Intermolecular Forces:

1
Chapter 12
Intermolecular Forces
Liquids, Solids, and Phase Changes
If you are doing this lecture online then print
the lecture notes available as a word document,
go through this ppt lecture, and do all the
example and practice assignments for discussion
time.
2
Why do gases differ from liquids solids?

• Gases are tiny particles far apart with no
  attraction for each other (ideally) and they are
  moving rapidly in random directions
• Gases obey a set of laws Ideal Gas Laws
• But liquids solids don't have a set of "laws"
  because....
• Liquids - condensed from gases
• - not compressible therefore not much space
• between molecules
• -moving randomly but more slowly, are
  attracted
• to each other!
• Solids - ordered fixed in place particles
• - close together
• - strong forces of attraction!

3
Interparticle Forces

• Interparticle forces of attraction between
  "particles" that affect physical state physical
  behavior
• Note particles can be ions, atoms or molecules
• True chemical bonding forces are intra-particle
  forces within a chemical substance
• They are atom to atom or ion to ion, not
  molecule to molecule ionic bonding, metallic
  bonding, covalent bonding, and (new to us)
  network covalent bonding
• Interparticle forces are between particles, not
  within them
• Special group of interparticle forces between
  molecules and atoms
• - called intermolecular forces
• - affect behavior of covalent compounds
• Next slide is table from your packet of handouts

4
(No Transcript)
5
Properties
High MP, brittle solid
Low MP, soft when solid
Range of MPs, malleable
ADD
Network Covalent Atom-Atom like cov
bond 500 Diamond
Very high MP, very hard solid
6
Network Covalent Bonding

• Make special note of the network covalent solids
  that I added to Table 12.2
• - covalent bonds are extended throughout the
  crystal solid
• - diamond and SiC, also SiO2
• Diamond is fcc with 4 C's in "holes" in unit
  cell, which we see later in chapter
• Graphite has unique structure (see diagram)

7



True Intermolecular Forces in pure substances
8
Ion-Dipole Attraction

• Interparticle forces involved in dissolving ionic
  solids in water or other polar solvents
• Ion-dipole attraction has to overcome ion-ion
  attraction in a solids crystal lattice
• This is why some compounds are soluble and some
  are not
• Called Hydration of an ion typically
  endothermic - takes energy to pull ions apart
• Some energy is gained back as Heat of Hydration -
  polar water molecules orient themselves and
  surround the individual ions
• More detail coming soon in Chapter 13

9
The Major InterMolecular (IM) Forces

• The three major IM Forces are very weak to
  moderate forces of attraction between molecules
  (or atoms)
• - London (dispersion) forces (LDF)
• - Dipole-dipole attractions
• - Special case of enhanced dipole attraction
  called Hydrogen-bonding
• - Dipole-induced dipole attraction, between
  different types of molecules or atoms

10
London Dispersion Forces (LDF)

• "Instantaneous dipole" causes neighboring
  electron clouds to also move to one side,
  inducing a dipole in them
• Leads to a small force of attraction between
  slight positive and slight negative ends of two
  different particles these attractions are
  called London dispersion forces
• Strength of LDF depends on (1) size of electron
  cloud in atom and/or (2) number of atoms in a
  molecule
• Polarizability increases down a group
• Polarizability decreases left to right across a
  period
• Look at the elemental halogens fluorine and
  chlorine are gases bromine is liquid and iodine
  is solid because of size of e- cloud around each
  molecule

11
Dispersion forces among nonpolar molecules.
Figure 12.14
instantaneous dipoles
separated Cl2 molecules
An instantaneous dipole in one Cl2 molecule will
induce a dipole in a nearby Cl2 molecule. The
partial charges attract the molecules together.
This process takes place with the particles
throughout the container.
12
Figure 12.15
Molar mass and boiling point.
The strength of LDF increases with the number of
electrons, which correlates with molar mass.
Therefore, LDF increases down a group in the
Periodic Table, which can be verified by the
increase in boiling points.
13
Molecular shape and boiling point.
Figure 12.16
fewer points for dispersion forces to act
more points for dispersion forces to act
Spherical molecular shapes make less contact with
each other than do cyllindrical shapes, so they
have a lower boiling point.
(Its just Pentane.)
2,2-dimethylpropane
I have corrected the organic names compare to
your textbook, which is using old naming rules.
14
Polar molecules and dipole-dipole forces.
Figure 12.11
solid
In a solid or a liquid, polar molecules are close
enough for the attraction to hold. Orientation
is more orderly in a solid because the average KE
of the particles is lower.
15
Dipole-dipole Attraction

• Dipoles involve polar molecules which are
  attracted to each other because of the slight
  positive and slight negative "poles" to the
  molecules
• Compare molecules of F2, HF, HCl, HBr, HI
• Boiling points F2ltHClltHBrltHIltHF

16
Data for IM Forces

• F2 HF HCl HBr HI
• Tot e-s 18 10 18 36 54
• MM 38 20 36.5 81 128
• DEN 0 1.8 1.0 0.8 0.5
• Dip Mom 0 1.4 1.1 0.8 0.4
• Disp 100 Low 81.4 94.5 99.5
• Dipole 0 High 18.6 5.5 0.5
• BP, K 85 291 188 206 238
• DHvap 6.86 High 16 18 20

17
Enhanced Dipole-Dipole or Hydrogen-Bonding?

• Dipole forces are decreasing down the
  hydrohalogen group because DEN is decreasing
• WHY is HF so very different in boiling point?
• HF represents a special case of dipole-dipole
  attraction called Hydrogen-bonding
• Occurs when H is bonded to a highly EN atom that
  is also very small H to F, O or N
• Size of EN atom is important also, allows H to
  get very close, as seen with radius in pm below
• N-70 O-73 F-72 Cl-100 (no H-bond)

18
SAMPLE PROBLEM 12.3
Drawing Hydrogen Bonds Between Molecules of a
Substance
SOLUTION
(a) C2H6 has no H bonding sites.
(c)
The N-H is also attracted to the N-H.
19
Hydrogen bonding and boiling point.
Figure 12.13
Boiling points of the binary covalent hydrides of
Groups 14 17 plotted agains Period digit.
Shows that H2O, HF and NH3 do not follow the
downward trend, as shown by the dashed line for
group 16.
20
Dipole-Induced Dipole between two different
compounds particles

• Dipole-induced dipole forces account for limited
  solubility of oxygen in water
• Ability to do this function of polarizability
  of molecule
• Compare H2 to I2 bigger molecule polarizes,
  soluble in water, which is demonstrated by the
  much greater solubility of I2 in water

21
SAMPLE PROBLEM 12.4
Predicting the Type and Relative Strength of
Intermolecular Forces
(a) MgCl2 or PCl3
(b) CH3NH2 or CH3F
(c) CH3OH or CH3CH2OH

• Bonding forces are stronger than
  nonbonding(intermolecular) forces.
• Hydrogen bonding is a strong type of
  dipole-dipole force.
• Dispersion forces are decisive when the
  difference is molar mass or molecular shape.

22
SAMPLE PROBLEM 12.4
Predicting the Type and Relative Strength of
Intermolecular Forces
continued
SOLUTION
(a) Mg2 and Cl- are held together by ionic
bonds while PCl3 is covalently bonded and the
molecules are held together by dipole-dipole
interactions. Ionic bonds are stronger than
dipole interactions and so MgCl2 has the higher
boiling point.
(b) CH3NH2 and CH3F are both covalent compounds
and have bonds which are polar. The dipole in
CH3NH2 can H-bond while CH3F is just
dipole-dipole. Therefore CH3NH2 has the stronger
interactions and the higher boiling point.
(c) Both CH3OH and CH3CH2OH can H bond but
CH3CH2OH has more CH for more London dispersion
force interaction. Therefore CH3CH2OH has the
higher boiling point.
(d) Hexane and 2,2-dimethylbutane are both
nonpolar with only London dispersion forces to
hold the molecules together. Hexane has the
larger surface area, thereby the greater
dispersion forces and the higher boiling point.
23
Practice with Intermolecular Forces explain the
forces behind this data

• 1. Butane (CH3CH2CH2CH3) melts at -138oC and
  boils at 0.5oC, while acetone (CH3COCH3) melts
  at -95oc and boils at 56oC, yet both weigh 58
  g/mol. Draw Lewis structures and explain the
  differences in MPs and BPs.
• 2. Guess BP order for CCl4, N2, Cl2, ClNO
  (chlorine-nitrogen-oxygen).

24
Figure 12.17 modified
Summary diagram for analyzing the interparticle
forces in a sample.
Metal atoms only
INTERACTING PARTICLES (atoms, molecules, ions)
METALLIC BONDING
ions present
ions not present
ions only IONIC BONDING (Section 9.2)
nonpolar molecules or atoms Only,
LONDON DISPERSION FORCES
polar molecules only DIPOLE-DIPOLE FORCES
ion polar molecule ION-DIPOLE FORCES
polar nonpolar molecules DIPOLE- INDUCED
DIPOLE FORCES
HYDROGEN BONDING
LONDON DISPERSION FORCES ALSO PRESENT IN ALL OF
ABOVE. NETWORK COVALENT BONDING POSSIBLE FOR VERY
FEW ATOMS.
25
Practice

• See handouts
• Also chapter problems 2, 29, 31, 33, 37, 39, 43,
  45
• 4th ed. 119 What forces are overcome when the
  following events occur (a) NaCl dissolves in
  water, (b) krypton boils, (c) water boils, (d)
  CO2 sublimes?

26
QUESTIONS TO ASK IN PREDICTING THE KINDS OF
INTERPARTICLE FORCES THAT WILL BE PRESENT IN A
SOLID OR A LIQUID

•  Start at the top with the first question, Is it
  metallic?. When you can answer yes, you are
  done. If the answer is no, keep going down the
  list. The it refers to whatever substance you
  are working with.
•  
• Question If yes, this force is present MP
  BP Examples (MP, BP)
•  Is it metallic? (ONLY metal present) Metallic
  bonding High Iron (1555, 3000)
•  
• Is it ionic? (cation anion present) Ionic
  bonding High NaCl(804, not defined)
•  
• Is it network covalent compound? Network covalent
  bonding High Diamond (3550, not def), SiC,
  SiO2
• Is the substance molecular? (covalent bonds
  present)
•  
• In the molecule, is H attached by a covalent
  Hydrogen bonding Medium Water (0.100)
• bond to F, O or N?
•  
• Does the molecule have a dipole moment?
  Dipole-dipole attraction Low HCl (-114, -85)
•  
• Is it a molecule with no dipole moment? Only
  London Forces Very low Hydrogen (-257, -253)
• Iodine (114, 183)
•  

27
PROPERTIES OF LIQUIDS INTERPARTICLE FORCES

• Why would a metal object with higher density than
  water float on water?
• Why can we fill a glass of water above its rim?
• Surface tension is related to strength of
  attractive forces in liquid the stronger the
  attractive forces the greater the surface tension
• Surface tension is the energy required to
  increase surface area by a unit amount units are
  J/m2

28
The molecular basis of surface tension.
Figure 12.18
hydrogen bonding occurs across the surface and
below the surface
the net vector for attractive forces is downward
Molecules in the interior of a liquid experience
IM forces in all directions. Molecules at the
surface experience a net attraction downward,
causing the liquid to minimize the number of
molecules at the surface, ergo surface tension.
hydrogen bonding occurs in three dimensions
29
Table 12.3
Surface Tension and Forces Between Particles
Surface Tension (J/m2) at 200C
Substance
Formula
Major Force(s)
diethyl ether
dipole-dipole dispersion
CH3CH2OCH2CH3
1.7x10-2
ethanol
H bonding
2.3x10-2
CH3CH2OH
1-butanol
H bonding dispersion
2.5x10-2
CH3CH2CH2CH2OH
water
H bonding
7.3x10-2
H2O
mercury
metallic bonding
48x10-2
Hg
30
Shape of water or mercury meniscus in glass.
Figure 12.19
capillarity
H2O
Hg
See note in box below or look in textbook.
31
Properties of Liquids

• Capillary action rising of a liquid through a
  narrow space against the force of gravity
• Viscosity resistance to flow, units in
  Newton-seconds/m2

32
Table 12.4 Viscosity of Water at Several
Temperatures
viscosity - resistance to flow
Viscosity (Ns/m2)
Temperature(0C)
20
1.00x10-3
40
0.65x10-3
0.47x10-3
60
80
0.35x10-3
The units of viscosity are Newton-seconds per
square meter.
33
Why water is special

• Water molecules are 80 H-bonded at normal
  conditions
• Molecules are so close together that you cannot
  tell which H's belong to which O in each molecule
• This is important to life on earth ( possibly
  elsewhere)
• Ice floats on liquid water because the solid
  (ice) is less dense than the liquid (good for
  fishies)
• Ice forms a crystal structure in tetrahedral
  arrangement
• Hydrogen-bonding also accounts for other physical
  properties
• Lower weight alcohols are very soluble in water
  because of the -OH functional group
• Great solvent properties because water is so
  polar
• Very high specific heat as noted back in Chapter
  6
• High surface tension

34
The H-bonding ability of the water molecule.
Figure 12.20
hydrogen bond donor
hydrogen bond acceptor
Because it has two O-H bonds and two lone pairs,
one water molecule can engage in as many as four
hydrogen-bonding attractions to surrounding water
molecules, which are arranged tetrahedrally.
35
The hexagonal structure of ice.
Figure 12.21
A. The geometric arrangement of the
hydrogen-bonding in water leads to open,
hexagonally shaped crystal structure of ice.
Thus, when water freezes, the volume increases.
B. The delicate six-pointed beauty of snowflakes
reflects the hexagonal crystal structure of ice.
36
SOLIDS CRYSTAL STRUCTURES

• SOLIDS fixed particles that cannot move with
  velocity, but do vibrate and rotate in position,
  so they do have KE
• Generally have long-range order - crystals have
  well-defined regular shapes, or if short-range
  order they are amorphous - no regular shape, like
  asphalt, wax, glass
• Crystal structure includes the four types of
  solids
• ionic (all cation-anion units)
• metallic (Cu, Zn, U, etc.)
• molecular/atomic (ice, I2, etc.)
• network covalent (diamond, SiC, SiO2)

37
General Properties of the Four Types of
Crystalline Solids

• 1. Ionic (KNO3, MgO) high MP/BP some
  water-solb, brittle, conduct only when molten or
  aqueous
• 2. Molecular (C10H8, I2) low MP/BP more solb
  in nonpolar nonconductors
• 3. Network Covalent (Cdiamond, SiC, SiO2) very
  high MP/BP insolb, brittle, non- or
  semi-conductor
• 4. Metallic (Cu, Fe, U) wide range of MPs
  insolb, malleable, ductile, elec conductor

38
The striking beauty of crystalline solids.
Figure 12.22
celestite
pyrite
amethyst
halite
Figure 12.22 in current 2nd edition has
wulfanite, barite, calcite, quartz as amethyst,
and beryl (emerald)
39
Crystal Structures

• Crystals have a crystal lattice arrangement of
  which smallest pieces are unit cell in 3-D,
  containing gt one formula unit
• Seven basic types cubic, tetragonal,
  orthorhombic, monoclinic, hexagonal,
  rhombohedral, triclinic
• See packet of handouts and Dry Lab VI(?) in Lab
  Manual for Crystal Structures and
  Characteristics
• Simplest are the cubic, of which there are three
  types
• Simple cubic (sc) metals and ionic cmpds
• Body-centered cubic (bcc) metals
• Face-centered cubic (fcc) metals and ionic cmpds

40
The crystal lattice and the unit cell.
Figure 12.23
A. The lattice is an array of points that defines
the positions of the particles in a crystal
structure. It is shown here as points connected
by lines. One unit cell is highlighted.
A checkerboard is a two-dimensional analogy for a
lattice.
41
The three cubic unit cells.
Figure 12.24 (1 of 3)
Simple Cubic
Atoms touch along edge of cube
Atoms/unit cell 1/8 8 1 Cell length 2r
coordination number 6
See notes box below slide.
42
The three cubic unit cells.
Figure 12.24 (2 of 3)
Body-centered Cubic
Atoms touch along main diagonal.
Atoms/unit cell (1/88) 1 2 Cell length
4r/(3)1/2
See notes box below slide.
43
The three cubic unit cells.
Figure 12.24 (3 of 3)
Face-centered Cubic
Atoms touch along face diagonal.
Atoms/unit cell (1/88)(1/26) 4 Cell length
4r/(2)1/2
See notes box below slide.
44
Cell length and cell volume

• See figure 12.28 for derivation of cell length
  based on which cubic structure makes up the unit
  cell. For metals cell length is sc length 2r
  bcc length 4r/(3)1/2 fcc length 4r/(2)1/2
• Cell length determination is different for ionic
  compounds, which are simple cubic or
  face-centered cubic sc length 2(rR)/(3)1/2
  fcc length 2(rR)
• Volume of any cube (length)3

45
Crystal Structures Practice

• See handouts and practice problems
• Also chapter problems 61, 64, 67, 73, 75
• 4th ed. 98 Polonium is a rare radioactive
  metal that is the only element with a crystal
  structure based on the simple cubic unit cell.
  If its density is 9.142 g/cm3, calculate an
  atomic radius for a polonium atom.

46
Crystal Structures Practice

• 4th ed. 101 Tantalum, with D 16.634 g/cm3,
  has a bcc structure with an edge length of 3.3058
  Angstroms. Use its molar mass and this data to
  prove Avogadros number.

47
Phase Changes
sublimination
vaporizing
melting
solid
liquid
gas
condensing
freezing
deposition
48
Vapor Pressure

• Evaporation/vaporization small fraction of
  molecules have high enough velocity to escape
  force of attraction at surface
• RATE OF EVAPORATION will increase with
  increasing T, since fraction of molecules with
  escape vel will increase
• In a closed system, a dynamic equilibrium will be
  reached
• Rate of evaporation rate of condensation
• Vapor pressure vapor molecules exert a partial
  pressure called vapor pressure

49
Liquid-gas equilibrium.
Figure 12.4
A. In a closed flask at const T, with air
removed, Pi 0. As molecules escape surface to
become vapor, P increases. B. At equilibrium,
of molecules escaping liquid of molecules
condensing, P is constant. C. Plot of P vs. time
shows P becomes constant.
50
The effect of temperature on the distribution of
molecular speed in a liquid.
Figure 12.5
With T1 lower than T2, most probable molecular
speed, u1, is less than u2. Fraction of
molecules with escape velocity is greater at
the higher temperature. At higher T, equilibrium
is reached with more molecules in the vapor
phase, therefore at a higher P.
51
Vapor Pressure Practice

• If 1.00 L of water is placed in 2.30x104 L closed
  room, will all the water evaporate? Given D
  0.997 g/mL, Vapor Pressure 23.8 torr at 25.0oC.
• Water will evap till room is at 23.8 torr partial
  pressure of water vapor.
• See how many moles at that point
• n PV/RT 29.4 mol
• How many moles in 1.0 L beaker? About 55 moles -
  won't all evaporate

52
Vapor Pressure vs. Temperature

• Boiling point occurs when you see bubbles of
  gas forming in the liquid and coming to surface
• Any pure liquid remains at constant T while
  boiling, since this is a change of state
• Definition
• BP is the Temperature at which VP barometric P
  
• Why does water boil at 100oC in Fairfield and at
  95oC in Denver?

53
Figures 12.6 and 12.7
Figure 12.7
A linear plot of vapor pressure- temperature
relationship.
Vapor pressure as a function of temperature and
intermolecular forces.
The Clausius-Clapeyron equation comes from this
graph y mx b
54
You practice drawing and labelling a generic VP
curve
55
Vapor pressure curves

• Initial "phase diagrams" incorporated into P/T
  diagrams that will include solid phase later
• Why can NH3 be condensed from gas to liquid at
  Room T by compression, but N2 can't?

56
Relative Humidity

• NOT IN TEXT
• Relative humidity as reported by weather
  forecasters
• water evapactual partial P/equil vapor P 100
• If actual is 12.8 and VP for given T is 21.1,
  relative humidity is 61

57
VP and DHvap

• DHvap is related to VP and T thru the
  Clausius-Clapeyron equation
• ln P (-DHvap/RT) C (where R 8.314 J/mol-K,
  T in K)
• If plotted on a graph, the slope is
• (ln p2 ln p1)/(1/T2 1/T1) -DHvap/R
• Rearranges to Clausius-Clapeyron Equation (next
  slide)

58
The Clausius-Clapeyron Equation
MEMORIZE!
Alternately, if you dont want to use the
negative sign ln (P2/P1) DHvap/R(1/T1-1/T2)
59
Clausius-Clapeyron equation examples

• Look at Sample Problem 12.2 in text.
• My Example hexane has DHvap 30.1 kJ/mol and
  at 25.0oC, VP 148 torr. What will VP be at
  50.0oC?
• ln (P2/148) (-30.1x103J/8.314 J/mol-K)(1/323.15
  1/298.15)
• ln (P2/148) 0.9394 (take antilog of both sides)
• P2/148 e0.9394 2.55
• P2 379 torr

60
Practice

• Chapter problems 17 18
• 17 A liquid has DHovap of 35.5 kJ/mol and a BP
  of 122oC at 1.00 atm. What is its VP at 113oC?
• 18 What is the DHovap of a liquid that has a VP
  of 641 torr at 85.2oC and a BP of 95.6oC at 1.00
  atm?

61
Iodine subliming.
Figure 12.8 4th ed., not in principles
62
Simple Phase Diagrams

• Are P vs. T diagrams showing three phases for
  pure elements or compounds, incorporates the VP
  curve
• Critical Point is where liquid and gas cannot be
  distinguished from each other
• Triple Point is where solid, liquid and gas
  phases meet and all three are present
• Example For water the Triple Point is 0.01oC
  and 4.58 torr
• For CO2 Triple Point is at -56.7oC and 5.1 atm

63
Phase diagrams for CO2 and H2O.
Figure 12.8
Each region depicts the T P under which the
phase is stable. Lines between regions show
conditions at which two phases exist in
equilibrium. The Critical Point shows conditions
beyond which liquid and gas cannot be
distinguished from each other. At the triple
point, three phases exist is equilibrium. CO2
phase diagram is typical with forward sloping
solid-liquid line solid is more dense than
liquid. H2O phase diagram is sloping backward
solid is less dense than liquid.
64
Phase Diagrams

• You must draw and label phase diagrams based on
  data given to you and determine the physical
  state of a substance from its placement on a
  phase diagram.
• Work on problems 20 22