Chapter 12
Intermolecular Forces
Earth: 15 ºC
Uranus: -214 ºC
What happens to molecules at the melting point?
Intramolecular forces
versus
Intermolecular forces
(aka. van der Waals forces)
Chemical Change
Breaks covalent, ionic
and metallic bonds
Physical Change
Electrostatic forces
between particles
Intramolecular forces
Intermolecular forces
Section 12.1: Physical States and Phase Changes
Kinetic-Molecular View of Three States of Matter
Increasing Energy
P.E. and K.E. together
determine the physical state
of any given substance.
Solid
Liquid
Gas
P.E. – draws molecules together
Coulomb’s Law (Chap2) – particles with
opposite charge attract each other.
The energy of attraction between two
particles is proportional to the product of
the charges and inversely proportional
to the distance between them.
K.E. – separates or disperses molecules
K.E. ~ f(absolute temperature)
(in Kelvins)
ºC = K - 273
Section 12.1: Characteristics of Physical States
Increasing Energy
Solid
• Definite shape.
• Definite volume.
• Particles fixed & close.
• Particle interaction v. strong.
• Particle movement v. slow.
Ex: ice, iron, table salt
Liquid
Gas
• Takes shape of container.
• Definite volume.
• Particles (molecules) random
& close.
• Particle interaction strong.
• Particle movement moderate.
Ex: water, oil, vinegar
• Takes shape of container.
• Fills container volume.
• Particles (molecules) random
and far apart.
• Essentially no interaction.
• Particle movement very fast.
Examples: water vapor, helium
gas
Section 12.1: Phase Changes
Condensation: Gas to liquid
Vaporization: Liquid to gas
Heat of vaporization
Freezing: Liquid to solid
Melting (Fusion): Solid to liquid
Heat of fusion
Sublimation: Solid to gas
Deposition: Gas to solid
Heat of sublimation
Section 12.1: Energy and Phase Changes
Enthalpy changes (∆H) accompany phase changes
Exothermic Phase Changes (-∆H)
Condensation: Gas to liquid
Freezing: Liquid to solid
Deposition: Gas to solid
Enothermic Phase Changes (+∆H)
Vaporization: Liquid to gas
Melting (Fusion): Solid to liquid
Sublimation: Solid to gas
Section 12.1: Energy and Phase Changes
Enthalpy change is different for different substances
For a pure substance: ∆H is measured in change per mole of the substance
and is specific to the pressure and temperature conditions
Pressure is usually 1 atm, Temperature is that at which the phase change occurs
Example: Phase changes of water
H2O (l)  H2O (g)
H2O (s)  H2O (l)
∆H = ∆H º vap = 40.7 kJ/mol (at 100 ºC)
∆H = ∆H º fus = 6.02 kJ/mol (at 0 ºC)
H2O (g)  H2O (l)
H2O (l)  H2O (s)
∆H = ∆H º vap = -40.7 kJ/mol (at 100 ºC)
∆H = ∆H º fus = -6.02 kJ/mol (at 0 ºC)
Why is ∆H º vap (40.7 kJ/mol) greater than ∆H º fus (6.02 kJ/mol)?
∆H º subl = ∆H º fus + ∆H º fus
Section 12.2: Quantifying Phase Changes
Temperature
The Heating-Cooling Curve – shows how the temperature of a substance changes as
heat is added or removed from a substance at a constant rate (at a constant P too)
Heat Removed
Interlude: Pressure Matters too
(but we assume 1 atm in this class for phase change calculations)
Interlude: Pressure Matters too
(but we assume 1 atm in this class for phase change calculations)
Methane hydrates
http://www.windows.ucar.edu/tour/link=/earth/Water/temp.html&edu=mid
Pressures ~ 1000 atm
CH4 freezing point: -182.5 ºC
http://pathways.fsu.edu/faculty/geeo/
Section 12.2: Quantifying Phase Changes
But back to temperature……….. 5 heat-releasing stages
Temp change: q = nC∆T
where q is heat, n is # of moles, ∆H is heat released/absorbed
Temperature
Temp constant: q = n∆H
where q is heat, n is # of moles, C is molar heat capacity
Heat Removed
Section 12.2: Quantifying Phase Changes
In class problem: 12.20
Suggested problem: 12.2. 12.3, 12.12, 12.19, 12.27
Section 12.2: Equilibrium and Phase Changes
In a closed system, phases changes of many substances reach equilibrium.
Open Container
Closed Container
Open system
– volume of liquid does not change
– net direction of molecule movement
is out of the liquid
Closed system
– volume of liquid does not change
– net direction of molecule movement
is out of the liquid
Your system is defined by you:
Systems can be closed to
some things, but not others
Is the ocean a closed system?
Closed Container
Heat source
Is Earth a closed system?
Concept of Open Systems and Steady-State
When matter/energy is leaving and entering an open system, it can reach Steady-state
Water in
Definition of steady-state: FluxIN = FluxOUT
Flux: Mass or Volume / time
Ex: 100 L H2O / hr
20 g CaCO3 / day
Water out
This system is not in steady-state if the volume changes with time.
Water in
Water in
Water in
Water out
Water out
Time
Water out
Section 12.2: Liquid-Gas Equilibrium and Vapor Pressure
Equilibrium vapor pressure – the pressure exerted by a vapor when it has reached
equilibrium in a system that is closed with respect to the vapor molecules
Vapor is stuck in the container and will accumulate,
putting pressure (P=Force/Area) on container walls.
When enough time passes,
the system will reach equilibrium
with respect to the vapor entering
and exiting the liquid.
Universal Concept: When a system at equilibrium is disturbed, it counteracts the
disturbance and eventually re-establishes a state of equilibrium
(For chemical reactions, called Le Châtelier’s principle  Chap17, CHEM 163)
Section 12.2: Liquid-Gas Equilibrium and Vapor Pressure
Higher T = Higher V.P.
Higher T  increases the fraction of molecules moving fast enough to escape the liquid
 decreases the fraction of molecules moving slow enough to be captured
Section 12.2: Liquid-Gas Equilibrium and Vapor Pressure
Clausius-Clapeyron equation - mathematical relationship between T and P
Nonlinear relationship between T and P
(in graph) expressed as linear relationship:
Know: P1,
T1, ∆Hvap
Section 12.2: Quantifying T – P Relationships
In class problem: 12.22, 12.24
Suggested problem: 12.21, 12.23
Section 12.2: Vapor Pressure and Boiling Point
Boiling point – temperature at which the vapor pressure equals the external pressure
Pressure exerted on the Earth by all the gas particles in the
Earth’s atmosphere.
Atmospheric Pressure
Pressure = Force / Area
From physics: F = ma
F  force
m  mass of particle
a  acceleration (= g, acceleration due to gravity)
Why is my soup not as hot at Camp Muir?!??!
Altitude = 10,000 feet
Atmospheric pressure lower = 590 mm Hg
Boiling Point = 90 °C
Altitude = Sea level
760 mm Hg (=1 atm)
Boiling Point = 100 °C
People living in Denver, CO use pressure cookers to cook food at higher temperature.
Section 12.3: Types of Intermolecular Forces
Bonding (Intramolecular) forces:
Relatively strong  involve large
charges that are closer together
Nonbonding (Intermolecular) forces:
Relatively weak  involve smaller
charges that are farther together
Section 12.3: Types of Intermolecular Forces
Why are bonding (intramolecular) forces stronger
than van der Waals (intermolecular forces)?
Periodic Table trends are similar to
those for bond length.
Section 12.3: Types of Intermolecular Forces
(1) Ion-dipole – an ion interacts with a partial charge
Example: NaCl (table salt) dissolves in water
Na
Cl
Na
Cl
Cl
Na
Cl
Na
Na
Cl
Na
Cl
Cl
Na
Cl
Na
O
+
Dissolution
(NaCl “dissociates”)
+
+
+
+
+
+
+
Cl
Na
+
-1
+
+
+
+
+
+
+
+
+1
H
H
+
Section 12.3: Types of Intermolecular Forces
(2) Dipole-dipole – polar molecules interact
The greater the dipole moment of a molecule, the great the dipole-dipole forces
between molecules of that type  more energy needed to separate them
Section 12.3: Types of Intermolecular Forces
(2) Dipole-dipole – polar molecules interact
Hydrogen bond – a special type of dipole-dipole force that arises between atoms that
have a H atom bonded to a small, highly electronegative atom with lone electron pairs
+
N, O, and F all fit this profile.
+
+
+
O
+
+
+
+
+
H
H
+
(3) Charge-induced dipoles – a molecule with a full or partial charge induces a
temporary dipole on a nonpolar molecule
(4) London (dispersion) forces – caused by momentary oscillations of e- charge
in atoms and, therefore, are present in all particles (atoms, ions, and molecules)
Section 12.3: Trends in Polarizability
Polarizability – the ease with which the e- cloud of a particle can be distorted
Smaller atoms (ions) are less
polarizable than larger ones  e-’s
closer to the nucleus and, therefore,
held more tightly
Polarizability
• Increases down a Group
• Decreases from L  R
• Cations less polarizable than their
original atoms
Anions are more polarizable than
original atoms
Why dry ice (solid
CO2) sublimates
Biodiesel Lab: Bomb Calorimeter
Combustion reaction – heat flows from the system to the surroundings = exothermic
Heat is lost to: (1) water in the calorimeter
(2) the calorimeter itself
Section 12.4: Zooming in on Liquids
Liquids are least understood at the molecular level.
Increasing Energy
Solid
Orderliness of particles 
Different regions identical
Liquid
Orderly &
random at
different times
Gas
Randomness of particles 
any region is pretty much
identical to any other
Macroscopic properties of liquids are well understood:
• Surface tension
• Capillarity
• Viscosoty
Section 12.4: Surface Tension
Intermolecular forces exert different effects on a molecule at the surface of a liquid
than at the interior: A liquid tends to minimize the # of molecules at the surface.
Interior molecules – attracted by water
molecules on all sides
Surface molecules –attracted to water
molecules below and on sides 
Experience a net downward attraction
How insects walk on water. (water strider)
• At the surface of a liquid, water molecules behave as a thin, elastic membrane or
“skin”  surface tension – energy required to increase the surface area (J/m2 of
surface area increased)
Section 12.4: Surface Tension
The stronger the forces are
between the particles in a liquid,
the greater the surface tension.
Surfactants (surface- active agents)—destroy surface tension by congregating at the
surface and disrupting the hydrogen bonds between water molecules
Example: Needle on water.
Example: Respiratory distress syndrome (RDS) in infants.
Occurs when infant
H-bonds break,
does not produce a
needle sinks.
surfactant that breaks
H-bonds and does not
allow O2 and CO2
exchange between
alveoli and capillaries in
the lungs.
H-bonds hold needle
on water surface.
Section 12.4: Capillarity
Capillary action – the rising of a liquid through a narrow space against the pull of
gravity  due to competition between intermolecular forces in a liquid (cohesive forces)
And those between the liquid and the tube walls (adhesive forces)
TLC and plant pigment lab
Meniscus on a test tube
Glass = SiO2
Mercury
Water
(H-bonds with SiO2) (Metallic bonds stronger than
any interaction with SiO2)
Section 12.4: Viscosity
A liquid’s resistance to flow  resistance decreases as Temp increases
Molecular shape plays a role – Biodiesel lab
triglyceride
+
methanol
Larger molecules – make
more contact = higher viscosity
3 methyl ester
+
glycerol
Smaller molecules –
make less contact = lower viscosity
Section 12.5: Uniqueness of Water
••
The water molecules is bent and highly polar 
due to this structure and charge distribution, water
can engage in four H bonds with its neighbors.
O
+
H
H
+
(1) Water is the “universal solvent” (solvent = the compound that does the dissolving)
Dissolves a range of solutes ( = the compounds that are dissolved)
Ionic substances
Na
Cl
Na
Cl
Cl
Na
Cl
Na
Na
Cl
Na
Cl
Cl
Na
Cl
Na
Polar Covalent
substances
Nonpolar Covalent
substances
CH3CH2OH
C6H12O6
N2 gas
Section 12.5: Uniqueness of Water
(2) Water has a high specific heat capacity (the measure of the heat absorbed by a
substance for a given rise in temperature – Section 6.3)
In other words, water can absorb a lot of heat with relatively small changes in temp.
Earth: Daily temperature changes = 40 ºC (in deserts – most extreme)
Waterless Moon: 250 ºC daily fluctuations
Water has a high heat of vaporization –
heat from Sun results in vaporization of
ocean water  heat stored in water vapor
carried poleward  heat released when
water vapor condenses back to liquid water
– called latent heat transport
oceanmotion.org/html/background/climate.htm
Section 12.5: Uniqueness of Water
(3) Surface properties are crucial to living things
Trees get water due to capillary action
in soils and in xylem (veins of trees)
Solids
(minerals)
Air
Water
Plant debris floating on water surface
provides shelter and nutrients
Section 12.5: Uniqueness of Water
(4) Density of solid and liquid water
Large spaces in the ice due to the hexagonal crystal structure result in solid water
being more dense than liquid  lake surfaces freeze in winter (organisms live below)
Section 12.5: Uniqueness of Water Summarized
Heating-Cooling Curve Practice
How much heat would need to be added to heat 50.0 g of water ice at -50.0 ºC to
water vapor at 135 ºC?
Given:
Cice = 37.6 J/mol ºC
Cliquid = 75.4J/mol ºC
Cgas = 33.1 j/mol ºC
∆Hfusion = 6.02 kJ/mol
∆Hvaporization = 40.7 kJ/mol
Answer: 1.59 x 105 J