Transcript Intermolecular Forces

Intermolecular Forces and
Liquids and Solids
Chapter 11
States of Matter
• Comparison of gases, liquids, and solids.
Gases are compressible fluids. Their
molecules are widely separated.
Liquids are relatively incompressible fluids.
Their molecules are more tightly packed.
Solids are nearly incompressible and rigid.
Their molecules or ions are in close contact
and do not move.
A phase is a homogeneous part of the system in
contact with other parts of the system but
separated from them by a well-defined boundary.
2 Phases
Solid phase - ice
Liquid phase - water
Intermolecular Forces
Intermolecular forces are attractive forces between molecules.
Intramolecular forces hold atoms together in a molecule.
Intermolecular vs Intramolecular
•
41 kJ to vaporize 1 mole of water (inter)
•
930 kJ to break all O-H bonds in 1 mole of water (intra)
“Measure” of intermolecular force
Generally,
intermolecular
forces are much
weaker than
intramolecular
forces.
boiling point
melting point
DHvap
DHfus
DHsub
Intermolecular Forces; Explaining
Liquid Properties
• The term van der Waals forces is a general
term including dipole-dipole and London
forces.
Van der Waals forces are the weak attractive
forces in a large number of substances.
Hydrogen bonding occurs in substances
containing hydrogen atoms bonded to certain
very electronegative atoms.
Van der Waals Forces
• Van der Waals Forces are types of
intermolecular forces.
• Consists of:
1 Dipole-dipole interactions
2 Ion-dipole interactions
3 London forces
7
Dipole-Dipole Forces
• Polar molecules can attract one another
through dipole-dipole forces.
The dipole-dipole force is an attractive
intermolecular force resulting from the tendency
of polar molecules to align themselves positive
end to negative end.
d+
H
Cl
d-
d+
H
Cl
d-
Intermolecular Forces
Dipole-Dipole Forces
Attractive forces between polar molecules
Orientation of Polar Molecules in a Solid
Intermolecular Forces
Ion-Dipole Forces
Attractive forces between an ion and a polar molecule
Ion-Dipole Interaction
London Forces
• London forces are the weak attractive
forces resulting from instantaneous dipoles
that occur due to the distortion of the
electron cloud surrounding a molecule.
London forces increase with molecular weight.
The larger a molecule, the more easily it can be
distorted to give an instantaneous dipole.
All covalent molecules exhibit some London force.
• London forces:
– exist between all molecules,
– is the only attractive force between
nonpolar atoms or molecules,
– and dipole-dipole attractions occur
between polar molecules.
• Electrons are in constant motion.
• Electrons can be, in an instant, arranged in such a way
that they have a dipole. (Instantaneous dipole)
• The temporary dipole interacts with other temporary
dipoles to cause attraction.
Intermolecular Forces
Dispersion Forces
Attractive forces that arise as a result of temporary
dipoles induced in atoms or molecules
ion-induced dipole interaction
dipole-induced dipole interaction
Induced Dipoles Interacting With Each Other
Intermolecular Forces
Dispersion Forces Continued
Polarizability is the ease with which the electron distribution
in the atom or molecule can be distorted.
Polarizability increases with:
•
greater number of electrons
•
more diffuse electron cloud
Dispersion
forces usually
increase with
molar mass.
What type(s) of intermolecular forces exist between
each of the following molecules?
HBr
HBr is a polar molecule: dipole-dipole forces. There are
also dispersion forces between HBr molecules.
CH4
CH4 is nonpolar: dispersion forces.
S
SO2
SO2 is a polar molecule: dipole-dipole forces. There are
also dispersion forces between SO2 molecules.
Hydrogen Bonding
• Hydrogen bonding is a force that exists
between a hydrogen atom covalently bonded to
a very electronegative atom, X, and a lone pair
of electrons on a very electronegative atom, Y.
:
:
:
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 exposed.
Hydrogen Bonding
• Hydrogen bonding:
– not considered a Van der Waals Force
– is a special type of dipole-dipole attraction
– is a very strong intermolecular attraction
causing higher than expected b.p. and m.p.
• Requirement for hydrogen bonding:
– molecules have hydrogen directly bonded to
O, N, or F
Examples of hydrogen bonding:
H2O
NH3
HF
Intermolecular Forces
Hydrogen Bond
The hydrogen bond is a special dipole-dipole interaction between
they hydrogen atom in a polar N-H, O-H, or F-H bond and an
electronegative O, N, or F atom.
A
H…B
or
A
A & B are N, O, or F
H…A
Hydrogen Bond
Hydrogen Bonding
• Molecules exhibiting hydrogen bonding have
abnormally high boiling points compared to
molecules with similar van der Waals forces.
For example, water has the highest boiling
point of the Group VI hydrides.
Similar trends are seen in the Group V and VII
hydrides.
Why is the hydrogen bond considered a “special”
dipole-dipole interaction?
Decreasing molar mass
Decreasing boiling point
Review: Hydrogen Bonding
• A hydrogen atom bonded to an electronegative
atom appears to be special.
The electrons in the O-H bond are drawn to the
O atom, leaving the dense positive charge of the
hydrogen nucleus exposed.
It’s the strong attraction of this exposed nucleus
for the lone pair on an adjacent molecule that
accounts for the strong attraction.
A similar mechanism explains the attractions in
HF and NH3.
Review: Hydrogen Bonding
O
H
O
H
H
O
O
H
H
H
H
H
Van der Waals Forces and the
Properties of Liquids
• In summary, intermolecular forces play a
large role in many of the physical properties
of liquids and gases. These include:
vapor pressure
boiling point
surface tension
viscosity
Van der Waals Forces and the
Properties of Liquids
• Surface tension increases with increasing
intermolecular forces.
Surface tension is the energy needed to
reduce the surface area of a liquid.
To increase surface area, it is necessary to pull
molecules apart against the intermolecular
forces of attraction.
Properties of Liquids
Surface tension is the amount of energy required to stretch or
increase the surface of a liquid by a unit area.
High surface
tension
Strong intermolecular
forces
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.
Properties of Liquids
Cohesion is the intermolecular attraction between like molecules
Adhesion is an attraction between unlike molecules
Adhesion
Cohesion
Liquid levels in capillaries.
Properties of Liquids; Surface
Tension and Viscosity
This explains why falling raindrops are nearly
spherical, minimizing surface area.
In comparisons of substances, as intermolecular
forces between molecules increase, the
apparent surface tension also increases.
Review: Surface Tension
• Surface tension - a measure of the attractive forces
exerted among molecules at the surface of a liquid.
• Surface molecules are surrounded and attracted by fewer
liquid molecules than those below.
• Net attractive forces on surface molecules pull them
downward.
– Results in “beading”
• Surfactant - substance added which decreases the surface
tension
– example: soap
• Surfactant - substance added which
decreases the surface tension
• The nonpolar end is hydrophobic.
• The polar end is hydrophilic.
• Surfactants are amphipathic.
Example - soap
DDBSA is the largest-volume synthetic
surfactant because of its relatively low cost,
good performance, the fact that it can be
dried to a stable powder. DDBSA is also
bio-degradable, thus providing
environmental friendliness. DDBSA is
mainly used to produce household
detergents including laundry powders,
laundry detergents, dishwashing liquids
and other household cleaners. It is used
as an emulsifier for applications like
agricultural produces and industrial
lubricants (Calcium Sulfonate Grease).
DDBSA as a detergent in water:
The Sinking Duck!!!!!!!
Why?
At the air-water interface the hydrophobic
end can escape the water, whereas the
hydrophilic end remains nestled in the
water’s surface. The concentration of
water at the interface is reduced.
Therefore, the attractive forces between
molecules at the interface is reduced,
thus the surface tension is reduced.
Van der Waals Forces and the
Properties of Liquids
• Viscosity increases with increasing
intermolecular forces because increasing
these forces increases the resistance to flow.
Other factors, such as the possibility of
molecules tangling together, affect viscosity.
Liquids with long molecules that tangle
together are expected to have high viscosities.
Intermolecular Forces; Explaining
Liquid Properties
• 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.
In comparisons of substances, as intermolecular
forces increase, viscosity usually increases.
Properties of Liquids
Viscosity is a measure of a fluid’s resistance to flow.
Strong
intermolecular
forces
High
viscosity
Kinetic Viscometery
Viscometers
The Solid State
• Particles highly organized and well defined
fashion
• Fixed shape and volume
• Properties of Solids:
– incompressible
– m.p. depends on strength of attractive force
between particles
– Crystalline solid - regular repeating structure
– Amorphous solid - no organized structure.
9
Solid State
• A solid is a nearly incompressible state of
matter with a well-defined shape. The units
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).
A crystalline solid possesses rigid and long-range order. In a
crystalline solid, atoms, molecules or ions occupy specific
(predictable) positions.
An amorphous solid does not possess a well-defined
arrangement and long-range molecular order.
A unit cell is the basic repeating structural unit of a crystalline
solid.
At lattice points:
lattice
point
Unit Cell
Unit cells in 3 dimensions
•
Atoms
•
Molecules
•
Ions
Shared by 8
unit cells
Shared by 2
unit cells
1 atom/unit cell
2 atoms/unit cell
4 atoms/unit cell
(8 x 1/8 = 1)
(8 x 1/8 + 1 = 2)
(8 x 1/8 + 6 x 1/2 = 4)
When silver crystallizes, it forms face-centered cubic
cells. The unit cell edge length is 409 pm. Calculate
the density of silver.
d=
m
V
V = a3 = (409 pm)3 = 6.83 x 10-23 cm3
4 atoms/unit cell in a face-centered cubic cell
1 mole Ag
107.9 g
-22 g
x
m = 4 Ag atoms x
=
7.17
x
10
mole Ag 6.022 x 1023 atoms
7.17 x 10-22 g
m
3
=
=
10.5
g/cm
d=
V
6.83 x 10-23 cm3
Crystal Defects
• There are principally two kinds of defects
that occur in crystalline substances.
Chemical impurities, such as in rubies,
where the crystal is mainly aluminum
oxide with an occasional Al3+ ion replaced
with Cr3+, which gives a red color.
Defects in the formation of the lattice.
Crystal planes may be misaligned, or sites
in the crystal lattice may remain vacant.
Extra distance = BC + CD = 2d sinq = nl
(Bragg Equation)
X rays of wavelength 0.154 nm are diffracted from a
crystal at an angle of 14.170. Assuming that n = 1,
what is the distance (in pm) between layers in the
crystal?
nl = 2d sin q
n=1
nl
q = 14.170 l = 0.154 nm = 154 pm
1 x 154 pm
=
= 314.0 pm
d=
2 x sin14.17
2sinq
Physical Properties
• Many physical properties of a solid can be
attributed to its structure.
Melting Point and Structure
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).
Physical Properties
• Many physical properties of a solid can be
attributed to its structure.
Melting Point and Structure
For ionic solids and covalent network solids
to melt, chemical bonds must be broken.
For that reason, their melting points are
relatively high.
Physical Properties
• Many physical properties of a solid can be
attributed to its structure.
Melting Point and Structure
Note that for ionic solids, melting points
increase with the strength of the ionic bond.
Ionic bonds are stronger when:
1. The magnitude of charge is high.
2. The ions are small (higher charge density).
Solid State
From this point of view, there are four
types of solids.
Ionic (Ionic bond)
Covalent (Covalent bond)
Molecular (Van der Waals forces)
Metallic (Metallic bond)
Types of Crystals
Ionic Crystals
• Lattice points occupied by cations and anions
• Held together by electrostatic attraction
• Hard, brittle, high melting point, boiling point
• Poor conductor of heat and electricity
• If dissolves in water, electrolytes
• NaCl
CsCl
ZnS
CaF2
Types of Crystals
Covalent Crystals
• Lattice points occupied by atoms
• Held together by covalent bonds
• Hard, high melting point, boiling point
• Poor conductor of heat and electricity
• Extremely Hard carbon
• Diamond
atoms
diamond
graphite
Types of Crystals
Molecular Crystals
• Lattice points occupied by molecules
• Held together by intermolecular forces
• Soft, low melting point
• Poor conductor of heat and electricity
• Often volatile
• Ice, dry ice
Types of Crystals
Metallic Crystals
• Lattice points occupied by metal atoms
• Held together by metallic bonds
• Soft to hard, low to high melting point
• Good conductors of heat and electricity
• Iron, copper, and silver
Cross Section of a Metallic Crystal
nucleus &
inner shell emobile “sea”
of e-
Crystal Structures of Metals
Physical Properties
• Many physical properties of a solid can be
attributed to its structure.
Melting Point and Structure
Metals often have high melting points, but there
is considerable variability.
Melting points are low for Groups IA and IIA but
increase as you move into the transition metals.
The elements in the middle of the transition
metals have the highest melting points.
Physical Properties
• Many physical properties of a solid can be
attributed to its structure.
Hardness and Structure
Hardness depends on how easily structural
units can be moved relative to one another.
Molecular solids with weak intermolecular
attractions are rather soft compared with
ionic compounds, where forces are much
stronger.
Physical Properties
• Many physical properties of a solid can be
attributed to its structure.
Hardness and Structure
Covalent network solids are quite hard
because of the rigidity of the covalent network
structure.
Diamond and silicon carbide (SiC), threedimensional covalent network solids, are
among the hardest substances known.
Physical Properties
• Many physical properties of a solid can be
attributed to its structure.
Hardness and Structure
Molecular and ionic crystals are generally
brittle because they fracture easily along
crystal planes.
Metallic solids, by contrast, are malleable.
Physical Properties
• Many physical properties of a solid can be
attributed to its structure.
Electrical Conductivity and Structure
Molecular and ionic solids are generally
considered nonconductors.
Ionic compounds conduct in their molten
state, as ions are then free to move.
Metals are all considered conductors.
Physical Properties
• Many physical properties of a solid can be
attributed to its structure.
Electrical Conductivity and Structure
Of the covalent network solids, only graphite
conducts electricity.
Types of Crystals
An amorphous solid does not possess a well-defined
arrangement and long-range molecular order.
A glass is an optically transparent fusion product of inorganic
materials that has cooled to a rigid state without crystallizing
Crystalline
quartz (SiO2)
Non-crystalline
quartz glass
Chemistry In Action: High-Temperature Superconductors
Chemistry In Action: And All for the Want of a Button
T < 13 0C
white tin
grey tin
stable
weak
Changes of State
• A change of state or phase transition is a
change of a substance from one state to
another.
gas
boiling
sublimation
condensation
liquid
freezing
melting
solid
condensation or
deposition
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
affected significantly by only large pressure
changes.
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 until it reaches
atmospheric pressure.
At this point, stable bubbles of vapor form
within the liquid. This is called boiling.
The normal boiling point is the boiling point at
1 atm.
The melting point of a solid or
the freezing point of a liquid is
the temperature at which the
solid and liquid phases coexist
in equilibrium
Freezing
H2O (l)
Melting
H2O (s)
• What happens when you go to a mountain where
the atmospheric pressure is lower than 1 atm?
(The boiling point lowers.)
T2 > T1
Condensation
Evaporation
Least
Order
Greatest
Order
Molar heat of sublimation
(DHsub) is the energy required to
sublime 1 mole of a solid.
DHsub = DHfus + DHvap
( Hess’s Law)
Deposition
H2O (g)
Sublimation
H2O (s)
Vapor Pressure
• Liquids are continuously vaporizing.
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.
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.
Liquids and solids with relatively high vapor
pressures at normal temperatures are said to
be volatile.
Measurement of
the vapor pressure
of water.
Vapor Pressure of a Liquid
• What happens when you put water in a sealed
container?
• Both liquid water and water vapor will exist in the
container.
• How does this happen below the boiling point?
• Kinetic Theory - Liquid molecules are in continuous
motion, with their average kinetic energy directly
proportional to the Kelvin temperature.
6
Variation of vapor
pressure with
temperature.
The equilibrium vapor pressure is the vapor pressure
measured when a dynamic equilibrium exists between
condensation and evaporation
H2O (l)
Dynamic Equilibrium
Rate of
Rate of
= evaporation
condensation
H2O (g)
Before
Evaporation
At
Equilibrium
The green line
represents the
minimum energy
required to break
the intermolecular
attractions.
Even at the cold
temp, some
molecules can be
converted.
energy + H2O(l)  H2O(g)
• Once there are molecules in the vapor phase, they can
be converted back to the liquid phase
H2O(g)  H2O(l) + energy
• evaporation - the process of conversion of liquid to gas,
at a temperature too low to boil
• condensation - conversion of the gas to the liquid state.
• When the rate of evaporation equals the rate of
condensation, the system is at equilibrium.
• Vapor pressure of a liquid - the pressure exerted by the
vapor at equilibrium
H2O(g)
H2O(l)
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 DHfus.
For ice, the heat of fusion is 6.01 kJ/mol.
H 2O(s )  H 2O(l );
DH fus  6.01 kJ
Heat of Phase Transition
• To boil a pure substance at its melting
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 DHvap.
For ice, the heat of vaporization is 40.66 kJ/mol.
H 2O(l )  H 2O(g );
DH vap  40.66 kJ
Heating Curve
A Problem to Consider
• The heat of vaporization of ammonia is
23.4 kJ/mol. How much heat is required to
vaporize 1.00 kg of ammonia?
First, we must determine the number of
moles of ammonia in 1.00 kg (1000 g).
1 mol NH 3
1.00  10 g NH 3 
 58.8 mol NH 3
17.0 g NH 3
3
A Problem to Consider
• The heat of vaporization of ammonia is
23.4 kJ/mol. How much heat is required to
vaporize 1.00 kg of ammonia?
Then we can determine the heat required
for vaporization.
58.8 mol NH 3  23.4 kJ/mol  1.38  10 kJ
3
Molar heat of vaporization (DHvap) is the energy required to
vaporize 1 mole of a liquid at its boiling point.
Clausius-Clapeyron Equation
DHvap
ln P = +C
RT
P = (equilibrium) vapor pressure
T = temperature (K)
R = gas constant (8.314 J/K•mol)
Vapor Pressure Versus Temperature
Clausius-Clapeyron Equation
• We noted earlier 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.
Consequently, the vapor pressure of a
liquid at two different temperatures is
described by:
P2 DH vap 1 1
ln 
P1
R
T1 T2
(
)
A Problem to Consider
• Carbon disulfide, CS2, has a normal boiling point
of 46oC (vapor pressure = 760 mmHg) and a heat
of vaporization of 26.8 kJ/mol. What is the vapor
pressure of carbon disulfide at 35oC?
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
A Problem to Consider
• Carbon disulfide, CS2, has a normal boiling point
of 46oC (vapor pressure = 760 mmHg) and a heat
of vaporization of 26.8 kJ/mol. What is the vapor
pressure of carbon disulfide at 35oC?
Taking the antiln we obtain:
P2
 antiln(-0.361)
(760 mm Hg)
P2  antiln(-0.361)  760 mm Hg
P2  530 mm Hg
The boiling point is the temperature at which the
(equilibrium) vapor pressure of a liquid is equal to the
external pressure.
The normal boiling point is the temperature at which a liquid
boils when the external pressure is 1 atm.
The Critical Phenomenon of SF6
T < Tc
T > Tc
T ~ Tc
T < Tc
Phase Diagrams
• A phase diagram is a graphical way to
summarize the conditions under which the
different states of a substance are stable.
The diagram is divided into three areas
representing each state of the substance.
The curves separating each area
represent the boundaries of phase
changes.
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
C
solid
D
liquid
.
A
temperature
gas
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
C
solid
liquid
.
D
A
temperature
gas
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
C
solid
liquid
.
D
A
temperature
gas
Phase Diagrams
• If a liquid is more dense than its solid, the
curve leans slightly to the left, causing the
melting point to decrease with pressure.
.
pressure
B
C
solid
liquid
.
D
A
temperature
gas
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.
.
pressure
B
C
solid
liquid
.
D
A
temperature
gas
Phase Diagrams
• Curve AC, which divides the liquid region
from the gaseous region, represents the boiling
points of the liquid for various pressures.
.
pressure
B
C
solid
liquid
.
D
A
temperature
gas
Phase Diagrams
• Curve AD, which divides the solid region from
the gaseous region, represents the vapor
pressures of the solid at various temperatures.
.
pressure
B
C
solid
liquid
.
D
A
temperature
gas
Phase Diagrams
• The curves intersect at A, the triple point,
which is the temperature and pressure where
three phases of a substance exist in equilibrium.
.
pressure
B
C
solid
liquid
.
D
A
temperature
gas
Phase Diagrams
• The curves intersect at A, the triple point,
which is the temperature and pressure where
three phases of a substance exist in equilibrium.
.
pressure
B
C
solid
liquid
.
D
A
temperature
gas
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
C
solid
liquid
.
D
A
temperature
gas
Tcrit
Phase Diagrams
• The vapor pressure at the critical temperature is
called the critical pressure. Note that curve AC
ends at the critical point, C.
.
B
pressure
Pcrit
C
solid
liquid
.
D
A
temperature
gas
Tcrit
The critical temperature (Tc) is the temperature above which the
gas cannot be made to liquefy, no matter how great the applied
pressure.
The critical pressure
(Pc) is the minimum
pressure that must be
applied to bring about
liquefaction at the
critical temperature.
A phase diagram summarizes the conditions at which a
substance exists as a solid, liquid, or gas.
Phase Diagram of Water
Phase Diagram of Carbon Dioxide
At 1 atm
CO2 (s)
CO2 (g)
Effect of Increase in Pressure on the Melting Point
of Ice and the Boiling Point of Water
Chemistry In Action: Liquid Crystals
WORKED
EXAMPLES
Worked Example 11.1
Worked Example 11.3a
Worked Example 11.3b
Worked Example 11.5
Worked Example 11.8