Lecture Notes Chem 150 - K. Marr

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Lecture Notes Chem 150 - K. Marr

• Chapter 12
• Intermolecular Attractions
• the Properties of Liquids Solids
• Silberberg 3 ed

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Intermolecular Forces Liquids, Solids, and
Phase Changes
12.1 An Overview of Physical States and Phase
Changes 12.2 Quantitative Aspects of Phase
Changes 12.3 Types of Intermolecular
Forces 12.4 Properties of the Liquid
State 12.5 The Uniqueness of Water 12.6 The
Solid State Structure, Properties, and
Bonding 12.7 Advanced Materials
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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
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Chapter 12 Intermolecular Attractions the
Properties of Liquids Solids

• Why do the gas laws work with almost any gas?
• Gases are alike
• Mainly empty space ? Weak intermolecular
  attractions

Liquid Conforms to shape of container volume
limited by surface
Solid Maintains its own shape and volume
Gas Conforms to shape and volume of container
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Why arent there liquid laws and solid laws?

• Little empty space between molecules
• Particles close together in Ss and Ls
• Stronger and quite varied intermolecular
  attractions in Ss Ls than in gases...Why?
• Attractions decrease as distance between
  molecules increase I.M.F. a 1/
  d2
• Attractions dependent on chemical composition
• Polar vs Nonpolar molecules
• e.g. Water vs Carbon Dioxide (sublimes. _at_ -58.5
  oC)

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Intermolecular Attractions Bonds between
Molecules

• Much weaker than Chemical Bonding within
  molecules
• Chemical Bonds (ionic and covalent) determine
  chemical properties
• Intermolecular bonds determine physical
  properties
• e.g. density, mp, bp, solubility, vapor
  pressure, etc.

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Kinds of Intermolecular Attractions

• Dipole-Dipole Attractions
• Hydrogen Bonds (H-FON Bonds)
• London Forces (dispersion forces)
• Ion-Dipole (e.g. spheres of hydration)? Chapter
  13
• Induced Dipole forces ?Chapter 13
• Ion induced
• Dipole induced

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How do IMFs affect Heats of Vaporization and
Fusion?
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Relative magnitudes of forces in Molecular
compounds

• Covalent bonds gtgtgtgtgt Hydrogen bonding gtgt
• Dipole-dipole interactions gtgtgtgtgt London forces

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What kind of IMF??
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Dipole-Dipole Attractions

• Dipoles are polar molecules
• Molecules w/ polar bonds and asymmetric
  distribution of charge
• What determines if a bond is nonpolar, polar or
  ionic?
• What determines if a molecule with polar bonds
  is polar or nonpolar
• Much weaker than covalent bonds
• Important in maintaining the shape of many
  biological molecules e.g. Proteins

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Hydrogen Bonds (H-FON Bonds)

• Special kind of dipole-dipole interaction
• Found in HF and molecules containing O-H or N-H
  bonds
• 5xs the strength of a typical dipole-dipole bond
• 5 the strength of a covalent bond
• Important in biological molecules
• e.g. DNA, proteins

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Myoglobin
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Hemoglobin
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Hydrogen bonding is responsible for the
expansion of water when it freezes.
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Exercise Which of the following molecules
display hydrogen bonding?

1. Methane, CH4
2. methyl ether, CH3OCH3
3. Hydrogen peroxide, H2O2
4. methyl alcohol, CH3OH

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London Forces Attractions between temporary
dipoles
Electrostatic Attraction
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Random movement of electrons may cause
temporary charge imbalances
London Force
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London (dispersion) forces between nonpolar
molecules
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Why are London forces the greatest in large
molecules?
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London Forces (dispersion forces) exist in all
molecules

• Ave. strength ltltlt Dipole-Dipole interactions
• Result from temporary charge imbalances
• Due to the random movement of electrons
• Nucleus of one atom attracts electrons from a
  neighboring atom.
• At the same time, the electrons in one particle
  repel the electrons in the neighbor and create a
  short lived charge imbalance.

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Relationship between atomic size and the strength
of London forces

• Greatest in large atoms
• Electron clouds more easily distorted
• Halogens and Noble Gases BP increase w/ molar
  mass
• Ion Induced Dipoles
• dipoles can be induced by ions
• attractions exist between ions and dipoles

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London Forces Effect of molecular surface area
Cyclopentane, BP 49.3 oC
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IMFs inNonpolar Organic Molecules

• What kind of attractive forces are present?
• What role do molecular size and surface area
  play?
• Linear molecules have more surface area than if
  they are folded into a sphere.
• Linear molecules have higher melting and boiling
  points because of the increased attractions.

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Predicting the Relative Boiling Points of
Substances

• The substance with the strongest intermolecular
  attractions will have the higher BP. Why?
• More energy needed to separate molecules ? higher
  boiling temperature
• e.g. Halogens and noble gases

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Cooling Curve H2O (g) ? H2O (s) What is
happening in to KE and PE at each part of the
curve?
DHfus 6.01 kJ/mol
DHvap 40.7 kJ/mol
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Molar Heat of Vaporization, DHvap

• Heat absorbed when one mole liquid is changed to
  one mole of vapor at constant T and P
• Depends on strength of IMFs
• Endothermic (results in PE elevation)
• For water DHvap 40.7 kJ/mol _at_ 100oC

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Molar Heat of fusion, DHvap

• Heat absorbed when 1 mole solid is changed to 1
  mole of liquid at constant T and P.
• Depends on strength of IMFs
• Endothermic (results in PE elevation)
• For water DHfus 6.01 kJ/mol

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Quantitative Aspects of Phase Changes Within a
phase A change in heat is associated with a
change in average KE and, therefore, a change in
temperature. q (mass)(Specific Heat)(Dt) During
a phase change A change in heat occurs at
constant temperature, which is associated with a
change in PE, as the average distance between
molecules changes--Bond IMF formation is
exothermic, IMF breaking is endothermic q
(moles of substance)(enthalpy of phase change)
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Calculation of the Heat of Fusion of Ice

• Use the data below to calculate the heat of
  fusion of water.
• A piece of ice at zero Celsius melts in 100.0 g
  water until the waters temperature also becomes
  zero.
• Initial water temp. 44.0 oC
• Mass of ice that melted 56.0 g.
• Specific heat of water, CH2O 4.184 J/g oC
• Calculate the error and explain the source of
  the error. DHfus H2O 6.01 kJ/mol

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Application Questions

• The molar heat of vaporization of water at 25 oC
  is 43.99 kJ/mol. How many kilojoules of heat
  would be required to vaporize 125 mL (125 g) of
  water at 25 oC?
• Answer 305kJ
• How much heat would be needed to convert 125 mL
  water at 25.oC to steam at 100.0 oC? The heat of
  vaporization of water at 100.0 oC is 40.657
  kJ/mol and the specific heat of water is 4.184
  J/goC.
• Answer 321 kJ

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General Properties of Liquids and Solids

• Macroscopic properties depend on Microscopic
  properties
• Microscopic properties of Ls and Ss
• Molecules tightly packed
• Strong intermolecular attractions

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Macroscopic properties of Ss and Ls

• Compressibility
• Little to none. Why?
• Diffusion (ability to flow) (T6)
• Slow in liquids
• Like moving in a crowded room
• Nonexistent in Solids
• Particles not free to move

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Macroproperties Liquids

• Why are raindrops spherical?
• Increases stability by maximizing the number of
  IMFs and decreasing surface tension
• Surface Tension
• E needed to increase the surface area of a liquid
  by a given amount (J/m2)
• Depends on nature of intermolecular forces

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The Molecular Basis of Surface Tension Surface
molecules experience fewer intermolecular forces
than interior molecules

• Liquids minimize surface area by forming
  spherical surfaces ? Lowers PE, thus increases
  stability ( e.g. Raindrops, overfilled glass)
• Surface molecules are at higher P.E. than
  interior molecules
• Recall Bond formation results in P.E. lowering

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Macroproperties Liquids

• Wetting of a Surface by a Liquid
• Spreading of a Liquid across a surface
• Caused by attraction of liq. molecules to surface
  molecules
• Why are Liquids with a Low Surface Tension Good
  Wetters?........
• Low Surface Tension means weak IM Forces
• Which wets solids surfaces better, hydrocarbons
  (gasoline/oil ) or water?.......

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Evaporation and Sublimation

• Where molecules leave surface and enter vapor
  space around them
• Evaporation L ? Vapor
• Sublimation S ? Vapor

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Evaporation L ?Vapor (T9)

• Factors that affect the rate of evaporation
• Surface area, Temp., Strength of IMFs
• Why does evaporation occur at temp.s below the
  BP?
• Why does an increase in temp. increase the rate
  of evaporation?
• Why does sweating cool you?
• Why does a fan cool you?

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Effect of Temperature on the Distribution of
Molecular Speeds in a Liquid
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Evaporation L ?Vapor Application
Questions (T10)

• If the liquids are water and ethanol, which one
  is liquid A? Liquid B?
• Why does the evaporation of ethanol from the skin
  feel cooler than the evap. of water?
• Which removes more K.E.?
• Which liquid will evaporate more quickly, acetone
  or ethanol?
• Which is more difficult to clean up, a marine
  spill of low or high MW crude oil?

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Application Questions (T12)

1. How can snow or ice cubes in the freezer
   disappear w/o melting?
2. Why do moth balls (naphthalene) need to be
   replaced periodically?

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Changes of State Dynamic Equilibrium

• Change of State (T11)
• Change from one physical state to another
• S L G or S
  G
• Occur under conditions of Dynamic Equilibriium
• Two opposing events occurring _at_ equal rates
• Application Questions
• Why does fanning cause you to feel cooler?
• Why does sweating in the tropics cool you less
  than sweating in the desert?

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Vapor Pressure of Solids Liquids

• Vapor Pressure
• Pressure exerted by a vapor in a closed flask in
  equilibrium w/ its liquid (T11)
• Measurement of vapor pressure (T 13a)

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Rateevap gt Ratecond
Rateevap Ratecond
Liquid Gas Equilibrium
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• Only Temperature and Intermolecular Forces affect
  Vapor Pressure
• The of molecules with sufficient K.E.to escape
  the liquid surface increases with Temperature

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Factors affecting Vapor Pressure

• Strength IMFs (T13b)
• Temperature
• At higher temps a higher of the molecules have
  suffiecient K.E. to escape the liquid surface
• Factors not Affecting VP
• Surface Area
• Increases the rate of evaporation and
  condensation equally
• Amount of Liquid
• Evaporation occurs at surface

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VP in Solids

• Due to vibrations of surface molecules ?
  Sublimation (T12)
• Examples
• Snow and ice cubes
• Dry ice (solid carbon dioxide)

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Boiling Points of Liquids

• Boiling Point
• Temp. at which VP of a liquid equals atmospheric
  pressure (T14)
• Depends on
• Strength of IMFs
• Atmospheric pressure
• Why do liquids Boil?
• Why do bubbles form on sides first?

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Application Questions

1. Explain why water can exist at temps above its
   normal BP in a cars radiator.
2. Explain why a pressure cooker can cook a beef
   stew faster than in a normal pot.
3. Explain why water, hydrogen fluoride and ammonia
   have much higher BPs than one might predict by
   size alone (T15)

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Clausius-Clapeyron Equation

• Relates the Vapor Pressure of a liquid to the
  heat of vaporation, -DHvap
• Ln P -DHvap/RT constant or
• Ln (P1/P2) -DHvap /R(1/T2 - 1/T1)
• Note P Vap Pressure R 8.314 J/mol K
• Our next lab experiment will involve the use of
  the Clausius-Clapeyron Equation to calculate the
  DHvap for methanol and ethanol

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Clausius-Clapeyron Eqn ln P -DHvap/RT
c P Vapor Pressure R 8.314 J/mol K
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Dynamic Equilibrium and Le Chateliers Principle

• Le Chateliers Principle When a stress is
  applied to a system at equilibrium, the
  equilibrium will shift to relieve that stress
  and, if possible, restore equilibrium.
• Position of Equilibrium Refers to relative
  amount of reactants and products
• Application Questions.................

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Le Chateliers Principle Application Questions

1. Use L.C.P. to predict how an increase in
   temperature will affect the vapor pressure of a
   solid.
2. Use L.C.P. to predict how a decrease in
   temperature will affect the vapor pressure of a
   liquid.
3. Use L.C.P. to predict how an increase in
   atmospheric pressure will affect the vapor
   pressure of a liquid.

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Phase Diagrams

• Objective
• Interpret phase diagrams and show how a phase
  diagram can be used to represent the
  thermodynamic relationship between the three
  states of matter for a particular substance.

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Phase Diagrams (T16)

• Lines represent equilibrium between phases
• Triple point
• T and P at which all three phases present _at_
  equilibrium
• Critical Temperature
• Temp at which liquid phase can not be
  distinguished from its vapor
• Supercritical Fluid
• Fluid at a Temp gt Critical Temp.
• e.g. Supercritical carbon dioxide
• Used to decaffeinate coffee and tea

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Application Questions

1. What phase will occur if water at 20. oC and
   2.15 torr is heated to 50. oC under constant
   pressure?
2. What phase will water be in if it is at a
   pressure of 330 torr and a temperature of 50 oC?
3. Use L.C.P. to explain why the MP of ice decreases
   as pressure increases.

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Application Questions (T17)

1. At what temperature does Dry ice sublime?
2. What effect does an increase in pressure have on
   the melting point of carbon dioxide. Use L.C.P.
   to explain why.

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Crystalline Solids

• Objective
• Describe how atoms, molecules, or ions are
  arranged in crystalline solids
• Unit Cells to Know
• Simple cubic (T18)
• Face-centered cubic (T19, 20, 21)
• Body-centered cubic (T21)

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X-Ray Diffraction

• Objective
• Describe the use of X-Ray diffraction to
  determine the structure of crystals (T23a)

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Physical Properties and Crystal Types

• Objective Relate the properties of solids to
  crystal type
• Understand Table 12.5, (T24a)
• Crystal Types
• Ionic
• Molecular
• covalent (Network)
• Metallic (T24c)

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Application Questions

1. Boron nitride, which has the empirical formula
   BN, melts under pressure at 3000 oC and is as
   hard as diamond. What is the probable crystal
   type of BN?
2. Crystals of elemental sulfur are easily crushed
   and melt at 113 oC to give a clear yellow liquid
   that does not conduct electricity. What is the
   probable crystal type for solid sulfur?

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Physical Properties of Graphite vs. Diamond
Property Graphite Diamond
Density (g/mL) 2.27 3.51
Hardness Very soft Very hard
Color Shiny black Colorless/transparent
Electrical Conductivity High None
DHcomb (kJ/mol) -393.5 -395.4
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Which Structure is Diamond? Graphite?
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Uses of Unit Cells

• Unit cells may be used to determine the
• Empirical formula of an ionic compound
• E.g. Problems 12.90 and 12.91, page 480
  Silberberg 3rd ed.
• Molar Mass of a substance
• E.g. Problem 12.93, page 480 Silberberg 3rd ed.
• Density of a substance
• (divide the mass of the atoms per unit cell by
  the volume of the unit cell)

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