General Chemistry

1. Liquids and Solids
AP Chem Unit 10

2. Sections
Intermolecular Forces
Liquid state
Solid Structures
Metal Structures
Carbon and Silicon Networks

3. Sections
Molecular Solids
Ionic Solids
Vapor pressure and State Change
Phase Diagrams

4. 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
compressibility
density
intermolecular forces

5. States of Matter
H2O(s)H2O(l) ΔH°fus = 6.02 kj/mol
H2O(l)H2O(g) ΔH°vap = 40.7 kj/mol
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

6. Intermolecular Forces
10.1

7. Intermolecular Forces
Electrons shared within the molecule are called
intramolecular bonding.
In the condensed states of matter the
attraction between molecules are called
intermolecular forces.

8. 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.
40.7kj needed to vaporize water
934kj to break the O-H bond

9. 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

10. Dipole-Dipole Forces
In a condensed
state, molecules
line up dipoles to
minimize repulsions
and maximize
attractions.

11. Dipole-Dipole Forces
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.
These strong attractions have a strong
impact on melting points and boiling
points.

12. Boiling Points of Covalent Hydrides

13. 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.

14. Hydrogen bonds

15. 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.

16. 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.

17. 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.

18. London Dispersion Forces
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.
larger atoms have more electrons and
increases the probability of a temporary
dipole.

19. London Dispersion Forces

20. London Dispersion Forces
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.

21. The Liquid State
10.2

22. Liquid Characteristics
lack of rigidity
low compressibility
high density
rounded droplets
capillary action
viscosity

23. 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.

24. Rounded Droplets
Molecules that are polar and have stronger
intermolecular forces have stronger surface
tensions.

25. Surface Tension

26. Capillary Action
Capillary action is the spontaneous rising of a
liquid in a narrow tube. This action is due to
two forces
cohesive forces- the intermolecular forces
among the molecules.
adhesive forces – the attractive forces
between the liquid and the container.

27. 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.

28. Adhesive forces
A nonpolar substance, such as mercury,
has a convex meniscus because the
cohesive forces are stronger than the
adhesive forces.

29. Meniscus: Water vs. Mercury

30. 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.

31. Viscosity
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.

32. Introduction to Structures
and Types of Solids
10.3

33. Types of Solids
Crystalline solids
Amorphous solids

34. Crystalline Solids
Crystalline solids have a regular arrangement
of components at a microscopic level and
produce beautiful, characteristic shapes of
crystals:

35. Crystalline Solids
The positions of components are usually
represented by a lattice.
lattice is a three dimensional system of
units repeating in a pattern. The smallest
repeating unit of the lattice is called the
unit cell.

36. Three types of Crystalline Solids

37. 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.

38. 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.

39. X-ray Analysis of Solids

40. X-ray Analysis of Solids
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.

41. X-ray Analysis of Solids
A diffractometer is a computer-controlled
instrument used for carrying out the X-ray
analysis of crystals
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.

42. X-ray Analysis of Solids
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)

43. X-ray Analysis of Solids

44. 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
2.33 Á

45. Types of Solids
Ionic solids
ionic solids are made of ions
Molecular solids
Molecular solids have small units of
covalently bonded molecules.
Atomic solids
Atomic solids are made of elements such
as carbon (graphite, diamond and the
fullerenes), boron, silicon, and all metals.

46. Fullerenes

47. Types of Solids

48. Atomic Solids
Atomic solids are broken down into subgroups
depending on the bond that exists in the solid:
Metallic solids
Has delocalized nondirectional covalent
bonding.
Network solids
atoms bond with strong directional covalent
bonding that lead to giant molecules and
networks

49. Atomic Solids
Group 8A solids
noble gases are attracted to each other with
London dispersion forces.

50. Classification of Solids

51. Structure and
Bonding in Metals
10.4

52. Metal Characteristics
Most of the properties that we see in metals
is due to the nondirectional covalent
bonding found in metal crystals.
High thermal conductivity
Electrical conductivity
Malleability
Ductility

53. Metallic Crystals
Metallic crystals can be pictured as containing
spherical atoms packed together that can
be bonded to each other equally in all
directions.
This arrangement is called closest packing.

54. 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.

55. Closest Packing
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.

56. Closest Packing

57. Closest Packing: Hexagonal

58. Closest Packed: Cubic

59. Closest Packing
Knowing the net number of atoms in a particular
unit cell is important for many applications
involving solids.

60. Closest Packing
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.

61. Closest Packing
Face – Centered Cubic Unit Cell

62. Cubic Substances
Metals that form cubic closest packed solids are:
aluminum
iron
copper
cobalt
nickel

63. Hexagonal Substances
Metals that form hexagonal closest packed
solids are:
magnesium
zinc

64. Other Metal Solids
Calcium and certain other metals can
crystallize in either cubic or hexagonal solids.
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.

65. Example Problem
Silver crystallizes in a cubic closest packed
structure. The radius of a silver atom is
144pm. Calculate the density of solid
silver?
10.6 g/cm3

66. 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.

67. 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.

68. Electron Sea Model

69. 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.

70. Band Model (MO Model)

71. Band Model
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.

72. Band Model

73. 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.
Interstitial alloy – is formed when some of the
interstices (holes) in the closest packed lattice are
occupied by smaller atoms.

74. Metal Alloys

75. Substitutional Alloys
Example: brass: 1/3 of copper metal atoms
are replaced by zinc atoms
Sterling silver- 93% silver and 7% copper.
Pewter- 85% tin, 7% copper, 6% bismuth and 2%
antimony.
Plumbers solder – 95% tin and 5% antimony

76. 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.

77. Interstitial Alloy
The amount of carbon directly affects the
properties of steel:
Mild steels- contains less than .2% carbon:
nails chains and cables.
Medium steels- contain .2-.6% carbon: rails
and structural steel
High-carbon steel – .6-1.5% carbon: springs,
tools and cutlery.

78. 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.

79. Carbon and Silicon
Network Atomic Solids
10.5

80. 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.

81. 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.

82. 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.

83. 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.

84. Carbon

85. Carbon: Graphite layers

86. 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.

87. 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.

88. Silica
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.

89. 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 produce 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.

90. Glass

91. 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.

92. Silicates

93. Silicates
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.

94. Clays
Clay comes from the weathering
of feldspar, an Aluminosilicate
(Na2O/K2OAl2O36SiO2). 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.

95. 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.

96. 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.

97. 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.

98. 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.

99. Energy Level Diagrams for N-type
and P-type Semiconductors.

100. 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.

101. P-N Junction
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.

102. P-N Junction
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.

103. P-N Junction
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.

104. P-N Junction
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.

105. Molecular Solids
10.6

106. Molecular Solids
Sometimes Molecular solids can have large
discrete molecular units in a lattice-type
position. These molecules have strong bonds
within the molecules but relatively weak
between the molecules.
Common examples: Ice, dry ice (solid
carbon dioxide), Sulfur (S8), Phosphorus (P4)

107. Molecular

108. Molecular Solids
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.

109. Ionic Solids
10.7

110. Ionic Solids
Ionic solids are stable, high melting
substances held together by the strong
electrostatic forces that exist between
oppositely charged ions.

111. Ionic Solids
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.

112. Ionic Solids
There are three types
of holes in closest
packed structures:
1.Trigonal holes are
formed by three
spheres in the same
layer

113. Ionic Solids
There are three types
of holes in closest
packed structures:
2.Tetrahedral holes
are formed when a
sphere sits in the
dimple of three
spheres in an
adjacent layer.

114. Ionic Solids
There are three types
of holes in closest
packed structures:
3.Octahedral holes are
formed between two
sets of three spheres
in adjoining layers of
the closest packed
structures.

115. Ionic Solids
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.

116. Ionic Solids
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.

117. Ionic Solids
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 face-
centered 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.

118. Example Problem
Determine the net number of Na+
and Cl-
ions
in the sodium chloride unit cell.
4 Na+
and 4 Cl-

119. Example Problem
Classify each of the following substances
according to the type of solid it forms:
gold
carbon dioxide
lithium fluoride
krypton
metallic
molecular
ionic
8a

120. Types and Properties of Solids

121. Vapor Pressure and
Changes of State
10.8

122. 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.
Water has strong intermolecular forces and
this increases the energy required to
vaporize. Also making it a great coolant.

123. Vaporization
The energy required to vaporize 1 mole of
liquid at 1 atm is called the heat of
vaporization or the enthalpy of
vaporization.
The symbol for this is ΔHvap.

124. Vapor Pressure
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.

125. Vapor Pressure
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.

126. Rates of Condensation and
Evaporation.

127. 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.

128. Vapor Pressure
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.

129. Vapor Pressure
When the system reaches equilibrium, the
vapor pressure can be determined from the
change in the height of the mercury column

130. Vapor Pressure
Patmosphere = Pvapor + PHg column
Pvapor = Patmosphere - Phg column

131. Vapor Pressure
The vapor pressure of liquids vary widely.
Liquids with high vapor pressures are
said to be volatile. They evaporate
rapidly in an open dish.

132. Vapor Pressure
The vapor pressure of a liquid is principally
determined by the size of the
intermolecular forces in the liquid.
Liquids with strong molecular forces have
relatively low vapor pressures because it
takes so much energy for the molecules to
escape.
In general, substances with large molar
masses have relatively low vapor pressures
due of large dispersion forces.

133. Vapor Pressure
Vapor pressure increases significantly with
temperature.
In order to break intermolecular forces, a
sufficient amount of kinetic energy is
needed.
As temperature of the liquid increases, so
does kinetic energy of the liquid.

134. Vapor Pressure of Water

135. Vapor Pressure
The nature of the temperature, vapor
pressure relationship is quadratic. Pvap vs.
1/T (Kelvin) gives a direct relationship.

136. Vapor Pressure
Vapor pressure equation:
R is the universal gas constant (8.3145 J/K). In
means natural logarithm.
ln
PvapT1
PvapT 2





÷÷=
∆Hvap
R
1
T2
−
1
T1





÷

137. Example Problem
Using the graph, determine whether water or diethyl
ether has the larger enthalpy of vaporization.
Ether has the
smaller slope
therefore
smaller ΔHvap

138. Example Problem
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.
93.7 torr

139. 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.

140. 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.

141. Heating Curve

142. 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

143. Heating Curve
The energy change that occurs at the
melting point when a solid melts is called
heat of fusion or enthalpy of fusion.
Note that changes of state are physical
changes, although intermolecular forces have
been broken, no chemical bonds have been
broken.

144. Enthalpy of Fusions
and Melting Points

145. Melting and Boiling
The melting and boiling points for a substance
are determined by the vapor pressure of the
solid and liquid states.

146. Melting and Boiling
At 0°C the vapor
pressure of ice is
less than that of
liquid. Vapor
pressure of ice
increases more
rapidly than water.

147. Melting and Boiling
A point is reached
when the liquid and
solids have
identical vapor
pressures. This is
the melting point.

148. 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.

149. Freezing Point
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.

150. Freezing Point
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.

151. Melting Point and Boiling Point
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.

152. 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.

153. Supercooled and Superheated

154. Supercooled and Superheated
Changes of state do not always occur exactly at
the boiling point or 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.

155. Supercooled and Superheated
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.

156. Supercooled and Superheated
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.

157. Phase Diagrams
10.9

158. Phase Diagram
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.

159. Phase Diagram

160. Phase Diagram
The solid/liquid boundary has a negative
slope.
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.

161. Boiling Points of Water vs. Pressure

162. Example 1
Pressure is 1 atm. Water
moves through the
changes of state
according to the vapor
pressure at the
corresponding
temperatures.

163. 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.

164. 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.

165. 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.

166. Example 4
The fluid region is
neither true liquid or
vapor. This unusual
behavior occurs because
the conditions are beyond
the critical point for
water.

167. 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.

168. Critical Point
The critical point for water is 374°C and
218 atm. Anything beyond this point,
involves the intermediate “fluid” region.

169. Phase Diagrams
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

170. THE END!!