AP Chapter 11
Intermolecular Forces / Liquids and
Solids
HW: 2 4 6 13 15 16 19 31 39 45 105
11.1 – A Comparison of the States of
Matter
The fundamental difference between states of
matter is the distance between particles.
Indefinite volume
Indefinite shape
Rotational, Vibrational and
Translational Particle Motion
Definite volume
Indefinite shape
Rotational and Vibrational
Definite volume
Definite shape
Rotational Motion
States of Matter
Because in the solid and liquid states particles are
closer together, we refer to them as condensed
phases.
Phase
• Homogeneous part of a system in contact with other
parts of the system but separated from them by a
well defined boundary
• The state/phase a substance is in at a particular
temperature and pressure depends on two
properties:
– The kinetic energy of the particles
– The strength of the attractions between the particles
11.2 - Intermolecular Forces
(External Bonds)
The attractions between molecules are not nearly
as strong as the intramolecular attractions
(covalent bonds) that hold compounds together.
Intermolecular Forces
They are, however, strong enough to control
physical properties such as boiling and melting
points, vapor pressures, and viscosities.
1. Ion-Dipole Interactions
• 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
• Hydration = Process of water molecules surrounding a
solute
2. Intermolecular Forces
Intermolecular forces between molecules are
referred to as van der Waals forces.
van der Waals Forces
• a. Dipole-dipole interactions
• b. London Dispersion forces
– Induced Dipoles (not a true London Dispersion
Force)
– Instantaneous Dipoles
a. Dipole-Dipole Interactions
• Molecules that have permanent
dipoles 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.
– The larger the dipole
moment, the stronger the
attraction
Dipole-Dipole Interactions
The more polar (larger resultant dipole moment)
the molecule, the higher is its boiling point
because the attractions are stronger.
Hydrogen Bonding
• Can be considered a unique kind of
dipole-dipole interaction
• Hydrogen atom in a polar bond
externally bonds to the non-bonding
electron pair (lone pair) on an
electronegative atom
– Polar bond usually the H connected to a
N, O, or F
• We call these interactions hydrogen
bonds.
• Often seen in biological systems.
Hydrogen Bonding
Hydrogen bonding arises in part from the high
electronegativity of nitrogen, oxygen, and
fluorine.
Because when hydrogen is bonded to one of those
very electronegative elements, the hydrogen nucleus is
exposed.
b. London Dispersion Forces
(occur between nonpolar substances/atoms)
-Induced Dipoles (not all books talk about these. Not
a TRUE London Dispersion Force)
• Occurs when a nonpolar atom/molecule
comes in proximity to charged
particle/molecule
• Induces a charge on the nonpolar molecule
+
->
And then
is
attracted
to dipole
London Dispersion Force
-Instantaneous Dipole
While the electrons in the 1s orbital of helium would
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.
Causes an instantaneous dipole. The same can occur
in any molecule…
Instantaneous Dipole
Looks like this
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
-A London Dispersion Force Occurs
-London dispersion forces, or dispersion forces,
are attractions between an instantaneous dipole
and an induced dipole.
London Dispersion Forces
• These forces are present in all molecules,
whether they are polar or nonpolar.
• The tendency of an electron cloud to distort in
this way is called polarizability.
Factors Affecting London Forces
• The shape of the molecule
affects the strength of
dispersion forces: 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.
(that’s why oil is thicker than water)
Factors Affecting London Forces
• The strength of dispersion forces tends to
increase with increased molecular weight.
• Larger atoms have larger electron clouds, which
are easier to polarize.
Dipole-Dipole Interactions vs.
London Forces
• Nonpolar molecules only have London Forces
• Dipole molecules have both Dipole-Dipole and
London Forces
– Which is more important to the property of the
molecule??
• When comparing two dipole molecules with
different BP, etc.:
– If size, shape and # of e-s of the two molecules
are very similar, the reason the BPs are different
are due to dipole-dipole interactions
– If the molecules are of very different size than
dispersion forces are the main explanation
Dispersion Forces vs. Dipole-Dipole Interactions
Why choose these two molecules to compare? Both have identical numbers of
electrons, and if you made models you would find that the sizes were similar - as
you can see in the diagrams. That means that the dispersion forces in both
molecules should be much the same.
The higher boiling point of fluoromethane is due to the large permanent dipole on
the molecule because of the high electronegativity of fluorine. However, even given
the large permanent polarity of the molecule, the boiling point has only been
increased by some 10°. (Dipole-Dipole Interactions don’t always make that much of
a difference).
Dispersion Forces vs. Dipole-Dipole Interactions
Here is another example showing the dominance of the dispersion forces.
Trichloromethane, CHCl3, is a highly polar molecule because of the
electronegativity of the three chlorines. There will be quite strong dipole-dipole
attractions between one molecule and its neighbours.
On the other hand, tetrachloromethane, CCl4, is non-polar. The outside
of the molecule is uniformly - in all directions. CCl4 has to rely only on
dispersion forces
So which has the highest boiling
point? CCl4 does, because it is a
The boiling points are:
bigger molecule with more
CHCl3
61.2°C
electrons. The increase in the
CCl4
76.8°C
dispersion forces more than
compensates for the loss of
dipole-dipole interactions.
11.3 - Physical Properties of Liquids
The strength of the
attractions between
particles can greatly
affect the properties
of a substance or
solution.
Viscosity
• Resistance of a liquid to
flow is called viscosity.
• It is related to the ease
with which molecules
can move past each
other.
• Viscosity increases with
stronger intermolecular
forces and decreases
with higher temperature.
Surface Tension
No molecules pull
upward on the
surface
Surface tension results
from the net inward
force experienced by
the molecules on the
surface of a liquid.
Capillary Action
• Liquid in a tube is not at the same level as the
liquid in the surrounding container
-Cohesion – Intermolecular attraction between liquid molecules
-Adhesion – Attraction between unlike molecules (ex – liquid and the
tube)
Adhesion > Cohesion
Cohesion > Adhesion
11.4 – Phase Changes
• Transfer from one phase to another by
addition or removal of energy
• Liquid – Vapor Equilibrium
– When molecules have enough energy they
escape the surface
– Higher T, Higher E, More Evaporation
Vapor Pressure
• At any temperature, some molecules in a liquid
have enough energy to escape.
• As the temperature rises, the fraction of
molecules that have enough energy to escape
increases.
Vapor Pressure
As more molecules escape the liquid, the
pressure they exert increases.
Vapor Pressure
The liquid and vapor
reach a state of
dynamic equilibrium:
liquid molecules
evaporate and vapor
molecules condense
at the same rate. (in a
closed container)
Vapor Pressure
• 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 its vapor pressure
is 760 torr.
• Molar Heat of
Vaporization = Energy
needed to vaporize one
mole of liquid.
Clausius-Clapeyron Equation (p. 456)
• Describes the relationship between P (vapor
pressure) and T
• Plot ln P on y-axis vs 1/T on x-axis
ln(Pvap) = -DHvap/(RT) + c
– Slope = -DHvap/R
– y-intercept = c
– R = 8.314 J/molK
• Comparing two Temps:
Sometimes shown without –
and T2 and T1 reversed…
Boiling Point
• Temp at which VP liquid = P external
• Water = Normal BP = 100oC
– Normal = P external = 1 atm
Critical Temperature and Pressure
• Gas -> Liquid occurs by either decrease
temp or increase pressure
• Critical Temp = Tc = “Last” temperature at
which a liquid can exist. At ALL pressures
above this temp, only a gas can exist
• Critical Pressure = Pc = Minimum pressure
that must be applied to bring about
liquification at the Critical Temperature
Liquid – Solid Equilibrium
• Freezing = l->s
• Melting/Fusion = s->l
• Melting Point /
Freezing Point =
Temperature at which
solid and liquid are at
equilibrium
• Water = Normal fp =
0oC
• Molar Heat of Fusion
= Energy required to
melt 1 mole of a solid
Heating / Cooling Curves:
Supercooling
• Liquid temporarily cooled below Freezing
Point
• Happens when removed from heat so
rapidly that molecules don’t have time to
form solid.
– Unstable
– Can solidify by seed crystal or scratch on side
of container
Solid-Vapor Equilibrium
• Sublimation = s->g
• Deposition = g->s
• Molar Heat of Sublimation = Energy
required to sublime 1 mole solid
Heat Calculations
• BEFORE Phase Changes (temp is changing):
q = msDT
• DURING Phase Changes (temp constant due
to all energy going into changing bonds):
q = m*1/M*DH
moles
11.6 - Phase Diagrams
Phase diagrams display the state of a substance at
various pressures and temperatures and the
places where equilibria exist between phases.
Phase Diagrams
• Triple Point = Conditions at which s, l, g all coexist
– Water = 0.01oC and 0.006 atm
• The AB line is the liquid-vapor interface.
• It starts at the triple point (A), the point at which all
three states are in equilibrium.
Phase Diagrams
It ends at the critical point (B); above this critical
temperature and critical pressure the liquid and
vapor are indistinguishable from each other.
Phase Diagrams
Each point along this line is the boiling point of
the substance at that pressure.
Phase Diagrams
• The AD line is the interface between liquid and
solid.
• The melting point at each pressure can be found
along this line.
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 Diagram of Water
• Note the high critical
temperature and
critical pressure:
– These are due to the
strong van der Waals
(both dipole-dipole
and London
Dispersion) forces
between water
molecules.
Structure and Properties of Water
• Water is a very unique
compound
• Has a VERY high
specific heat (can
absorb large amounts
of heat but with small
change in temperature)
• Solid is less dense than
the liquid
• Maximum density
occurs at 4oC
Phase Diagram of Water
• The slope of the
solid–liquid line is
negative.
– This means that as
the pressure is
increased at a
temperature just
below the melting
point, water goes
from a solid to a
liquid.
– RARE
Phase Diagram of Carbon Dioxide
Carbon dioxide cannot
exist in the liquid state
at pressures below
5.11 atm; CO2
sublimes at normal
pressures.
Phase Diagram of Carbon Dioxide
• Solid – Liquid line has
a + slope
• Cannot liquify a solid
using pressure
*ice skating
11.7 – Crystal Structure of Solids
• We can think of solids
as falling into two
groups:
– Crystalline—particles
are in highly ordered
arrangement.
Solids
– Amorphous—no
particular order in the
arrangement of
particles.
Crystalline Solids
Because of the order
in a crystal, we can
focus on the
repeating pattern of
arrangement called
the unit cell.
Unit Cells
• There are 7 basic types of unit cells
11.8 – Bonding in Solids
Molecular Crystals
• Molecules held together by external bonds
– Ex: Ice, Iodine (s), S8, P4
11.8 – Bonding in Solids
Covalent-Network
• Diamonds are an example of a covalent-network
solid in which MANY atoms are covalently bonded
to each other.
Covalent-Network
• Graphite is an example in which atoms are held
together by covalent bonds and the layers are
held with van der Waals forces.
11.8 – Bonding in Solids
Ionic
– Ions at the lattice pts
– Cations and anions in alternating pattern
CsCl
ZnS
CaF2
11.8 – Bonding in Solids
Metallic
• Metal Ions in an electron sea
Types of Crystalline Solids
Amorphous Solids
-No pattern
-Glass
-Quickly cooled liquid->solid