Chapter 12 – States of Matter
12.1 Gases [KMT, diffusion/effusion,
pressure]
12.2 Forces of Attraction
12.3 Liquids and Solids
12.4 Phase Changes
Section 12.1 Gases
Gases expand, diffuse, exert pressure,
and can be compressed because they
are in a low density state consisting of
tiny, constantly-moving particles.
• Use the kinetic-molecular theory to explain the
behavior of gases.
• Describe how mass affects the rates of diffusion and
effusion.
• Explain how gas pressure is measured and calculate
the partial pressure of a gas.
Section 12.1 Gases
Key Concepts
• The kinetic-molecular theory explains the properties of
gases in terms of the size, motion, and energy of their
particles.
• Dalton’s law of partial pressures is used to determine
the pressures of individual gases in gas mixtures.
• Graham’s law is used to compare the diffusion rates
of two gases.
Kinetic-Molecular Theory
Liquids and solids have wide range of
physical properties at room
temperature
Gases at room temperature usually
display similar physical properties
despite having different compositions
Why?
Kinetic-Molecular Theory
KMT describes gas behavior by making
assumptions about the size, motion,
and energy of gas particles
Kinetic  “to move”
Objects in motion have kinetic energy
KE = ½ m v2 = kinetic energy of particle
• m = mass, v = velocity
KMT - General
Explains ideal gas behavior
Assumptions simplify the theory
• don’t work in real gases
General picture for gas particles
Small - can ignore their volume
In constant motion
• their collisions cause pressure
KMT Assumptions (1)
Small particles, far apart
• Gas mostly empty space
• Reason for compressibility of gases
No interactions between particles
• No attractive or repulsive forces
between molecules, except for direct
collision
KMT Assumptions (2)
Constant straight-line motion until
collision with wall or other molecule
Collisions are elastic
• No kinetic energy lost during collision
• KE1(b) + KE2(b) = KE1(a) + KE2(a)
1,2 = particles 1 & 2
(b) = before collision (a) = after collision
• KE1(b) = KE1(a) for collision with wall
KMT Assumptions (3)
Temperature (T) is a measure of the
average kinetic energy of gas particles
At a given temperature, all gases have
the same average KE
KE(avg) = ½ m v2 = f(T) only
Raise T, KE(avg) & speed increase
Lower T, KE(avg) & speed decrease
KMT – Gas velocities
KE(avg) = ½ m v2 = f(T) only
At 25 C, H2 & N2 have same KE(avg)
m(N2) / m(H2) = 28 / 2 = 14
[v2(H2)]avg / [v2(N2)]avg = 14
[v(H2)]avg = 14 [v(N2)]avg = 3.74 [v(N2)]avg
Lighter gases move faster than heavier
ones (at same temperature)
Range of velocities
Temperature is an average
Molecules of many speeds in the
average
Shown on a graph called a velocity
distribution
Number of Particles
273 K
1273 K
2000 K
Molecular Velocity
Maxwell-Boltzmann Distribution Curve
Lighter molecules move faster on average
O2 - heavy
Note how fast
they move
H2 - light
Gas Molecule Velocities
http://hyperphysics.phy-astr.gsu.edu/hbase/kinetic/eqpar.html#c2
http://hyperphysics.phy-astr.gsu.edu/hbase/sound/souspe3.html#c1
At 25C, vavg for H2 = 1776 m/s
= 3974 mi/hr
At 25C, vavg for Xe = 219 m/s
= 491 mi/hr
For comparison, the speed of sound in
dry air at 25C vsound = 346 m/s
= 776 mi/hr
KMT & Gas Behavior
Low density
Chlorine gas 2.95  10-3 g/mL @ 20C
Cl2 molecular mass = 70.9 amu
Liquid water 1.0 g/mL @ 20C
H2O molecular mass = 18.0
 Lots of space between gas
molecules relative to a liquid
KMT & Gas Behavior
Compression – Expansion
Large amount of space allows
compression
Constant motion of particles explains
expansion
Expand
Low Density
Low P
Compress
High Density
High P
KMT & Diffusion/Effusion
Diffusion – the movement of one
material through another
• One gas moves through another gas –
no interactions
Movement of odor molecules (perfume)
Move from high to low areas of
concentration
Lighter gases diffuse faster
KMT & Diffusion/Effusion
Effusion – escape of gas through small
opening
Thomas Graham – various gases
effusing into vacuum
Graham’s Law of Effusion
Rate  1/  molar mass
Also can apply to diffusion rates
rate A / rate B = mass B/  mass A
Gas Effusion Through Pinhole
Vacuum
Sealed
Separation
With Pinhole
in Separation
Graham’s Law: Sample Problem 12.1
Ratio of diffusion rate of ammonia
(MM=17.0 g/mol) to HCl (MM=36.5 g/mol)?
rate A / rate B = (mass B / mass A)
rate NH3 / rate HCl =
(36.5 g/mol / 17.0 g/mol)
rate NH3 / rate HCl = 1.47
Practice – Effusion/Diffusion
Problems 1- 3 page 405
Problems 42 - 44 page 434
Pressure
Pressure is force per unit area
P= F
A
SI unit = pascal (Pa) = 1 N/m2
Gas molecules fill container
Molecules move around and hit sides
Collisions with sides provide force
Container has area
Measured with a barometer
Atmospheric Pressure
Column of
air
extending
from sea
level to the
upper
atmosphere
Pressure – Altitude Dependence
Standard atmospheric pressure refers
to pressure at sea level
As the altitude increases, the pressure
decreases
Pressure – Altitude Dependence
Vacuum
P exerted by
Hg column
1 atm
Pressure
Barometer
Pressure of
atmosphere at
sea level will
hold a column of
mercury (Hg)
760 mm high
1 atm =
760 mm Hg =
101.3 kPa
Open end
Manometer
h
Gas
Column of mercury
to measure
pressure
h is how much
lower the pressure
is than outside
Manometer
Open end
h
Gas
Column of mercury
to measure
pressure
h is how much
higher the pressure
is than outside
Units of Pressure (1)
1 atmosphere = 1 atm
(average sea level pressure at 0C)
Column of mercury related units
• useful for mercury barometers and
manometers
1 atm = 760 mm Hg
1 mm Hg = 1 torr
1 atm = 760 torr
Units of Pressure (2)
Units directly in terms of pressure
1 atm = 101,325 Pascals
= 101.325 kPa
Pascal (Pa) = SI unit of pressure
• Pa = N/m2 (force of 1 Newton per m2)
1 atm = 14.7 psi (pounds force per in2)
• Common US engineering unit
• We will deal primarily with SI unit
Partial Pressures
Each gas in a mixture exerts pressure
independently of other gases present
Partial Pressure (PP)– portion of the
total pressure contributed by a single
gas
PP depends upon
• # moles, volume, temperature
PP independent of identity of gas
Dalton’s Law of Partial Pressures
Total pressure of gas mixture (Ptotal) =
sum of partial pressures (Pi)
Ptotal = P1 + P2 + … + Pn
+
1 mol He
P1
1 mol H2
P2
2 mol gas
Ptotal=P1+P2
Dalton’s Law of Partial Pressures
Water has a vapor pressure which
depends upon temperature [f(T)]
Water molecules have kinetic energy
and can escape from liquid
(see following slide)
Energy Distribution of Molecules in a Liquid
Minimum KE
needed for
vaporization
Kinetic Energy (KE)
Vapor Pressue
Kinetic energy (or velocity) distribution
of molecules means that a fraction of
molecules have sufficient energy to
overcome intermolecular forces and
escape the liquid
Origin of a Vapor Pressure
Open
Closed
H2O
Vapor
Liquid
H 2O
Saturated Vapor Pressure – H2O
Generating and Collecting a Gas
Decomposing KClO3(s) to produce O2(g)
Collecting a Gas Over Water
Dalton’s Law of Partial Pressures
If have a gas collected over water, total
pressure given by
Ptotal = Pgas + Pw
Pw = vapor pressure of water at a given
temperature
Dalton’s Law of Partial Pressures
Pw = vapor pressure of water at a given
temperature
• If T = 100C, then Pw=1 atm
What happens to PP of N2, O2 (air)
over boiling water at sea level?
Dalton’s Law: Example Problem 12.2
Mixture of O2, CO2 and N2 gases has total
pressure of 0.97 atm. PP of O2 if PP(CO2)
= 0.70 atm and PP(N2) = 0.12 atm?
Ptotal = P1 + P2 + … + Pn
Ptotal = PN2 + PCO2 + PO2
PO2 = Ptotal  PN2  PCO2
PO2 = 0.97 atm  0.12 atm  0.70 atm
PO2 = 0.15 atm
Practice (Partial Pressure, Units)
Problems 4 - 7, 409
Problems 45 - 46 page 434 (Dalton’s)
Problems 47- 49 page 434 (P units)
Chapter 12 – States of Matter
12.1
12.2
12.3
12.4
Gases
Forces of Attraction
Liquids and Solids
Phase Changes
Section 12.2 Forces of Attraction
Intermolecular forces—including
dispersion forces, dipole-dipole forces,
and hydrogen bonds—determine a
substance’s state at a given
temperature.
• Describe intramolecular forces.
• Compare and contrast intermolecular forces.
• Identify the intermolecular forces that will act on a
specific molecule.
• Relate the strength of intermolecular forces to
properties such as melting and boiling points.
Section 12.2 Forces of Attraction
Key Concepts
• Intramolecular (bonding) forces are stronger than
intermolecular forces.
• London dispersion (aka induced dipole – induced
dipole) forces are intermolecular forces between
temporary dipoles.
• London dispersion forces increase with molecular size
and with the amount of surface contact with
surrounding molecules.
• Dipole-dipole forces occur between polar molecules.
• For hydrogen bonding to occur, a hydrogen atom must
be bonded directly to F, N, or O.
• Hydrogen bonding is the strongest of the
intermolecular forces.
Molecular Forces Compared
Forces between molecules
are intermolecular forces
Bond is
intramolecular force
Intramolecular Forces (Bonds)
Intermolecular forces
Inside molecules, atoms are bonded to
each other
• Bonds are intramolecular forces
Intermolecular refers to forces between
molecules
• Hold molecules together in condensed
states (liquid, solid)
Intramolecular (Bonding) vs
Intermolecular Forces
Relative Strength of Forces
Strong (intramolecular)
• covalent bonding
• ionic bonding
Weak (intermolecular)
• Dipole dipole
• London dispersion forces
During phase changes the molecules stay
intact
• Energy used to overcome intermolecular
forces
Relative Strength of Forces
Intermolecular - between molecules
Intramolecular – hold atoms together
• 41 kJ vaporize 1 mole H2O (inter)
• 930 kJ to break all O-H bonds in 1
mole of water (intra)
BP, MP, VP measures of strength of
intermolecular forces
• Higher boiling point  stronger forces
Induced Dipoles - Example
d+
d-
H
H
d+
d
H
H
Induces
Instantaneous
Two
second
nonpolar
instantaneous
dipole
molecules
created
dipole
London Dispersion Force
Non-polar molecules also exert forces
on each other
• Otherwise, no solids or liquids
Electrons are not evenly distributed at
every instant in time
Have an instantaneous dipole
Induces a dipole in the atom next to it
Induced dipole- induced dipole
interaction
London Dispersion Force
Due to induced temporary dipoles
• Act between all molecules
• Only force between nonpolar molecules
and noble gas atoms
Strength depends primarily on molecular
size
• Very weak for small molecules
• Fairly strong for large ones
Molecular shape also plays a role
• Spherical vs elongated molecules
Dispersion Force & Molecular Size
Dispersion Force – Molecular Shape
n-Pentane C5H12
BP = 309.4 K
Neopentane C5H12
BP = 282.7 K
Dispersion Force & Molecular Shape
higher boiling point
weaker dispersion forces;
lower boiling point
Long skinny molecule
…
compact
molecule
Other Kinds of Induced Dipoles
Dispersion Forces (Ar, Cl2)
Ion - induced dipole
Dipole –
induced dipole
Polarity and Dipole Moment
+
Dipole
_
Dipole moment is a vector pointing from
center of - charge to center of + charge
Magnitude proportional to size of charges
and to separation distance
All polar covalent bonds have a dipole
moment
Dipole-Dipole Interactions
Dipole – Dipole Interactions
Orientation of Polar Molecules in a solid
Dipole-Dipole Interactions
The more polar the molecule, the
higher the boiling point
Ion – Dipole Interactions
Attractive forces between an ion and a polar
molecule
Ion-Dipole Interactions
Hydrogen Bonding
Especially strong dipole-dipole forces when
H is attached to F, O, or N
These three because• Have high electronegativity
• Are small enough to get close
Effects BP & other properties affected by
intermolecular forces
Hydrogen Bonding in Water
Each O has 2
lone pairs;
interacts with
two other H2Os
100
Boiling Points
H 2O
0ºC HF
Hydrogen bonding
NH3
H2Se
H2S
HCl
-100
PH3
SiH4
-200
CH4
AsH3
HBr
GeH4
H2Te
SbH3
HI
SnH4
Nonpolar (tetrahedral)
Properties of 3 Molecular Compounds
Polar,
multiple
H-bonds
Nonpolar
-164
Polar,
multiple
H-bonds
-33.4
Intermolecular Hydrogen Bonds
Intermolecular hydrogen bonds give
proteins their secondary shape, forcing
protein molecules into particular
orientations, like a folded sheet
Intramolecular Hydrogen Bonds
Intramolecular
hydrogen
bonds can
cause
proteins to
take a helical
shape
Hydrogen Bonding in Nylon
Hydrogen bonding helps make nylon strong
Inter- and Intramolecular Hydrogen
Bonding in Cellulose
Ion-Molecule & Intermolecular Forces
Summary
Hydrogen Bond
Ion Dipole
Dipole – Induced
Ion – Induced Dipole
Dipole
Dipole-Dipole
Dispersion
What type(s) of intermolecular
forces exist between each of the
following molecules?
HBr
HBr polar: dipole-dipole forces
Also dispersion forces
CH4
CH4 nonpolar: dispersion forces
SO2
S
Bent shape  polar: dipole-dipole + disp.
Chapter 12 – States of Matter
12.1
12.2
12.3
12.4
Gases
Forces of Attraction
Liquids and Solids
Phase Changes
Section 12.3 Liquids and Solids
The particles in solids and liquids have
a limited range of motion and are not
easily compressed.
• Contrast the arrangement of particles in liquids and
solids.
• Describe the factors that affect surface tension,
viscosity, capillary action, beading, and the formation
of a meniscus.
• Distinguish between adhesion and cohesion.
• Explain how the unit cell and crystal lattice are
related.
Section 12.3 Liquids and Solids
• Explain why ice floats in water but almost all other
solids sink in their liquids
• Describe the factors that are different in the 7 different
crystal categories.
• Draw the simple, body centered, and face centered
cubic unit cells.
• Identify the 5 types of crystalline solids, give an
example of each type and be able to rank the order of
properties such as melting point among the various
types.
• Distinguish among crystalline, amorphous, and
quasi-crystalline solids.
Section 12.3 Liquids and Solids
Key Concepts
• The kinetic-molecular theory explains the behavior of
solids and liquids.
• Intermolecular forces in liquids affect viscosity, surface
tension, cohesion, and adhesion.
• Crystalline solids can be classified by their shape and
composition.
Three States of Matter
Liquids – Density/Compressibility
At 25°C and 1 atm pressure, liquids
much denser than gases
• Liquid water 1250 x denser than vapor
• Both have ~ same kinetic energy
Liquids not very compressible
• High P needed to reduce V by a few %
Liquids – Fluidity
Diffusion of liquid molecules through a
liquid much slower than gas through
gas
Liquids in this sense are “less fluid”
than gases
Liquids – Viscosity
Measure of resistance of liquid to flow
High viscosity  high flow resistance
• Corn syrup higher viscosity than water
High viscosity  high intermolecular
forces
• Strong hydrogen bonding

Glycerol – 3 OH groups per molecule
• Long chains as opposed to compact

Oils – typically long carbon chains C18
Glycerol – Multiple Hydrogen Bonding
Viscosity and Temperature
Viscosity decreases with increasing
temperature
• Added kinetic energy makes it easier for
molecules to overcome intermolecular
forces
Viscosity and Temperature
Issue for motor oils
• Winter need low viscosity for flow
• Summer need higher viscosity to
provide adequate lubrication properties
Multi-grade oils have chemical
additives for automatic viscosity control
• Low T – compact, no strong effect
• High T – uncoil, tangle with oil, increase
viscosity
Liquids
Additional properties due to internal
attraction of molecules
• Surface tension
• Capillary action
• Beading
Stronger intermolecular forces cause
each of these to increase
Surface Tension
Definition: amount of
energy required to
stretch or increase
surface of liquid by a
unit area
Strong intermolecular
forces  high surface
tension
Molecules at
surface pulled
toward interior
Interior
Surface
molecules
area
pulled in all
minimized
directions
Surface Tension & Detergents (Surfactants)
Surfactants are
compounds that
lower surface
tension (typically of
water) – all
detergents contain
surfactants
Typical detergent
molecule: sodium
lauryl sulfate –
charged (polar)
head & nonpolar tail
Surface Tension & Detergents
Surfactant forms micelle or monolayer in water
Former good for removing oily dirt
Latter changes surface tension
Air
Monolayer
Bulk Water
Micelle
Surface Tension on a Needle
Needle floats, even
though density of
steel much higher
than density of water
Needle actually rests
in small depression
in liquid surface
Vertical components
of force balance
weight
Cohesive vs Adhesive Forces
Intermolecular forces are cohesive,
connecting like things
• Main cohesive force acting in liquid
water is hydrogen bond
Adhesive forces refer to forces between
the molecule and something else, most
often a solid surface
Liquids in Contact with Solid
Surface – Case 1
Adhesive forces
greater than
cohesive forces
Liquid clings to
walls of
container
Liquid “wets”
surface
Water
Glass
Liquids in Contact with Solid
Surface – Case 2
Cohesive
forces greater
than adhesive
forces
Liquid curves
downward
Liquid does not
“wet” surface
Mercury
Glass
Capillary Action
Liquids spontaneously rise in a narrow
tube
Glass is polar – attracts water molecule
Strong adhesive forces between water
and glass surface
Capillary Action / Meniscus
Cohesion Adhesion
 Adhesive Forces
Capillary Action
Result of surface
tension and adhesive
forces
Liquid rises when
adhesive forces greater
than cohesive forces
At point of contact
between liquid and
solid, upward forces are
as shown in diagram
Capillary Action
Cohesive forces
greater than
adhesive forces
Level of fluid in tube
will be below
surface of
surrounding fluid
Contact Angle 
Paraffin – not wetted
Left: Φ > 90°; cohesive forces > adhesive
forces
Right: Φ < 90°; adhesive forces > cohesive
forces
Beading
Liquid spreads;
adhesive forces
are comparable
in strength to
cohesive forces
Liquid “beads up.”
Which forces are
stronger, adhesive
or cohesive?
Beading
If polar substance
placed on non-polar
surface:
• There are cohesive,
• But no adhesive
forces
Beading - Mercury
Solids
KMT predicts that at same temperature,
solid particles have same amount of KE
as liquid
• Strong forces between particles
• Must be in motion – vibrations
Higher degree of order in solid than in
liquid
Solids are not “fluid”
(Exception – fluidized beds)
Solids - Density
Most solids more dense than their liquid
phase
• Typically 10% higher
• Water an exception – solid structure
less closely packed than liquid
Thermal expansion solid-liquid-gas
Density
Normally, density
(r) changes as
solid
liquid
gas
Temperature
Thermal Expansion of Water
1.0004
Density (kg/m3)
Density of ice less
than water at 0C
• Icebergs float
Density of water is
maximum at 4°C
• Nearly frozen
water floats to top
of lake and hence
freezes at surface
1.0002
1.0000
0
4
8
Temperature (°C)
Hexagonal
Structure
of
Ordinary
Ice
Holes in structure gives ice
a density less than water
Crystalline Solids
Atoms regularly arranged in
crystal lattice
Unit cell - smallest
arrangement of connected
points that can be repeated
in 3 dimensions to form
lattice (smallest arrangement
of atoms in a crystal lattice
that has same symmetry as
whole crystal)
Unit Cell in Crystal Lattice
Two Different Unit Cells in NaCl
Crystal Categories (Types of Unit Cells)
7 categories based on overall shape
Difference based on relative lengths of
sides and angles between sides of unit cell
7 Crystal Categories - Summary
4 Types of Unit Cell
(Orthorhombic Crystal Class – only one with all 4)
Body
Primitive
Centered
Face
Side
Centered
Centered
Cubic Lattice Types (Unit Cells)
Simple
Cubic
Body-Centered
Cubic
Face-Centered
Cubic
Cubic Unit Cells
Simple Cubic
Body-Centered Face-Centered
NaCl: Anions in FCC Unit Cell
Only ions connected by lines actually touch
Cl- in contact with center Na+ - 6 coordinate
Unit Cell in NaCl
Slice Through Atoms
Cubic
Unit
Cells
Types of Crystalline Solids
(Table 12.5, p 422)
ions
Types of Crystalline Solids (Tab. 12.5)
Atomic (noble gases only)
• Ar
Molecular
• Ice (individual H2O molecules)
Covalent network (limited category)
• Diamond, quartz (SiO2), SiC,BN
Ionic
• NaCl
Metallic
• Cu
(ions in lattice – book incorrect p. 422)
Molecular Solids
Held together by
• Dispersion forces

Larger for large molecules
• Dipole-Dipole forces
• Hydrogen Bonds
Physical Properties – Molecular Solids
Melting/boiling points low compared to
ionic substances
Many are gases or volatile liquids
• O2, CO2, H2S
Form relatively soft solids
• Wax (typically a large hydrocarbon)
Poor conductors of heat and electricity
• No ions or delocalized electrons
Covalent Network Solids
Interconnected covalent
bonds – essentially one
gigantic molecule
Brittle, hard,
nonconductors
Diamond
SiO2 (quartz)
SiC
BN
Crystal Structure of Diamond
3-dimensional
network extremely strong,
rigid
What kind of
forces must be
overcome to
melt diamond?
Crystal Structure of Graphite
Hexagons of
sp2hybridized
carbon
(graphene)
Relatively
weak forces
between
layers
Covalent Network Solids
Network for graphite in-plane only
Diamond
Graphite
Substance MP (C) Substance MP (C)
Diamond
3550
Fe
1536
W
3422
Fe2O3
1460
MgO
2900
NaCovalent
890
2SO4
Network
SiC
2700
NaCl Solids
808
Transition
CaO
2600
CaCl2
782
Al2O3
2040
MgCl2
714
Metals
Representative
ZnO
1975
Al Metal 660
Metal
Oxides 275
Alkali
Li2O
1700
ZnCl
2
Metal
SiO
1700
K
63
2
Amorphous Solids
Non-crystalline solids
• No regular arrangement in space
Formed when material cooled quickly
• Atoms don’t have time to get into
ordered arrangement
Glass, rubber, plastics, obsidian (lava)
Quasicrystals (NIB)
Although textbook lists only crystalline and
amorphous solids, a third arrangement exists –
quasicrystalline solids
Atoms are arranged in regular patterns like in a
normal crystal but the pattern never repeats itself
Quasiperiodic crystal (quasicrystal) - structure
that is ordered but not periodic
Quasicrystalline pattern can continuously fill all
available space, but it lacks translational
symmetry (shifted copy will never match exactly
with its original)
http://en.wikipedia.org/wiki/Quasicrystal#cite_note-0
Quasicrystal
Quasicrystals exist universally in many metallic
alloys and some polymers.
Quasicrystals are found most often in aluminium
alloys but numerous other compositions are also
known.
http://en.wikipedia.org/wiki/File:Penrose_Tiling_%28Rhombi%29.svg
Penrose Tiling
Illustrates Type of Arrangement in a
Quasicrystal (regular but non repeating)
http://www.livescience.com/16393-nobel-prize-chemistry-quasicrystals.html
Silver Aluminum Quasicrystal
Regular patterns that follow mathematical rules but don't
repeat themselves. CREDIT: AMES Lab, U.S. DOE
Practice
Problems 17-19, 21-23 page 403
Problems 43-51 page 414
Chapter 12 – States of Matter
12.1
12.2
12.3
12.4
Gases
Forces of Attraction
Liquids and Solids
Phase Changes
Section 12.4 Phase Changes
Matter changes phase when energy is
added or removed.
• Explain how the addition and removal of energy can
cause a phase change.
• Name the six types of phase changes and specify if the
change is endo- or exothermic.
• Explain the terms superheated and supercooled as
applied to liquids
• Interpret vapor pressure and its dependence on
temperature in terms of kinetic molecular theory and
the strength of intermolecular forces.
Section 12.4 Phase Changes
• Interpret a phase diagram including the special points
(triple and critical).
• Explain how the critical temperature and pressure
relate to the ability to liquefy a substance.
Section 12.4 Phase Changes
Key Concepts
• States of a substance are referred to as phases when
they coexist as physically distinct parts of a mixture.
• Energy changes occur during phase changes.
• Phase diagrams show how different temperatures and
pressures affect the phase of a substance.
Phase Changes and Energy
Heat is transfer of energy from higher
temperature (T2) object to lower
temperature (T1) object
In phase changes, energy used to
overcome (disrupt) intermolecular
forces between molecules
For phase changes at equilibrium,
T2 = T1 (both phases at same T)
Phase Changes and Energy
Require energy (endothermic process)
• Vaporization
• Sublimation
• Melting
Release energy (exothermic process)
• Condensation
• Deposition
• Freezing
Melting
Amount of energy required to melt
substance depends on strength of
intermolecular forces
Melting point (MP) - temperature at
which solid and liquid phases coexist
For amorphous solids, exact melting
point more difficult to determine
Supercooled Liquids
Under certain conditions (generally,
pure liquids that lack nucleation sites),
liquids can be cooled below their
freezing points and still remain liquid
Called a supercooled liquid
Shock or introduction of suitable nuclei
can cause rapid freezing
Superheated Liquids
Under certain conditions (generally,
pure liquids that lack nucleation sites –
smooth surface cups), liquids can be
heated above their boiling points and
still remain liquid
Called a superheated liquid
Shock or introduction of suitable nuclei
can cause rapid boiling
Types of Crystalline Solids
(Table 12.5, p 422)
ions
Melting & Types of Solids (Tbl 12.5)
Ionic, covalent network, and many
metallic solids generally have high MPs
– forces involved are bonding forces
Atomic solids low MPs
Molecular solids low to moderate MPs
• H-bonding solids requires more energy
than most polar solids because of
stronger dipole-dipole forces
Vaporization
Process by which liquid changes to gas
(vapor)
Kinetic energy (or velocity) distribution
means that a fraction of molecules
have sufficient energy to overcome
intermolecular forces and escape the
liquid – see following slide
Energy Distribution of Molecules in a Liquid
Minimum KE
needed for
vaporization
Kinetic Energy (KE)
Vaporization
Vaporization occurs at surface – fewer
forces acting upon a surface molecule
Called evaporation
Those molecules that do escape exert
a gas pressure (partial pressure) that is
equal to the vapor pressure
• True even at temperatures near the
freezing point
Vapor Pressure
Liquid in Closed Container
Vapor Pressure – Approach to Equilbirium
Rate
Rate of evaporation
Rate of Equal rates –
equilibrium
conden
sation vapor pressure
attained
Time
Measuring Vapor Pressure
Before Evaporation
At Equilibrium
Vapor Pressure – Differing Liquids
0 torr
Vacu
um
24 torr
H2O
65 torr 545 torr
Etha
nol
Diethyl
ether
Vapor Pressure
Fraction of molecules having sufficient
energy to overcome intermolecular
forces at 2 different temperatures
Vapor Pressure
Effect of Temperature
Vapor Pressure – T Dependence
Boiling
Temperature at which vapor pressure =
atmospheric pressure (normal BP)
Can superheat if no nuclei present
Boiling (as with melting) requires
energy
Temperature of liquid same as
temperature of vapor produced during
boiling process
Sublimation
Solid changes directly to gas without
first becoming liquid
• Iodine
• Carbon dioxide
• Napthalene (moth balls)
• Solid air fresheners
• Ice at low pressures (freeze drying)
• Water in your freezer
Energy-Releasing Transformations
Condensation (opposite of evaporation)
Deposition (opposite of sublimation)
• High altitude snowflake formation
Freezing (opposite of melting)
Phase Diagrams
Pressure vs Temperature plots with
phases identified
Lines on phase diagram indicate where
phase changes can occur
Can find MP, BP, VP as f(T), etc
Triple point – all 3 phases
simultaneously present
• Multiple solid phases  multiple triple pts.
Phase Diagram
Pressure
Liquid
Solid
Gas
Temperature
Phase Diagram - Critical Point
High pressures can liquify a gas
Critical Temperature – temperature
beyond which cannot liquify a vapor
regardless of the pressure applied
• Liquid and gas phases indistinguishable
Critical Pressure – the vapor pressure
at the critical temperature
Critical Point – the point defined by the
critical temperature and pressure
Phase Diagram - Water
Pressure (atm)
X
Temperature (C)
Pressure (atm)
Phase Diagram – Carbon Dioxide
Temperature (C)
Phase Diagram - Sulfur
With 2 separate solid phases (monoclinic,
rhombic), have 3 triple points
The Critical Point
At higher temp At room temp more vapor & its
relatively little vapor
density increases …
at low density
… while density of liquid
decreases; molecular
motion increases
At Tc,
densities of
liquid &
vapor equal
-single
phase
Phase Diagram - Critical Point
Critical Temperatures and Pressures
These 4 gases can’t be liquefied at room temperature,
regardless of applied pressure. Why not?
Practice (Phase Changes)
Problems 27-33 page 430
Problems 75-82 page 435
End of Chapter