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Chapter 11 Intermolecular Forces

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Title: Chapter 11 Intermolecular Forces


1
Chapter 11 Intermolecular Forces
2
11.1 Intermolecular Forces (IMF)
  • IMF lt intramolecular forces (covalent, metallic,
    ionic bonds)
  • IMF strength solids gt liquids gt gases
  • Boiling points and melting points are good
    indicators of relative IMF strength.

3
11.2 Types of IMF
  • Electrostatic forces act over larger distances
    in accordance with Coulombs law
  • Ion-ion forces strongest found in ionic
    crystals (i.e. lattice energy)

4
  1. Ion-dipole between an ion and a dipole (a
    neutral, polar molecule/has separated partial
    charges)
  • Increase with increasing polarity of molecule and
    increasing ion charge.

Ex Compare IMF in Cl- (aq) and S2- (aq).
lt
5
  1. Dipole-dipole weakest electrostatic force exist
    between neutral polar molecules
  • Increase with increasing polarity (dipole moment)
    of molecule

Ex What IMF exist in NaCl (aq)?
6
  • Hydrogen bonds (or H-bonds)
  • H is unique among the elements because it has a
    single e- that is also a valence e-.
  • When this e- is hogged by a highly EN atom (a
    very polar covalent bond), the H nucleus is
    partially exposed and becomes attracted to an
    e--rich atom nearby.

7
  • H-bonds form with H-XX', where X and X' have
    high EN and X' possesses a lone pair of e-
  • X F, O, N (since most EN elements) on two
    molecules

F-H O-H N-H
F O N
8
  • There is no strict cutoff for the ability to
    form H-bonds (S forms a biologically important
    hydrogen bond in proteins).
  • Hold DNA strands together in double-helix

Nucleotide pairs form H-bonds
DNA double helix
9
  • H-bonds explain why ice is less dense than water.

10
Ex Boiling points of nonmetal hydrides
  • Conclusions
  • Polar molecules have higher BP than nonpolar
    molecules
  • ? Polar molecules have stronger IMF
  • BP increases with increasing MW
  • ? Heavier molecules have stronger IMF

Boiling Points (ºC)
  • NH3, H2O, and HF have unusually high BP.
  • ? H-bonds are stronger than dipole-dipole IMF

11
Inductive forces
  • Arise from distortion of the e- cloud induced by
    the electrical field produced by another particle
    or molecule nearby.
  • London dispersion between polar or nonpolar
    molecules or atoms
  • Proposed by Fritz London in 1930
  • Must exist because nonpolar molecules form liquids

Fritz London(1900-1954)
12
  • How they form
  • Motion of e- creates an instantaneous dipole
    moment, making it temporarily polar.
  • Instantaneous dipole moment induces a dipole in
    an adjacent atom
  • Persist for about 10-14 or 10-15 second
  • Ex two He atoms

13
Geckos!
  • Geckos feet make use of London dispersion forces
    to climb almost anything.
  • A gecko can hang on a glass surface using only
    one toe.
  • Researchers at Stanford University recently
    developed a gecko-like robot which uses synthetic
    setae to climb walls

http//en.wikipedia.org/wiki/Van_der_Waals27_forc
e
14
  • London dispersion forces increase with
  • Increasing MW, of e-, and of atoms
    (increasing of e- orbitals to be distorted)
  • Boiling points
  • Effect of MW Effect of atoms
  • pentane 36ºC Ne 246C
  • hexane 69ºC CH4   162C
  • heptane 98ºC
  • ??? effect
  • H2O 100C
  • D2O 101.4C
  • Longer shapes (more likely to interact with
    other molecules)
  • C5H12 isomers 2,2-dimethylpropane 10C
  • pentane 36C

15
Summary of IMF
Van der Waals forces
16
Ex Identify all IMF present in a pure sample of
each substance, then explain the boiling points.
BP(C) IMF Explanation
HF 20
HCl -85
HBr -67
HI -35
Lowest MW/weakest London, but most polar/strongest dipole-dipole and has H-bonds
Low MW/weak London, moderate polarity/dipole-dipole and no H-bonds
Medium MW/medium London, moderate polarity/dipole-dipole and no H-bonds
Highest MW/strongest London, but least polar bond/weakest dipole-dipole and no H-bonds
London, dipole-dipole, H-bonds
London, dipole-dipole
London, dipole-dipole
London, dipole-dipole
17
11.3 Properties resulting from IMF
  • Viscosity resistance of a liquid to flow
  • Viscosity depends on
  • -the attractive forces between molecules
  • -the tendency of molecules to become entangled
  • -the temperature

18
11.3 Properties resulting from IMF
  • Surface tension energy required to increase the
    surface area of a liquid

19
  • 3. Cohesion attraction of molecules for other
    molecules of the same compound
  • 4. Adhesion attraction of molecules for a
    surface

20
  • Meniscus curved upper surface of a liquid in a
    container a relative measure of adhesive and
    cohesive forces
  • Ex

Hg
H2O
(cohesion rules)
(adhesion rules)
21
Phase Changes
  • Surface molecules are only attracted inwards
    towards the bulk molecules.
  • Sublimation solid ? gas.
  • Vaporization liquid ? gas.
  • Melting or fusion solid ? liquid.
  • Deposition gas ? solid.
  • Condensation gas ? liquid.
  • Freezing liquid ? solid.
  • Energy Changes Accompanying Phase Changes
  • Energy change of the system for the above
    processes are

22
Intermolecular Forces Bulk and Surface
23
Phase Changes
  • Energy Changes Accompanying Phase Changes
  • Sublimation ?Hsub gt 0 (endothermic).
  • Vaporization ?Hvap gt 0 (endothermic).
  • Melting or Fusion ?Hfus gt 0 (endothermic).
  • Deposition ?Hdep lt 0 (exothermic).
  • Condensation ?Hcon lt 0 (exothermic).
  • Freezing ?Hfre lt 0 (exothermic).
  • Generally heat of fusion (enthalpy of fusion) is
    less than heat of vaporization
  • it takes more energy to completely separate
    molecules, than partially separate them.

24
Phase Changes
  • Energy Changes Accompanying Phase Changes
  • All phase changes are possible under the right
    conditions (e.g. water sublimes when snow
    disappears without forming puddles).
  • The sequence
  • heat solid ? melt ? heat liquid ? boil ? heat gas
  • is endothermic.
  • The sequence
  • cool gas ? condense ? cool liquid ? freeze ? cool
    solid
  • is exothermic.

25
Phase Changes
Energy Changes Accompanying Phase Changes
26
Phase Changes
  • Heating Curves
  • Plot of temperature change versus heat added is a
    heating curve.
  • During a phase change, adding heat causes no
    temperature change.
  • These points are used to calculate ?Hfus and
    ?Hvap.
  • Supercooling When a liquid is cooled below its
    melting point and it still remains a liquid.
  • Achieved by keeping the temperature low and
    increasing kinetic energy to break intermolecular
    forces.

27
Phase Changes
Heating Curves
28
Heating Curve Illustrated
29
Phase Changes
  • Critical Temperature and Pressure
  • Gases liquefied by increasing pressure at some
    temperature.
  • Critical temperature the minimum temperature for
    liquefaction of a gas using pressure.
  • Critical pressure pressure required for
    liquefaction.

30
Critical Temperature, Tc
31
Transition to Supercritical CO2
32
Supercritical CO2 Used to Decaffeinate Coffee
33
Vapor Pressure
  • Explaining Vapor Pressure on the Molecular Level
  • Some of the molecules on the surface of a liquid
    have enough energy to escape the attraction of
    the bulk liquid.
  • These molecules move into the gas phase.
  • As the number of molecules in the gas phase
    increases, some of the gas phase molecules strike
    the surface and return to the liquid.
  • After some time the pressure of the gas will be
    constant at the vapor pressure.

34
Gas-Liquid Equilibration
35
Vapor Pressure
  • Explaining Vapor Pressure on the Molecular Level
  • Dynamic Equilibrium the point when as many
    molecules escape the surface as strike the
    surface.
  • Vapor pressure is the pressure exerted when the
    liquid and vapor are in dynamic equilibrium.

36
Vapor Pressure
  • Volatility, Vapor Pressure, and Temperature
  • If equilibrium is never established then the
    liquid evaporates.
  • Volatile substances evaporate rapidly.
  • The higher the temperature, the higher the
    average kinetic energy, the faster the liquid
    evaporates.

37
Liquid Evaporates when no Equilibrium is
Established
38
Vapor Pressure
Volatility, Vapor Pressure, and Temperature
39
Vapor Pressure
  • Vapor Pressure and Boiling Point
  • Liquids boil when the external pressure equals
    the vapor pressure.
  • Temperature of boiling point increases as
    pressure increases.
  • Two ways to get a liquid to boil increase
    temperature or decrease pressure.
  • Pressure cookers operate at high pressure. At
    high pressure the boiling point of water is
    higher than at 1 atm. Therefore, there is a
    higher temperature at which the food is cooked,
    reducing the cooking time required.
  • Normal boiling point is the boiling point at 760
    mmHg (1 atm).

40
Phase Diagrams
  • Phase diagram plot of pressure vs. Temperature
    summarizing all equilibria between phases.
  • Given a temperature and pressure, phase diagrams
    tell us which phase will exist.
  • Features of a phase diagram
  • Triple point temperature and pressure at which
    all three phases are in equilibrium.
  • Vapor-pressure curve generally as pressure
    increases, temperature increases.
  • Critical point critical temperature and pressure
    for the gas.
  • Melting point curve as pressure increases, the
    solid phase is favored if the solid is more dense
    than the liquid.
  • Normal melting point melting point at 1 atm.

41
Phase Diagrams
  • Any temperature and pressure combination not on a
    curve represents a single phase.

42
Phase Diagrams
  • The Phase Diagrams of H2O and CO2
  • Water
  • The melting point curve slopes to the left
    because ice is less dense than water.
  • Triple point occurs at 0.0098?C and 4.58 mmHg.
  • Normal melting (freezing) point is 0?C.
  • Normal boiling point is 100?C.
  • Critical point is 374?C and 218 atm.
  • Carbon Dioxide
  • Triple point occurs at -56.4?C and 5.11 atm.
  • Normal sublimation point is -78.5?C. (At 1 atm
    CO2 sublimes it does not melt.)
  • Critical point occurs at 31.1?C and 73 atm.

43
Phase Diagrams
The Phase Diagrams of H2O and CO2
44
11.4 Phase Changes
  • Processes
  • Endothermic melting, vaporization, sublimation
  • Exothermic condensation, freezing, deposition

I2 (s) and (g)
Microchip
45
Water Enthalpy diagram or heating curve
46
11.5 Vapor pressure
Pressure cooker 2 atm
Normal BP 1 atm
10,000 elev 0.7 atm
29,029 elev (Mt. Everest) 0.3 atm
  • A liquid will boil when the vapor pressure equals
    the atmospheric pressure, at any T above the
    triple point.

47
11.6 Phase diagrams CO2
  • Lines 2 phases exist in equilibrium
  • Triple point all 3 phases exist together in
    equilibrium (X on graph)
  • Critical point, or critical temperature
    pressure highest T and P at which a liquid can
    exist (Z on graph)

Temp (ºC)
  • For most substances, inc P will cause a gas to
    condense (or deposit), a liquid to freeze, and a
    solid to become more dense (to a limit.)

48
Phase diagrams H2O
  • For H2O, inc P will cause ice to melt.

49

50

51
11.7-8 Structures of solids
  • Amorphous without orderly structure
  • Ex rubber, glass
  • Crystalline repeating structure have many
    different stacking patterns based on chemical
    formula, atomic or ionic sizes, and bonding

52
Cubic Unit Cells in Crystalline Solids
  • Primitive-cubic shared atoms are located only at
    each of the corners. 1 atom per unit cell.
  • Body-centered cubic 1 atom in center and the
    corner atoms give a net of 2 atoms per unit cell.
  • Face-centered cubic corner atoms plus half-atoms
    in each face give 4 atoms per unit cell.

53
Common Lattice Structures
54
Types of Crystalline Solids
Type Particles Forces Notable properties Examples
Atomic Atoms London dispersion Poor conductors Very low MP Ar (s),Kr (s)
A small (2 cm long) piece of rapidly melting
argon ice (the liquid is flowing off at the
bottom) which has been frozen by allowing a slow
stream of the gas to flow into a small graduated
cylinder which was immersed into a cup of liquid
nitrogen
55
Molecular crystals Molecules (polar or non-polar) London dispersion, dipole-dipole, H-bonds Poor conductors Low to moderate MP SO2(s) CO2 (s), C12H22O11, H2O (s)
Sucrose (liq at room T)
Ice(liq at room T)
Carbon dioxide, dry ice(g at room T)
56
Covalent (a.k.a. covalent network) Atoms bonded in a covalent network Covalent bonds Very hard Very high MP Generally insoluble Variable conductivity C (diamond graphite) SiO2 (quartz) Ge, Si, SiC, BN
Diamond
Graphite
SiO2
57
Ionic Anions and cations Crystals shatter! Ion-ion (ionic bonding) High Lattice Energy Hard brittle High MP,BP Poor conductors Some solubility in H2O NaCl, Ca(NO3)2
58
Metallic Metal cations in a diffuse, delocalized e- cloud Metallic bonds Usually face-centered or body centered Excellent conductors Malleable Ductile High but wide range of MP Cu, Al, Fe (hard) Alloys Pb, Au, Na (soft)
59
Overall
  • Physical properties depend on these forces. The
    stronger the forces between the particles,
  • (a) the higher the melting point.
  • (b) the higher the boiling point.
  • (c) the lower the vapor pressure (partial
    pressure of vapor in equilibrium with liquid or
    solid in a closed container at a fixed
    temperature).
  • (d) the higher the viscosity (resistance to
    flow).
  • (e) the greater the surface tension (resistance
    to an increase in surface area).

60
Practice
  • Determine the type of solid and the forces
    holding the particles together
  • SiO2 Covalent Network Covalent Bonds
  • NaNO3 Ionic Electrostatic Att.
  • C2H6 Molecular Dispersion
  • CH3OH Molecular Dispersion, Dipole-Dipole,
    H-Bond
  • C(diamond) Covalent Network Covalent Bonds
  • Al Metallic Metallic
  • Kr Atomic (Molecular) Dispersion
  • H2O Molecular Dispersion, Dipole-Dipole, H-Bond

61
Extra Material
  • The following pages contain some additional
    material and review items

62
Examples
63
Ionic Solids
  • stable, high melting points
  • held together by strong electrostatic forces
    between oppositely charged ions
  • larger ions are arranged in closest packing
    arrangement
  • smaller ions fit in the holes created by the
    larger ions

64
Cubic Unit Cells in Crystalline Solids
  • Primitive-cubic shared atoms are located only at
    each of the corners. 1 atom per unit cell.
  • Body-centered cubic 1 atom in center and the
    corner atoms give a net of 2 atoms per unit cell.
  • Face-centered cubic corner atoms plus half-atoms
    in each face give 4 atoms per unit cell.

65
Common Lattice Structures
66
Calculations involving the Unit Cell
  • The density of a metal can be calculated if we
    know the length of the side of a unit cell.
  • The radius of an metal atom can be determined if
    the unit cell type and the density of the metal
    known
  • Relationship between length of side and radius of
    atom
  • Primitive 2r l FCC BCC
  • E.g. Polonium crystallizes according to the
    primitive cubic structure. Determine its density
    if the atomic radius is 167 pm.
  • E.g.2 Calculate the radius of potassium if its
    density is 0.8560 g/cm3 and it has a BCC crystal
    structure.

67
Figure 11.31
  • Length of sides a, b, and c as well as angles a,
    b, g vary to give most of the unit cells. Return
    to unit cells

68
Unit Cells in Crystalline Solids
  • Metal crystals made up of atoms in regular arrays
    the smallest of repeating array of atoms is
    called the unit cell.
  • There are 14 different unit cells that are
    observed which vary in terms of the angles
    between atoms some are 90, but others are not.
    Go to Figure 11.31

69
Packing of Spheres and the Structures of Metals
  • Arrays of atoms act as if they are spheres. Two
    or more layers produce 3-D structure.
  • Angles between groups of atoms can be 90 or can
    be in a more compact arrangement such as the
    hexagonal closest pack (see below) where the
    spheres form hexagons.
  • Two cubic arrays one directly on top of the other
    produces simple cubic (primitive) structure.
  • Each atom has 6 nearest neighbors (coordination
    number of 6) nearest neighbor is where an atom
    touches another atom.
  • 54 of the space in a cube is used.
  • Offset layers produces a-b-a-b arrangement since
    it takes two layers to define arrangement of
    atoms.
  • BCC structure an example.
  • Coordination is 8.

70
Packing of Spheres and the Structures of Metals
  • FCC structure has a-b-c-a-b-c stacking. It takes
    three layers to establish the repeating pattern
    and has 4 atoms per unit cell and the
    coordination number is 12.

71
Metallic Crystals
  • can be viewed as metals atoms (spheres) packed
    together in the closest arrangement possible
  • closest packing- when each sphere has 12
    neighbors
  • 6 on the same plane
  • 3 in the plane above
  • 3 in the plane below

72
Bonding of Metals
  • the highest energy level for most metal atoms
    does not contain many electrons
  • these vacant overlapping orbitals allow outer
    electrons to roam freely around the entire metal

73
Bonding of Metals
  • these roaming electrons
  • form a sea of electrons
  • around the metal atoms
  • malleability and ductility
  • bonding is the same in every direction
  • one layer of atoms can slide past another without
    friction
  • conductivity
  • from the freedom of electrons to move around the
    atoms

74
Metal Alloys
  • substance that is a mixture of elements and has
    metallic properties
  • substitutional alloy
  • host metal atoms are replaced by other metal
    atoms
  • happens when they have similar sizes
  • interstitial alloy
  • metal atoms occupy spaces created between host
    metal atoms
  • happens when metal atoms have large difference in
    size

75
Examples
  • Brass
  • substitutional
  • 1/3 of Cu atoms replaced by Zn
  • Steel
  • interstitial
  • Fe with C atoms in between
  • makes harder and less malleable

76
Chapter 11 Overview
  • Changes of State
  • Phase transitions
  • Phase Diagrams
  • Liquid State
  • Properties of Liquids Surface tension and
    viscosity
  • Intermolecular forces explaining liquid
    properties
  • Solid State
  • Classification of Solids by Type of Attraction
    between Units
  • Crystalline solids crystal lattices and unit
    cells
  • Structures of some crystalline solids
  • Calculations Involving Unit-Cell Dimensions
  • Determining the Crystal Structure by X-ray
    Diffraction

Exam on Friday We will begin Chp 14 Thursday
77
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.

78
Phase Transitions
  • Melting change of a solid to a liquid.
  • Freezing change a liquid to a solid.
  • Vaporization change of a solid or liquid to a
    gas. Change of solid to vapor often called
    sublimation.
  • Condensation change of a gas to a liquid or
    solid. Change of a gas to a solid often called
    deposition.

H2O(s) ? H2O(l) H2O(l) ? H2O(s) H2O(l) ? H2O(g)
or H2O(s) ? H2O(g) H2O(g) ? H2O(l) or H2O(g) ?
H2O(s)
79
Vapor Pressure
  • In a sealed container, some of a liquid
    evaporates to establish a pressure in the vapor
    phase.
  • Vapor pressure partial pressure of the vapor
    over the liquid measured at equilibrium and at
    some temperature.
  • Dynamic equilibrium

80
Temperature Dependence of Vapor Pressures
  • The vapor pressure above the liquid varies
    exponentially with changes in the temperature.
  • The Clausius-Clapeyron equation shows how the
    vapor pressure and temperature are related. It
    can be written as

81
Clausius Clapeyron Equation
  • A straight line plot results when ln P vs. 1/T is
    plotted and has a slope of ?Hvap/R.
  • Clausius Clapeyron equation is true for any two
    pairs of points.
  • Write the equation for each and combine to get

82
Using the Clausius Clapeyron Equation
  • Boiling point the temperature at which the vapor
    pressure of a liquid is equal to the pressure of
    the external atmosphere.
  • Normal boiling point the temperature at which the
    vapor pressure of a liquid is equal to
    atmospheric pressure (1 atm).

E.g. Determine normal boiling point of chloroform
if its heat of vaporization is 31.4 kJ/mol and it
has a vapor pressure of 190.0 mmHg at
25.0C. E.g.2. The normal boiling point of
benzene is 80.1C at 26.1C it has a vapor
pressure of 100.0 mmHg. What is the heat of
vaporization?
83
Energy of Heat and Phase Change
  • Heat of vaporization heat needed for the
    vaporization of a liquid.
  • H2O(l) ?H2O(g) DH 40.7 kJ
  • Heat of fusion heat needed for the melting of a
    solid.
  • H2O(s) ?H2O(l) DH 6.01 kJ
  • Temperature does not change during the change
    from one phase to another.

E.g. Start with a solution consisting of 50.0 g
of H2O(s) and 50.0 g of H2O(l) at 0C. Determine
the heat required to heat this mixture to 100.0C
and evaporate half of the water.
84
Phase Diagrams
  • Graph of pressure-temperature relationship
    describes when 1,2,3 or more phases are present
    and/or in equilibrium with each other.
  • Lines indicate equilibrium state two phases.
  • Triple point- Temp. and press. where all three
    phases co-exist in equilibrium.
  • Critical temp.- Temp. where substance must always
    be gas, no matter what pressure.
  • Critical pressure- vapor pressure at critical
    temp.
  • Critical point- point where system is at its
    critical pressure and temp.

85
Properties of Liquids
  • Surface tension the energy required to increase
    the surface area of a liquid by a unit amount.
  • Viscosity a measure of a liquids resistance to
    flow.
  • Surface tension The net pull toward the interior
    of the liquid makes the surface tend to as small
    a surface area as possible and a substance does
    not penetrate it easily.
  • Viscosity Related to mobility of a molecule
    (proportional to the size and types of
    interactions in the liquid).
  • Viscosity decreases as the temperature increases
    since increased temperatures tend to cause
    increased mobility of the molecule.

86
Intermolecular Forces
  • Intermolecular forces attractions and repulsions
    between molecules that hold them together.
  • Intermolecular forces (van der Waals forces) hold
    molecules together in liquid and solid phases.
  • Ion-dipole force interaction between an ion and
    partial charges in a polar molecule.
  • Dipole-dipole force attractive force between
    polar molecules with positive end of one molecule
    is aligned with negative side of other.
  • London dispersion Forces interactions between
    instantaneously formed electric dipoles on
    neighboring polar or nonpolar molecules.
  • Polarizability ease with which electron cloud of
    some substance can be distorted by presence of
    some electric field (such as another dipolar
    substance). Related to size of atom or molecule.
    Small atoms and molecules less easily polarized.

87
Boiling Points vs. Molecular Weight
  • Hydrogen bonds the interaction between hydrogen
    bound to an electronegative element (N, O, or F)
    and an electron pair from another electronegative
    element. Hydrogen bonding is the dominate force
    holding the two DNA molecules together to form
    the double helix configuration of DNA.

88
Comparisonof Energies for Intermolecular Forces
Interaction Forces Approximate Energy
Intermolecular
London 1 10 kJ
Dipole-dipole 3 4 kJ
Ion-dipole 5 50 kJ
Hydrogen bonding 10 40 kJ
Chemical bonding
Ionic 100 1000 kJ
Covalent 100 1000 kJ
89
Structure of Solids
  • Types of solids
  • Crystalline a well defined arrangement of
    atoms this arrangement is often seen on a
    macroscopic level.
  • Ionic solids ionic bonds hold the solids in a
    regular three dimensional arrangement.
  • Molecular solid solids like ice that are held
    together by intermolecular forces.
  • Covalent network a solid consists of atoms held
    together in large networks or chains by covalent
    networks.
  • Metallic similar to covalent network except
    with metals. Provides high conductivity.
  • Amorphous atoms are randomly arranged. No
    order exists in the solid. Example glass
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