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Lecture Notes Chem 150 - K. Marr. Chapter 12 Intermolecular Attractions & the Properties of Liquids & Solids Silberberg 3 ed. Intermolecular Forces: Liquids, Solids, and Phase Changes. 12.1 An Overview of Physical States and Phase Changes 12.2 Quantitative Aspects of Phase Changes

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lecture notes chem 150 k marr
Lecture Notes Chem 150 - K. Marr

Chapter 12

Intermolecular Attractions &

the Properties of Liquids & Solids

Silberberg 3 ed

slide2

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

slide3

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

chapter 12 intermolecular attractions the properties of liquids solids
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

why aren t there liquid laws and solid laws
Why aren’t there liquid laws and solid laws?
  • Little empty space between molecules:

Particles close together in S’s and L’s

  • Stronger and quite varied intermolecular attractionsin S’s & L’s 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. @ -58.5 oC)

intermolecular attractions bonds between molecules
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.

kinds of intermolecular attractions
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
relative magnitudes of forces in molecular compounds
Relative magnitudes of forces in Molecular compounds

Covalent bonds>>>>>Hydrogen bonding>>

Dipole-dipole interactions>>>>> London forces

dipole dipole attractions
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
hydrogen bonds h fon bonds
Hydrogen Bonds (H-FON Bonds)
  • Special kind of dipole-dipole interaction
    • Found in HF and molecules containing O-H or N-H bonds
  • 5x’s the strength of a typical dipole-dipole bond
  • ~ 5% the strength of a covalent bond
  • Important in biological molecules

e.g. DNA, proteins

exercise which of the following molecules display hydrogen bonding
Exercise: Which of the following molecules display hydrogen bonding?
  • Methane, CH4
  • methyl ether, CH3OCH3
  • Hydrogen peroxide, H2O2
  • methyl alcohol, CH3OH
london forces dispersion forces exist in all molecules
London Forces (dispersion forces) exist in all molecules
  • Ave. strength <<< 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.
relationship between atomic size and the strength of london forces
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
imf s in nonpolar organic molecules
IMF’s 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.
predicting the relative boiling points of substances
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

slide32

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

molar heat of vaporization d h vap
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 IMF’s
  • Endothermic (results in PE elevation)
  • For water: DHvap = + 40.7 kJ/mol @ 100oC
molar heat of fusion d h vap
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 IMF’s
  • Endothermic (results in PE elevation)
  • For water: DHfus = + 6.01 kJ/mol
slide36

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)

calculation of the heat of fusion of ice
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 water’s 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
application questions
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
general properties of liquids and solids
General Properties of Liquids and Solids
  • Macroscopic properties depend on Microscopic properties
  • Microscopic properties of L’s and S’s
    • Molecules tightly packed
    • Strong intermolecular attractions
macroscopic properties of s s and l s
Macroscopic properties of S’s and L’s
  • 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
macroproperties liquids
Macroproperties Liquids

Why are raindrops spherical?

  • Increases stability by maximizing the number of IMF’s 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
slide43

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
macroproperties liquids48
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?.......
evaporation and sublimation
Evaporation and Sublimation

Where molecules leave surface and enter vapor space around them

  • Evaporation: L  Vapor
  • Sublimation: S  Vapor
evaporation l vapor t9
Evaporation: L Vapor (T9)
  • Factors that affect the rate of evaporation
    • Surface area, Temp., Strength of IMF’s
  • 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?
slide51

Effect of Temperature on the

  • Distribution of Molecular Speeds in a Liquid
evaporation l vapor application questions t10
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?
application questions t12
Application Questions (T12)
  • How can snow or ice cubes in the freezer disappear w/o melting?
  • Why do moth balls (naphthalene) need to be replaced periodically?
changes of state dynamic equilibrium
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 @ 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?
vapor pressure of solids liquids
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)
slide56

Rateevap > Ratecond

  • Rateevap = Ratecond
  • Liquid – Gas Equilibrium
slide57

Only Temperature and Intermolecular Forces affect Vapor Pressure

  • The % of molecules with sufficient K.E.to escape the liquid surface increases with Temperature
factors affecting vapor pressure
Factors affecting Vapor Pressure
  • Strength IMF’s (T13b)
  • Temperature
    • At higher temps a higher % of the molecules have suffiecient K.E. to escape the liquid surface
  • Factors notAffecting VP
    • Surface Area
      • Increases the rate of evaporation and condensation equally
    • Amount of Liquid
      • Evaporation occurs at surface
vp in solids
VP in Solids
  • Due to vibrations of surface molecules Sublimation (T12)

Examples

        • Snow and ice cubes
        • Dry ice (solid carbon dioxide)
boiling points of liquids
Boiling Points of Liquids
  • Boiling Point
    • Temp. at which VP of a liquid equals atmospheric pressure (T14)
    • Depends on
        • Strength of IMF’s
        • Atmospheric pressure
  • Why do liquids Boil?
  • Why do bubbles form on sides first?
application questions61
Application Questions
  • Explain why water can exist at temps above its normal BP in a car’s radiator.
  • Explain why a pressure cooker can cook a beef stew faster than in a normal pot.
  • Explain why water, hydrogen fluoride and ammonia have much higher BP’s than one might predict by size alone (T15)
clausius clapeyron equation
Clausius-Clapeyron Equation
  • Relates the Vapor Pressure of a liquid to the heat of vaporation, -DHvap

Ln P = -DHvap/RT + constantor

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
slide63

Clausius-Clapeyron Eqn:

ln P = -DHvap/RT + c

P = Vapor Pressure R = 8.314 J/mol K

dynamic equilibrium and le chatelier s principle
Dynamic Equilibrium and Le Chatelier’s Principle
  • Le Chatelier’s 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.................

le chatelier s principle application questions
Le Chatelier’s Principle : Application Questions
  • Use L.C.P. to predict how an increase in temperature will affect the vapor pressure of a solid.
  • Use L.C.P. to predict how a decrease in temperature will affect the vapor pressure of a liquid.
  • Use L.C.P. to predict how an increase in atmospheric pressure will affect the vapor pressure of a liquid.
phase diagrams
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.

phase diagrams t16
Phase Diagrams(T16)
  • Lines represent equilibrium between phases
  • Triple point
    • T and P at which all three phases present @ equilibrium
  • Critical Temperature
    • Temp at which liquid phase can not be distinguished from its vapor
  • Supercritical Fluid
    • Fluid at a Temp > Critical Temp.
      • e.g. Supercritical carbon dioxide

Used to decaffeinate coffee and tea

application questions69
Application Questions
  • What phase will occur if water at –20. oC and 2.15 torr is heated to 50. oC under constant pressure?
  • What phase will water be in if it is at a pressure of 330 torr and a temperature of 50 oC?
  • Use L.C.P. to explain why the MP of ice decreases as pressure increases.
application questions t17
Application Questions (T17)
  • At what temperature does Dry ice sublime?
  • What effect does an increase in pressure have on the melting point of carbon dioxide. Use L.C.P. to explain why.
crystalline solids
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)
x ray diffraction
X-Ray Diffraction

Objective:

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

physical properties and crystal types
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)
application questions76
Application Questions
  • 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?
  • 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?
uses of unit cells
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)