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Okay, let's break this down step-by-step: * We are given: 1 mole of ice at -25°C * Heat of fusion of ice = 6.01 kJ/mol * Heat of vaporization of water = 40.7 kJ/mol * Specific heat of ice = 2.09 J/g°C * Specific heat of water = 4.18 J/g°C 1) Heat ice from -25°C to 0°C: Q = m * c * ΔT Q = (18 g) * (2.09 J/g°C) * (25°C) = 903 J 2) Heat

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1 Chapter 10 Liquids and Solids
3 Van Der Waals Forces • These are intermolecular forces of attraction between neutral molecules. • The Nobel Prize in Physics 1910 (Johannes van der Waals) • "for his work on the equation of state for gases and liquids"
4 intER vs. intRA molecular forces • Intramolecular forces are the forces within a molecule or ionic compound Example: Individual therapy NaCl Ionic bond between atom of Na and atom of Cl • Intermolecular forces are the forces between molecules or ions and molecules Example: couples therapy Solid liquid gas
5 Intramolecular forces Intermolecular forces
Strength Intramolecular bonds > intermolecular forces Intramolecular bonds are stronger because it would take a lot more energy to overcome covalent bonds and break apart the molecule than to overcome intermolecular forces in between the atoms (to make it become a liquid or gas). 6
Phase Changes • When a substance changes from solid to liquid to gas, the molecules remain intact. • The changes in state are due to changes in the forces among molecules rather than in those within the molecules. 7
10 Intermolecular Force Model Basis of Attraction Energy (kJ/mol) Example
11 3 Types of van der Waals Forces • Dipole-Dipole forces • London Dispersion forces • Hydrogen bonding
12
13 DIPOLE-DIPOLE FORCES • These are forces of attraction that occur between polar molecules. (big difference in electron negativity) • These forces are effective only when polar molecules are very close. As distance increase strength of bond decreases. • For molecules of approximately equal mass and size, the strength of force of attraction increases as the polarity increases. • Radius have an effect on strength of dipole.
Dipole-Dipole Forces Copyright © Cengage Learning. All rights reserved 14 To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE
15 DIPOLE-DIPOLE FORCES + _ _ _ + + + _
16 FYI
17 DIPOLE-DIPOLE FORCES • Molecules with larger dipole moments have higher melting and boiling points (hard to break) than those with small dipole moments. • Dipole attractions are relatively weak and tend to be liquids or gas at room temperature.
18 HYDROGEN BONDING • A special type of dipole-dipole interaction between the hydrogen atom in a polar bond and an unshared electron pair of an element that is very electronegative usually a F, O, or N atom on another molecule • (note that all of these have very high EN’s and small atomic radii).
Hydrogen Bonding Copyright © Cengage Learning. All rights reserved 19 To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE
20
21 HYDROGEN BONDING • These types of bonds are super- humanly strong. • (unusually strong dipole dipole 4X stronger that diopole dipole)
22 HYDROGEN BONDING
23 WHY HYDROGEN BONDING IS EFFECTIVE • F, O, & N are extremely small and very electronegative atoms. • Hydrogen atoms are very small and have no inner core of electrons, therefore, the positive side of the bond dipole has the concentrated charge of the partially exposed, nearly bare proton of the nucleus. • …in other words, the atoms have a large difference in electronegativity and their nuclei can get really close.
24 IMPORTANCE OF HYDROGEN BONDING • Are important biologically, in stabilizing proteins and keeping DNA together. • Also explains why ice is less dense than water (see text).
25
26 LONDON DISPERSION FORCES • Fritz London • These are forces that arise as a result of temporary dipoles induced in the atoms or molecules.( it’s a temporary accident!) • All molecules have some degree of LD forces
London Dispersion Forces Copyright © Cengage Learning. All rights reserved 27 To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE
29 LONDON DISPERSION FORCES • LD forces occur between neutral non-polar molecules. (nobles gases and nonpolar compounds) • Occurs in all molecules, including nonpolar ones. • LD forces are weak • The greater the number of electrons the greater the LD force. (ie the greater the melting and boiling pt.). Say This: The larger the electron cloud the more polarizable an the greater the strength of the interaction • LD force molecules have Low melting and boiling pts
30 See Graphic on next slide • The motion of electrons in an atom or molecule can create an instantaneous dipole moment. • EX: in a collection of He (g) the average distribution of electrons about a nucleus is spherical, the molecules are non-polar and there is no attraction.
31 INSTANTANEOUS AND INDUCED DIPOLES Pg 454- 455 in text
32 LONDON DISPERSION FORCES (CONT) • These forces tend to increase in strength with an increase in molecular weight (The size of the molecule generally increases with mass and the electrons are less tightly held…allows the electron cloud to be more easily distorted. • These forces are stronger in linear molecules than comparable “bunched up” molecules.
33 LONDON DISPERSION FORCES LD forces are generally the WEAKEST intermolecular forces. Molecules with more electrons will experience more LD forces
Copyright © Cengage Learning. All rights reserved 34 Which molecule is capable of forming stronger intermolecular forces? N2 H2O Explain. CONCEPT CHECK!
Copyright © Cengage Learning. All rights reserved 35 Draw two Lewis structures for the formula C2H6O and compare the boiling points of the two molecules. C H H C H H H O H C H H C H H H O H CONCEPT CHECK!
Copyright © Cengage Learning. All rights reserved 36 Which gas would behave more ideally at the same conditions of P and T? CO or N2 Why? CONCEPT CHECK!
37
38 n-pentane vs neopentane • BP = 309.4 K BP = 282.7 K • Same atomic masses different structure
39 Generalizations Regarding Relative Strengths of IM Forces • If molecules have comparable molecular weights and shapes, dispersion forces are approximately equal. Any difference in attractive forces is due to dipole-dipole attractions. • If molecules differ widely in molecular weight, dispersion forces are the decisive factor. The most massive molecule has the strongest attractions.
40 Because melting points (MPs) and boiling points (BPs) of covalent molecules increase with the strengths of the forces holding them together, it is common to use MPs and BPs as a way to compare the strengths of intermolecular forces. This is shown below, with the molecular formulas, molar masses and normal BPs of the first five straight-chain hydrocarbons. Molecular Formula Molar Mass Normal BP (C) CH4 16 - 161.5 C2H6 30 - 88.5 C3H8 44 - 42.1 C4H10 58 - 0.5 C5H12 72 36.1
41 Which noble gas element has the lowest boiling point? He Ne Ar Kr Xe
42 The chemical forces between HCl is/are • Dispersion • Covalent bond • Hydrogen bond • Dipol-dipole • Two of the above All Molecules Have Not symmetrical Polar
43 Consider the following list of compounds. How many of these have hydrogen bonding as their principle IMF HCl NH3 CH3OH H2S CH4 PH3 Hydrogen Bonding is between H and highley EN atoms such as N, O, F, and H
44 Which of the following statements are false or correct and why? O2 is dipole dipole HCl is hydrogen bonding CO2 is dipole dipole NH3 is hydrogen FALSE Dipole dipole not symmetric/polar FALSE London Dispersion symmetric/nonpolar TRUE H + N,O, or F FALSE London Dispersion symmetric/nonpolar
45 ION-DIPOLE FORCES • Attraction between an ion and the partial charge on the end of a polar molecule.
46 ION-DIPOLE FORCES (CONT) • The magnitude of attraction increases as either the charge of the ion increases or magnitude of the dipole moment increases. • Ion-dipole forces are important in solutions of ionic substances in polar liquids (e.g. water) • Stronger than Hydrogen bonding
47 ION-DIPOLE FORCES AND THE SOLUTION PROCESS
48
49 ANSWER
50 A.Identify the types of bonds in 1. Glucose 2. Cyclohexane B.Glucose is soluble in water but cyclohexane is not. Why?
51 A.Identify the types of bonds in 1. Glucose H, LD, VanderWal, Dip-dip 2. Cyclohexane LD only B.Glucose is soluble in water but cyclohexane is not. Why? Glucose is polar and cyclohexane is nonpolar. Polar compounds are soluble in polar solvent and visversa.
52
53 Homework • Pg 504-505 #’s : 35, 36, 37, 39 (you may need to read 10.1 for this part esp. LD portion)
55 CHARACTERISTICS OF LIQUIDS • Surface tension • Capillary action • Viscosity
56 COHESIVE FORCES • Intermolecular forces that bind like molecules to one another (e.g. hydrogen bonding).
• Which force dominates alongside the glass tube – cohesive or adhesive forces? cohesive forces Convex Meniscus Formed by Nonpolar Liquid Mercury Copyright © Cengage Learning. All rights reserved 57
58 ADHESIVE FORCES • Intermolecular forces that bind a substance to a surface.
Concave Meniscus Formed by Polar Water • Which force dominates alongside the glass tube – cohesive or adhesive forces? adhesive forces Copyright © Cengage Learning. All rights reserved 59
60 SURFACE TENSION • A measure of the inward forces that must be overcome in order to expand the surface area of a liquid and resist and external force. • The greater the forces of attraction between molecules (IMF) of the liquid, the greater the surface tension.
61 Surface Tension Cont. • Surface tension of a liquid decreases with increasing temperature. • The stronger the intermolecular forces the stronger the surface tension. Water has a high surface tension do to hydrogen bonding.
62 CAPILLARY ACTION • Another way surface tension manifests. • The rise of liquids up very narrow tubes. This is limited by adhesive and cohesive forces.
63 Formation of meniscus • Water : adhesive forces are greater than cohesive forces • Mercury: Cohesive are greater than adhesive forces.
64 VISCOSITY • The resistance of a liquid to flow. • The less “tangled” a molecule is expected to be, the less viscous it is. Water = less Viscosity syrup = high Viscosity Larger molecules stronger IM
65 Viscosity Cont. • Viscosity decreases with increasing temperature (molecules gain kinetic energy and can more easily overcome forces of attraction). • Viscosity Increases as pressure increases. • Liquids with strong IMF have a higher viscosity.
66 Homework • Pg 505 #’s 43-45 all
Four types of solids 67
Ionic Solids 68 • Formed by cation and anion. Typicaly metals and non metals. • Crystalline solids • High melting points and Boiling point due to strong attractions • Poor conductors in solid form good conductors in solution. • Brittle
Copyright © Cengage Learning. All rights reserved 69 To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE
Molecular Solids • Neutral molecules that form molecular lattice structures • Low MP and BP • Non conductors in all states 70
Copyright © Cengage Learning. All rights reserved 71 To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE
Covalent network solids • Distinct atoms all bound covalently. • High MP and BP • Made of carbons and Si Ge and B • Poor conductors except for graphite sp2 hybridization and delocalized electrons 72 Diamond
Network Solids Copyright © Cengage Learning. All rights reserved 73 To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE
Metallic Solids • Metallic bonding • Great conductors or heat and electricity • Ductile and malleable • MP and BP vary 74
ID The type of Solid based on the formula 75
76
77 Malleability Brittle Nothing flexible polymer
92 Distillation Vaporization AKA: steam Condensation Takes advantage differences in IMF and thus vapor pressures and therefore their BP . Higher VP = Lower BP and weaker IMF
93 Changes in state • Liquid  Gas Vaporization Endothermic • Gas  Liquid Condensation Exothermic
94 • Solid  Gas Sublimation Endothermic • Gas  Solid Deposition Exothermic
95 • Solid  Liquid Melting Endothermic • Liquid  Solid Freezing Exothermic
96 Changes of state • The energy involved it phase changes is calculated using – Heat of fusion (solid  liquid or liquid solid) – Heat of vaporization (liquid gas or gas liquid)
97 Energy Changes and Phase Changes Heat of Vaporization: Vaporization is an endothermic process ( it requires heat). Energy is required to overcome intermolecular forces to turn liq to gas. (AKA evaporate) Hvap is an Indicator of strength of IMF Larger molecule…greater IMF…greater Hvap Methane Propane CH4 C3H8 9.2 kJ/mol 18.1 kJ/mol
98 Question How much energy does it take to vaporizer 111 g of water? Given: Hvap water= 40.67 kJ/mol 111 g H2O 1 mol x 40.6kJ = 250kJ 18g 1mol
99 • Heat of Fusion: the enthalpy change associated with melting. (Solid to liquid.) • Hfusion water= 6.01 kJ/mol NOTE: heat of fusion is always smaller than heat of vaporization. This makes sense think about the level of “order” in the molecules in these phases.
Which is larger for a given substance: ΔHvap or ΔHfus? Explain why. Copyright © Cengage Learning. All rights reserved 100 CONCEPT CHECK!
Changes of State Copyright © Cengage Learning. All rights reserved 101 To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE
102 Heating Curve • A plot of the temperature versus time
103 Heat of Vaporization Heat of Fusion
104 Example Calculate the enthalpy change associated with converting 1.00 mole of ice -25ºC to water 150ºC at 1 atm. Specific heat of ice, water, and steam are 2.09 J/g ºC and 4.184 J/g ºC, 1.84 J/g ºC . The heat of fusion of ice is 6.01 kJ/mol and heat of vaporization of water is 40.67 kJ/mol
105 ICE -25 °C q1 Ice water 0°C q2 Water 0°C q3 Water vapor 100°C q4 Vapor 150°C
106 1 mol ice  1 mol ice 1 mol water  1 mol water  1 mol steam T= -25ºC 0ºC 0ºC 100ºC 100ºC -> 150 qtotal = q1 + q2 + q3 + q4 + q5 1.)q = 2.09(18g)(-25-0) 2.) q = 6.02 KJ/mol (convert heat of fusion) 3.) q = 4.184(18g) (100-0) 4.) q = 40.7 KJ/mol (convert heat of vaporization) 5. ) q = 1.84(18 g) (150-100)
107 Critical Stuff • Critical Temperature: The temperature above which it is impossible to liquefy the gas under study no matter how high the applied pressure. • Critical Pressure: The pressure required to liquefy a gas as at its critical temperature NOTE: the critical temp of a gas gives an indication of the strength of the IMF of that gas. A substance with weak attractive forces would have a low critical temp.
108 Which gas can be liquefied at 25ºC Gas Critical Temp ºC Critical Pressure atm Ammonia 132 112 Ethanol 158 78 Argon -186 6 Critical Temp above 25ºC Critical Temp under 25ºC
109 Vapor Pressure (vp) Vapor Pressure: Pressure exerted by molecules that have enough energy to escape the surface. As T ↑ VP ↑evaporation ↑ Liquids with high VP are volatile (alcohol evaporates easily) Liquids that have strong IMF have low vapor pressures. (take a lot of energy to overcome IMF so it can evaporate)
Kinetic energy % o f M o l e c u l e s T2 • At higher temperature more molecules have enough energy • Higher vapor pressure.
111 • Liquids with high VP are volatile (alcohol evaporates easily) • Liquids that have strong IMF have low vapor pressures. • (take a lot of energy to overcome IMF so it can evaporate) substance vapor pressure at 25oC diethyl ether C4H10O 0.7 atm Bromine Br2 0.3 atm ethyl alcohol C2H5OH 0.08 atm Water H2O 0.03 atm
Evaporation • Molecules at the surface break away and become gas. • Only those with enough KE escape • Evaporation is a cooling process. • It requires heat. • Endothermic.
Condensation Change from gas to liquid Achieves a dynamic equilibrium with vaporization in a closed system. What the heck is a “dynamic equilibrium?”
Liquid/Vapor Equilibrium Copyright © Cengage Learning. All rights reserved 114 To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE
When first sealed the molecules gradually escape the surface of the liquid As the molecules build up above the liquid some condense back to a liquid. Dynamic equilibrium
As time goes by the rate of vaporization remains constant but the rate of condensation increases because there are more molecules to condense. Equilibrium is reached when Rate of Vaporization = Rate of Condensation Dynamic equilibrium
117 VP example In a closed container the number of partials changing from liquid  vapor will eventually equal the number changing from vapor  liquid.
118 Boiling Point Note: The normal boiling point of water is 100oC. The term normal refers to standard pressure or 1 atm, or also 101.3 kPa. The vapor pressure of the liquid = air pressure above the liquid
119 Boiling Pts. of H2O at Various Elevations Altitude compared to Sea Level (m) Boiling Point (°C) 1609 98.3 177 100.3
120 How to make something boil 1. Increase the VP of the liquid (heat it) so that the VP of the liquid is > that of the atmosphere. 2. Lower the atmospheric pressure (pressure above the liquid) (put a lid on it)
121 Boiling Point ↑ boiling pt by ↑ in IMF Or ↓ VP At high altitudes (low air pressure) water boils at a lower temperature
122 Normal Boiling Point • Temperature at which something boils when the vp =1 atm • Note the lower the external pressure the lower the boiling point.
123 Freezing point/melting point • They are the same but in opposite directions. • When heated the particles vibrate more rapidly until they shake themselves free of each other. • Ionic solids have strong intermolecular forces so a high mp. • Covalent/molecular solids have weak intermolecular forces so a low mp.
124 Phase Diagram • A graphical way to summarize the conditions under which equilibrium exists between different states of matter. • Allows you to predict the phase of a substance that is stable at a given temperature and pressure
125 Triple point = three phase are in equilibrium with each other at the same time Boiling Point Melting Point 1 atm Critical point
126 Critical Point: The temp beyond which the ,molecules of a substance have to much kinetic energy to stick together to form a liquid.
127 Phase diagrams of substances other than water the slope of the solid liquid line slopes forward. (positive) In water the slope of the solid-liquid lines slopes downward. (negative) Water Not Water
128 Homework • Pg 508 • #’s : 85, 87,89, 91,

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intern_forc_biceps2.ppt

  • 1.
  • 2.
    3 Van Der WaalsForces • These are intermolecular forces of attraction between neutral molecules. • The Nobel Prize in Physics 1910 (Johannes van der Waals) • "for his work on the equation of state for gases and liquids"
  • 3.
    4 intER vs. intRAmolecular forces • Intramolecular forces are the forces within a molecule or ionic compound Example: Individual therapy NaCl Ionic bond between atom of Na and atom of Cl • Intermolecular forces are the forces between molecules or ions and molecules Example: couples therapy Solid liquid gas
  • 4.
  • 5.
    Strength Intramolecular bonds >intermolecular forces Intramolecular bonds are stronger because it would take a lot more energy to overcome covalent bonds and break apart the molecule than to overcome intermolecular forces in between the atoms (to make it become a liquid or gas). 6
  • 6.
    Phase Changes • Whena substance changes from solid to liquid to gas, the molecules remain intact. • The changes in state are due to changes in the forces among molecules rather than in those within the molecules. 7
  • 7.
    10 Intermolecular Force Model Basis ofAttraction Energy (kJ/mol) Example
  • 8.
    11 3 Types ofvan der Waals Forces • Dipole-Dipole forces • London Dispersion forces • Hydrogen bonding
  • 9.
  • 10.
    13 DIPOLE-DIPOLE FORCES • Theseare forces of attraction that occur between polar molecules. (big difference in electron negativity) • These forces are effective only when polar molecules are very close. As distance increase strength of bond decreases. • For molecules of approximately equal mass and size, the strength of force of attraction increases as the polarity increases. • Radius have an effect on strength of dipole.
  • 11.
    Dipole-Dipole Forces Copyright ©Cengage Learning. All rights reserved 14 To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE
  • 12.
  • 13.
  • 14.
    17 DIPOLE-DIPOLE FORCES • Moleculeswith larger dipole moments have higher melting and boiling points (hard to break) than those with small dipole moments. • Dipole attractions are relatively weak and tend to be liquids or gas at room temperature.
  • 15.
    18 HYDROGEN BONDING • Aspecial type of dipole-dipole interaction between the hydrogen atom in a polar bond and an unshared electron pair of an element that is very electronegative usually a F, O, or N atom on another molecule • (note that all of these have very high EN’s and small atomic radii).
  • 16.
    Hydrogen Bonding Copyright ©Cengage Learning. All rights reserved 19 To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE
  • 17.
  • 18.
    21 HYDROGEN BONDING • Thesetypes of bonds are super- humanly strong. • (unusually strong dipole dipole 4X stronger that diopole dipole)
  • 19.
  • 20.
    23 WHY HYDROGEN BONDINGIS EFFECTIVE • F, O, & N are extremely small and very electronegative atoms. • Hydrogen atoms are very small and have no inner core of electrons, therefore, the positive side of the bond dipole has the concentrated charge of the partially exposed, nearly bare proton of the nucleus. • …in other words, the atoms have a large difference in electronegativity and their nuclei can get really close.
  • 21.
    24 IMPORTANCE OF HYDROGEN BONDING •Are important biologically, in stabilizing proteins and keeping DNA together. • Also explains why ice is less dense than water (see text).
  • 22.
  • 23.
    26 LONDON DISPERSION FORCES •Fritz London • These are forces that arise as a result of temporary dipoles induced in the atoms or molecules.( it’s a temporary accident!) • All molecules have some degree of LD forces
  • 24.
    London Dispersion Forces Copyright© Cengage Learning. All rights reserved 27 To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE
  • 25.
    29 LONDON DISPERSION FORCES •LD forces occur between neutral non-polar molecules. (nobles gases and nonpolar compounds) • Occurs in all molecules, including nonpolar ones. • LD forces are weak • The greater the number of electrons the greater the LD force. (ie the greater the melting and boiling pt.). Say This: The larger the electron cloud the more polarizable an the greater the strength of the interaction • LD force molecules have Low melting and boiling pts
  • 26.
    30 See Graphic onnext slide • The motion of electrons in an atom or molecule can create an instantaneous dipole moment. • EX: in a collection of He (g) the average distribution of electrons about a nucleus is spherical, the molecules are non-polar and there is no attraction.
  • 27.
  • 28.
    32 LONDON DISPERSION FORCES (CONT) •These forces tend to increase in strength with an increase in molecular weight (The size of the molecule generally increases with mass and the electrons are less tightly held…allows the electron cloud to be more easily distorted. • These forces are stronger in linear molecules than comparable “bunched up” molecules.
  • 29.
    33 LONDON DISPERSION FORCES LDforces are generally the WEAKEST intermolecular forces. Molecules with more electrons will experience more LD forces
  • 30.
    Copyright © CengageLearning. All rights reserved 34 Which molecule is capable of forming stronger intermolecular forces? N2 H2O Explain. CONCEPT CHECK!
  • 31.
    Copyright © CengageLearning. All rights reserved 35 Draw two Lewis structures for the formula C2H6O and compare the boiling points of the two molecules. C H H C H H H O H C H H C H H H O H CONCEPT CHECK!
  • 32.
    Copyright © CengageLearning. All rights reserved 36 Which gas would behave more ideally at the same conditions of P and T? CO or N2 Why? CONCEPT CHECK!
  • 33.
  • 34.
    38 n-pentane vs neopentane •BP = 309.4 K BP = 282.7 K • Same atomic masses different structure
  • 35.
    39 Generalizations Regarding Relative Strengthsof IM Forces • If molecules have comparable molecular weights and shapes, dispersion forces are approximately equal. Any difference in attractive forces is due to dipole-dipole attractions. • If molecules differ widely in molecular weight, dispersion forces are the decisive factor. The most massive molecule has the strongest attractions.
  • 36.
    40 Because melting points(MPs) and boiling points (BPs) of covalent molecules increase with the strengths of the forces holding them together, it is common to use MPs and BPs as a way to compare the strengths of intermolecular forces. This is shown below, with the molecular formulas, molar masses and normal BPs of the first five straight-chain hydrocarbons. Molecular Formula Molar Mass Normal BP (C) CH4 16 - 161.5 C2H6 30 - 88.5 C3H8 44 - 42.1 C4H10 58 - 0.5 C5H12 72 36.1
  • 37.
    41 Which noble gaselement has the lowest boiling point? He Ne Ar Kr Xe
  • 38.
    42 The chemical forcesbetween HCl is/are • Dispersion • Covalent bond • Hydrogen bond • Dipol-dipole • Two of the above All Molecules Have Not symmetrical Polar
  • 39.
    43 Consider the followinglist of compounds. How many of these have hydrogen bonding as their principle IMF HCl NH3 CH3OH H2S CH4 PH3 Hydrogen Bonding is between H and highley EN atoms such as N, O, F, and H
  • 40.
    44 Which of thefollowing statements are false or correct and why? O2 is dipole dipole HCl is hydrogen bonding CO2 is dipole dipole NH3 is hydrogen FALSE Dipole dipole not symmetric/polar FALSE London Dispersion symmetric/nonpolar TRUE H + N,O, or F FALSE London Dispersion symmetric/nonpolar
  • 41.
    45 ION-DIPOLE FORCES • Attractionbetween an ion and the partial charge on the end of a polar molecule.
  • 42.
    46 ION-DIPOLE FORCES (CONT) •The magnitude of attraction increases as either the charge of the ion increases or magnitude of the dipole moment increases. • Ion-dipole forces are important in solutions of ionic substances in polar liquids (e.g. water) • Stronger than Hydrogen bonding
  • 43.
    47 ION-DIPOLE FORCES ANDTHE SOLUTION PROCESS
  • 44.
  • 45.
  • 46.
    50 A.Identify the typesof bonds in 1. Glucose 2. Cyclohexane B.Glucose is soluble in water but cyclohexane is not. Why?
  • 47.
    51 A.Identify the typesof bonds in 1. Glucose H, LD, VanderWal, Dip-dip 2. Cyclohexane LD only B.Glucose is soluble in water but cyclohexane is not. Why? Glucose is polar and cyclohexane is nonpolar. Polar compounds are soluble in polar solvent and visversa.
  • 48.
  • 49.
    53 Homework • Pg 504-505 #’s: 35, 36, 37, 39 (you may need to read 10.1 for this part esp. LD portion)
  • 50.
    55 CHARACTERISTICS OF LIQUIDS •Surface tension • Capillary action • Viscosity
  • 51.
    56 COHESIVE FORCES • Intermolecularforces that bind like molecules to one another (e.g. hydrogen bonding).
  • 52.
    • Which forcedominates alongside the glass tube – cohesive or adhesive forces? cohesive forces Convex Meniscus Formed by Nonpolar Liquid Mercury Copyright © Cengage Learning. All rights reserved 57
  • 53.
    58 ADHESIVE FORCES • Intermolecularforces that bind a substance to a surface.
  • 54.
    Concave Meniscus Formedby Polar Water • Which force dominates alongside the glass tube – cohesive or adhesive forces? adhesive forces Copyright © Cengage Learning. All rights reserved 59
  • 55.
    60 SURFACE TENSION • Ameasure of the inward forces that must be overcome in order to expand the surface area of a liquid and resist and external force. • The greater the forces of attraction between molecules (IMF) of the liquid, the greater the surface tension.
  • 56.
    61 Surface Tension Cont. •Surface tension of a liquid decreases with increasing temperature. • The stronger the intermolecular forces the stronger the surface tension. Water has a high surface tension do to hydrogen bonding.
  • 57.
    62 CAPILLARY ACTION • Anotherway surface tension manifests. • The rise of liquids up very narrow tubes. This is limited by adhesive and cohesive forces.
  • 58.
    63 Formation of meniscus •Water : adhesive forces are greater than cohesive forces • Mercury: Cohesive are greater than adhesive forces.
  • 59.
    64 VISCOSITY • The resistanceof a liquid to flow. • The less “tangled” a molecule is expected to be, the less viscous it is. Water = less Viscosity syrup = high Viscosity Larger molecules stronger IM
  • 60.
    65 Viscosity Cont. • Viscositydecreases with increasing temperature (molecules gain kinetic energy and can more easily overcome forces of attraction). • Viscosity Increases as pressure increases. • Liquids with strong IMF have a higher viscosity.
  • 61.
  • 62.
    Four types ofsolids 67
  • 63.
    Ionic Solids 68 • Formedby cation and anion. Typicaly metals and non metals. • Crystalline solids • High melting points and Boiling point due to strong attractions • Poor conductors in solid form good conductors in solution. • Brittle
  • 64.
    Copyright © CengageLearning. All rights reserved 69 To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE
  • 65.
    Molecular Solids • Neutralmolecules that form molecular lattice structures • Low MP and BP • Non conductors in all states 70
  • 66.
    Copyright © CengageLearning. All rights reserved 71 To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE
  • 67.
    Covalent network solids •Distinct atoms all bound covalently. • High MP and BP • Made of carbons and Si Ge and B • Poor conductors except for graphite sp2 hybridization and delocalized electrons 72 Diamond
  • 68.
    Network Solids Copyright ©Cengage Learning. All rights reserved 73 To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE
  • 69.
    Metallic Solids • Metallicbonding • Great conductors or heat and electricity • Ductile and malleable • MP and BP vary 74
  • 70.
    ID The typeof Solid based on the formula 75
  • 71.
  • 72.
  • 73.
    92 Distillation Vaporization AKA: steam Condensation Takes advantagedifferences in IMF and thus vapor pressures and therefore their BP . Higher VP = Lower BP and weaker IMF
  • 74.
    93 Changes in state •Liquid  Gas Vaporization Endothermic • Gas  Liquid Condensation Exothermic
  • 75.
    94 • Solid Gas Sublimation Endothermic • Gas  Solid Deposition Exothermic
  • 76.
    95 • Solid Liquid Melting Endothermic • Liquid  Solid Freezing Exothermic
  • 77.
    96 Changes of state •The energy involved it phase changes is calculated using – Heat of fusion (solid  liquid or liquid solid) – Heat of vaporization (liquid gas or gas liquid)
  • 78.
    97 Energy Changes andPhase Changes Heat of Vaporization: Vaporization is an endothermic process ( it requires heat). Energy is required to overcome intermolecular forces to turn liq to gas. (AKA evaporate) Hvap is an Indicator of strength of IMF Larger molecule…greater IMF…greater Hvap Methane Propane CH4 C3H8 9.2 kJ/mol 18.1 kJ/mol
  • 79.
    98 Question How much energydoes it take to vaporizer 111 g of water? Given: Hvap water= 40.67 kJ/mol 111 g H2O 1 mol x 40.6kJ = 250kJ 18g 1mol
  • 80.
    99 • Heat ofFusion: the enthalpy change associated with melting. (Solid to liquid.) • Hfusion water= 6.01 kJ/mol NOTE: heat of fusion is always smaller than heat of vaporization. This makes sense think about the level of “order” in the molecules in these phases.
  • 81.
    Which is largerfor a given substance: ΔHvap or ΔHfus? Explain why. Copyright © Cengage Learning. All rights reserved 100 CONCEPT CHECK!
  • 82.
    Changes of State Copyright© Cengage Learning. All rights reserved 101 To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE
  • 83.
    102 Heating Curve • Aplot of the temperature versus time
  • 84.
  • 85.
    104 Example Calculate the enthalpychange associated with converting 1.00 mole of ice -25ºC to water 150ºC at 1 atm. Specific heat of ice, water, and steam are 2.09 J/g ºC and 4.184 J/g ºC, 1.84 J/g ºC . The heat of fusion of ice is 6.01 kJ/mol and heat of vaporization of water is 40.67 kJ/mol
  • 86.
    105 ICE -25 °C q1 Icewater 0°C q2 Water 0°C q3 Water vapor 100°C q4 Vapor 150°C
  • 87.
    106 1 mol ice 1 mol ice 1 mol water  1 mol water  1 mol steam T= -25ºC 0ºC 0ºC 100ºC 100ºC -> 150 qtotal = q1 + q2 + q3 + q4 + q5 1.)q = 2.09(18g)(-25-0) 2.) q = 6.02 KJ/mol (convert heat of fusion) 3.) q = 4.184(18g) (100-0) 4.) q = 40.7 KJ/mol (convert heat of vaporization) 5. ) q = 1.84(18 g) (150-100)
  • 88.
    107 Critical Stuff • CriticalTemperature: The temperature above which it is impossible to liquefy the gas under study no matter how high the applied pressure. • Critical Pressure: The pressure required to liquefy a gas as at its critical temperature NOTE: the critical temp of a gas gives an indication of the strength of the IMF of that gas. A substance with weak attractive forces would have a low critical temp.
  • 89.
    108 Which gas canbe liquefied at 25ºC Gas Critical Temp ºC Critical Pressure atm Ammonia 132 112 Ethanol 158 78 Argon -186 6 Critical Temp above 25ºC Critical Temp under 25ºC
  • 90.
    109 Vapor Pressure (vp) VaporPressure: Pressure exerted by molecules that have enough energy to escape the surface. As T ↑ VP ↑evaporation ↑ Liquids with high VP are volatile (alcohol evaporates easily) Liquids that have strong IMF have low vapor pressures. (take a lot of energy to overcome IMF so it can evaporate)
  • 91.
    Kinetic energy % o f M o l e c u l e s T2 • Athigher temperature more molecules have enough energy • Higher vapor pressure.
  • 92.
    111 • Liquids withhigh VP are volatile (alcohol evaporates easily) • Liquids that have strong IMF have low vapor pressures. • (take a lot of energy to overcome IMF so it can evaporate) substance vapor pressure at 25oC diethyl ether C4H10O 0.7 atm Bromine Br2 0.3 atm ethyl alcohol C2H5OH 0.08 atm Water H2O 0.03 atm
  • 93.
    Evaporation • Molecules atthe surface break away and become gas. • Only those with enough KE escape • Evaporation is a cooling process. • It requires heat. • Endothermic.
  • 94.
    Condensation Change from gasto liquid Achieves a dynamic equilibrium with vaporization in a closed system. What the heck is a “dynamic equilibrium?”
  • 95.
    Liquid/Vapor Equilibrium Copyright ©Cengage Learning. All rights reserved 114 To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE
  • 96.
    When first sealedthe molecules gradually escape the surface of the liquid As the molecules build up above the liquid some condense back to a liquid. Dynamic equilibrium
  • 97.
    As time goesby the rate of vaporization remains constant but the rate of condensation increases because there are more molecules to condense. Equilibrium is reached when Rate of Vaporization = Rate of Condensation Dynamic equilibrium
  • 98.
    117 VP example In aclosed container the number of partials changing from liquid  vapor will eventually equal the number changing from vapor  liquid.
  • 99.
    118 Boiling Point Note: Thenormal boiling point of water is 100oC. The term normal refers to standard pressure or 1 atm, or also 101.3 kPa. The vapor pressure of the liquid = air pressure above the liquid
  • 100.
    119 Boiling Pts. ofH2O at Various Elevations Altitude compared to Sea Level (m) Boiling Point (°C) 1609 98.3 177 100.3
  • 101.
    120 How to makesomething boil 1. Increase the VP of the liquid (heat it) so that the VP of the liquid is > that of the atmosphere. 2. Lower the atmospheric pressure (pressure above the liquid) (put a lid on it)
  • 102.
    121 Boiling Point ↑ boilingpt by ↑ in IMF Or ↓ VP At high altitudes (low air pressure) water boils at a lower temperature
  • 103.
    122 Normal Boiling Point •Temperature at which something boils when the vp =1 atm • Note the lower the external pressure the lower the boiling point.
  • 104.
    123 Freezing point/melting point •They are the same but in opposite directions. • When heated the particles vibrate more rapidly until they shake themselves free of each other. • Ionic solids have strong intermolecular forces so a high mp. • Covalent/molecular solids have weak intermolecular forces so a low mp.
  • 105.
    124 Phase Diagram • Agraphical way to summarize the conditions under which equilibrium exists between different states of matter. • Allows you to predict the phase of a substance that is stable at a given temperature and pressure
  • 106.
    125 Triple point =three phase are in equilibrium with each other at the same time Boiling Point Melting Point 1 atm Critical point
  • 107.
    126 Critical Point: Thetemp beyond which the ,molecules of a substance have to much kinetic energy to stick together to form a liquid.
  • 108.
    127 Phase diagrams ofsubstances other than water the slope of the solid liquid line slopes forward. (positive) In water the slope of the solid-liquid lines slopes downward. (negative) Water Not Water
  • 109.
    128 Homework • Pg 508 •#’s : 85, 87,89, 91,