1. States of Matter

2. Kinetic Molecular Theory Particles are always in motion. Temperature is a measure of average kinetic energy of particles. Intermolecular forces hold particles together. Stronger forces require more energy (higher temp.) to overcome.

3. Intermolecular Forces Always weaker than chemical bond Affect structure and state of matter

4. Dipole-Dipole Forces Positive and negative ends of polar molecules attract each other. About 1% as strong as covalent or ionic bonds Weaken as distance between molecules increases

5. Hydrogen Bonding Especially strong dipole-dipole force Occurs when H bonds to a strongly electronegative atom—O, N, or F Very strong because 1) molecule is very polar & 2) small size of H

6. H Bonding Example—water More pronounced in molecules formed from small atoms (dipoles can come closer) High boiling point

7. London Dispersion Forces Forces that exist in all atoms and molecules but that are significant only among Noble gases and nonpolar molecules Result from temporary dipoles formed when electrons distribute themselves unevenly—can induce a dipole in a neighboring atom VERY WEAK

8. London Dispersion Forces Stronger in larger atoms or molecules due to the greater chance of the formation of instantaneous dipoles.

9. Physical Properties Melting and boiling points are higher when IM attractions are stronger More energy required to separate molecules

10. Which would have the higher boiling point & Why? Cl 2 or F 2 ? H 2 O or H 2 S? SiBr 2 or SBr 2 ? CH 4 or C 10 H 22 ? O 2 or NO?

11. States of Matter Gases—weak IM forces (like London Dispersion Forces) Liquids—intermediate IM forces Solids—strong IM forces

12. Liquids & IM Forces Surface tension—result of IM forces that resist an increase in surface area Capillary action—result of cohesive forces within liquid and adhesive forces between liquid and tube

13. Liquids (cont’d) Viscosity—the ability of a liquid to resist flow (resist change in shape) All effects are higher with more polar molecules.

14. Solids Amorphous—without definite structure Crystalline—definite structures

15. Solids (Crystals) Ionic solids—made of charged particles; ions at lattice points Molecular solids—made of neutral particles; molecules at lattice points Atomic solids—made of neutral particles; atoms at lattice points; 3 types

16. An atomic solid, an ionic solid & a molecular solid

17. Ionic Solids Ions at lattice points Closest packed spheres Arranged to minimize repulsions and maximize attractions Conducts only when melted

18. Molecular Solids Lattice positions occupied by molecules Internal covalent bonds are strong, but intermolecular forces are weak IM force: dipole/dipole if polar covalent bond; London dispersion forces (larger in larger molecules)

19. Atomic Solids 1. Network—directional covalent bonds; forms giant molecules (diamond, graphite,and silicon); highest melting points 2. Metallic—delocalized covalent bonds; atoms have closest packing structure; high melting points

20. Atomic Solids 3. Group 8A—Noble gases— London dispersion forces only; low melting points.

21. Network Atomic Solids Strong, directional bonds Form giant “molecules” Typically brittle & poor conductors Examples—carbon and silicon

22. Carbon Network Follows a molecular orbital (not atomic orbital) model

23. Diamond Tetrahedral--sp 3 hybridized bonds stabilize structure Large gaps exist between filled and unfilled molecular orbitals—hard for electrons to move—no conductivity

24. Graphite Fused carbon rings form sheets Trigonal planar—sp 2 hybridized (1 p orbital remains unhybridized) Delocalized electrons in orbital causes graphite to be conductive

25. Figure 10.22: The structures of diamond and graphite. In each case only a small part of the entire structure is shown.

26. Closest Packed Solids aba pattern—alternating layers— atoms in 3 rd layer lie directly above atoms in 1 st layer—hexagonal unit cell—body centered abca pattern—atoms in 1 st and 4 th layers are in line; 2 nd & 5 th layer; 3 rd & 6 th layer—face-centered cubic cell

28. aba Packing

29. abca Packing

30. Density of Closest Packed Solids To calculate density, you need to know: MASS VOLUME

31. Mass Figure out how many atoms in one unit cell Multiply by molar mass Divide by Avogadro’s number You now know a mass in grams

33. Face-Centered Cubic Unit Cell If these are atoms of calcium, what is the mass of the cell?

35. Volume Determine the length of one side of the cube by using the atomic radius (varies depending on type of unit cell) Cube the side length.

36. Simple Cubic--aaa If the atomic radius of this atom is 122 pm, what is the volume?

38. Body-Centered Cubic--aba If the atomic radius of the atom is 246 pm, what is the volume of the cell?

40. Face-Centered Cubic--abca If the atomic radius is 291 pm, what is the volume of the cell?

42. Sample Problem Silver crystallizes in a face-centered cubic closest packed structure. The radius of the silver atom is 144 pm. Calculate the density of silver.

43. Bonding in Metals Strong, non-directional bonds Atoms are hard to separate but easy to move. “Electron sea” model Mobile electrons carry heat or electricity easily

44. Band Model or Molecular Orbital Model Electrons travel around metal crystal in a molecular (instead of atomic) orbitals Result is a continuum of levels that eventually merge to form a band.

45. Figure 10.19: The molecular orbital energy levels produced when various numbers of atomic orbitals interact.

46. Band or MO Model Empty orbitals close in energy exist. Electrons are very mobile into and out of these similar-energy orbitals--CONDUCTIVITY.

47. Semiconductors Some electrons can cross the “energy gap” between molecular orbitals—somewhat conductive Higher temperatures result in more electrons’ being able to reach conductive bands

48. Doping Adding other elements with one more or one less electron than a semiconductor can increase conductivity

49. n-type semiconductor An element with one more valence electron is added More valence electrons are available to move into conduction bands What could be used to dope Si?

50. p-type semiconductor An element with one less valence electron is added The absence of a valence electron creates a hole through which electrons can travel

51. Alloys Introducing other elements into metallic structure is easy. Substitutional—some “host” atoms are replaced by atoms of similar size (Brass = copper + zinc) Interstitial—small atoms occupy spaces between metal atoms (such as carbon used to harden steel)

52. Phase Changes

53. Condensation/Evaporation Boiling point Freezing/Melting Melting point Deposition/Sublimation

54. Energy of Phase Changes Heating/Cooling Lab

55. Energy of Phase Changes Heating Curve Temperature changes when only one phase is present No temperature change during phase change

56. For Water Heat of fusion—334 J/g Heat of vaporization—2260 J/g Specific heat 1.84 J/g o C for gas 4.18 J/g o C for liquid 2.1 J/g o C for solid

57. What amount of energy is needed to completely melt 45 g of ice at its melting point?

58. What amount of energy is needed to completely vaporize 45 g of water at its boiling point?

59. Calculate the amount of energy required to raise the temperature of 125 g of ice from -12 o C to 75 o C.

60. What is the final temperature/state of 95 g of ice at -15 o C if 65 kJ of energy are added?

61. The heat of vaporization for carbon dioxide is 571 J/g. How much energy would be required to change 6.92 g of dry ice to carbon dioxide gas?

62. Vapor Pressure In a closed system, # of liquid molecules decreases as they enter the gas phase. Eventually, an equilibrium is reached—constant number of molecules in both phases.

63. Equilibrium

64. Vapor Pressure At equilibrium evaporation rate is exactly the same as condensation rate. Vapor pressure = the pressure of the vapor present at equilibrium

65. Measuring Vapor Pressure Difference between atmospheric pressure with a vacuum above and with vapor above.

66. Effect of IM Forces Low IM forces lead to high vapor pressure; less attraction allows molecules to vaporize more easily High IM forces lead to low vapor pressure; molecules are tightly held in liquid

67. Enthalpy (  H) of Vaporization Change in energy when a liquid vaporizes Vapor pressure varies with temperature. (Why??) Note: Function is not linear

68. Linear Function Graphing ln of vapor pressure and 1/T gives a linear graph. Can use slope- intercept form of equation to solve for points.

69. Line Equation ln(P vap ) =-  vap (1) + C R (T) Format is y = mx + b R is the universal gas constant. C is a constant (unique for each substance)

70. Clausius-Clapeyron Equation Since C is temperature independent, by measuring vapor pressure at several different temperatures, heat of vaporization can be calculated.

73. Problem 2 Ethanol’s vapor pressure at 30. o C is 100. torr. At 62 o C the vapor pressure has increased to 400. torr. What is  H for ethanol?

75. In Your Scientist’s Notebook What is ethanol’s vapor pressure at 15 o C?

76. Heating Curve Changes of state are evident on a heating curve:

77. Heating Curve As heat is added temperature typically increases. At the plateaus, all energy is going into the phase change, so temperature is constant.

78. Enthalpy (  H) of Fusion Enthalpy change that occurs at the melting point when a solid melts or freezes

79. Problem The enthalpy of fusion for water is 6.02 kJ per mole. How much energy would be required to change 43.2 g of ice to liquid water?

80. What is the enthalpy change when 52.8 g of CCl 4 freezes? Enthalpy of fusion is 2.51 kJ/mol.

81. Boiling Point The temperature at which the vapor pressure of a liquid is equal to the atmospheric pressure Varies with pressure (Think about a pressure cooker.)

82. Phase Change Enthalpy The amount of heat gained or lost—  H Endothermic: boiling (evaporating), melting, sublimating Exothermic: condensing, freezing, depositing

83. Phase Change Entropy The amount of disorder in a system Increasing entropy: boiling, melting, sublimating Decreasing entropy: condensing, freezing, depositing

84. Phase Diagram Lines show equilibrium boundaries for phase changes Equilibrium—the system shows no net change Dynamic equilibrium—two processes of change occur at exactly the same rate

85. Typical Phase Diagram

86. Points of Interest Critical temperature—the temp. above which a gas cannot be liquefied at any pressure Critical pressure—the pressure required to liquefy the gas at critical temperature

87. Points of Interest Triple point—the point at which solid, liquid & gas have equal vapor pressure and exist in equilibrium “Normal” boiling or melting point—read at 1 atm pressure

89. What is the enthalpy / entropy change when… Evaporation occurs? Sublimation occurs? Freezing occurs? Melting occurs? Deposition occurs?

90. Calculations with Enthalpy To calculate total energy change, consider each stage in the process. Ice at -10 o C to steam at 110 o C: Heating ice from -10 o to 0 o C Melting Heating water from 0 o to 100 o C Boiling Heating steam from 100 o to 110 o C

91. Finding the Numbers Melting or Freezing: heat of fusion (  H fus ) Vaporizing or Condensing: heat of vaporization (  H vap ) Temperature changes: specific heats (q = mc  T) Sign depends on whether heat is gained or lost in the process.

92. Example What is the total energy required to change 45.0 g of ice at -5 o C to water at 50 o C? Specific heat of ice: 2.1 J/g o C Specific heat of water: 4.21 J/g o C  H fus : 6.02 kJ/mol  H vap : 40.7 kJ/mol