are
attractions between an instantaneous dipole and an
induced dipole
• LDF present in all molecules, whether polar or nonpolar
• The tendency of an electron cloud to distort in this way is
called polarizability
Factors Affecting London Forces
(polarizability)
• The shape of the molecule affects
the strength: long, skinny molecules
(like n-pentane) tend to have
stronger dispersion forces than
short fat ones (like neopentane)
• This is due to the increased surface
area in n-pentane
• The strength of dispersion forces tends
to increase with molecular weight.
• Larger atoms have larger electron
clouds which are easier to polarize.
Dipole-Dipole Interactions
• Molecules that have permanent dipoles (polar molecules)
are attracted to each other.
• The positive end of one is
attracted to the negative
end of the other and
vice-versa
• These forces are only
important when the
molecules are close to
each other
Dipole-Dipole Interactions
• The more polar the molecule, the greater the dipole-dipole
intereractions, so its boiling point is higher
Which Have a Greater Effect: DipoleDipole or LDF?
• If two molecules are of comparable size and shape,
dipole-dipole interactions will likely be the dominating
force.
• If one molecule is much larger than another, dispersion
forces will likely determine its physical properties
Boiling Points as a Function of Molecular
Weight
• The BP of the group 4A
(bottom) and 6A (top)
hydrides are shown as a
function of molecular
weight.
• In general, the BP increase
with increasing molecular
weight because of the
increasing strength of
dispersion forces
How Do We Explain This?
• The nonpolar series SnH4
to CH4 follow the expected
trend.
• The polar series follows the
trend from H2Te through
H2S, but water is an
anomaly
• This is due to…
Hydrogen Bonding
• The dipole-dipole interactions
experienced when H is bonded to
N, O, or F are unusually strong.
• We call these interactions hydrogen
bonds.
• Hydrogen bonding arises from the
high electronegativity of N, O, and
F.
• Also when hydrogen is bonded to
one of those elements, the
hydrogen nucleus is exposed (very
positive side of compound)
Ion-Dipole Interactions
• A fourth type of force, ion-dipole interactions, are an
important force in solutions of ions.
• The strength of these forces are what make it possible for
ionic substances to dissolve in polar solvents.
Summarizing Intermolecular Forces
Predicting the Types and Relative
Strengths of Intermolecular Forces
• List the forces involved in each substance then list in
order of increasing boiling points:
• Barium Chloride
• Hydrogen molecule
• Carbon Monoxide
• Hydrogen Fluoride
• Neon
Solution
• H2<Ne<CO<HF<BaCl2
• Attractive forces are stronger for ionic substances than
molecular.
• BaCl2 should be highest
• The intermolecular forces of the remaining substances
depend on molecular weight, polarity, and hydrogen
bonding.
• H2 lowest –nonpolar and lowest molecular weight (LDF)
• Ne – nonpolar, higher molecular weight than H2 (LDF)
• CO, HF, molecular weights are roughly the same. (LDF, D-D)
• HF also has hydrogen bonding (LDF, D-D, HB)
Practice
Substance
Boiling Point
Water, H2O
100C
Ammonia, NH3
-33C
Methane, CH4
-164C
In terms of intermolecular forces, account for the differences in the
boiling points in the above table.
Water and ammonia have higher boiling points than methane
because water and ammonia form hydrogen bonds and methane
does not. Hydrogen bonds are relatively strong intermolecular
bonds and require more energy to separate them. Water has a
much higher boiling point than ammonia because water has two
nonbonding pairs of electrons per molecule versus one
nonbonding pair for ammonia. Also, oxygen is more
electronegative than nitrogen. Therefore, the hydrogen bonding in
water is much stronger than that in ammonia.
Intermolecular Forces Affect Many
Physical Properties
• The strength of the
attractions between
particles can greatly
affect the properties of a
substance or solution.
• Viscosity
• Surface tension
• Boiling point
Viscosity
• Resistance of a liquid
to flow
• Related to the ease
with which molecules
can move past each
other.
• Viscosity increases
with stronger IMF
and decreases with
higher temperature.
Surface Tension
• Results from the net inward
force experienced by the
molecules on the surface of
a liquid
• These are COHESIVE
forces
• Cohesive forces are forces
bind similar molecules to
one another
• Adhesive forces bind a
substance to a surface
• Capillary action
• miniscus
Phase Changes
Energy Changes Associated with
Changes of State
• Heat of Fusion: Energy required to change a solid at its
melting point to a liquid. ∆Hfus
• Heat of Vaporization: Energy required to change a liquid
at its boiling point to a gas. ∆Hfus
• Notice that the heat of vaporization is always larger than
the heat of fusion
1. Heating Curves
• The heat added to the
system at the MP and
BP goes into pulling the
molecules farther apart
from each other.
• The temperature of the
substance does not rise
during the phase change
• For Water:
• ∆Hfus = 6.01 kJ/mol
• ∆Hvap = 40.7 kJ/mol
Calculating ∆H for Temperature and Phase
Changes
• Calculate the enthalpy change upon converting 1 mol of
ice at -25.0 C to water vapor at 125 C under a constant
pressure of 1 atm.
• The specific heats of ice, water, and steam are 2.09 J/g-K, 4.18 J/g-
K, and 1.84 J/g-K respectively.
• For water, ∆Hfus = 6.01 kJ/mol and ∆Hvap = 40.67 kJ/mol
• Heating stages: q = cm∆T/1000
• Phase changes: q = (∆Hfus/vap) (mol)
Answer
• Solid phase: q = (2.09 J/g-k) (18.01g) (25 C) / 1000
= 0.941 kJ
S-L phase change: q = (6.01 kJ/mol) (1 mol)
= 6.01 kJ
Liquid phase: q = (4.18 J/g-K) (18.01g) (100 C) / 1000
= 7.528 kJ
L-G phase change: q = (40.67 kJ/mol) (1 mol)
= 40.67 kJ
gas phase: q = (1.84 J/g-K) (18.01g) (25 C) / 1000
= 0.828 kJ
Total = 0.941 + 6.01 + 7.528 + 40.67 + 0.828 = 56.0 kJ
Practice
• What is the enthalpy change during the process in which
100 g of water at 50.0⁰C is cooled to ice at -30⁰C?
Vapor Pressure
• The vapor pressure of a liquid is the pressure exerted by
its vapor when the liquid and vapor are in dynamic
equilibrium
• Molecules reach dynamic equilibrium when evaporation
and condensation occur at equal rates.
• It may appear that nothing is occurring because there is no net
change,. However, molecules are constantly passing from the liquid
to gas phase and from the gas to the liquid phase.
2. Effect of Temp on Kinetic Energy
• When vaporization occurs in an open container,
equilibrium never occurs because the vapor continues to
form until the liquid completely evaporates.
• Liquids with high vapor pressure evaporate more quickly
than those with low vapor pressure. These are called
volatile liquids
• More volatile liquids have
more molecules with
sufficient energy to escape
a liquid’s surface.
• The weaker the
intermolecular forces, the
higher the vapor pressure
and more easily molecules
can escape
3. Vapor Pressure and BP
• The boiling point of a liquid
is the temperature at which
its vapor pressure equals
atmospheric pressure
• The normal boiling point is
the temperature at which
vapor pressure is 760 torr
(1 atm)
• Estimate the boiling point of diethyl ether under an external
pressure of 0.80 atm.
• At what external pressure will ethanol have a boiling point
of 60⁰C?
4. Critical Temperature and Pressure
• A gas normally liquefies at
some point when pressure
is applied to it.
• The highest temperature at
which a distinct liquid
phase can form is referred
to as the critical
temperature
• The critical pressure is the
pressure required to bring
about liquefaction at this
critical temperature
Phase Diagrams
• Below A the substance cannot exist in the liquid state
• Along the AC line the solid and gas phases are in
equilibrium; the sublimation point at each pressure is
along this line
Phase Diagrams of H2O and CO2
• Water’s melting line slants
left with increasing pressure
• This indicates that for water
the melting point decreases
with increasing pressure
• The occurs with water
because the liquid form is
more compact the the solid
form.
• CO2 does not exist at a
liquid at 1 atm
• Solid Cos sublimes at 1atm
to form a gas
• CO2 does not have a normal
melting point, but instead a
normal sublimation point
Liquid Crystals
• Liquid crystals do not pass directly from the solid to liquid
phase. They pass through an intermediate liquid crystal
phase that has some of the structure of solids and some
of the freedom of motion of liquids.
Liquid Crystals
• Nematic liquid crystal
• Molecules are aligned so their long axes to point in same direction
but ends are not aligned
• Smectic A liquid crystal
• Molecules are aligned along the long acess and along their ends
• Smectic C liquid crystal
• Same as A except the are inclined with respect to their layer planes
LCD Displays
• Liquid Crystal displays (SCDs) are used in watches,
calculators, and computer screens
• A thin liquid crystalline material is placed between 2
electrically conducting glass electronde
• When electricity passes through the electrodes the
polarization of the light can “twist” which will allow light to
flow through it.