States of Matter and Intermolecular Forces

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States of Matter and Intermolecular Forces
Chapter Eleven
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Chapter Preview

• Intramolecular forces (bonds) govern molecular
  properties such as molecular geometries and
  dipole moments.
• Intermolecular forces determine the macroscopic
  physical properties of liquids and solids.
• This chapter
• describes changes from one state of matter to
  another.
• explores the types of intermolecular forces that
  underlie these and other physical properties of
  substances.

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Molecular Forces Compared
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States of Matter Compared
Intermolecular forces are very important.
Intermolecular forces are of little significance
why?
Intermolecular forces must be considered.
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Vaporization and Condensation

• Vaporization is the conversion of a liquid to a
  gas.
• The enthalpy of vaporization (DHvapn) is the
  quantity of heat that must be absorbed to
  vaporize a given amount of liquid at a constant
  temperature.
• Condensation is the reverse of vaporization. The
  enthalpy of condensation (DHcondn) accompanies
  this change of a gas to a liquid.
• Enthalpy is a function of state therefore, if a
  liquid is vaporized and the vapor condensed at
  constant temperature, the total DH must be zero
• DHvapn DHcondn 0
• DHcondn DHvapn

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• Example 11.1
• How much heat, in kilojoules, is required to
  vaporize 175 g methanol, CH3OH, at 25 C?
• Example 11.2 An Estimation Example
• Without doing detailed calculations, determine
  which liquid in Table 11.1 requires the greatest
  quantity of heat for the vaporization of 1 kg of
  liquid.

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Vapor Pressure

• The vapor pressure of a liquid is the partial
  pressure exerted by the vapor when it is in
  dynamic equilibrium with the liquid at a constant
  temperature.

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LiquidVapor Equilibrium
until equilibrium is attained.
More vapor forms rate of condensation of that
vapor increases
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Vapor pressure increases with temperature why?
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Vapor Pressure Curves
What is the vapor pressure of H2O at 100 C,
according to this graph? What is the significance
of that numeric value of vapor pressure?
Which of the five compounds has the strongest
intermolecular forces? How can you tell?
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• Example 11.3
• Suppose the equilibrium illustrated in Figure
  11.3 is between liquid hexane, C6H14, and its
  vapor at 298.15 K. A sample of the equilibrium
  vapor is found to have a density of 0.701 g/L.
  What is the vapor pressure of hexane, in mmHg, at
  298.15 K?

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Boiling Point and Critical Point

• Boiling point the temperature at which the vapor
  pressure of the liquid equals the external
  pressure.
• Normal boiling point boiling point at 1 atm.
• Critical temperature (Tc) the highest
  temperature at which a liquid can exist.
• The critical pressure, Pc, is the vapor pressure
  at the critical temperature.
• The condition corresponding to a temperature of
  Tc and a pressure of Pc is called the critical
  point.

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The Critical Point
At Tc, the densities of liquid and vapor are
equal a single phase.
At room temperature there is relatively little
vapor, and its density is low.
At higher temperature, there is more vapor, and
its density increases
while the density of the liquid decreases
molecular motion increases.
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These four gases cant be liquefied at room
temperature, no matter what pressure is applied
why not?
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• Example 11.4 A Conceptual Example
• To keep track of how much gas remains in a
  cylinder, we can weigh the cylinder when it is
  empty, when it is full, and after each use. In
  some cases, though, we can equip the cylinder
  with a pressure gauge and simply relate the
  amount of gas to the measured gas pressure.
  Which method should we use to keep track of the
  bottled propane, C3H8, in a gas barbecue?

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Phase Changes Involving Solids

• Melting (fusion) transition of solid ? liquid.
• Melting point temperature at which melting
  occurs.
• Same as freezing point!
• Enthalpy of fusion, DHfusion, is the quantity of
  heat required to melt a set amount (one gram, one
  mole) of solid.
• Sublimation transition of solid ? vapor.
• Example Ice cubes slowly disappear in the
  freezer.
• Enthalpy of sublimation, DHsubln, is the sum of
  the enthalpies of fusion and vaporization.
• Triple point all three phasessolid, liquid,
  vaporare in equilibrium.

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Some Enthalpies of Fusion
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Cooling Curve for Water
Once all of the liquid has solidified, the
temperature again drops.
The liquid water cools until
the freezing point is reached, at which time
the temperature remains constant as solid forms.
If the liquid is cooled carefully, it can
supercool.
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Heating Curve for Water
until the solid begins to melt, at which time
the temperature remains constant
The temperature of the solid increases as it is
heated
until all the solid is melted, at which time
the temperature again rises.
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Phase Diagrams
AD, solid-liquid equilibrium.
AC, liquid-vapor equilibrium.

• A phase diagram is a graphical representation of
  the conditions of temperature and pressure under
  which a substance exists as a solid, liquid, a
  gas, or some combination of these in equilibrium.

AB, solid-vapor equilibrium.
Triple point
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Phase Diagram for HgI2
HgI2 has two solid phases, red and yellow.
As the vessel is allowed to cool, will the
contents appear more red or more yellow?
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Phase Diagram for CO2
Note that at 1 atm, only the solid and vapor
phases of CO2 exist.
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Phase Diagram for H2O
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• Example 11.5 A Conceptual Example
• In Figure 11.14, 50.0 mol H2O(g) (steam) at 100.0
  C and 1.00 atm is added to an insulated cylinder
  that contains 5.00 mol H2O(s) (ice) at 0 C. With
  a minimum of calculation, use the data provided
  to determine which of the following will describe
  the final equilibrium condition

(a) ice and liquid water at 0 C, (b) liquid
water at 50 C, (c) steam and liquid water at
100 C, (d) steam at 100 C. Data you will need
are ?Hfusion 6.01 kJ/mol, ?Hvapn 40.6 kJ/mol
(at 100 C), molar heat capacity of H2O(l) 76 J
mol1 C1.
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Intermolecular Forces

• are forces between molecules.
• They determine melting points, freezing points,
  and other physical properties.
• Types of intermolecular forces include
• dispersion forces
• Dipoledipole forces
• hydrogen bonding

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Dispersion Forces

• exist between any two particles.
• Also called London forces (after Fritz London,
  who offered a theoretical explanation of these
  forces in 1928).
• Dispersion forces arise because the electron
  cloud is not perfectly uniform.
• Tiny, momentary dipole moments can exist even in
  nonpolar molecules.

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Dispersion Forces Illustrated (1)
At a given instant, electron density, even in a
nonpolar molecule like this one, is not perfectly
uniform.
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Dispersion Forces Illustrated (2)
the other end of the molecule is slightly ().
The region of (momentary) higher electron density
attains a small () charge
When another nonpolar molecule approaches
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Dispersion Forces Illustrated (3)
this molecule induces a tiny dipole moment
in this molecule.
Opposite charges ________.
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Strength of Dispersion Forces

• Dispersion force strength depends on
  polarizability the ease with which the electron
  cloud is distorted by an external electrical
  field.
• The greater the polarizability of molecules, the
  stronger the dispersion forces between them.
• Polarizability in turn depends on molecular size
  and shape.
• Heavier molecule gt more electrons gt a more-
  polarizable molecule.
• As to molecular shape

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Molecular Shape and Polarizability
can have greater separation of charge along its
length. Stronger forces of attraction, meaning
Long skinny molecule
higher boiling point.
giving weaker dispersion forces and a lower
boiling point.
In the compact isomer, less possible separation
of charge
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DipoleDipole Forces

• A polar molecule has a positively charged end
  (d) and a negatively charged end (d).
• When molecules come close to one another,
  repulsions occur between like-charged regions of
  dipoles. Opposite charges tend to attract one
  another.
• The more polar a molecule, the more pronounced is
  the effect of dipoledipole forces on physical
  properties.

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DipoleDipole Interactions
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Predicting Physical Properties of Molecular
Substances

• Dispersion forces become stronger with increasing
  molar mass and elongation of molecules. In
  comparing nonpolar substances, molar mass and
  molecular shape are the essential factors.
• Dipoledipole and dipole-induced dipole forces
  are found in polar substances. The more polar the
  substance, the greater the intermolecular force
  is expected to be.
• Because they occur in all substances, dispersion
  forces must always be considered. Often they
  predominate.

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• Example 11.6
• Arrange the following substances in the expected
  order of increasing boiling point carbon
  tetrabromide, CBr4 butane, CH3CH2CH2CH3
  fluorine, F2 acetaldehyde, CH3CHO.

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Homology

• A series of compounds whose formulas and
  structures vary in a regular manner also have
  properties that vary in a predictable manner.
• This principle is called homology.
• Example both densities and boiling points of the
  straight-chain alkanes increase in a continuous
  and regular fashion with increasing numbers of
  carbon atoms in the chain.
• Trends result from the regular increase in molar
  mass, which produces a fairly regular increase in
  the strength of dispersion forces.

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• Example 11.7 An Estimation Example
• The boiling points of the straight-chain alkanes
  pentane, hexane, heptane, and octane are 36.1,
  68.7, 98.4, and 125.7 C, respectively. Estimate
  the boiling point of the straight-chain alkane
  decane.

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Hydrogen Bonds

• A hydrogen bond is an intermolecular force in
  which
• a hydrogen atom that is covalently bonded to a
  (small, electronegative) nonmetal atom in one
  molecule
• is simultaneously attracted to a (small,
  electronegative) nonmetal atom of a neighboring
  molecule.

Y H - - - Z
When Y and Z are small and highly electronegative
(N, O, F)
this force is called a hydrogen bond a
special, strong type of dipoledipole force.
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Hydrogen Bonds in Water
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Hydrogen Bonding in Ice
Hydrogen bonding arranges the water molecules
into an open hexagonal pattern.
Hexagonal is reflected in the crystal
structure. Open means reduced density of the
solid (vs. liquid).
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Hydrogen Bonding in Acetic Acid
Hydrogen bonding occurs between molecules.
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Hydrogen Bonding inSalicylic Acid
Hydrogen bonding occurs within the molecule.
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Intermolecular Hydrogen Bonds
Intermolecular hydrogen bonds give proteins their
secondary shape, forcing the protein molecules
into particular orientations, like a folded sheet

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Intramolecular Hydrogen Bonds
while intramolecular hydrogen bonds can cause
proteins to take a helical shape.
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• Example 11.8
• In which of these substances is hydrogen bonding
  an important intermolecular force N2, HI, HF,
  CH3CHO, and CH3OH? Explain.

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Liquids and Intermolecular Forces

• Much behavior and many properties of liquids can
  be attributed to intermolecular forces.
• Surface tension (g) is the amount of work
  required to extend a liquid surface and is
  usually expressed in J/m2.
• Adhesive forces are intermolecular forces between
  unlike molecules.
• Cohesive forces are intermolecular forces between
  like molecules.
• A meniscus is the interface between a liquid and
  the air above it.
• Viscosity is a measure of a liquids resistance
  to flow.

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Surface Tension
To create more surface, the molecules at the
surface must be separated from one another.
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Adhesive and Cohesive Forces
The liquid spreads, because adhesive forces are
comparable in strength to cohesive forces.
The liquid beads up. Which forces are stronger,
adhesive or cohesive?
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Meniscus Formation
What conclusion can we draw about the cohesive
forces in mercury?
Water wets the glass (adhesive forces) and its
attraction for glass forms a concave-up surface.
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Capillary Action
The adhesive forces wet more and more of the
inside of the tiny tube, drawing water farther up
the tube.
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Comparing Viscosity
Which oil flows more readily? Which oil has
stronger intermolecular forces between its
molecules? Oil is mostly hydrocarbons what kind
of forces are these?
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Types of Solids

• Amorphous solids have no significant long-range
  order.
• Crystalline solids have atoms/ions/molecules
  arranged in a regular pattern. Types of
  crystalline solids include
• Molecular solids, containing molecules held to
  one another by dispersion/dipoledipole/ hydrogen
  bonding forces.
• Network (covalent) solids.
• Ionic solids.
• Metallic solids (metals).

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Network Covalent Solids

• Network solids have a network of covalent bonds
  that extend throughout the solid, holding it
  firmly together.
• The allotropes of carbon provide good examples.
• Diamond has each carbon bonded to four other
  carbons in a tetrahedral arrangement using sp3
  hybridization.
• Graphite has each carbon bonded to three other
  carbons in the same plane using sp2
  hybridization.
• Fullerenes are roughly spherical collections of
  carbon atoms in the shape of a soccer ball.
• A nanotube can be thought of as a plane of
  graphite rolled into a tube.

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Crystal Structure of Diamond
Three-dimensional network is extremely strong,
rigid.
What kind of forces must be overcome to melt
diamond?
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Crystal Structure of Graphite
Hexagons of sp2-hybridized carbon atoms.
Forces between layers are relatively weak.
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Structure of a Buckyball
C60 molecule
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Structure of a Nanotube
A nanotube can be thought of as a sheet of
graphite, rolled into a tube, capped with half of
a buckyball.
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Ionic Bonds asIntermolecular Forces

• There are no molecules in an ionic solid, and
  therefore there cant be any intermolecular
  forces.
• The attractions are electrostatic interionic
  attractions.
• Lattice energy (Chapter 9) is a measure of the
  strength of interionic attraction.
• The attractive force between a pair of oppositely
  charged ions increases
• as the charges on the ions increase.
• as the ionic radii decrease.
• Lattice energies increase accordingly.

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Interionic Forces of Attraction
Mg2 and O2 have much stronger forces of
attraction for one another than do Na and Cl.
Melting point of MgO is about 2800 oC.
Melting point of NaCl is about 801 oC.
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• Example 11.9
• Arrange the ionic solids MgO, NaBr, and NaCl in
  the expected order of increasing melting point.

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Crystal Lattices

• To describe crystals, three-dimensional views
  must be used.
• The repeating unit of the lattice is called the
  unit cell.
• There are a number of different types of unit
  cell hexagonal, rhombic, cubic, etc.
• The three types of cubic unit cells are simple
  cubic, body-centered cubic (bcc), face-centered
  cubic (fcc).

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Cubic Unit Cells
The unit cell is a cube in each case.
Whole atoms shown for clarity.
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Close-Packed Structures
A close packed structure in two dimensions.
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Close Packing in Three Dimensions
Third layer directly above first layer HCP
Two layers, stacked, give two different locations
for the third layer
Third layer over the octahedral holes in the
second layer CCP
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• Example 11.10
• Copper crystallizes in the cubic close-packed
  arrangement. The atomic (metallic) radius of a Cu
  atom is 127.8 pm. (a) What is the length, in
  picometers, of the unit cell in a sample of
  copper metal? (b) What is the volume of that unit
  cell, in cubic centimeters? (c) How many atoms
  belong to the unit cell?

Example 11.11 Use the results of Example 11.10,
the molar mass of copper, and Avogadros number
to calculate the density of metallic copper.
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Ionic Crystal Structures

• Ionic crystals have two different types of
  structural unitscations and anions.
• The cations and anions ordinarily are different
  sizes.
• Smaller cations can fill the voids between the
  larger anions.
• Where the cations go depends on the size of the
  cations and on the size of the voids.
• The smallest voids are the tetrahedral holes,
  then the octahedral holes, and finally the holes
  in a cubic structure. Therefore

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Ionic Crystal Structures

• Tetrahedral hole filling occurs when the cations
  are small when the radii ratio is
• 0.225 lt rc/ra lt 0.414
• Octahedral filling occurs with larger cations,
  when the radii ratio is
• 0.414 lt rc/ra lt 0.732
• The arrangement is cubic if rc/ra gt 0.732.

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Unit Cell of Cesium Chloride
How many cesium ions are inside this unit cell?
How many chloride ions?
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Unit Cell of Sodium Chloride
How many sodium ions are inside this unit cell?
How many chloride ions?
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Experimental Determinationof Crystal Structures
X rays
Angle of diffraction can be used to find distance
d, using simple trigonometry.
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• Cumulative Example
• Here are some data about an organic compound
  Its normal boiling point is a few degrees below
  that of H2O(l), and its vapor density at 99.0 C
  and 752 Torr is within 5 of 2.0 g/L. Using these
  data and other information from this and previous
  chapters, determine which one of the following
  compounds it is most likely to be
• (CH3)2O
• CH3CH2CH2OH
• CH3CH2OCH3
• HOCH2CH2OH