Ch. 12: Liquids, Solids, and
Intermolecular Forces
Dr. Namphol Sinkaset
Chem 200: General Chemistry I
I. Chapter Outline
I.
II.
III.
IV.
V.
VI.
VII.
Introduction
Solids, Liquids, and Gases
Intermolecular Forces
Intermolecular Forces in Action
Vaporization and Vapor Pressure
Sublimation and Fusion
Heating Curves
I. Condensed States
• Liquids and solids are the condensed
states because of the close proximity of
atoms/molecules to one another.
• This proximity leads to much more
frequent interactions than in gases.
• Structure determines properties via the
nature of the interactions that occur.
I. Ethanol vs. Dimethyl Ether
• How can the same 9 atoms form two
compounds with such different boiling points?
• It’s because of the structure of the molecules!
I. Electrostatic Forces
• Every molecule in a sample of matter
experiences two types of electrostatic
forces.
 Intramolecular forces: the forces that exist
within the molecule (bonding). These
forces determine chemical reactivity.
 Intermolecular forces: the forces that exist
between molecules. These forces
determine physical properties.
II. Solids, Liquids, and Gases
II. Two Types of Solids
II. Differences
Between the States
• A few things stand out when we look at
the table.
 Densities of the condensed phases are
much higher.
 Distances between particles much smaller
in the condensed phases.
 More freedom of movement in liquid than
in solid.
II. Properties of States of Matter
• The state of matter is a “battle” between IM
forces and thermal energy (temperature).
II. Phase Changes
• Phase changes can be induced by
adding/removing heat or increasing/decreasing
pressure.
III. Intermolecular Forces
• IM forces originate from interactions between
charges, partial charges, and temporary charges
on molecules.
• IM forces are relatively weak because of smaller
charges and the distance between molecules.
III. Types of IM Forces
• There are different kinds of IM forces,
each with a different level of strength.




Dispersion force
Dipole-dipole force
*Hydrogen “bonding”
Ion-dipole force
III. Dispersion Force
• Dispersion force (London force) is present in
all molecules and atoms and results from
changes in e- locations.
III. Instantaneous Dipoles
• Charge separation in one creates charge
separation in the neighbors.
III. Dispersion Force Strength
• The ease with which e-’s can move in
response to an external charge is
known as polarizability.
• Large atoms with large electron clouds
tend to have stronger dispersion forces.
• Larger molecules tend to have stronger
dispersion forces.
III. Noble Gas Boiling Points
III. Dispersion Force and Shape
• Molecular size is not the only factor…
III. Dispersion Force and Shape
• Shape influences how the molecules interact
with one another…structure determines
properties.
III. Dispersion Force in a Family
III. Dipole-Dipole Force
• Occurs in polar molecules which have permanent
dipoles, so attraction is always present.
III. Effect of Dipole-Dipole Force
• Polar molecules have dispersion forces and
dipole-dipole forces.
• Effects can be seen in boiling and melting points.
III. Increasing Polarity
• If we increase
the polarity, but
keep molar mass
approximately
the same…
III. “Like Dissolves Like”
• Polar liquids are miscible (mix without
separating) with other polar liquids, but not
with nonpolar liquids.
• Can be explained with intermolecular forces.
III. Hydrogen “Bonding”
• This IM force is a misnomer since it’s
not an actual bond.
• Occurs between molecules in which H
is bonded to a highly electronegative
element (N, O, F), leading to high partial
positive and partial negative charges.
• It’s a “super” dipole-dipole force.
III. H “Bonding” Ethanol
III. Ethanol vs. Dimethyl Ether
• We compared these earlier.
• Hydrogen “bonding” is so much stronger than
dipole-dipole that one is a liquid at room
temp. while the other is a gas.
III. Effect of H “Bonding”
• Hydrogen
“bonding” is a
very strong
intermolecular
force.
• Without
hydrogen
“bonding” life as
we know it could
not exist!
III. Ion-Dipole Force
• Present in mixtures
of ionic compounds
& polar compounds.
• Example: NaCl(s)
dissolved in water.
III. Summary of IM Forces
III. Sample Problem
• Which substance has the highest boiling
point and why?
a) CH3OH
b) CO
c) N2
IV. Effect of IM Forces
• IM forces lead to a variety of different
physical properties.
 Without IM forces, only gas phase would
exist.
 Surface Tension
 Viscosity
 Capillary Action
IV. Solid, Liquid, or Gas?
• Whether a substance exists as a solid,
liquid, or gas depends on the
relationship between the intermolecular
attractions and the kinetic energy of the
molecules.
 It’s a battle – which dominates? The KE or
the IM attractions?
 Recall that the average KE of a sample is
related to its temperature.
IV. KE vs. IM Forces
• Gas: the kinetic energy of the molecules is
much greater than the intermolecular
attractions.
• Liquid: the kinetic energy of the molecules
is moderately greater than the
intermolecular attractions.
• Solid: the kinetic energy of the molecules
is less than the intermolecular attractions.
IV. More Dense Objects Can
Float
IV. Surface Tension
• IM force is an
attractive force.
• More surrounding
molecules means
more attractive
forces and thus
lower PE.
• Everyone wants to
be in the interior.
IV. Benefit of Being Spherical
• If all molecules want
to be interior, then
substance will seek
lowest surface area.
• Surface tension acts
as a “skin,” resisting
any increase in
surface area because
that requires energy.
IV. Viscosity
• Viscosity is the resistance of a liquid to
flow.
• The unit for viscosity is poise (P) which
is defined as 1 g/cm·s.
• Stronger IM forces increase viscosity
(particles more “sticky”), but temp. can
decrease viscosity.
• Longer molecules are more viscous
because they can become tangled.
IV. Sample Viscosities
IV. Capillary Action
• Capillary action is the ability of a liquid
to flow against gravity up a narrow tube.
• It is the result of two forces:
 Cohesive forces (IM forces)
 Adhesive forces (attractions between
molecules and surface of tube).
• Need adhesive > cohesive for liquid to
flow up the tube.
IV. Cohesive vs. Adhesive
V. Vaporization and IM Forces
• From experience,
we know that water
evaporates in an
open container.
• What factors
influence rate of
vaporization?
V. Vaporization Variables
• Temperature
• Surface area
• IM forces
V. Energetics of Vaporization
• As molecules evaporate, what happens
to the temperature of the samples left in
the beaker?
• Vaporization is an endothermic process
– it’s the reason why we sweat when we
get too hot.
• Condensation is an exothermic process.
V. Heat of Vaporization
• The energy needed to vaporize 1 mole of a
liquid to gas is the heat of vaporization, ΔHvap.
• Can be thought of the energy needed to
overcome IM forces of the liquid.
V. Dynamic Equilibrium
• In an open
flask, a liquid
will eventually
evaporate
away.
• What about a
closed flask?
V. Dynamic Equilibrium
• As evaporation occurs,
headspace fills with gas
molecules.
• Gas molecules condense
back to liquid phase.
• Eventually, rates become
equal.
• Pressure of gas at
dynamic equilibrium is
called the vapor
pressure.
V. Dynamic Equilibrium
• Systems at dynamic equilibrium will seek to
return to dynamic equilibrium when disturbed.
V. Vapor Pressure and Temp.
• Vapor pressure
depends on
temperature and
IM forces.
• Why?
V. Boiling Point
• When T is increased, the vapor pressure
increases due to the higher # of molecules
that can break away and enter gas phase.
• What if all molecules have necessary
thermal energy?
• At this point, vapor pressure = external
pressure, and boiling point is reached.
• The temperature at which vapor pressure
equals 1 atm is the normal boiling point.
V. Boiling Point vs. Altitude
V. Boiling Point
• At the boiling point, those aren’t air bubbles!
V. Pvap – T Relationship
V. Clausius-Clapeyron Equation
• This equation is in linear form, y = mx + b.
• The heat of vaporization can be found using
graphical analysis.
• Use R = 8.314 J/mole·K.
V. Graphical Analysis
V. Clausius-Clapeyron
Equation, 2-point Form
• If you have two sets of pressure, temperature
data for a liquid, the more convenient 2-point
form of the Clausius-Clapeyron equation can
be used.
V. Sample Problem
• Propane has a normal boiling point of
-4.20 °C and a heat of vaporization of
19.04 kJ/mole. What is the vapor
pressure of propane at 25.0 °C?
V. The Critical Point
• In a sealed container, as T of liquid is
heated, more and more vapor is formed,
and P increases.
• At the critical temperature, a supercritical
fluid forms; liquid can’t exist above this
temperature.
VI. Other Phase Changes
• Sublimation is the direct conversion of
particles from the solid phase to the gas
phase.
 Average KE is low, but always some that
have enough KE to break away.
• Fusion is the conversion of solid to
liquid.
• Also have deposition and freezing.
VI. Energetics of Fusion
• Different compounds
have different heats of
fusion.
• Notice they are much
lower than heats of
vaporization – why?
VII. Energies of Phase Changes
• The enthalpies involved in a phase
change depends on the amount of
substance and the substance itself.
• We look at a heating curve for 1.00
moles of H2O at 1.00 atm pressure.
• Note that there are sloping regions and
flat regions in the curve. (Why?)
VII. Heating Curve for H2O
VII. Heating Curve, Segment 1
• At this stage, we are heating ice from -25 °C
to 0 °C, increasing KE (vibrational motions).
• The heat required depends on the specific
heat capacity of ice.
VII. Heating Curve, Segment 2
• Here, the temperature stays the same, so the
average KE stays the same.
• Thus, the PE must be increasing.
• The heat gained is a factor of the heat of
fusion, the heat needed to melt 1 mole of
solid.
VII. Heating Curve, Segment 3
• During this stage, water is being heated from
0 °C to 100 °C; again, KE is increasing.
• The heat gained depends on the specific heat
capacity of water.
VII. Heating Curve, Segment 4
• Again, the temperature stays the same, so
the average KE stays the same.
• PE must be increasing.
• The heat gained is a factor of the heat of
vaporization.
VII. Heating Curve, Segment 5
• During this stage, steam is heated from 100
°C to 125 °C; average KE is increasing.
• The heat gained depends on the specific heat
capacity of steam.