Intermolecular Forces, Liquids, and Solids
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Intermolecular Forces, Liquids, and Solids
Some Characteristic Properties of the States of Matter
Assumes both the volume and shape of the container
Is Compressible
Diffusion within a gas occurs rapidly
Flows readily
Liquid Assumes the shape of the portion of the container it occupie
Does not expand to fill the container
Is virtually incompressible
Diffusion within a gas occurs slowly
Flows readily
Solid Retains its own shape and volume
Is virtually incompressible
Diffusion within a solid occurs extremely slowly
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Does not flow
Gas:
The strength of intermolecular forces vary over a wide range,
but are generally weaker than covalent or ionic bonds
16 kJ will overcome the intermolecular attraction leading to vaporization
431 kJ of energy is required to break the covalent bond between HCl3
Dipole Dipole forces exist between neutral polar molecules, are
effective only when polar molecules are very close together, and are
generally weaker then ion-dipole interactions
For two particles of
equal mass and size,
the strengths of the
intermolecular
attractions increase
with increasing
polarity
Substance Dipole moment Bp
Propane
0.1 231
Dimethyl ether 1.3 249
Methyl chloride 2.0 249
Acetaldehyde
2.7 293
Molecules that are attracting one another spend more time near each other
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than those molecules which repel. Thus the overall effect is a net attraction
Ion-Dipole Forces exist between an ion and the partial
charge on the end of a polar molecule
Na+
Cl-
The magnitude of the interaction depends upon the charge of the
ion, the dipole moment of the polar molecule and the distance from
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the center of the ion to the midpoint of the dipole
London Dispersion Forces can exist between non-polar atoms and
molecules where the movement of electrons create an instantaneous
dipole moment
The instantaneous distribution at any given moment can be
different from the average distribution producing an
instantaneous dipole moment
Because electrons repel, the motions of electrons
in one atom influence
the motions of electrons on its near neighbors
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London Dispersion Forces can exist between non-polar atoms and
molecules where the movement of electrons create an instantaneous
dipole moment
The ease with which the charge distribution in a molecule can be
distorted by an external force is called its polarizability. The greater
the polarizability of a molecule, the more easily its electron cloud can
be distorted to give a momentary dipole.
In general, larger molecules tend to have greater polarizability
because their electrons are farther from the nucleus
Because molecular size and molecular mass tend to parallel one
another, dispersion forces tend to increase in strength with
increasing molecular weight.
Halogen Boiling pt (K) Noble gas Boiling point (K)
F2
85.1
He
4.6
Cl2
238.6
Ne
27.3
Br2
332.0
Ar
87.5
I2
457.6
Kr
120.9
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London Dispersion Forces can exist between non-polar atoms and
molecules where the movement of electrons create an instantaneous
dipole moment
The shapes of molecules can play a role in the magnitude of dispersion forces.
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London Dispersion Forces can exist between non-polar atoms and
molecules where the movement of electrons create an instantaneous
dipole moment
London dispersion forces may also operate within polar molecules
and contribute to the overall attractive forces between the molecules;
some times more than the dipole-dipole forces
HBr: bp = 206.2 K
HCl: bp = 189.5 K
Dipole Moment = 1.03 D Dipole Moment = 0.79 D
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Hydrogen bonding is a special type of intermoleculer attraction that
exists between the hydrogen atom in a polar bond (e.g., H F, H O,
or H N) and an unshared electron pair on a nearby electronegative
atom (usually an F, O, or N atom on another molecule)
Because F, N, and O are so electronegative, a bond between hydrogen10and
any of these elements is quite polar, with hydrogen at the positive end.
Metallic solids consist of metal atoms. Bonding is due to valence
electrons that are delocalized throughout the entire solid. We can
visualize these electrons as an array of positive ions immersed in a
sea of delocalized electrons.
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Intermolecular Forces, Liquids, and Solids
Metallic bond
examples: Cu,
Fe, or alloy
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Intermolecular Forces, Liquids, and Solids
Properties of Liquids: Viscosity and Surface Tension
The resistance of liquids to flow is called their viscosity. The greater the viscosity, the
more slowly the liquid flows. Viscosity decreases with increasing temperature.
Viscosity is related to the ease with which individual molecules of a liquid can move
with respect to one another. It thus depends on the attractive forces between molecules
and and their structural character.
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Intermolecular Forces, Liquids, and Solids
Properties of Liquids: Viscosity and Surface Tension
Molecules in the interior are attracted equally in
all directions, whereas those at the surface
experience a net inward force
This net inward force pulls molecules from the
surface into interior, thereby reducing the surface
area
Surface tension is therefore defined as the energy
required to increase the surface of a liquid by a
unit amount
Forces that bind like molecules to one another are
called cohesive forces.
Forces that bind molecules to a surface are called
adhesive forces
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Changes of State
Whenever a change of state involves going to a less ordered state, energy
must be supplied in order to overcome intermolecular forces
ENDOTHERMIC
The heat needed for vaporization is called the heat of vaporization
(e.g. 40.67 kJ/mol for water)
Melting is also called fusion. The enthalpy of change associated15with
melting is called heat of fusion. (e.g. 6.01 kJ/mol for water)
Changes of State
Enthalpy and Temperature Changes Accompanying Heating
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Calculate the enthalpy change associated with converting 1.00 mol
of ice at -25°C to water vapor at 125 °C and 1 atm. The heat
capacities per gram (specific heats) of ice, water, and steam are 2.09
J/g-°C , 4.18 J/g-°C , and 1.84 J/g-°C, respectively. The heat of
fusion of ice is 6.01 kJ/mol, and the heat of vaporization of water is
40.67 kJ/mol
•Heat required to bring the
ice from -25 °C to 0 °C :
•Heat of fusion:
•Heat required to bring liquid
• water from 0 °C to 100 °C :
Heat of vaporization
Heat required to raise the
temperature of the vapor
to 125 °C
1.00 mol
H2O
1.00 mol
H2O
1.00 mol
H2O
1.00 mol
H2O
1.00 mol
H2O
18.0 g.
1.00 mol
H2O
6.01 kJ
mol H2O
18.0 g.
1.00 mol
H2O
40.67 kJ
mol H2O
18.0 g.
1.00 mol
H2O
2.09 J
g -°C
25 °C = 940 J
= 6.01 kJ
4.18 J
g -°C
100 °C =7520 J
= 40.67 kJ
1.84 J
g -°C
25°C
=830 J
Ht = .940 kJ + 6.01 kJ + 7.52 kJ + 40.67 kJ + .830 kJ = 55.97 kJ17
Changes of State
Vapor Pressure: pressure exerted by evaporating liquid in the space
above that liquid once dynamic equilibrium has been established
Substances with a high vapor pressure evaporate more rapidly than
thosewith low vapor pressures. Theses substances are more volatile
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The vapor pressure of a substance increases as the temperature increases
Changes of State
Vapor Pressure and Boiling Point: A liquid boils when its vapor
pressure equals the external pressure acting on the surface of the liquid
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The higher the outside pressure, the higher the boiling point
A phase diagram is a graphical way to summarize the conditions
under which equilibria exist between the difference states of matter.
Line A-D represents the
change in melting point
of the solid under
increased pressure
Line A-B is the vapor
curve of the liquid
The line A-C represents the variation in
vapor pressure during sublimation.
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A phase diagram is a graphical way to summarize the conditions
under which equilibria exist between the difference states of matter.
Note that the solid phase is the
stable phase under condition of
high pressure and low temperature
Note that the gas phase is the
stable phase under condition of
low pressure and high temperature
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Why are the fusion-freezing lines different for H2O and CO2?
Why does CO2 sublime under normal conditions?
How is freeze-dry food prepared?
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AS THE TEMPERATURE IS INCREASED : At point 6-A, the H2O exists entirely as a
solid. When the temperature reaches point 4-B, the solid begins to melt, and an
equilibrium condition occurs between the solid and the liquid. At a yet high
temperature, point 7-C, the solid has been converted entirely to a liquid. When point
8-D is encountered, vapor forms, and a liquid-vapor equilibrium is achieved. Upon
further heating, to point 9-E, the H2O is converted entirely to the vapor phase.
5
6
4
77
8 9
3
2
1
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A crystalline solid is a solid whose atoms, ions, or molecules are
ordered in well-defined arrangements. Amorphous solids have no
structure.
A unit cell is the repeating unit that makes up the crystalline
solid. The array of repeating points is called the crystal lattice
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The lattices of crystalline solids can be described in terms of seven
basic types. The following three types are the simplest of the seven.
Coordination number = # of nearest neighbors
# of atoms in unit cell
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The Crystalline (Body-Centered) Lattice of Sodium Chloride
Na+
Cl-
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The total cation-anion ratio in each unit cell must be the same as
that for the entire crystal. Therefore a unit cell of sodium chloride
must have equal number of Na+ and Cl-
(¼ Na+ per edge ) (12 edges) = 3 Na+
Na+
(1 Na+ per center) (1 center ) = 1 Na+
(1/8 Cl- per corner ) (8 corners) = 1 ClCl(1/2 Cl- per face) (6 faces ) = 3 Cl-
Na+
Cl-
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Aluminum crystallizes in a facecentered arrangement
a.how many atoms are in the unit
cell?
b.what is the coordination
number?
c.each atom has diameter of 2.86
Å, what is the length of a side?
d. what is the density?
°
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The Close Packing of Spheres: A Study of Metallic Solids
Lets make an assumption using Methane as a example that many
molecules may be approximated as being roughly spherical 29
The Close Packing of Spheres: A Study of Metallic Solids
The coordination number of close packing structures = 12
hexagonal close pack
cubic close pack =
face-centered cubic
The coordination number of body-centered cubic structures = 8
The coordination number of primitive cubic structures = 6
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Using X-Ray Diffraction to Determine the Structure of Crystalline Solids
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Covalent Network Solids are Held together in large networks or
chains by covalent bonds. These solids are much harder and have
higher melting points than molecular solids
Molecular solids consist of atoms or molecules held together by
dipole-dipole or London dispersion forces or hydrogen bonds 32
Ionic solids consist of ions held together by ionic bonds. The strength
of the bonds depend upon the charges of the ions. Melting point
increases as bond strength increases
NaCl
The structure of the solid is dependent upon the charges and
relative sizes of the ions.
The coordination number increases as the ratio of the cation
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radius to the anion radius increases.
• Atomic solids – (london) eg. He, Ar etc.
atoms exist at lattice pts
• molecular solids (London, dipole, H-Bond) CH4,
NH3
molecules exist at lattice pts
• Ionic solids (electrostatic attractions) NaCl
cations & anions exist at lattice pts
• Metallic solids- (Cations held by a sea of
electrons) Cu, Fe
cations exist at lattice pts
• Network solids – (3 dimensional covalent
bonds) diamond (C), quartz (SiO2)
atoms exist at lattice pts 34