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Topic 11
Liquids and Solids
1
States of Matter
Comparison of gases, liquids, and solids.
– Gases are compressible fluids. Their molecules are
widely separated with no volume or shape.
– Liquids are relatively incompressible fluids. Their
molecules are more tightly packed and touching with
volume but no shape.
– Solids are nearly incompressible and rigid. Their
molecules or ions are in close contact and do not move
with volume and shape.
2
Changes of State
A change of state or phase transition is a change of a
substance from one state to another.
The energy associated with changing states is equal for both
directions but opposite in sign meaning exothermic vs
endothermic.
gas
boiling or
vaporization
condensation
exo
endo
sublimation
liquid
endo
freezing or
crystallization
melting or
fusion
endo
condensation or
deposition
exo
exo
solid
3
Vapor Pressure
Liquids are continuously vaporizing even if their
not at their boiling point.
– If a liquid is in a closed vessel with space above
it, a partial pressure of the vapor state builds up
in this space.
– The vapor pressure of a liquid is the partial
pressure of the vapor over the liquid, measured
at equilibrium at a given temperature.
– The more gas molecules (vapor), the higher the
vapor pressure.
4
Vapor Pressure
The vapor pressure of a liquid
depends on its temperature.
– As the temperature increases, the
kinetic energy of the molecular
motion becomes greater, and vapor
pressure increases. More energy
available to convert liquid molecules
to gas; hence, more gas molecules
and higher vapor pressure.
– Liquids with relatively high vapor pressures at normal
temperatures are said to be volatile. The higher the vapor
pressure, the easier it goes from a liquid to a gas.
5
– i.e. acetone has a high vapor pressure.
Boiling Point
The temperature at which the vapor pressure
of a liquid equals the pressure exerted on the
liquid is called the boiling point.
– As the temperature of a liquid increases, the vapor
pressure increases (more gas molecules formed) until it
reaches atmospheric pressure and begins to boil.
– At this point, stable bubbles of vapor (same species) form
within the liquid. This is called boiling.
– The normal boiling point is the boiling point at 1 atm.
– Because atmospheric pressure varies with altitude and
weather conditions, the boiling point of a liquid does as
well.
as atm pressure decreases, boiling point decreases
– Boiling point of water at 1 atm is 100oC but approximately
6
o
71 C on Mount Everest at 8850m elevation (0.33 atm).
Clausius-Clapeyron Equation
We noted that vapor pressure was a
function of temperature.
– It has been demonstrated that the logarithm of
the vapor pressure of a liquid varies linearly
with absolute temperature, K.
– Consequently, the vapor pressure of a liquid at
two different temperatures is described by:
P2 H vap 1 1
ln 

P1
R
T1 T2
(
)
– Equation allows you to determine the vapor
pressure of substance at any temperature if you
know the vapor pressure at the normal boiling
point.
7
Carbon disulfide, CS2, has a normal boiling point of 46°C 319 K
(vapor pressure = 760 mmHg) and a heat of vaporization of
26.8 kJ/mol. What is the vapor pressure of carbon disulfide at
35°C? 308 K
P H vap 1 1
ln
2
P1

R
(T  T )
1
2
Substituting into the Clausius-Clapeyron equation, we obtain:
P2
26.8  103 J/mol
1
1
ln


(760 mm Hg)
8.31 J/(mol  K) 319 K 308 K
(
)
 (3225 K)  (-1.12  10-4 K -1 )  0.361
Taking the antiln we obtain:
P2
 antiln(-0.361)
(760 mm Hg)
P2  antiln(-0.361)  760 mm Hg
P2  530 mm Hg
Note:
Temp decreased
and
8
VP decreased
Freezing Point
The temperature at which a pure liquid
changes to a crystalline solid, or freezes,
is called the freezing point.
– The melting point is identical to the freezing
point and is defined as the temperature at which
a solid becomes a liquid.
– Unlike boiling points, melting points are not
affected significantly by pressure changes;
however, large pressure changes may have
some affect.
melting, +H endothermic
S
L
freezing, -H exothermic
9
Heat of Phase Transition
To melt a pure substance at its melting
point requires an extra boost of energy to
overcome lattice energies.
– The heat needed to melt 1 mol of a pure
substance is called the heat of fusion and
denoted Hfus.
– For ice, the heat of fusion is 6.01 kJ/mol.
H 2O( s ) 
 H 2O( l );
 H 2O (l );
H 2O ( s ) 
H fus  6.01 kJ
H crystal  6.01 kJ
10
Heat of Phase Transition
To boil a pure substance at its boiling
point requires an extra boost of energy to
overcome intermolecular forces.
– The heat needed to boil 1 mol of a pure substance
is called the heat of vaporization and denoted
Hvap.
– For liquid water, the heat of vaporization is 40.66
kJ/mol.
H 2 O( l ) 
 H 2O( g );
H vap  40.66 kJ
H 2O (l ) 
 H 2O ( g );
H cond  40.66 kJ
11
Figure :
Heating
curve for
water.
Note:
Temp does not
change while
changing states;
heat is used to
overcome forces
of attraction
Hvap = 40.66 kJ 
Hcond = -40.66 kJ 
Hfus = 6.01 kJ 
Hcrystal = -6.01 kJ 
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Phase Diagrams
A phase diagram is a graphical way to
summarize the conditions under which the
different states of a substance are stable for
different temperatures and pressures.
– The diagram is divided into three areas
representing each state of the substance.
– The curves separating each area represent the
boundaries of phase changes.
13
Phase Diagrams
Below is a typical phase diagram. It consists of
three curves that divide the diagram into
regions labeled “solid, liquid, and gas”.
.
pressure
B
solid
D
C
liquid
.
gas
A
temperature
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Phase Diagrams
Curve AB, dividing the solid region from the
liquid region, represents the conditions under
which the solid and liquid are in equilibrium.
.
pressure
B
solid
liquid
.
D
C
gas
A
temperature
15
Phase Diagrams
Usually, the melting point is only slightly
affected by pressure. For this reason, the
melting point curve, AB, is nearly vertical.
.
pressure
B
solid
liquid
.
D
C
gas
A
temperature
16
Phase Diagrams
If a liquid is more dense than its solid (i.e.
water), the curve leans slightly to the left (points
toward less dense state), causing the melting
point to decrease with pressure.
.
pressure
B
solid
less dense
liquid
more dense
.
D
C
gas
A
temperature
17
Phase Diagrams
If a liquid is less dense than its solid, the curve
leans slightly to the right, causing the melting
point to increase with pressure. Most solids
are more dense than liquids.
.
pressure
B
liquid
less dense
solid
more dense
.
D
C
gas
A
temperature
18
Phase Diagrams
Curve AC, which divides the liquid region from
the gaseous region, represents the boiling
points of the liquid for various pressures. Note
that pressure has a major affect on curve and
the boiling point.
.
pressure
B
solid
liquid
.
D
C
gas
A
temperature
19
Phase Diagrams
Curve AD, which divides the solid region from
the gaseous region, represents the vapor
pressures of the solid at various temperatures.
An example of a solid that goes directly to a gas
is dry ice, CO2.
.
pressure
B
solid
liquid
.
D
C
gas
A
temperature
20
Phase Diagrams
The curves intersect at A, the triple point,
which is the temperature and pressure where
three phases (solid, liquid, gas) of a substance
exist in equilibrium (i.e. water occurs at 273.15K
and 4.58 Torr).
.
pressure
B
solid
less dense
liquid
s
.
L
g
D
C
gas
A
temperature
21
Phase Diagrams
The temperature above which the liquid state of
a substance no longer exists regardless of
pressure is called the critical temperature.
.
pressure
B
solid
liquid
.
D
C
A
temperature
gas
Tcrit
22
Phase Diagrams
The vapor pressure at the critical temperature is called the
critical pressure. Note that curve AC ends at the critical
point, C.
No liquefaction is observed above the critical temperature.
When the pressure of the gas above the critical temperature
is increased beyond the critical pressure, we have what is
called a supercritical fluid, SCF.
SCF
B
Pcrit
C
pressure
.
solid
liquid
.
D
A
temperature
gas
Tcrit
23
Intermolecular Forces
In general, properties like boiling point, melting point,
viscosity, vapor pressure, and surface tension depend
on the strength of the attractive forces among the
molecules.
The stronger the attractive forces, the higher the boiling
point, viscosity and surface tension of a liquid; the
higher the melting point of a solid; and the lower the
vapor pressure of a liquid.
Molecules gain more freedom of movement as a solid
melts or as a liquid vaporizes. The amount of energy
that they need to overcome the forces of attraction
among them increases the stronger the forces are.
24
Surface Tension
Surface tension is the energy required to increase the
surface area of a liquid by a unit amount.
A molecule within a liquid is pulled in all
directions, whereas a molecule on the surface
is only pulled to the interior.
As a result, there is a tendency for the surface
area of the liquid to be minimized.
To increase the surface area of a liquid requires movement
of molecules within the interior, where they experience
stronger attractions, to the surface. This requires energy.
The stronger the intermolecular (attractive) forces
between molecules, the higher the surface tension.
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Viscosity
Viscosity is the resistance to flow exhibited by all liquids
and gases.
Viscosity can be illustrated by measuring the time required
for a steel ball to fall through a column of the liquid.
Even without such measurements, you know that syrup
has a greater viscosity than water (thicker the fluid, the
more resistant to flow).
The stronger the intermolecular (attractive) forces
among the molecules of the gas or liquid, the more
resistant they would be to flow, the higher the
viscosity.
You can lower the viscosity of a substance by increasing
the temperature giving the system more energy to
overcome the attractive forces thereby increasing flow.
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Intermolecular Forces
The attractive forces between two molecules
can be classified into two major types:
1.) van der Waals forces which are the weak
attractive forces in a large number of substances (all
covalent bonded).
2.) hydrogen bonding interactions which occurs
in substances containing hydrogen atoms bonded to
certain very electronegative atoms (O, N, & F).
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Van der Waals Forces
Van der Waals forces can be further classified into:
1.) London dispersion forces (occurs between any
pair of molecules)
2.) dipole-induced dipole interaction (occurs
between a nonpolar molecule and a polar
molecule)
3.) dipole-dipole interaction (occurs between two
polar molecules)
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London Dispersion Forces
London forces refers to the force of attraction that exists
between any pair of molecules and is the predominant
interaction among most molecules.
It is due to temporary molecular polarizations, which occur
because electrons are always moving causing a distortion of
the electron cloud surrounding a molecule.
The larger the molecule, the more frequently the polarization
occur.
Therefore, we expect attractions to be stronger among larger
molecules. This means London forces increase with
molecular weight. The larger a molecule, the more easily
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the electron cloud can be distorted.
London Dispersion Forces
Let’s look at N2 and O2 as pure substances.
Both are nonpolar molecules that only have London
dispersion forces of attraction. Since they only have
London forces, the molecule with the larger molar mass
will have more polarization and a higher boiling point,
viscosity, and surface tension as well as a lower vapor
pressure.
N2
O2
Mm
28g/mol
32g/mol
BP
-196oC
-183oC
Since O2 has a higher molar mass, we expect it to
have a higher boiling point than N2.
30
Dipole-Dipole Interactions
Dipole-dipole interactions refers to the force of
attraction that exists between polar molecules.
The dipole-dipole interaction is an attractive
intermolecular force resulting from the tendency of polar
molecules to align themselves positive end to negative
end.
d+
d
d+
d
H Cl
H Cl
There is a higher electron density in the Cl end of the
polar molecule; this end, we say, is partially negative.
The H end of the polar molecule is partially positive
and is attracted to the partially negative end of a nearby
HCl molecule which is a dipole-dipole interaction.
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Dipole-Dipole Interactions
Let’s look at N2, NO, and O2 pure substances.
Both N2 and O2 are nonpolar molecules that only have
London dispersion forces of attraction. However, NO is a
polar molecule which has dipole-dipole interactions in
addition to London forces to overcome which affects its
boiling point, viscosity, surface tension, and vapor pressure.
The boiling point of NO will be the highest among these
substances because of the additional dipole-dipole
interactions in the polar molecule.
N2
NO
O2
Mm
28g/mol
30g/mol
32g/mol
BP
-196oC
-152oC
-183oC
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Dipole-induced Dipole Interactions
Dipole-induced dipole interaction refers to the force
of attraction that exists between a polar molecule and a
nonpolar molecule.
When a nonpolar molecule comes close to the positive
end of a polar molecule, its electrons would be attracted
toward the polar molecule causing temporary (or
induced) polarization.
Similarly, when a molecule comes close to the negative
end of a polar molecule, its electrons would be repelled,
again causing temporary (or induced) polarization.
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Hydrogen Bonding
Hydrogen bonding is a force that exists between a
hydrogen atom covalently bonded to a very
electronegative atom (O, N, F).
:
:
:
– To exhibit hydrogen bonding, one of the following
three structures must be present.
H N
H O
H F
– Only N, O, and F are electronegative enough to
leave the hydrogen nucleus almost stripped bare
of electrons making it strongly attracted to a lone
pair of a highly electronegative atom (O, N, F) in
a nearby molecule.
34
Hydrogen Bonding in H2O
The H atoms in water are bonded to a highly
electronegative O atom.
Because H is almost stripped bare of electrons, it is
strongly attracted to the lone pair of the O atom in the
neighboring water molecule
O
H
O
H
H
O
O
H
H
H
H
H
35
Hydrogen Bonding
Hydrogen bonding interaction is much stronger than dipoledipole interaction and for small molecules stronger than
London forces.
Hydrogen bonding accounts for the unusually high boiling
point of water. Water molecules are small and a liquid at
room temperature; substances made of molecules of
comparable size are gaseous at room temperature.
O2 is a gas at room temperature, while H2O is a liquid even
though H2O molecules are smaller than O2 molecules.
Water is a polar molecule and capable of extensive
hydrogen bonding thereby raising its boiling point
considerably.
36
Hydrogen Bonding
Molecules exhibiting hydrogen bonding
have abnormally high boiling points
compared to molecules with similar van
der Waals forces.
Which of the following are capable of exhibiting
hydrogen bonding?
N2
CH4
HI
HF
(CH3)2O
C6H5OH
CH3OH
NH3
H2S
Within the Lewis structure of the molecule, H must be
attached to O, N, or F for hydrogen bonded to occur.
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Let’s look at H2O, H2S, H2Se, and H2Te as pure substances.
If you draw the Lewis structures and examine the VSEPR geometry of these
molecules, you would determine that all of these molecules are polar with a same
bent geometry.
This means that all of the molecules have London dispersion and dipole-dipole
interactions. Since they are all polar molecules, their London dispersion forces will
dictate their boiling points which will vary based on their molar mass (larger molar
mass, stronger London forces).
Based on molar mass, we would predict H2Te to have the highest boiling point
because it has the largest molar mass.
If we extrapolate the boiling point of H2O based on the other actual boiling points,
the boiling point of H2O should be -68oC.
However, since H2O has very strong hydrogen bonding to overcome as well, it’s
boiling point is actually extremely high, 100oC, as compared to the other
substances despite it’s low molar mass.
H 2O
Mm
18.02
due to hydrogen
100ooCC
BP bonding
-68
H 2S
34.08
-60.33oC
H2Se
H2Te
80.98
129.63g/mol
-41.3oC -2oC
38
For which of the following pairs of molecules do we expect
London dispersion forces, dipole-dipole, dipole-induced, and
hydrogen bonding?
First, we must realize through Lewis structures and VSEPR
that CO2, CH4 are nonpolar while H2O, HCl, NH3 are polar.
all have
CH4 and CH4
London
H2O and H2O
London
H2O and CO2
London
NH3 and NH3
London
HCl and HCl
polar-polar
dipole-dipole
polar-nonpolar
H–O, N, F
hydrogen bonding
dipole-induced
dipole-dipole
hydrogen bonding
39
London
dipole-dipole
Which species has the higher boiling point CS2 or CCl4?
If you draw the Lewis structures and examine the
VSEPR geometry of these molecules, you would
determine that both of these molecules are nonpolar
with CS2 having a linear geometry and CCl4 having a
tetrahedral geometry.
Since both are nonpolar molecules, they only have
London dispersion forces of attraction which means
the molecule with the larger molar mass (CCl4) will
have more polarization and a higher boiling point. 40
Which species has the higher boiling point H2O or CO?
If you draw the Lewis structures and examine the
VSEPR geometry of these molecules, you would
determine that H2O is a polar molecule with a bent
geometry and CO is a polar molecule with a linear
geometry.
Since both are polar molecules, they have London
dispersion and dipole-dipole interactions. Usually in
this instance, the species with the larger molar mass
would have the stronger London forces and higher
boiling point; however, H2O also has hydrogen
bonding causing the boiling point to be much higher
despite the lower molar mass.
41
Intermolecular Forces
In summary, intermolecular forces play a
large role in many of the physical properties
of liquids and gases. These include:
– vapor pressure
as intermolecular forces increase, vapor pressure decreases
– boiling point
as intermolecular forces increase, boiling point increases
– surface tension
as intermolecular forces increase, surface tension increases
– viscosity
as intermolecular forces increase, viscosity decreases
42
Crystalline Solids
The regular arrangement of particles in a crystalline solid
leads to the minimization of total potential energy of
interactions of the particles and the most stable arrangement.
A crystalline structure is said to have a long-range order.
The overall structure can be thought of in terms of a repeating
pattern, called a unit cell. The unit cells making up the solid
are in close contact and in fixed positions.
Solids are characterized by the type of force holding the
structural units together. In some cases, these forces are
intermolecular, but in others they are chemical bonds
(metallic, ionic, or covalent).
43
Crystalline Solid
Properties of crystals depend on the type of particles in
the lattice. Crystals can be classified as ionic, molecular,
or atomic.
– Ionic crystals tend to have very high melting points
due to strong attractions among ions. They are brittle
due to the strong repulsions that result when ions of
like charges are momentarily brought closer together
as ions are slightly displaced from their locations when
the crystal is, say, hit by a hammer.
– Molecular crystals tend to have low melting or
sublimation points. The attractive forces among the
molecules are relatively weak (mainly van der Waals
44
forces).
Crystalline Solid
Atomic crystals can be classified as nonbonding,
metallic, or covalent network.
–
Nonbonding atomic crystals are formed when noble gases are
frozen to very low temperatures. These atoms are held together
by very weak London dispersion forces.
–
Metallic crystals are made up of atoms of metallic elements. If
more than one element is present, the solid is a solution and is
called an alloy. A strong metallic bond is the reason metals have
high melting points and boiling points; most metals are solids at
room temperature. Because metal atoms can readily slip and roll
over each other without breaking the metallic bond, metals are
malleable and ductile.
–
Covalent network crystal can be thought of as one giant
molecule; the atoms are held together by very strong covalent
bonds. Thus, a network covalent crystal like diamond has a very
45
high melting point and is among the hardest material known.
Physical Properties
Many physical properties of a solid can be
attributed to its structure and forces of
attraction called crystal lattice energy or
ion-ion intermolecular forces.
– For a solid to melt, the forces holding the
structural units together must be overcome.
– For a molecular solid, these are weak
intermolecular attractions.
– Thus, molecular solids tend to have low
melting points (below 300oC).
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Physical Properties
– For ionic solids and covalent network solids
to melt, chemical bonds must be broken.
– For that reason, their melting points are
relatively high.
– Note that for ionic solids, melting points increase
with the strength of the ionic bond while solubility
decreases.
– Ionic bonds are stronger when:
1. The magnitude of charge is high. Higher the charge, the
stronger the attraction, the more energy needed to
overcome attraction; therefore, MP increases and solubility
decreases.
2. The ions are small (higher charge density). Smaller the
radius, the closer the opposite charges and larger attraction,
the more energy needed to overcome attraction; therefore,
47
MP increases and solubility decreases.
Summary: The attractive forces (crystal lattice energy)
between a pair of oppositely charged ions increases
(stronger bond) as the charges on the ions increases and as
ionic size decreases; hence higher MP and lower solubility.
Which of the following has the higher melting point and lower
solubility?
MgO
CaBr2
vs.
vs.
HW 69
code: liquids
NaCl
CaCl2
Since magnesium oxide involves higher
charges (+2, -2) than sodium chloride (+1, -1),
MgO will have the higher MP and lower
solubility.
Since both species have the same charges (+2,
-1), the size of the ions (anion in this case) will
affect the properties. Calcium chloride has the
smaller anion; therefore, it will have the
stronger attraction to calcium and have the 48
higher MP and lower solubility.