Ch10 z5e liq solids

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  • Z5e 10.1 p. 452
  • Rf Z5e Ch10 early slide
  • Z5e Fig 10.2 p. 453 Use this to predict relative boiling points.
  • HRW Top are polar Bottom are non-polar
  • Re Z5e ch 10 early slides
  • Z5e Fig. 10.4 p. 454 Non-polar tetrahedral hydrides of Gp 4A show steady increase in b.p. w/molar mass But, for other gps, lightest member has unexpectedly high b.p. d/t H-bonding (polar X-H) AND because small size of 1st element in each group allows close approach of dipoles, further strengthening intermolecular forces.
  • Z5e 453 rf Fig. 10.3
  • Rf Z5e Ch10 early slides
  • Z53 fig. 10.5 p. 455 Describes (a) instantaneous dipoles on atoms vs. (b) that on molecules
  • As number of e- increase, greater chance of momentary dipole interactions. So, importance of London forces increases as size of atom increases.
  • Z5e Section 10.2 p. 456
  • Rf. Z5e fig. 10.6 p. 456
  • SS look at Fig. 10.7 Z5e p. 456 for actual menisci of Hg and H2O
  • Z5e 456
  • Z5e 457
  • Z5e 457-458
  • Section 10.3 Z5e p. 458
  • Z5e 10.3 An Introduction to Structures and Types of Solids
  • Z5e 461-462.
  • Aka Simple cubic (rf. Fig. 10.9 Z5e 459)
  • Z5e459 Fig. 10.9. Three cubic unit cells and their corresponding lattices.
  • Z5e 462 Fig. 10.12
  • Z5e 462 -466
  • Section 10.4 Z5e 463
  • Hrw 181
  • Z53 468 and fig. 10.21
  • Z5e 468 Two Types of Alloys.
  • Z5e Section 10.5 p. 470
  • Rf. Fig. 10.22(a) Z5e 470
  • Rf. Fig. 10.22(b) Z5e 470.
  • Z5e 471 Fig. 10.24 - an extensive Pi-bonding network. Discuss Semiconductors Z5e p. 476; add figures
  • Z5e 471. Fig. 10.24
  • Z5e 470 Fig. 10.22.
  • Z5e 476
  • Z5e 476-477 & figs. 10.29 -10.31
  • Energy level diagrams for (a) an n-type semiconductor and (b) a p-type semiconductor
  • Section 10.6 Z5e 478
  • Z5E 479
  • Section 10.7 Z5e 479 Do Z5e SE 10.4 p. 483 Students copy Table 10.7 and I review!!!
  • Section 10.8 Z5e 484
  • Z5e Fig. 10.37 485 Rate of evaporation remains constant Rate of condensation increases as # molecules in vapor phase increase, until the two rates become equal
  • Z5e 484
  • Rf. Fig 10.38 Z5e 485
  • Fig 10.38 Z5e 485 Ether is most volatile. In each case a little liquid remains floating at top.
  • Rf. Fig 10.39 Z5e 486
  • Fig 10.40 Z5e 487. (b) plot used in Lab 9 (Vapor Pressure and Enthalpy of Vaporization in H2O).
  • Fig. 10.42 Z5e 490 Bp plateau longer than mp because almost 7x energy to vaporize vs. melt. Slopes of other lines are different because different states of water have different molar heat capacities.
  • Merrill 435 SP
  • Merrill 435 SP
  • Rf. #87 Z5e 508 Ans. 1680 kJ (see solutions guide 247).
  • Z5e 491-492
  • Z5e 492.
  • Section 10.9 Z5e 493
  • Rf. HRW Tr 63
  • Ask students what state exists for each letter
  • Fig 10.47 Z5e 493 Be able to draw, given a set of data and to describe, including meaning of (-) slope.
  • Fig. 10.52 Z5e 500 Be able to explain for test!!
  • Ch10 z5e liq solids

    1. 1. Liquids and solids pp Chapter 10
    2. 2. 10.1 Intermolecular Forces <ul><li>Liquids & solids are similar to each other compared to gases </li></ul><ul><li>They are incompressible . </li></ul><ul><li>Their density doesn’t change (much) with temperature. </li></ul><ul><li>These similarities are due </li></ul><ul><ul><li>to the molecules being close together in solids and liquids </li></ul></ul><ul><ul><li>and far apart in gases </li></ul></ul><ul><li>What holds them close together? </li></ul>
    3. 3. Intermolecular forces <ul><li>Within the molecules ( intra molecular) the atoms are bonded to each other . </li></ul><ul><li>Inter molecular refers to the forces between the molecules. </li></ul><ul><li>The intermolecular forces are what hold the molecules together in the condensed states . </li></ul>
    4. 4. Intermolecular forces <ul><li>Strong (these are “int ra ”) </li></ul><ul><ul><li>covalent bonding </li></ul></ul><ul><ul><li>ionic bonding </li></ul></ul><ul><li>Weak (“int er ”) </li></ul><ul><ul><li>Dipole dipole (including H-bonding). </li></ul></ul><ul><ul><li>London dispersion forces. </li></ul></ul><ul><li>During phase changes the molecules stay intact (but separate from each other). </li></ul><ul><li>Energy used to overcome forces. </li></ul>
    5. 5. Dipole - Dipole <ul><li>Remember where the polar definition came from? </li></ul><ul><li>Molecules line up in the presence of a electric field. The opposite ends of the dipole can attract each other so the molecules stay close together. </li></ul><ul><li>1% as strong as covalent bonds </li></ul><ul><li>Weaker with greater distance. </li></ul><ul><li>Small role in gases. </li></ul>
    6. 6. + - + - + - + - + - + - + - + - + - + -
    7. 7. How to show a bond is polar <ul><li>Not a whole charge, just a partial charge </li></ul><ul><li> means a partially positive </li></ul><ul><li> means a partially negative </li></ul><ul><li>The Cl pulls harder on the electrons </li></ul><ul><li>The electrons spend more time near the Cl </li></ul>H Cl  
    8. 8. Polar Molecules <ul><li>Molecules with a slightly positive and a slightly negative end </li></ul><ul><li>Requires two things to be true </li></ul><ul><li>The molecule must contain polar bonds </li></ul><ul><li>(This can be determined from differences in electronegativity). </li></ul><ul><li>Symmetry cannot cancel out the effects of the polar bonds. </li></ul><ul><li>(Must determine geometry first.) </li></ul>
    9. 9. Polar Molecules <ul><li>98% of the time, the molecule is polar if: </li></ul><ul><li>The molecule has polar bonds and </li></ul><ul><li>There is at least one lone pair on the central atom. </li></ul><ul><li>Example of an exception to this rule of thumb: CH 2 O polar because of C=O bond, but no lone pairs </li></ul>
    10. 10. Fig 6-26, p.191 Predicting Polarity Tr 37
    11. 11. Is it polar? <ul><li>HF </li></ul><ul><li>H 2 O </li></ul><ul><li>NH 3 </li></ul><ul><li>CCl 4 </li></ul><ul><li>CO 2 </li></ul><ul><li>Yes </li></ul><ul><li>Yes </li></ul><ul><li>Yes </li></ul><ul><li>No </li></ul><ul><li>No </li></ul>
    12. 12. Hydrogen Bonding <ul><li>An especially strong dipole-dipole forces when H is attached to F, O, or N </li></ul><ul><li>Occurs with these three because: </li></ul><ul><ul><li>They have high electronegativity . </li></ul></ul><ul><ul><li>They are small enough to get close . </li></ul></ul><ul><li>The hydrogen partially shares with the lone pair in the molecule next to it. </li></ul><ul><li>Affects boiling point. </li></ul><ul><li>The strongest of the intermolecular forces. </li></ul>
    13. 13. Hydrogen Bonding H H O  +  -  + H H O  +  -  +
    14. 14. Hydrogen bonding H H O H H O H H O H H O H H O H H O H H O
    15. 15. Boiling Points 0ºC 100 -100 -200 CH 4 SiH 4 GeH 4 SnH 4 PH 3 NH 3 SbH 3 AsH 3 H 2 O H 2 S H 2 Se H 2 Te HF HI HBr HCl
    16. 16. Water  +  -  +
    17. 17. London Dispersion Forces <ul><li>Non polar molecules and noble gases also exert forces on each other. </li></ul><ul><li>Otherwise, no solids or liquids. </li></ul><ul><li>Electrons are not evenly distributed at every instant in time. </li></ul><ul><li>Have an instantaneous dipole. </li></ul><ul><li>Induces a dipole in the atom next to it. </li></ul><ul><li>Induced dipole-induced dipole interaction. LD2: 6.69 Induced dipole in an oxygen molecule. </li></ul>
    18. 18. Example H H H H H H H H  +  - H H H H  +  -  + 
    19. 19. Figure 10.5 London Dispersion Forces
    20. 20. London Dispersion Forces <ul><li>Weak, short lived. </li></ul><ul><li>Lasts longer at low temperature. </li></ul><ul><li>Eventually long enough to make liquids. </li></ul><ul><li>More electrons means more polarizable. </li></ul><ul><li>Bigger molecules - higher melting & boiling points. </li></ul><ul><li>Much, much weaker than other forces. </li></ul><ul><li>Also called Van der Waal’s forces. </li></ul><ul><li>LD 2: 6.72 London Dispersion Forces </li></ul>
    21. 21. 10.2 Liquids <ul><li>Many of the properties due to internal attraction of atoms. </li></ul><ul><ul><li>Beading </li></ul></ul><ul><ul><li>Surface tension </li></ul></ul><ul><ul><li>Capillary action </li></ul></ul><ul><li>Stronger intermolecular forces cause each of these to increase . </li></ul>
    22. 22. Surface tension <ul><li>Molecules in the middle are attracted in all directions. </li></ul><ul><li>Molecules at the top are only pulled inside . </li></ul><ul><li>Minimizes surface area. </li></ul>
    23. 23. Capillary Action <ul><li>Liquids spontaneously rise in a narrow tube. </li></ul><ul><li>Inter molecular forces are cohesive , connecting like things. </li></ul><ul><li>Adhesive forces connect to something else . (Cohesive - between same molecules). </li></ul><ul><li>Glass is polar. </li></ul><ul><li>So, glass attracts water molecules (adhesive). </li></ul>
    24. 25. Beading <ul><li>If a polar substance is placed on a non-polar surface. </li></ul><ul><ul><li>There are cohesive , </li></ul></ul><ul><ul><li>But no adhesive forces. </li></ul></ul><ul><li>And visa versa </li></ul>
    25. 26. Viscosity <ul><li>How much a liquid resists flowing. </li></ul><ul><li>Large forces means more viscous. </li></ul><ul><li>Large molecules can get tangled up . </li></ul><ul><li>Cyclo hexane has a lower viscosity than hexane because . . . </li></ul><ul><li>It is a circle - more compact. </li></ul>
    26. 27. How much of these? <ul><li>Stronger forces  bigger effect. </li></ul><ul><ul><li>Hydrogen bonding </li></ul></ul><ul><ul><li>Polar bonding </li></ul></ul><ul><ul><li>LDF </li></ul></ul>
    27. 28. Model for Liquids <ul><li>Can’t see molecules so picture them as: </li></ul><ul><li>In motion but attracted to each other </li></ul><ul><li>With regions arranged like solids but </li></ul><ul><ul><li>with higher disorder. </li></ul></ul><ul><ul><li>with fewer holes than a gas . </li></ul></ul><ul><ul><li>Highly dynamic , regions changing between types. </li></ul></ul>
    28. 29. Phases <ul><li>The phase of a substance is determined by three things : </li></ul><ul><li>The temperature . </li></ul><ul><li>The pressure . </li></ul><ul><li>The strength of intermolecular forces . </li></ul>
    29. 30. 10.3 Solids <ul><li>Two major types: </li></ul><ul><li>Amorphous - those with much disorder in their structure. </li></ul><ul><li>Crystalline - have a regular arrangement of components in their structure. </li></ul>
    30. 31. Crystals <ul><li>Lattice - a three dimensional grid that describes the locations of the pieces in a crystalline solid. </li></ul><ul><li>Unit Cell -The smallest repeating unit in of the lattice. </li></ul><ul><li>Three common types . . . </li></ul>
    31. 32. Cubic
    32. 33. Body-Centered Cubic
    33. 34. Face-Centered Cubic
    34. 35. Figure 10.9 Three Cubic Unit Cells & the Corresponding Lattices
    35. 36. Solids <ul><li>There are many amorphous solids. </li></ul><ul><li>Like glass. </li></ul><ul><li>We tend to focus on crystalline solids. </li></ul><ul><li>Three types: </li></ul><ul><li>Ionic solids have ions at the lattice points. </li></ul><ul><li>Molecular solids have molecules. </li></ul><ul><ul><li>Can be covalent or network covalent </li></ul></ul><ul><li>Sugar vs. Salt. </li></ul><ul><li>Metallic (atomic) </li></ul>
    36. 37. Figure 10.12 Examples of Three Types of Crystalline Solids Atomic, Ionic and Molecular Solids
    37. 38. The book drones on about <ul><li>Using diffraction patterns to identify crystal structures. </li></ul><ul><li>Talks about metals and the closest packing model. </li></ul><ul><li>It is interesting, but trivial. </li></ul><ul><li>We need to focus on metallic bonding . </li></ul><ul><li>Why do metal atoms stay together? </li></ul><ul><li>How does their bonding affect their properties? </li></ul>
    38. 39. 10.4 Metallic Bonding <ul><li>The way metal atoms are held together in the solid. </li></ul><ul><li>Metals hold onto their valence electrons very weakly . </li></ul><ul><li>Think of the metals as positive ions floating in a sea of negative electrons . </li></ul>
    39. 40. 10.4 Metallic bonding 1s 2s 2p 3s 3p 12+ 12+ 12+ 12+ 12+ 12+ Filled Molecular Orbitals Empty Molecular Orbitals Magnesium Atoms
    40. 41. Filled Molecular Orbitals Empty Molecular Orbitals The 1s, 2s, and 2p electrons are close to nucleus, so they are not able to move around. 1s 2s 2p 3s 3p 12+ 12+ 12+ 12+ 12+ 12+ Magnesium Atoms
    41. 42. Filled Molecular Orbitals Empty Molecular Orbitals 1s 2s 2p 3s 3p 12+ 12+ 12+ 12+ 12+ 12+ Magnesium Atoms The 3s and 3p orbitals overlap and form molecular orbitals.
    42. 43. Filled Molecular Orbitals Empty Molecular Orbitals 1s 2s 2p 3s 3p 12+ 12+ 12+ 12+ 12+ 12+ Magnesium Atoms Electrons in these energy levels can travel freely throughout the crystal.  
    43. 44. Filled Molecular Orbitals Empty Molecular Orbitals 1s 2s 2p 3s 3p 12+ 12+ 12+ 12+ 12+ 12+ Magnesium Atoms This makes metals conductors Malleable because the bonds are flexible .  
    44. 45. Sea of Electrons <ul><li>Electrons are free to move through the solid. </li></ul><ul><li>So, metals conduct electricity. </li></ul>+ + + + + + + + + + + +
    45. 46. Metals are malleable & ductile <ul><li>Malleable - can be hammered into shape (they bend). </li></ul><ul><li>Ductile - can be drawn into wires. </li></ul>
    46. 47. Malleable + + + + + + + + + + + +
    47. 48. Malleable <ul><li>Electrons allow atoms to slide by (like ball bearings). </li></ul>+ + + + + + + + + + + +
    48. 49. Shiny <ul><li>Metals contain many orbitals (d-orbitals) separated by small energy differences. </li></ul><ul><li>So, can absorb a wide range of light frequencies. </li></ul><ul><li>Electrons then get excited to different levels because of the wide range of light frequencies. </li></ul><ul><li>When de-excited they give off light in that wide range - shiny. </li></ul>
    49. 50. Metal Alloys <ul><li>Alloy : contains mixture of elements and has metallic properties. </li></ul><ul><li>Substitutional : some host metal atoms replaced by similar size metal atoms (e.g., Cu/Zn in brass). </li></ul><ul><li>Interstitial : Interstices (holes) in closest packed metal structure occupied by small atoms (e.g., C in Fe). </li></ul>
    50. 51. Fig. 10.21p. 468 Two Types of Alloys <ul><li>Substitutional </li></ul><ul><li>Interstitial </li></ul>
    51. 52. Metal Alloys <ul><li>1. Substitutional Alloy : some metal atoms replaced by others of similar size. </li></ul><ul><li>brass = Cu/Zn </li></ul>Substances that have a mixture of elements and metallic properties.
    52. 53. Metal Alloys (continued) <ul><li>2. Interstitial Alloy : Interstices (holes) in closest packed metal structure are occupied by small atoms. </li></ul><ul><li>steel = iron + carbon </li></ul><ul><li>3. Both types : Alloy steels contain a mix of interstitial (carbon) and substitutional (Cr, Mo) alloys. </li></ul>
    53. 54. 10.5 Carbon & Silicon: Network Atomic Solids <ul><li>There are three types of solid carbon. </li></ul><ul><li>Amorphous - coal uninteresting. </li></ul><ul><li>Diamond - hardest natural substance on earth, insulates both heat and electricity. </li></ul><ul><li>Graphite - slippery, conducts electricity. </li></ul><ul><li>How the atoms in these network solids are connected explains why they have different properties. </li></ul>
    54. 55. Diamond- each Carbon is sp3 hybridized, connected to four other carbons. <ul><li>Carbon (diamond) atoms are locked into tetrahedral shape. </li></ul><ul><li>Strong  bonds give the huge molecule its hardness. </li></ul>
    55. 56. Why is diamond an insulator? <ul><li>The space between orbitals make it impossible for electrons to move around. </li></ul>E Empty MOs Filled MOs
    56. 57. <ul><li>Each carbon is connected to three other carbons and sp 2 hybridized . </li></ul><ul><li>So, the molecule is flat with 120º angles in fused 6-member rings. </li></ul><ul><li>The  bonds extend above and below the plane. </li></ul>Graphite is different.
    57. 58. This  bond overlap forms a huge  bonding network. <ul><li>Electrons are free to move through out these delocalized orbitals. </li></ul><ul><li>So, the layers slide by each other. </li></ul>
    58. 59. Figure 10.24 The p Orbitals
    59. 60. Figure 10.22 p. 470 The Structures of Diamond and Graphite
    60. 61. Semiconductors. <ul><li>Like diamond, silicon has an energy gap between its filled & empty molecular orbitals. </li></ul><ul><li>But, it’s gap is smaller so e - s can cross it at room temperature (25 degrees C). </li></ul><ul><li>Still not great, so it is a semi conductor . </li></ul>
    61. 62. Semiconductors <ul><li>The space between orbitals is closer than with carbon so slightly easier for electrons to move around. </li></ul>E Empty MOs Filled MOs
    62. 63. Semiconductors <ul><li>Conductivity is enhanced by doping with group 13 or group 15 elements. </li></ul>A substance in which some electrons can cross the band gap.
    63. 64. Semiconductors <ul><li>n-type - doping with atoms with one more electron ( e.g., As) </li></ul><ul><li>The extra e - s lie close in energy to the conduction bands and can be easily excited into those levels where they conduct electricity. </li></ul><ul><li>Hint: “ n -type” as in “ negative ” because have one more e - . </li></ul>
    64. 65. Semiconductors <ul><li>p-type - doping with atoms with one less electron ( e.g., B) </li></ul><ul><li>An e - vacancy (hole) is created so as another e - enters this hole, it leaves a new hole, etc. </li></ul><ul><li>The “hole” advances the opposite direction to the movement of e - s. </li></ul><ul><li>The movement of e - s causes conduction. </li></ul><ul><li>“ p ” as in “ positive ” because one less e - . </li></ul>
    65. 66. Fig 10.29 p. 476 Silicon Crystal Doped with (a) Arsenic and (b) Boron <ul><li>Arsenic has one more valence electron than Si </li></ul><ul><li>Boron has one less valence electron than Si </li></ul><ul><li>Electrons must be in singly occupied molecular orbitals to conduct a current. </li></ul><ul><li>As an electron fills a hole it leaves a new hole. </li></ul>
    66. 67. Figure 10.30 Energy Level Diagrams for (a) an n-Type Semiconductor and (b) a p-Type Semiconductor
    67. 68. Figure 10.31 p. 477 The p-n Junction <ul><li>p-n junction involves contact of p-type & n-type semiconductor. </li></ul><ul><li>Charge carriers of p-type region are holes. </li></ul><ul><li>No current flows (reverse bias) - note the terminals have been switched. </li></ul><ul><li>Current readily flows (forward bias). </li></ul><ul><li>Note: each electron that crosses the boundary leaves a hole behind. Thus, the electrons and the holes move in opposite directions. </li></ul>
    68. 69. 10.6 Molecular solids. <ul><li>Molecules occupy the corners of the lattices . </li></ul><ul><li>Different molecules have different forces between them. </li></ul><ul><li>These forces depend on the size of the molecule. </li></ul><ul><li>They also depend on the strength and nature of dipole moments . </li></ul>
    69. 70. Those without dipoles. <ul><li>Most are gases at 25ºC. </li></ul><ul><li>The only forces are London Dispersion Forces (extremely weak). </li></ul><ul><li>These depend on size of atom. </li></ul><ul><li>So, large molecules (such as I 2 ) can be solids even without dipoles (lots of e - s to generate LDFs). </li></ul><ul><li>I 2 (solid) Br 2 (liquid) Cl 2 (gas) large medium small </li></ul>
    70. 71. Those with dipoles. <ul><li>Dipole-dipole forces are generally stronger than L.D.F. </li></ul><ul><li>Hydrogen-bonding is stronger than Dipole-dipole forces. </li></ul><ul><li>No matter how strong the intermolecular force, it is always much, much weaker than the forces in bonds (1%). </li></ul><ul><li>Stronger forces lead to higher melting and freezing points. </li></ul>
    71. 72. Water is special <ul><li>Each molecule has two polar O-H bonds. </li></ul>H O H      -
    72. 73. Water is special <ul><li>Each molecule has two polar O-H bonds. </li></ul><ul><li>Each molecule has two lone pair on its oxygen. </li></ul>H O H    
    73. 74. Water is special <ul><li>Each molecule has two polar O-H bonds. </li></ul><ul><li>Each molecule has two lone pair on its oxygen. </li></ul><ul><li>Each oxygen can interact with 4 hydrogen atoms. </li></ul>H O H    
    74. 75. Water is special <ul><li>This gives water an especially high melting and boiling point. </li></ul>H O H     H O H     H O H    
    75. 76. 10.7 Ionic Solids <ul><li>The extremes in dipole-dipole forces-atoms are actually held together by opposite charges . </li></ul><ul><li>So, huge melting and boiling points. </li></ul><ul><li>Atoms are locked in lattice, so the solid is hard and brittle. </li></ul><ul><li>Every e 1- is accounted for, so ionic solids are poor conductors - good insulators. </li></ul>
    76. 77. Using Table 10.7 p. 458, what type of solids do the following form? <ul><li>Diamond </li></ul><ul><li>PH 3 </li></ul><ul><li>Mg </li></ul><ul><li>NH 4 NO 3 </li></ul><ul><li>Ar </li></ul><ul><li>atomic, covalent network </li></ul><ul><li>Molecular </li></ul><ul><li>Atomic, metallic </li></ul><ul><li>Ionic </li></ul><ul><li>Atomic, group 8A or 18 </li></ul>
    77. 78. 10.8 Vapor Pressure & Changes of State <ul><li>Vaporization - change from liquid to gas at boiling point. </li></ul><ul><li>Evaporation - change from liquid to gas below boiling point </li></ul><ul><li>Heat (or Enthalpy) of Vaporization (  H vap ) - the energy required to vaporize 1 mol at 1 atm. </li></ul>
    78. 79. Vaporization <ul><li>Vaporization is an endothermic process - it requires heat. </li></ul><ul><li>Energy is required to overcome intermolecular forces. </li></ul><ul><li>Responsible for cool earth. </li></ul><ul><li>Why we sweat. (Never let them see you.) </li></ul>
    79. 80. Condensation <ul><li>Change from gas to liquid . </li></ul><ul><li>Achieves a dynamic equilibrium with vaporization in a closed system. </li></ul><ul><li>What is a closed system? </li></ul><ul><li>A closed system means matter can’t go in or out. </li></ul><ul><li>Put a cork in it. </li></ul><ul><li>What the heck is a “ dynamic equilibrium ?” </li></ul>
    80. 81. Dynamic equilibrium <ul><li>When first sealed the molecules gradually escape the surface of the liquid. </li></ul>
    81. 82. Dynamic equilibrium <ul><li>When first sealed the molecules gradually escape the surface of the liquid. </li></ul><ul><li>As the molecules build up above the liquid some condense back to a liquid. </li></ul>
    82. 83. Dynamic equilibrium <ul><li>When first sealed the molecules gradually escape the surface of the liquid. </li></ul><ul><li>As the molecules build up above the liquid some condense back to a liquid. </li></ul><ul><li>As time goes by the rate of vaporization remains constant but the rate of condensation increases because there are more molecules to condense. </li></ul>
    83. 84. Dynamic equilibrium <ul><li>When first sealed the molecules gradually escape the surface of the liquid </li></ul><ul><li>As the molecules build up above the liquid some condense back to a liquid. </li></ul><ul><li>As time goes by the rate of vaporization remains constant but the rate of condensation increases because there are more molecules to condense. </li></ul><ul><li>Equilibrium is reached when: </li></ul>
    84. 85. <ul><li>Rate of Vaporization = Rate of Condensation </li></ul><ul><li>Molecules are constantly changing phase “ Dynamic ” </li></ul><ul><li>The total amount of liquid and vapor remains constant “ Equilibrium ” </li></ul>Dynamic equilibrium
    85. 86. Figure 10.37 The Rates of Condensation and Evaporation <ul><li>Rate of evaporation stays constant </li></ul><ul><li>Rate of condensation increases as as # vapor molecules increase, until the two rates are equal. </li></ul>
    86. 87. Vapor pressure <ul><li>The pressure above the liquid at equilibrium. </li></ul><ul><li>Liquids with high vapor pressures evaporate easily. They are called volatile . </li></ul><ul><li>Decreases with increasing intermolecular forces because . . . </li></ul><ul><ul><li>Bigger molecules (bigger LDF) </li></ul></ul><ul><ul><li>More polar molecules (dipole-dipole) </li></ul></ul>
    87. 88. Vapor pressure <ul><li>Increases with increasing temperature. </li></ul><ul><li>Easily measured in a barometer. </li></ul>
    88. 89. <ul><li>A barometer will hold a column of mercury 760 mm high at one atm </li></ul>Dish of Hg Vacuum P atm = 760 torr
    89. 90. <ul><li>A barometer will hold a column of mercury 760 mm high at one atm. </li></ul><ul><li>If we inject a volatile liquid in the barometer it will rise to the top of the mercury. </li></ul>Dish of Hg Vacuum P atm = 760 torr
    90. 91. <ul><li>A barometer will hold a column of mercury 760 mm high at one atm. </li></ul><ul><li>If we inject a volatile liquid in the barometer it will rise to the top of the mercury. </li></ul><ul><li>There it will vaporize and push the column of mercury down. </li></ul>Dish of Hg P atm = 760 torr Water
    91. 92. <ul><li>The mercury is pushed down by the vapor pressure. </li></ul><ul><li>P atm = P Hg + P vap </li></ul><ul><li>P atm - P Hg = P vap </li></ul><ul><li>760 - 736 = 24 torr </li></ul>Dish of Hg 736 mm Hg Water Vapor
    92. 93. Figure 10.38 Vapor Pressure:Which is most volatile?
    93. 94. Temperature Effect Kinetic energy # of molecules T 1 Energy needed to overcome intermolecular forces
    94. 95. <ul><li>At higher temperature more molecules have enough energy - higher vapor pressure. </li></ul>Kinetic energy # of molecules T 1 Energy needed to overcome intermolecular forces T 1 T 2 Energy needed to overcome intermolecular forces
    95. 96. Mathematical relationship <ul><li>ln is the natural logarithm </li></ul><ul><ul><li>ln = Log base e </li></ul></ul><ul><ul><li>e = Euler’s number an irrational number like  </li></ul></ul><ul><li> H vap is the heat of vaporization in J/mol </li></ul>P vapT1  H vap 1 1 ln (-----) = ---- (--- - ---) P vapT2 R T 2 T 1
    96. 97. Mathematical relationship <ul><li>R = 8.3145 J /K mol (use correct units). </li></ul><ul><li>Graph = straight line (y = mx + b) </li></ul><ul><li>Called Clausius-Clapeyron equation. </li></ul><ul><li>Be able to plot this in a free response question. </li></ul>P vapT1  H vap 1 1 ln (-----) = ---- (--- - ---) P vapT2 R T 2 T 1
    97. 98. Figure 10.40 The Vapor Pressure of Water
    98. 99. Changes of state <ul><li>The graph of temperature versus heat applied is called a heating curve . </li></ul><ul><li>The temperature a solid turns to a liquid is the melting point . </li></ul><ul><li>The energy required to change solid to liquid is called the Heat (or Enthalpy) of Fusion  H fus </li></ul><ul><li>The energy to change liquid to gas is the Heat of Vaporization  H vap </li></ul>
    99. 100. Changes of state <ul><li>The graph of temperature versus heat applied is called a heating curve . </li></ul><ul><li>Use q = m  T C p when the temperature is changing within a phase. </li></ul><ul><li>Use q = m  H fus when melting (no change in temperature) </li></ul><ul><li>Use q = m  H vap when vaporizing . </li></ul><ul><li>Add them up to get the overall Heat of Vaporization  H vap </li></ul>
    100. 101. Heating Curve for Water Ice Water and Ice Water Water and Steam Steam
    101. 102. Heating Curve for Water Heat of Fusion Heat of Vaporization Slope is Heat Capacity
    102. 103. Figure 10.42 Heating Curve for Water Bp plateau longer than Mp because takes almost 7x the energy to vaporize as melt. Slopes of other lines different because different states of water have different molar heat capacities
    103. 104. Example pp <ul><li>How much heat does it take to convert 10.0 g of ice at -10  C to steam at 150  C? </li></ul><ul><li>Raise ice temperature to its melting point. </li></ul><ul><li>q = heat = m x ∆T x C p = (10.0g)(10ºC)(2.06 J/g  C) = 206 J. </li></ul><ul><li>Melt the ice. </li></ul><ul><li>q = m x ∆H f = (10.0 g)(334 J/g) = 3340 J. </li></ul><ul><li>Raise H 2 O temperature to its boiling point. </li></ul><ul><li>q = heat = m x ∆T x C p = (10.0g)(100ºC)(4.18 J/g  C) = 4180 J. </li></ul>
    104. 105. Example Continued pp <ul><li>How much heat does it take to convert 10.0 g of ice at -10  C to steam at 150  C? </li></ul><ul><li>Vaporize the water. </li></ul><ul><li>q = m x ∆H v = (10.0 g)(2260 J/g) = 22 600 J. </li></ul><ul><li>Raise steam temperature to 150  C. </li></ul><ul><li>q = heat = m x ∆T x C p = (10.0g)(50ºC)(2.02 J/g  C) = 1010 J. </li></ul><ul><li>Total heat to convert 10.0 g of ice at -10  C to steam at 150  C is the sum of all. </li></ul><ul><li>206 + 3340 + 4180 + 22 600 + 1010 = 31 300 J </li></ul><ul><ul><ul><ul><li>= 31.3 kJ </li></ul></ul></ul></ul>
    105. 106. Another Phase Change Problem <ul><li>How much energy does it take to convert 0.500 kg ice at -20.0°C to steam at 250.°C? C p ice = 2.10 J/g°C; steam 2.00 J/g°C; liquid = 4.20 J/g°C  H fus = 6.02 kJ/mol  H vap = 40.7 kJ/mol The answer is . . . </li></ul><ul><li>1680 kJ (z5e508 #87; Sol 247) </li></ul>
    106. 107. Melting Point <ul><li>Melting point is determined by the vapor pressure of the solid and the liquid . </li></ul><ul><li>At the melting point the vapor pressure of the solid = vapor pressure of the liquid </li></ul><ul><li>Normal melting point is when the solid & liquid vapor pressures are the same and total pressure = 1 atmosphere </li></ul>
    107. 108. Solid Water Liquid Water Water Vapor Vapor
    108. 109. <ul><li>If the vapor pressure of the solid is higher than that of the liquid the solid will release molecules to achieve equilibrium. </li></ul>Solid Water Liquid Water Water Vapor Vapor
    109. 110. <ul><li>While the molecules of water vapor condense to a liquid . </li></ul>Solid Water Liquid Water Water Vapor Vapor
    110. 111. <ul><li>This can only happen if the temperature is above the freezing point since the solid is turning to liquid. </li></ul>Solid Water Liquid Water Water Vapor Vapor
    111. 112. <ul><li>If the vapor pressure of the liquid is higher than that of the solid, the liquid will release molecules to achieve equilibrium. </li></ul>Solid Water Liquid Water Water Vapor Vapor
    112. 113. <ul><li>While the molecules condense to a solid . </li></ul>Solid Water Liquid Water Water Vapor Vapor
    113. 114. <ul><li>The temperature must be below the freezing point since the liquid is turning to a solid. </li></ul>Solid Water Liquid Water Water Vapor Vapor
    114. 115. <ul><li>If the vapor pressure of the solid and liquid are equal , the solid and liquid are vaporizing and condensing at the same rate: The melting point . </li></ul>Solid Water Liquid Water Water Vapor Vapor
    115. 116. Boiling Point <ul><li>Reached when the vapor pressure equals the external pressure. </li></ul><ul><li>Normal boiling point is the boiling point at 1 atm pressure. </li></ul><ul><li>Superheating - Heating above the boiling point. </li></ul><ul><li>Supercooling - Cooling below the freezing point. </li></ul>
    116. 117. 10.9 Phase Diagrams. <ul><li>A plot of temperature versus pressure for a closed system, with lines to indicate where there is a phase change. </li></ul>
    117. 118. Phase Diagram for Water What happens to BP of water as pressure increases? To FP? BP increases, FP decreases What is the definition of triple point? The T and P at which all 3 states are in equilibrium
    118. 119. Temperature 1 Atm D D D Pressure D Solid Liquid Gas A A B B C C
    119. 120. Temperature Pressure Solid Liquid Gas Triple Point Critical Point
    120. 121. <ul><li>This is the phase diagram for water. </li></ul><ul><li>The density of liquid water is higher than solid water (negative slope). </li></ul><ul><li>Normal boiling point is the boiling point at 1 atm pressure. </li></ul>Temperature Pressure Solid Solid Liquid Gas
    121. 122. Phase Diagram for Carbon Note: There are 4 phases: Diamond, graphite, liquid, vapor. There are two “triple points”: Higher one is between Diamond, graphite & liquid Lower one is between graphite, liquid & vapor There can be different phases between solids if they have a different crystalline structure.
    122. 123. <ul><li>This is the phase diagram for CO 2 </li></ul><ul><li>The solid is more dense than the liquid (positive slope) </li></ul><ul><li>The solid sublimes at 1 atm. </li></ul>Solid Liquid Gas 1 Atm Temperature Pressure
    123. 124. Z5e Fig 10.52 Phase Diagram for CO 2 Solid/Liquid line has positive slope - Why? Density of solid carbon dioxide is greater than that of liquid.

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