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States of Matter and Intermolecular Forces

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Title: States of Matter and Intermolecular Forces


1
States of Matter and Intermolecular Forces
Chapter Eleven
2
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.

3
Molecular Forces Compared
4
States of Matter Compared
Intermolecular forces are very important.
Intermolecular forces are of little significance
why?
Intermolecular forces must be considered.
5
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

6
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7
  • 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.

8
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.

9
LiquidVapor Equilibrium
until equilibrium is attained.
More vapor forms rate of condensation of that
vapor increases
10
Vapor pressure increases with temperature why?
11
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?
12
  • 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?

13
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.

14
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.
15
These four gases cant be liquefied at room
temperature, no matter what pressure is applied
why not?
16
  • 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?

17
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.

18
Some Enthalpies of Fusion
19
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.
20
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.
21
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
22
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?
23
Phase Diagram for CO2
Note that at 1 atm, only the solid and vapor
phases of CO2 exist.
24
Phase Diagram for H2O
25
  • 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.
26
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

27
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.

28
Dispersion Forces Illustrated (1)
At a given instant, electron density, even in a
nonpolar molecule like this one, is not perfectly
uniform.
29
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
30
Dispersion Forces Illustrated (3)
this molecule induces a tiny dipole moment
in this molecule.
Opposite charges ________.
31
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

32
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
33
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.

34
DipoleDipole Interactions
35
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.

36
  • Example 11.6
  • Arrange the following substances in the expected
    order of increasing boiling point carbon
    tetrabromide, CBr4 butane, CH3CH2CH2CH3
    fluorine, F2 acetaldehyde, CH3CHO.

37
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.

38
  • 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.

39
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.
40
Hydrogen Bonds in Water
41
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).
42
Hydrogen Bonding in Acetic Acid
Hydrogen bonding occurs between molecules.
43
Hydrogen Bonding in Salicylic Acid
Hydrogen bonding occurs within the molecule.
44
Intermolecular Hydrogen Bonds
Intermolecular hydrogen bonds give proteins their
secondary shape, forcing the protein molecules
into particular orientations, like a folded sheet

45
Intramolecular Hydrogen Bonds
while intramolecular hydrogen bonds can cause
proteins to take a helical shape.
46
  • Example 11.8
  • In which of these substances is hydrogen bonding
    an important intermolecular force N2, HI, HF,
    CH3CHO, and CH3OH? Explain.

47
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.

48
Surface Tension
To create more surface, the molecules at the
surface must be separated from one another.
49
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?
50
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.
51
Capillary Action
The adhesive forces wet more and more of the
inside of the tiny tube, drawing water farther up
the tube.
52
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?
53
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).

54
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.

55
Crystal Structure of Diamond
Three-dimensional network is extremely strong,
rigid.
What kind of forces must be overcome to melt
diamond?
56
Crystal Structure of Graphite
Hexagons of sp2-hybridized carbon atoms.
Forces between layers are relatively weak.
57
Structure of a Buckyball
C60 molecule
58
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.
59
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60
Ionic Bonds as Intermolecular 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.

61
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.
62
  • Example 11.9
  • Arrange the ionic solids MgO, NaBr, and NaCl in
    the expected order of increasing melting point.

63
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).

64
Cubic Unit Cells
The unit cell is a cube in each case.
Whole atoms shown for clarity.
65
Close-Packed Structures
A close packed structure in two dimensions.
66
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
67
  • 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.
68
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

69
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.

70
Unit Cell of Cesium Chloride
How many cesium ions are inside this unit cell?
How many chloride ions?
71
Unit Cell of Sodium Chloride
How many sodium ions are inside this unit cell?
How many chloride ions?
72
Experimental Determination of Crystal Structures
X rays
Angle of diffraction can be used to find distance
d, using simple trigonometry.
73
  • 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
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