Ch 12

1
Chapters Eleven and Thirteen Intermolecular
Forces, Liquids and Solids and Solutions
2
Intermolecular Forces

• The covalent bond holding a molecule together is
  an intramolecular forces.
• The attraction between molecules is an
  intermolecular force.
• Intermolecular forces are much weaker than
  intramolecular forces (e.g. 16 kJ/mol vs. 431
  kJ/mol for HCl).
• When a substance melts or boils the
  intermolecular forces are broken (not the
  covalent bonds).
• When a substance condenses intermolecular forces
  are formed.

3
Types of Intermolecular Forces - Bonding and
Nonbonding
4
Types of Intermolecular Forces - Bonding and
Nonbonding
5
Intermolecular Forces

• Dipole-Dipole Forces
• Interaction between an ion (e.g. Na) and a
  dipole (e.g. water).
• Dipole-dipole forces exist between neutral polar
  molecules.
• Polar molecules need to be close together.
• Weaker than ion-dipole forces
• Q1 and Q2 are partial charges.

6
Intermolecular Forces

• Dipole-Dipole Forces
• There is a mix of attractive and repulsive
  dipole-dipole forces as the molecules tumble.
• If two molecules have about the same mass and
  size, then dipole-dipole forces increase with
  increasing polarity.

7
Orientation of polar molecules because of
dipole-dipole forces
8
Dipole moment and boiling point
9
Intermolecular Forces

• Hydrogen Bonding
• Special case of dipole-dipole forces.
• By experiments boiling points of compounds with
  H-F, H-O, and H-N bonds are abnormally high.
• Intermolecular forces are abnormally strong.
• H-bonding requires H bonded to an electronegative
  element (most important for compounds of F, O,
  and N).
• Electrons in the H-X (X electronegative
  element) lie much closer to X than H.
• H has only one electron, so in the H-X bond, the
  ? H presents an almost bare proton to the ?- X.
• Therefore, H-bonds are strong.

10
The Hydrogen Bond
A special dipole-dipole interaction occurs when a
H atom is covalently bonded to a small
electronegative atom, i.e. N, O, or F. The
Hydrogen Bond is a through space bond between a H
atom that is covalently bonded to one of the
electronegative atoms to another of the
electronegative atoms. H-F-----H-O-H
H2O------H-O-O
11
SAMPLE PROBLEM 12.2
Drawing Hydrogen Bonds Between Molecules of a
Substance
SOLUTION
(a) C2H6 has no H bonding sites.
(c)
12
Hydrogen bonding and boiling point
13
The H-bonding abilitiy of the water molecule
14
Intermolecular Forces

• Hydrogen Bonding
• Hydrogen bonds are responsible for
• Ice Floating
• Solids are usually more closely packed than
  liquids
• therefore, solids are more dense than liquids.
• Ice is ordered with an open structure to optimize
  H-bonding.
• Therefore, ice is less dense than water.
• In water the H-O bond length is 1.0 Å.
• The OH hydrogen bond length is 1.8 Å.
• Ice has waters arranged in an open, regular
  hexagon.
• Each ? H points towards a lone pair on O.
• Ice floats, so it forms an insulating layer on
  top of lakes, rivers, etc. Therefore, aquatic
  life can survive in winter.

15
Intermolecular Forces

• Hydrogen Bonding
• Hydrogen bonds are responsible for
• Protein Structure
• Protein folding is a consequence of H-bonding.
• DNA Transport of Genetic Information

16
Intermolecular Forces

• London Dispersion Forces
• Weakest of all intermolecular forces.
• It is possible for two adjacent neutral molecules
  to affect each other.
• The nucleus of one molecule (or atom) attracts
  the electrons of the adjacent molecule (or atom).
• For an instant, the electron clouds become
  distorted.
• In that instant a dipole is formed (called an
  instantaneous dipole).

17
Intermolecular Forces
London Dispersion Forces
18
Intermolecular Forces

• London Dispersion Forces
• One instantaneous dipole can induce another
  instantaneous dipole in an adjacent molecule (or
  atom).
• The forces between instantaneous dipoles are
  called London dispersion forces.
• Polarizability is the ease with which an electron
  cloud can be deformed.
• The larger the molecule (the greater the number
  of electrons) the more polarizable.

19
DISPERSION(London) FORCES
Molecular shape and boiling point
20
Intermolecular Forces
London Dispersion Forces
21
Intermolecular Forces

• London Dispersion Forces
• London dispersion forces increase as molecular
  weight increases.
• London dispersion forces exist between all
  molecules.
• London dispersion forces depend on the shape of
  the molecule.
• The greater the surface area available for
  contact, the greater the dispersion forces.
• London dispersion forces between spherical
  molecules are lower than between sausage-like
  molecules.

22
separated Cl2 molecules
DISPERSION(London) FORCES among nonpolar
molecules
instantaneous dipoles
23
DISPERSION(London) FORCES
Effect of Molar Mass and boiling point
24
SAMPLE PROBLEM 12.3
Predicting the Type and Relative Strength of
Intermolecular Forces
PROBLEM
For each pair of substances, identify the
dominant intermolecular forces in each substance,
and select the substance with the higher boiling
point.
(a) MgCl2 or PCl3
(b) CH3NH2 or CH3F
(c) CH3OH or CH3CH2OH
PLAN

• 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.

25
SAMPLE PROBLEM 12.3
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 that in CH3F cannot.
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 dispersion force
interaction. Therefore CH3CH2OH has the higher
boiling point.
(d) Hexane and 2,2-dimethylbutane are both
nonpolar with only dispersion forces to hold the
molecules together. Hexane has the larger
surface area, thereby the greater dispersion
forces and the higher boiling point.
26
Intermolecular Forces
Comparing Intermolecular Forces
27
Table 12.1
A Macroscopic Comparison of Gases, Liquids, and
Solids
State
Shape and Volume
Compressibility
Ability to Flow
Gas
Conforms to shape and volume of container
high
high
Liquid
Conforms to shape of container volume limited by
surface
very low
moderate
Solid
Maintains its own shape and volume
almost none
almost none
28
Types of Phases Changes
A liquid changing into a gas - vaporizationthe
reverse process - condensation A solid changing
into a liquid - fusion (melting)the reverse
process - freezing (solidification) A solid
changing directly into a gas - sublimationthe
reverse process - deposition Enthalpy changes
accompany phase changes. Vaporization, fusion,
and sublimation areEXOTHERMIC the reverse
processes ENDOTHERMIC
29
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.

30
Heats of vaporization and fusion for several
common substances.
31
Phase changes and their enthalpy changes
32
Quantitative Aspects of Phase Changes
Energy changes result in a change in temperature
and/or change in phase.
Within a phase, a change in heat is accompanied
by a change in temperature which is associated
with a change in average Ek as the most probable
speed of the molecules changes.
q (amount)(molar heat capacity)(DT)
During a phase change, a change in heat occurs at
a constant temperature, which is associated with
a change in Ep, as the average distance between
molecules changes.
q (amount)(enthalpy of phase change)
33
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.

34
Phase Changes
Energy Changes Accompanying Phase Changes
35
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.

36
Phase Changes
Heating Curves
37
A cooling curve for the conversion of gaseous
water to ice
Heat Removed
38
Calculating the Loss of Heat - Cooling steam at
110o C down to ice at -10o C
q (amount)(molar heat capacity)(DT) -
change of temp
q (amount)(enthalpy of phase change) - change
of phase
q n Cwater(g) (100-110) q n (-?HOvap)
q n Cwater(l) (0-100) q n (-?HOfus)
q n Cwater(s) (-10-0)
39
Vapor pressure as a function of temperature and
intermolecular forces
A linear plot of vapor pressure- temperature
relationship
40
The Clausius-Clapeyron Equation
Subtraction two equations for two temperatures.
41
SAMPLE PROBLEM 12.1
Using the Clausius-Clapeyron Equation
SOLUTION
34.90C 308.0K

T2 350K 770C
42
Phase diagrams for CO2 and H2O
43
A Molecular Comparison of Liquids and Solids

• Physical properties of substances understood in
  terms of kinetic molecular theory
• Gases are highly compressible, assumes shape and
  volume of container
• Gas molecules are far apart and do not interact
  much with each other.
• Liquids are almost incompressible, assume the
  shape but not the volume of container
• Liquids molecules are held closer together than
  gas molecules, but not so rigidly that the
  molecules cannot slide past each other.
• Solids are incompressible and have a definite
  shape and volume
• Solid molecules are packed closely together. The
  molecules are so rigidly packed that they cannot
  easily slide past each other.

44
A Molecular Comparison of Liquids and Solids
45
A Molecular Comparison of Liquids and Solids

• Converting a gas into a liquid or solid requires
  the molecules to get closer to each other
• cool or compress.
• Converting a solid into a liquid or gas requires
  the molecules to move further apart
• heat or reduce pressure.
• The forces holding solids and liquids together
  are called intermolecular forces.

46
Some Properties of Liquids

• Viscosity
• Viscosity is the resistance of a liquid to flow.
• A liquid flows by sliding molecules over each
  other.
• The stronger the intermolecular forces, the
  higher the viscosity.
• Surface Tension
• Bulk molecules (those in the liquid) are equally
  attracted to their neighbors.

47
Some Properties of Liquids
Surface Tension
48
Some Properties of Liquids

• Surface Tension
• Surface molecules are only attracted inwards
  towards the bulk molecules.
• Therefore, surface molecules are packed more
  closely than bulk molecules.
• Surface tension is the amount of energy required
  to increase the surface area of a liquid.
• Cohesive forces bind molecules to each other.
• Adhesive forces bind molecules to a surface.

49
Some Properties of Liquids
Surface Tension
50
Some Properties of Liquids

• Surface Tension
• Meniscus is the shape of the liquid surface.
• If adhesive forces are greater than cohesive
  forces, the liquid surface is attracted to its
  container more than the bulk molecules.
  Therefore, the meniscus is U-shaped (e.g. water
  in glass).
• If cohesive forces are greater than adhesive
  forces, the meniscus is curved downwards.
• Capillary Action When a narrow glass tube is
  placed in water, the meniscus pulls the water up
  the tube.

51
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.

52
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.

53
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.

54
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.

55
Vapor Pressure
Volatility, Vapor Pressure, and Temperature
56
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).

57
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.

58
Phase Diagrams

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

59
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.

60
Phase Diagrams
The Phase Diagrams of H2O and CO2
61
Bonding in Solids

• There are four types of solid
• Molecular (formed from molecules) - usually soft
  with low melting points and poor conductivity.
• Covalent network (formed from atoms) - very hard
  with very high melting points and poor
  conductivity.
• Ions (formed form ions) - hard, brittle, high
  melting points and poor conductivity.
• Metallic (formed from metal atoms) - soft or
  hard, high melting points, good conductivity,
  malleable and ductile.

62
Bonding in Solids
63
Bonding in Solids

• Molecular Solids
• Intermolecular forces dipole-dipole, London
  dispersion and H-bonds.
• Weak intermolecular forces give rise to low
  melting points.
• Room temperature gases and liquids usually form
  molecular solids and low temperature.
• Efficient packing of molecules is important
  (since they are not regular spheres).

64
Bonding in Solids

• Covalent Network Solids
• Intermolecular forces dipole-dipole, London
  dispersion and H-bonds.
• Atoms held together in large networks.
• Examples diamond, graphite, quartz (SiO2),
  silicon carbide (SiC), and boron nitride (BN).
• In diamond
• each C atom has a coordination number of 4
• each C atom is tetrahedral
• there is a three-dimensional array of atoms.
• Diamond is hard, and has a high melting point
  (3550 ?C).

65
Bonding in Solids

• Covalent Network Solids

66
Bonding in Solids

• Covalent Network Solids
• In graphite
• each C atom is arranged in a planar hexagonal
  ring
• layers of interconnected rings are placed on top
  of each other
• the distance between C atoms is close to benzene
  (1.42 Å vs. 1.395 Å in benzene)
• the distance between layers is large (3.41 Å)
• electrons move in delocalized orbitals (good
  conductor).

67
Bonding in Solids

• Ionic Solids
• Ions (spherical) held together by electrostatic
  forces of attraction
• The higher the charge (Q) and smaller the
  distance (d) between ions, the stronger the ionic
  bond.
• There are some simple classifications for ionic
  lattice types

68
Bonding in Solids

• Metallic Solids
• Metallic solids have metal atoms in hcp, fcc or
  bcc arrangements.
• Coordination number for each atom is either 8 or
  12.
• Problem the bonding is too strong for London
  dispersion and there are not enough electrons for
  covalent bonds.
• Resolution the metal nuclei float in a sea of
  electrons.
• Metals conduct because the electrons are
  delocalized and are mobile.

69
Bonding in Solids
Metallic Solids
70
The Solution Process

• A solution is a homogeneous mixture of solute
  (present in smallest amount) and solvent (present
  in largest amount).
• Solutes and solvent are components of the
  solution.
• In the process of making solutions with condensed
  phases, intermolecular forces become rearranged.
• Consider NaCl (solute) dissolving in water
  (solvent)
• the water H-bonds have to be interrupted,
• NaCl dissociates into Na and Cl-,
• ion-dipole forces form Na ?-OH2 and Cl-
  ?H2O.
• We say the ions are solvated by water.
• If water is the solvent, we say the ions are
  hydrated.

71
The Solution Process
72
The Solution Process

• Energy Changes and Solution Formation
• There are three energy steps in forming a
  solution
• separation of solute molecules (?H1),
• separation of solvent molecules (?H2),
  andformation of solute-solvent interactions
  (?H3).
• We define the enthalpy change in the solution
  process as
• ?Hsoln ?H1 ?H2 ?H3.
• ?Hsoln can either be positive or negative
  depending on the intermolecular forces.

73
The Solution Process
Energy Changes and Solution Formation
74
The Solution Process

• Energy Changes and Solution Formation
• Breaking attractive intermolecular forces is
  always endothermic.
• Forming attractive intermolecular forces is
  always exothermic.
• To determine whether ?Hsoln is positive or
  negative, we consider the strengths of all
  solute-solute and solute-solvent interactions
• ?H1 and ?H2 are both positive.
• ?H3 is always negative.
• It is possible to have either ?H3 gt (?H1 ?H2)
  or ?H3 lt (?H1 ?H2).

75
The Solution Process
Energy Changes and Solution Formation
76
The Solution Process

• Energy Changes and Solution Formation
• Examples
• NaOH added to water has ?Hsoln -44.48 kJ/mol.
• NH4NO3 added to water has ?Hsoln 26.4 kJ/mol.
• Rule polar solvents dissolve polar solutes.
  Non-polar solvents dissolve non-polar solutes.
  Why?
• If ?Hsoln is too endothermic a solution will not
  form.
• NaCl in gasoline the ion-dipole forces are weak
  because gasoline is non-polar. Therefore, the
  ion-dipole forces do not compensate for the
  separation of ions.
• Water in octane water has strong H-bonds. There
  are no attractive forces between water and octane
  to compensate for the H-bonds.

77
The Solution Process

• Solution Formation, Spontaneity, and Disorder
• A spontaneous process occurs without outside
  intervention.
• When energy of the system decreases (e.g.
  dropping a book and allowing it to fall to a
  lower potential energy), the process is
  spontaneous.
• Some spontaneous processes do not involve the
  system moving to a lower energy state (e.g. an
  endothermic reaction).
• If the process leads to a greater state of
  disorder, then the process is spontaneous.

78
The Solution Process

• Solution Formation, Spontaneity, and Disorder
• Example a mixture of CCl4 and C6H14 is less
  ordered than the two separate liquids.
  Therefore, they spontaneously mix even though
  ?Hsoln is very close to zero.
• There are solutions that form by physical
  processes and those by chemical processes.

79
The Solution Process
Solution Formation, Spontaneity, and Disorder
80
The Solution Process

• Solution Formation and Chemical Reactions
• Example a mixture of CCl4 and C6H14 is less
  ordered
• Consider
• Ni(s) 2HCl(aq) ? NiCl2(aq) H2(g).
• Note the chemical form of the substance being
  dissolved has changed (Ni ? NiCl2).
• When all the water is removed from the solution,
  no Ni is found only NiCl2.6H2O. Therefore, Ni
  dissolution in HCl is a chemical process.

81
The Solution Process
Solution Formation and Chemical Reactions
82
The Solution Process

• Solution Formation and Chemical Reactions
• Example
• NaCl(s) H2O (l) ? Na(aq) Cl-(aq).
• When the water is removed from the solution, NaCl
  is found. Therefore, NaCl dissolution is a
  physical process.

83
Ways of Expressing Concentration

• All methods involve quantifying amount of solute
  per amount of solvent (or solution).
• Generally amounts or measures are masses, moles
  or liters.
• Qualitatively solutions are dilute or
  concentrated.
• Definitions

84
Ways of Expressing Concentration

• Parts per million (ppm) can be expressed as 1 mg
  of solute per kilogram of solution.
• If the density of the solution is 1g/mL, then 1
  ppm 1 mg solute per liter of solution.
• Parts per billion (ppb) are 1 ?g of solute per
  kilogram of solution.

85
Ways of Expressing Concentration

• Mole Fraction, Molarity, and Molality
• Recall mass can be converted to moles using the
  molar mass.
• Recall
• Recall

86
Ways of Expressing Concentration

• Mole Fraction, Molarity, and Molality
• We define
• Converting between molarity (M) and molality (m)
  requires density.

87
Ways of Expressing Concentration
Mole Fraction, Molarity, and Molality
88
Saturated Solutions and Solubility

• Mole Fraction, Molarity, and Molality
• Dissolve solute solvent ? solution.
• Crystallization solution ? solute solvent.
• Saturation crystallization and dissolution are
  in equilibrium.
• Solubility amount of solute required to form a
  saturated solution.
• Supersaturated a solution formed when more
  solute is dissolved than in a saturated solution.

89
Factors Affecting Solubility

• Solute-Solvent Interactions
• Polar liquids tend to dissolve in polar solvents.
• Miscible liquids mix in any proportions.
• Immiscible liquids do not mix.
• Intermolecular forces are important water and
  ethanol are miscible because the broken hydrogen
  bonds in both pure liquids are re-established in
  the mixture.
• The number of carbon atoms in a chain affect
  solubility the more C atoms the less soluble in
  water.

90
Factors Affecting Solubility
Solute-Solvent Interactions
91
Factors Affecting Solubility

• Solute-Solvent Interactions
• The number of -OH groups within a molecule
  increases solubility in water.

92
Factors Affecting Solubility

• Solute-Solvent Interactions
• Generalization like dissolves like.
• The more polar bonds in the molecule, the better
  it dissolves in a polar solvent.
• The less polar the molecule the less it dissolves
  in a polar solvent and the better is dissolves in
  a non-polar solvent.
• Network solids do not dissolve because the strong
  intermolecular forces in the solid are not
  re-established in any solution.

93
Factors Affecting Solubility

• Pressure Effects
• Solubility of a gas in a liquid is a function of
  the pressure of the gas.
• The higher the pressure, the more molecules of
  gas are close to the solvent and the greater the
  chance of a gas molecule striking the surface and
  entering the solution.
• Therefore, the higher the pressure, the greater
  the solubility.
• The lower the pressure, the fewer molecules of
  gas are close to the solvent and the lower the
  solubility.

94
Factors Affecting Solubility
Pressure Effects
95
Factors Affecting Solubility

• Pressure Effects
• Henrys Law
• Cg is the solubility of gas, Pg the partial
  pressure, k Henrys law constant.
• Carbonated beverages are bottled under gt
  1 atm. As the bottle is opened,
  decreases and the solubility of CO2 decreases.
  Therefore, bubbles of CO2 escape from solution.

96
Factors Affecting Solubility

• Temperature Effects
• Experience tells us that sugar dissolves better
  in warm water than cold.
• As temperature increases, solubility of solids
  generally increases.
• Sometimes, solubility decreases as temperature
  increases (e.g. Ce2(SO4)3).

97
Factors Affecting Solubility
Temperature Effects
98
Factors Affecting Solubility

• Temperature Effects
• Experience tells us that carbonated beverages go
  flat as they get warm.

99
Factors Affecting Solubility

• Temperature Effects
• Experience tells us that carbonated beverages go
  flat as they get warm.
• Gases are less soluble at higher temperatures.
• Thermal pollution if lakes get too warm, CO2 and
  O2 become less soluble and are not available for
  plants or animals.

100
Colligative Properties

• Colligative properties depend on quantity of
  solute molecules. (E.g. freezing point
  depression and melting point elevation.)
• Lowering the Vapor Pressure
• Non-volatile solvents reduce the ability of the
  surface solvent molecules to escape the liquid.
• Therefore, vapor pressure is lowered.
• The amount of vapor pressure lowering depends on
  the amount of solute.

101
Colligative Properties
Lowering the Vapor Pressure
102
Colligative Properties

• Raoults Law
• Raoults Law PA is the vapor pressure with
  solute, PA? is the vapor pressure without
  solvent, and ?A is the mole fraction of A, then
• Recall Daltons Law

103
Colligative Properties

• Raoults Law
• Ideal solution one that obeys Raoults law.
• Raoults law breaks down when the solvent-solvent
  and solute-solute intermolecular forces are
  greater than solute-solvent intermolecular
  forces.
• Boiling-Point Elevation
• Goal interpret the phase diagram for a solution.
• Non-volatile solute lowers the vapor pressure.
• Therefore the triple point - critical point curve
  is lowered.

104
Colligative Properties
Boiling-Point Elevation
105
Colligative Properties

• Boiling-Point Elevation
• At 1 atm (normal boiling point of pure liquid)
  there is a lower vapor pressure of the solution.
  Therefore, a higher temperature is required to
  teach a vapor pressure of 1 atm for the solution
  (?Tb).
• Molal boiling-point-elevation constant, Kb,
  expresses how much ?Tb changes with molality, m

106
Colligative Properties

• Freezing-Point Depression
• At 1 atm (normal boiling point of pure liquid)
  there is no depression by definition
• When a solution freezes, almost pure solvent is
  formed first.
• Therefore, the sublimation curve for the pure
  solvent is the same as for the solution.
• Therefore, the triple point occurs at a lower
  temperature because of the lower vapor pressure
  for the solution.
• The melting-point (freezing-point) curve is a
  vertical line from the triple point.

107
Colligative Properties

• Freezing-Point Depression
• The solution freezes at a lower temperature (?Tf)
  than the pure solvent.
• Decrease in freezing point (?Tf) is directly
  proportional to molality (Kf is the molal
  freezing-point-depression constant)

108
Colligative Properties
Freezing-Point Depression
109
Colligative Properties

• Osmosis
• Semipermeable membrane permits passage of some
  components of a solution. Example cell
  membranes and cellophane.
• Osmosis the movement of a solvent from low
  solute concentration to high solute
  concentration.
• There is movement in both directions across a
  semipermeable membrane.
• As solvent moves across the membrane, the fluid
  levels in the arms becomes uneven.

110
Colligative Properties

• Osmosis
• Eventually the pressure difference between the
  arms stops osmosis.

111
Colligative Properties

• Osmosis
• Osmotic pressure, ?, is the pressure required to
  stop osmosis

112
Colligative Properties

• Osmosis
• Osmotic pressure, ?, is the pressure required to
  stop osmosis
• Isotonic solutions two solutions with the same ?
  separated by a semipermeable membrane.
• Hypotonic solutions a solution of lower ? than a
  hypertonic solution.
• Osmosis is spontaneous.
• Red blood cells are surrounded by semipermeable
  membranes.

113
Colligative Properties

• Osmosis
• Crenation
• red blood cells placed in hypertonic solution
  (relative to intracellular solution)
• there is a lower solute concentration in the cell
  than the surrounding tissue
• osmosis occurs and water passes through the
  membrane out of the cell.
• The cell shrivels up.

114
Colligative Properties

• Osmosis
• Crenation and Hemolysis

115
Colligative Properties

• Osmosis
• Hemolysis
• red blood cells placed in a hypotonic solution
• there is a higher solute concentration in the
  cell
• osmosis occurs and water moves into the cell.
• The cell bursts.
• To prevent crenation or hemolysis, IV
  (intravenous) solutions must be isotonic.
• Examples of osmosis
• Cucumber placed in NaCl solution loses water to
  shrivel up and become a pickle.

116
Colligative Properties

• Osmosis
• Limp carrot placed in water becomes firm because
  water enters via osmosis.
• Salty food causes retention of water and swelling
  of tissues (edema).
• Water moves into plants through osmosis.
• Salt added to meat or sugar to fruit prevents
  bacterial infection (a bacterium placed on the
  salt will lose water through osmosis and die).
• Active transport is the movement of nutrients and
  waste material through a biological system.
• Active transport is not spontaneous.

117
Colloids

• Colloids are suspensions in which the suspended
  particles are larger than molecules but too small
  to drop out of the suspension due to gravity.
• Particle size 10 to 2000 Å.
• There are several types of colloid
• aerosol (gas liquid or solid, e.g. fog and
  smoke),
• foam (liquid gas, e.g. whipped cream),
• emulsion (liquid liquid, e.g. milk),
• sol (liquid solid, e.g. paint),
• solid foam (solid gas, e.g. marshmallow),
• solid emulsion (solid liquid, e.g. butter),
• solid sol (solid solid, e.g. ruby glass).

118
Colloids

• Tyndall effect ability of a Colloid to scatter
  light. The beam of light can be seen through the
  colloid.

119
Colloids

• Hydrophilic and Hydrophobic Colloids
• Focus on colloids in water.
• Water loving colloids hydrophilic.
• Water hating colloids hydrophobic.
• Molecules arrange themselves so that hydrophobic
  portions are oriented towards each other.
• If a large hydrophobic macromolecule (giant
  molecule) needs to exist in water (e.g. in a
  cell), hydrophobic molecules embed themselves
  into the macromolecule leaving the hydrophilic
  ends to interact with water.

120
Colloids
Hydrophilic and Hydrophobic Colloids
121
Colloids

• Hydrophilic and Hydrophobic Colloids
• Typical hydrophilic groups are polar (containing
  C-O, O-H, N-H bonds) or charged.
• Hydrophobic colloids need to be stabilized in
  water.
• Adsorption when something sticks to a surface we
  say that it is adsorbed.
• If ions are adsorbed onto the surface of a
  colloid, the colloids appears hydrophilic and is
  stabilized in water.
• Consider a small drop of oil in water.
• Add to the water sodium stearate.

122
Colloids
Hydrophilic and Hydrophobic Colloids
123
Colloids

• Hydrophilic and Hydrophobic Colloids
• Sodium stearate has a long hydrophobic tail
  (CH3(CH2)16-) and a small hydrophobic head
  (-CO2-Na).
• The hydrophobic tail can be absorbed into the oil
  drop, leaving the hydrophilic head on the
  surface.
• The hydrophilic heads then interact with the
  water and the oil drop is stabilized in water.

124
Colloids
Hydrophilic and Hydrophobic Colloids
125
Colloids

• Hydrophilic and Hydrophobic Colloids
• Most dirt stains on people and clothing are
  oil-based. Soaps are molecules with long
  hydrophobic tails and hydrophilic heads that
  remove dirt by stabilizing the colloid in water.
• Bile excretes substances like sodium stereate
  that forms an emulsion with fats in our small
  intestine.
• Emulsifying agents help form an emulsion.

126
Colloids

• Removal of Colloidal Particles
• Colloid particles are too small to be separated
  by physical means (e.g. filtration).
• Colloid particles are coagulated (enlarged) until
  they can be removed by filtration.
• Methods of coagulation
• heating (colloid particles move and are attracted
  to each other when they collide)
• adding an electrolyte (neutralize the surface
  charges on the colloid particles).
• Dialysis using a semipermeable membranes
  separate ions from colloidal particles.