Liquids and Solids: AP Chemistry Unit 10
Sections
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Intermolecular Forces
Liquid state
Solid Structures
Metal Structures
Carbon and Silicon Networks
Molecular Solids
Ionic Solids
Vapor pressure and State Change
Phase Diagrams
States of Matter
When considering the three states of matter, properties of gases are strikingly
different than solids and liquids. Liquids and solids share many similar
characteristics
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compressibility
density
intermolecular forces
H2O(s)
 H2O(l)
ΔH°fus = 6.02 kj/mol
H2O(l)

ΔH°vap = 40.7 kj/mol
H2O(g)
Water densities:
25°C and 1atm
.99707g/cm3
25°C and 1065 atm
1.046g/cm3
400°C and 1atm
3.26x10-4 g/cm3
400°C and 242 atm
.157g/cm3
1
Intermolecular Forces 10.1

Electrons shared within the molecule are called intramolecular bonding.

In the condensed states of matter the attraction between molecules are
called intermolecular forces.
It is important to realize that when a molecule changes state, the molecule stays
intact. The changes in state are due to the change in forces surrounding the
molecule not from changes within the molecule.
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
40.7kj needed to vaporize water
934kj to break the O-H bond
Dipole–Dipole Forces
Dipole-dipole forces occur when polar molecule (molecules with dipole moments)
electrostatically attract each other by lining up the positive and negative ends of the
dipoles.

Dipole-dipole forces are about 1% as strong as a covalent or ionic bond and
rapidly become weaker when distances between the dipoles increases. The
distances in a gas make these attractions relatively unimportant

In a condensed state, molecules line up dipoles to minimize repulsions and
maximize attractions.
2
Dipole-Dipole Forces continued…
Some dipole-dipole forces are unusually strong. These usually form between H and
another very electronegative atom.

These are stronger due to the high polarity of the bond and the closeness of
the dipoles between the atoms.
o These strong attractions have a strong impact on melting points and
boiling points.
Boiling Points of Covalent Hydrides
Hydrogen bonds
Hydrogen bonds are the strongest in the smallest and lightest of the covalent
molecules. This is primarily due to two factors:

large difference in electronegativities

small size of the atoms allows for close dipole interactions.
3
Hydrogen bonds continued….
Hydrogen Bonds and Organics
Methanol (CH3OH) and ethanol (CH3CH2OH) have much higher boiling points than
would be expected from their molar masses because of the O-H bonds that produce
hydrogen bonding.
London Dispersion Forces
Even without dipoles, molecules exert forces on each other.

The forces that exist among noble gas atoms and nonpolar molecules are
called London dispersion forces.
Usually it is assumed that electron dispersion is uniform throughout the molecule,
but this is not always the case.

Since the movements of the electrons around the nucleus are somewhat
random, a momentary nonsymmetrical electron distribution can develop
that creates a temporary dipolar arrangement of charge.

This temporary change in polarity can, in turn, temporarily change the
distribution of the neighboring molecule.

This phenomenon leads to an inter-atomic attraction that is relatively weak
and short-lived, but can be significant in larger atoms at lower temperatures.
o larger atoms have more electrons and increases the probability of a
temporary dipole.
4
London Dispersion Forces continued…
Polarizability is the ease at which an electron cloud can be distorted into a
temporary dipole.

large atoms have a larger polarizability than smaller atoms

This also applies to molecules like H2, CH4, CCl4 and CO2; smaller molecules,
but nonpolar.
The Liquid State 10.2
Liquid Characteristics
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lack of rigidity
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low compressibility
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high density
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rounded droplets
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capillary action
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viscosity
Rounded Droplets

Occur due to the intermolecular forces of the liquid. The liquid molecules are
subject to attraction from the side and from below, so liquid tends to form a
shape with the minimum surface area – sphere.

The resistance of a liquid to increase surface area is from the energy that it
takes to overcome intermolecular forces. This resistance is called surface
tension.

Molecules that are polar and have stronger intermolecular forces have
stronger surface tensions.
5
Capillary Action
Capillary action is the spontaneous rising of a liquid in a narrow tube. This action is
due to two forces
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cohesive forces- the intermolecular forces among the molecules.

adhesive forces – the attractive forces between the liquid and the container.
Adhesive forces
Adhesive forces happen when bonds within the container have polar bonds

For example: glass has O atoms that carry a partial negative charge that
attracts the partial positive charge of the hydrogen in water. This balance
between the strong cohesive forces and the strong adhesive forces produce a
meniscus.

A nonpolar substance, such as mercury, has a convex meniscus because the
cohesive forces are stronger than the adhesive forces.
Meniscus: Water vs. Mercury
Viscosity
Viscosity is a fluids resistance to flow.

liquids with strong cohesive forces tend to be highly viscous.

Example: glycerol is highly viscous because of its ability to create hydrogen
bonds.
6
Viscosity continued…

Molecular complexity also can affect viscosity because they can become
entangled in each other.

Example: Gasoline has carbon chains from 3-8C long and is nonviscous.
Grease is 20-25C long and is very viscous.
Introduction to Structures and Types of Solids 10.3
Types of Solids
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Crystalline solids
Amorphous solids
Crystalline Solids
Crystalline solids have a regular arrangement of components at a microscopic level
and produce beautiful, characteristic shapes of crystals:
The positions of components are usually represented by a lattice.
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lattice is a three dimensional system of units repeating in a pattern. The
smallest repeating unit of the lattice is called the unit cell.
Amorphous Solids
Amorphous solids have considerable disorder in their structures.

Example: Common glass looks like a solution frozen in place. It has a rigid
shape but a great deal of disorder within its structure.
7
X-ray Analysis of Solids
The structures of crystalline solids are commonly determine by X-ray diffraction.

This type of diffraction occurs when beams of light are scattered as they go
through spaces between substances. Light scatters when the size of the
spaces are similar to the wavelength of light.
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A single wavelength is directed at the crystal and a diffraction pattern is
obtained. The diffraction pattern is a series of light and dark areas on a
photographic plate from constructive and destructive interference from
waves of light.

The diffraction pattern can then be used to determine the interatomic
spacings.

A diffractometer is a computer-controlled instrument used for carrying out
the X-ray analysis of crystals
o It rotates the crystal with respect to the X-ray beam and collects the
data produced by the scattering. The techniques have been refined to
the point that very complex structures can be determined, such as
large biological enzymes.
The Bragg equation combines trigonometry and physics to determine the atomic
spaces between crystals:

nλ = 2d sin θ

d is the distance between atoms and θ is the angle of incidence and reflection
of the light. n is an integer, most commonly 1. (n is usually given)
8
X-ray Analysis of Solids continued….
Example Problem: X-rays of wavelength 1.54 Â were used to analyze an
aluminum crystal. A reflection was produced at θ = 19.3°. Assuming n=1,
calculate the distance d between the planes of atoms producing this
reflection
Types of Solids
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Ionic solids
o ionic solids are made of ions

Molecular solids
o Molecular solids have small units of covalently bonded molecules.

Atomic solids
o Atomic solids are made of elements such as carbon (graphite,
diamond and the fullerenes), boron, silicon, and all metals.
Atomic Solids
Atomic solids are broken down into subgroups depending on the bond that exists in
the solid:

Metallic solids
o Has delocalized nondirectional covalent bonding.

Network solids
o atoms bond with strong directional covalent bonding that lead to
giant molecules and networks
9
Atomic Solids continued….
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Group 8A solids
o noble gases are attracted to each other with London dispersion forces.
Classification of Solids
Structure and Bonding in Metals 10.4
Metal Characteristics
Most of the properties that we see in metals is due to the nondirectional covalent
bonding found in metal crystals.
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High thermal conductivity
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Electrical conductivity
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Malleability
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Ductility
Metallic Crystals
Metallic crystals can be pictured as containing spherical atoms packed together that
can be bonded to each other equally in all directions.
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This arrangement is called closest packing.
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Closest Packing

The spheres pack in layers. Each sphere is surrounded by six others. These
layers do not lie directly over those in the first layer, instead they fill the
indentations of the layer below. The third layer is in the same position as the
first. This is called aba arrangement.
o The aba arrangement has the hexagonal unit cell and the resulting
structure is called the hexagonal closest packed (hcp) structure.

The abc arrangement has a face-centered cubic unit cell and the resulting
structure is called the cubic closest packed (ccp) structure. This has a
repeating vertical placement every fourth layer.
Closest Packing: Hexagonal
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Closest Packing: Cubic
Knowing the net number of atoms in a particular unit cell is important for many
applications involving solids.
Example: A face centered cube (unit cell) is defined by the centers of the
spheres on the cube’s corners. Therefore 8 cubes share a given corner
sphere, so 1/8 of this sphere lies inside the unit cell. (8 corners x 1/8 sphere
= 1sphere). The sphere at the center of each face is shared by two cubes. (6
faces x ½ sphere = 3 spheres). The total number of spheres for a face
centered cube is 4.
Face – Centered Cubic Unit Cell
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Cubic Substances
Metals that form cubic closest packed solids are:
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aluminum
iron
copper
cobalt
nickel
Hexagonal Substances
Metals that form hexagonal closest packed solids are:
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magnesium
zinc
Other Metal Solids
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Calcium and certain other metals can crystallize in either cubic or hexagonal
solids.
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Some metals, including many alkali metals, have structures that are
characterized by a body-centered cubic (bcc) unit cell. In this structure, each
sphere has 8 neighbors.
Example Problem: Silver crystallizes in a cubic closest packed structure. The
radius of a silver atom is 144pm. Calculate the density of solid silver?
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Bonding Models for Metals
In order to determine bonding for metals, one must account for the typical
properties: durable, high melting point, malleable, ductile, and efficient in uniform
conduction of heat and electricity in all directions.

These characteristics indicate that the bonds are strong and nondirectional.
In other words, it is not easy to separate metal atoms but easy to move them.
Electron Sea Model
Metal cations ‘swim’ in a sea of valence electrons that are mobile and shared.

This accounts for conduction and malleability and ductility.
Band Model (MO Model)
In this model, the electrons are assumed to travel around the metal crystal in
molecular orbitals formed from the valence atomic orbitals of the metal atoms.

When metals atoms interact, the large number of resulting molecular orbitals
become more closely spaced and finally form a virtual continuum of levels,
called bands.
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Band Model continued…

The electrons in partially filled MO’s are mobile. These conduction electrons
are free to travel throughout the metal crystal. The MO occupied by these
conducting electrons are called conduction bands.
Metal Alloys
An alloy is best defined as a substance that contains a mixture of elements and has
metallic properties. There are two types of alloys:

Substitutional alloy– some of the host metal atoms are replaced by other
metal atoms of similar size.
15

Interstitial alloy – is formed when some of the interstices (holes) in the
closest packed lattice are occupied by smaller atoms.
Substitutional Alloy
Example: brass: 1/3 of copper metal atoms are replaced by zinc atoms
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Sterling silver- 93% silver and 7% copper.
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Pewter- 85% tin, 7% copper, 6% bismuth and 2% antimony.

Plumbers solder – Plumbers solder – 95% tin and 5% antimony
Interstitial Alloy
Example: Steel contains carbon atoms in the holes of an iron crystal. The
presence of the interstitial atoms changes the properties of the host metal.
Iron is relatively soft, ductile and malleable, but when carbon (which forms
directional bonds), is introduced into the crystal, it makes the iron bonds
stronger and less ductile.
The amount of carbon directly affects the properties of steel:
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Mild steels- contains less than .2% carbon: nails chains and cables.
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Medium steels- contain .2-.6% carbon: rails and structural steel
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High-carbon steel – .6-1.5% carbon: springs, tools and cutlery.
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Mixed Alloys
Some steels contain elements in addition to iron and carbon. These are called
alloy steels and are viewed as being mixed interstitial and substitutional alloys.

Bicycle frames are usually constructed from a wide variety of alloy steels.
Carbon and Silicon Network Atomic Solids 10.5
Network Solids
Many atomic solids contain strong directional covalent bonds to form a solid that
might be viewed as a “giant molecule.” These materials are typically brittle and do
not efficiently conduct heat and electricity. Two examples of these network solids
are carbon and silicon.
Carbon
Two most common forms of carbon are diamond and graphite. They are typical
network solids.

Diamond is the hardest naturally occurring substance.

Graphite is slippery, black and a conductor.
Diamond

Each carbon is surrounded by a tetrahedral arrangement of other carbon
atoms to form a large molecule. Diamond is an insulator not a conductor.
Each carbon is sp3 hybridized with localized bonding and therefore does not
conduct.

Diamonds are often used for industrial cutting implements.

The application of 150,000 atm at 2800°C can break graphite bonds and
rearrangement into a diamond structure.
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Graphite


The structure of graphite is based on layers of carbon atoms arranged in
fused 6 C rings. The unhybridized p orbitals allow for delocalized electrons
and therefore conductivity.
Graphite is used as a industrial lubricant. Because graphite has strong bonds
within the layers and weak bonding between the layer, the layers slide past
one another readily.
Carbon: Graphite layers
Silicon
Silicon is an important constituent of the compounds that make up the earth’s crust.
Silicon is to geology what carbon is to biology and is fundamental to most rocks,
sands and soils found in the earth’s crust.

Carbon compounds typically have long strings of C-C bonds

Silicon compounds typically involve chains of Si-O bonds.
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Silica

The fundamental silicon-oxygen compound is silica, which has the empirical
formula SiO2. The structure that is formed is based on a network of SiO4
tetrahedra with shared oxygen atoms rather than smaller SiO2 units.
Silica continued…
When silica is heated above its melting point (1600°c) and cooled rapidly, an
amorphous solid called glass results. Glass has a lot of disorder as opposed to
the crystalline nature of quartz. Glass, also homogeneous, more closely
resembles a very viscous solution than it does a crystalline solid.
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Glass
The properties of glass can vary greatly depending on the additives.

Common glass results when substances like Na2CO3 are added to the silica
melt.

B2O3 produces borosilicate glass which does not expand and contract during
large temperature changes. (Pyrex)

K2O produces especially hard glass that can be ground into shapes for lenses
and contacts.
Silicates
Compounds closely related to silica and found in most rocks, soils and clays
are the silicates. Like silica, the silicates are based on interconnected SiO4
tetrahedra, but instead of a O/Si ratio of 2:1, the ratio is typically higher. This
higher ratio tends to make silicon-oxygen anions.
20
Ceramics are typically made from clays (which contain silicates) and hardened
by firing at high temperatures. They tend to be strong, brittle and heat and
chemical resistant.

Ceramic is heterogeneous and contain two phases: minute crystals of
silicates that are suspended in a glassy cement.
Clays
Clay comes from the weathering of feldspar, an Aluminosilicate
(Na2O/K2OŸAl2O3Ÿ6SiO2). This weathering produces kaolinite, that consists
of tiny thin platelets of Al2Si2O5(OH)4. When dry these platelets cling
together and lock into place; when wet they can slide over one another.
During firing, these platelets bind and form a glass.
Ceramics
Ceramics constitute one of the most important classes of ‘high-tech”
materials. Their stability at high temperatures and resistance to corrosion,
make them an obvious choice for constructing jet and car engines.
Organoceramics are taking form by the addition of organic polymers to
ceramics. This reduces some of the brittle nature of ceramics and allows
them to be used for things such as flexible superconducting wire,
microelectronic devices, prosthetic devices and artificial bones.
Semiconductors
Elemental silicon has the same structure as diamond. The structure is different in
that the energy gap between filled and empty MO’s is not as large and electrons can
delocalize and make silicon a semi-conductor. At higher temperatures, more
electrons get excited in the conduction bands and the conductivity of silicon
increases.
N-type Semiconductor

When small fraction of silicon atoms are replaced by arsenic atoms (one
more valence electron), extra electrons become available for conduction and
produce an n-type semi-conductor. These can conduct an electric current.
21
P-type Semiconductor

When small fraction of silicon atoms are replaced by boron atoms (one less
valence electron), an electron ‘vacancy’ is made. As electrons move, the fill
the ‘hole’ and make a new one. This movement of electrons can therefore
carry a current. This type of conductor (less electrons) is called a p-type
semiconductor.
Energy Level Diagrams for N-type and P-type Semiconductors.
P-N Junction
Most important applications of semiconductors involve connection of a p-type and
an n-type to form a p-n junction.

The red dots represent excess electrons in the n-type semiconductor and the
white circles represent holes (electron vacancies.

At the junction a small number of electrons migrate from the n-type region
into the p-type region. The effect of these migrations is to place a negative
charge on the p-type region and a positive charge on the n-type region.

This charge buildup, called the contact potential or junction potential,
prevents further migration of electrons. This transfer of electrons is
therefore a ‘one-way’ transfer and under an external battery source will
allow flow of electrons from the n to the p type regions.

When current is opposed it is said to be under reverse bias. When current
flows easily, the junction is said to be under forward bias.

A p-n junction is a good rectifier, a device that produces a pulsating direct
current from an alternating current.
22

When placed in a circuit where the current is constantly reversing, a p-n
junction only transmits current under forward bias. Radios, computers and
other electronic devices all use this rectifiers. This p-n junction
revolutionized electronics.
Molecular Solids 10.6
Network Solids
Sometimes network solids can be considered to be one giant molecule or have large
discrete molecular units in a lattice-type position. These molecules have strong
bonds within the molecules but relatively weak between the molecules.

Ice, dry ice (solid carbon dioxide), Sulfur (S8), Phosphorus (P4)
23
Network Solids continued…

When molecules do have dipole moments, their intermolecular forces are
significantly greater, especially when hydrogen bonding is possible.

Water not only has polar bonds, a dipole moment, has hydrogen bonds, but it
also can have a total of four hydrogens associated with every oxygen atom.
Ionic Solids 10.7
Ionic Solids
Ionic solids are stable, high melting substances held together by the strong
electrostatic forces that exist between oppositely charged ions.

Most binary ionic solids can be explained by the closest packing of spheres.
Typically the larger ions, usually anions, are packed in one of the closest
packed arrangements (hcp and ccp).

The smaller cations fit into the holes among the closest packed anions. This
packing maximizes the electrostatic attractions among oppositely charged
ions and minimizes the repulsion of like charges.
There are three types of holes in closest packed structures:
1. Trigonal holes are formed by three spheres in the same layer
2. Tetrahedral holes are formed when a sphere sits in the dimple of three
spheres in an adjacent layer.
3. Octahedral holes are formed between two sets of three spheres in
adjoining layers of the closest packed structures.
24
Ionic Solids continued...

The holes increase in size in the order:
trigonal < tetrahedral < octahedral

The trigonal holes are so small that they are never occupied in binary ionic
compounds. Tetrahedral and octahedral holes are occupied if the relative
sizes of the ions allow.
Example: Zinc Sulfide (ZnS) creates a ccp structure. The Zn2+ has a radius of
70pm and the S2- ion has an ionic radius of 180pm. There are 4 spheres
(atoms/anions) in a face-centered cubic unit cell and 8 tetrahedral holes. So
only half of the holes in the ccp unit are filled with cations.
Example: Sodium chloride can be described in terms of a ccp structure. Na+
resides in octahedral holes. The locations of the octahedral holes in the facecentered cubic unit is marked by X. The number of spheres (anions) in the
structure is the same number of octahedral holes. Since NaCl is a 1:1 binary
compound. All octahedral holes are used.
Example: Determine the net number of Na+ and Cl- ions in the sodium chloride
unit cell.
Example: Classify each of the following substances according to the type of solid
it forms:




gold
carbon dioxide
lithium fluoride
krypton
25
Types and Properties of Solids
Vapor Pressure and Changes of State 10.8
Vaporization
Vaporization, or evaporation, is the process of liquid molecules escaping the liquid’s
surface and forming a gas.

Vaporization is endothermic because energy is required to overcome the
relatively strong intermolecular forces in the liquid.
o Water has strong intermolecular forces and this increases the energy
required to vaporize. Also making it a great coolant.

The energy required to vaporize 1 mole of liquid at 1 atm is called the heat of
vaporization or the enthalpy of vaporization.
o The symbol for this is ΔHvap.
Condensation is the process by which vapor molecules re-form a liquid.

The evaporation process occurs at a constant rate at a given temperature,
and once an equilibrium has been reached, the rate of condensation will
equal the rate of evaporation.
26
Vapor Pressure continued…

Molecules in a given system are constantly escaping from and entering the
liquid at high rate. However, there is not net change because the two
opposite processes just balance each other.
Rates of Condensation and Evaporation
Vapor Pressure
The pressure of the vapor present at equilibrium is called the equilibrium vapor
pressure, or more commonly, the vapor pressure of the liquid.

A simple barometer can measure the vapor pressure of a liquid.

Liquid is injected at the bottom of the tube of mercury and floats to the
surface. A portion of the liquid evaporates at the top of the column,
producing a vapor whose pressure pushes some mercury out of the tube.
27
Vapor Pressure …
When the system reaches equilibrium, the vapor pressure can be determined from
the change in the height of the mercury column

Patmosphere = Pvapor + PHg column

Pvapor = Patmosphere - Phg column
The vapor pressure of liquids vary widely. Liquids with high vapor pressures are
said to be volatile. They evaporate rapidly in an open dish.

The vapor pressure of a liquid is principally determined by the size of the
intermolecular forces in the liquid.
o Liquids with strong molecular forces have relatively low vapor
pressures because it takes so much energy for the molecules to
escape.
o In general, substances with large molar masses have relatively low
vapor pressures due of large dispersion forces.

Vapor pressure increases significantly with temperature.
o In order to break intermolecular forces, a sufficient amount of kinetic
energy is needed.
o As temperature of the liquid increases, so does kinetic energy of the
liquid.
28
Vapor Pressure of Water

The nature of the temperature, vapor pressure relationship is quadratic. Pvap
vs. 1/T (Kelvin) gives a direct relationship.
Vapor pressure equation:

R is the universal gas constant (8.3145 J/K), and C is a constant characteristic
of a given liquid (y=intercept). In means natural logarithm.
Example: Using the graph, determine whether water or diethyl ether has the
larger enthalpy of vaporization.
29
Example: The vapor pressure of water at 25°C is 23.8 torr, and the heat of
vaporization of water at 25°C is 43.9 kj/mol. Calculate the vapor pressure of
water at 50°C.
Sublimation
Like liquids, solids have vapor pressures. When a solid sublimes, it goes directly
from the solid to the gaseous state without passing through the liquid state.

Example: Dry Ice.
Changes of State
Typically when a solid is heated, it will form a liquid and then boil to form a vapor.
This process can be represented by a heating curve.

Temperature vs. time when energy is added a constant rate.
Heating Curve
30
Heating Curve
The plateaus in the heating curve represent the positions of phase change.

At the melting point, the temperature remains constant until the solid has
completely changed to liquid.

At the boiling point, the temperature remains constant as the added energy is
used to vaporize the liquid

The energy change that occurs at the melting point when a solid melts is
called heat of fusion or enthalpy of fusion.
o Note that changes of state are physical changes, although
intermolecular forces have been broken, no chemical bonds have been
broken.
Enthalpy of Fusions and Melting Points
Melting and Boiling
The melting and boiling points for a substance are determined by the vapor
pressure of the solid and liquid states.

At 0°C the vapor pressure of ice is less than that of liquid. Vapor pressure of
ice increases more rapidly than water.

A point is reached when the liquid and solids have identical vapor pressures. This is
the melting point.
31
Melting and Boiling…
Freezing Point

At a temperature at which the vapor pressure of the solid is greater than that
of the liquid, the solid would sublime and the vapor would be added to the
water.

At a temperature at which the vapor pressure of the solid is less than that of
the liquid, the liquid would evaporate and the vapor would be added to the
ice.

At a temperature at which the vapor pressures of the solid and liquid are
identical, the vapor is in equilibrium. This is the freezing point of water.
32
Melting and Boiling…


The normal melting point of a liquid is the temperature at which the solid and liquid
states have the same vapor pressure under conditions where the total pressure is 1
atmosphere.
The normal boiling point of a liquid is the temperature at which the vapor
pressure of the liquid is exactly 1 atmosphere.
Supercooled and Superheated
Changes of state do not always occur exactly at the boiling point or melting point.

Water can be supercooled below 0°C at 1 atm and remain in the liquid state.
At some point the correct ordering of molecules occurs and ice forms,
releasing energy in the exothermic process and bringing the temperature
back up to the melting point.

A liquid can also be superheated, or raised to temperatures above its boiling
point, especially if it is heated rapidly. Boiling requires high-energy
molecules to gather in the same vicinity for bubble formation. This may not
happen at the boiling point.

Once a bubble does form, when a liquid is superheated, its internal pressure
is greater than the atmospheric pressure. This bubble can burst before rising
to the surface, blowing the surrounding liquid out of the container. This is
called bumping and is a common experimental problem.
o Boiling chips are often added to prevent bumping. These are bits of
porous ceramic material containing trapped air that escapes on
heating, forming tiny bubbles that act as ‘starters’ for the vapor
bubble formation. This allows for smooth onset of boiling.
33
Phase Diagrams
Phase Diagrams
A phase diagram is a convenient way of representing the phases of a substance as a
function of temperature and pressure. It shows which state exists at a given
temperature and pressure.

Conditions for these phase diagrams are assumed to be a closed system and
is plotted as temperature vs pressure.

The solid/liquid boundary has a negative slope.
o Melting point of ice decreases as external pressure increases.

This is different for most substances other than water because the density of
ice is less than that of liquid water at the melting point.
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Example 1: Pressure is 1 atm. Water moves through the changes of state
according to the vapor pressure at the corresponding temperatures.
Example 2: Pressure is 2 torr. Water will sublime at -10°C. This is when the
vapor pressure of the ice is equal to the external pressure of 2 torr. Vapor
pressure of liquid water is always greater than 2 torr and therefore will not
form.
Example 3: Pressure is 4.58 torr. When temperature reaches .01°C
(273.16K), water reaches the triple point. Solid and liquid water have
identical vapor pressures and all three states of water exist. This is the only
condition in a closed system that allows this.
Example 4: Pressure is 225 atm. Liquid water can be present at this
temperature because of the high external pressure. As temperature
increases, liquid gradually turns to vapor, but goes through a ‘fluid’ region.

The fluid region is neither true liquid or vapor. This unusual behavior
occurs because the conditions are beyond the critical point for water.
Critical Point

The critical temperature can be defined as the temperature above which
the vapor cannot be liquefied no matter what pressure is applied.

The critical pressure is the pressure required to produce liquefaction at
the critical temperature

Together, the critical temperature and critical pressure define at the
critical point.
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Critical Point continued…
 The critical point for water is 374°C and 218 atm. Anything beyond this
point, involves the intermediate “fluid” region.
Phase Diagrams: CO2
The phase diagram for CO2 shows the liquid state does not exist at a pressure of 1
atm. The solid/liquid line has a positive slope, since the density of solid CO2 is
greater than that of liquid CO2
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