1. I. Properties, Phase Changes & Diagrams Ch. 11 - Liquids & Solids & IMF’s

2. A. Inter… and Intra…  Intermolecular Forces  Intramolecular Forces

3. A. Inter… and Intra…  Attractive forces between molecules.  Chemical bonds within molecules.

4. Types of IMF(InterMF)

5. B. Liquids vs. Solids LIQUIDS Stronger than in gases Yes high Somewhat slower than in gases SOLIDS Very strong No (?alloys) Very high No extremely slow IMF Strength Fluid Density Compressible Diffusion

6. Summarizing Intermolecular Forces

7. H–Bonding in Water and Ethanol

8. H–Bonds  Substances that can hydrogen bond will have higher boiling points and melting points than similar substances that cannot.  But hydrogen bonds are not nearly as strong as chemical bonds. 2–5% the strength of covalent bonds

9. Effect of H–Bonding on Boiling Point

10. HF, H 2 O, and NH 3 have hydrogen bonds. Therefore, they have higher boiling points than would be expected from the general trends. For nonpolar molecules, such as the hydrides of group 4, the intermolecular attractions are due to dispersion forces. Therefore, they increase down the column, causing the boiling point to increase. Polar molecules, such as the hydrides of groups 5–7, have both dispersion forces and dipole–dipole attractions. Therefore, they have higher boiling points than the corresponding group 4 molecules. Boiling Points of Group 4A and 6A Compounds

11. Ion–Dipole Attraction  In a mixture, ions from an ionic compound are attracted to the dipole of polar molecules.  The strength of the ion– dipole attraction is one of the main factors that determines the solubility of ionic compounds in water.

12. Summary  Dispersion forces are the weakest of the intermolecular attractions.  Dispersion forces are present in all molecules and atoms.  The magnitude of the dispersion forces increases with molar mass.  Polar molecules also have dipole– dipole attractive forces.

13. Summary (cont.)  Hydrogen bonds are the strongest of the intermolecular attractive forces a pure substance can have.  Hydrogen bonds will be present when a molecule has H directly bonded to either O, N, or F atoms. The only example of H bonded to F is HF.  Ion–dipole attractions are present in mixtures of ionic compounds with polar molecules.  Ion–dipole attractions are the strongest intermolecular attraction.  Ion–dipole attractions are especially important in aqueous solutions of ionic compounds.

14. B. Liquid Properties  Surface Tension (linked video) Surface Tension (linked video) attractive force between particles in a liquid that minimizes surface areaattractive force between particles in a liquid that minimizes surface area

15. B. Liquid Properties  Capillary Action attractive force between the surface of a liquid and the surface of a solid watermercury

16. B. Liquid Properties  Viscosity – resistance to motion between molecules of a liquid Molasses and water – high viscosity because of strong IMF (hydrogen bonds) Viscosity increases as temperature decreases

17. C. Properties of Water  Strong IMF (hydrogen bonds) Strong IMF (hydrogen bonds)  Most abundant substance on earth  High boiling point  Density of solid < density of liquid Density of solid < density of liquid  High surface tension  High heat of vaporization

18. D. Phase Changes

19.  Evaporation molecules gain enough energy to overcome IMF (without boiling)  Volatility (volatile) measure of evaporation rate ie. Gas, paint thinner, alcohol are very volatile (can smell the vapors)

20. D. Phase Changes Kinetic Energy # of Particles p. 480 Boltzmann Distribution tempvolatilityIMFvolatility

21. D. Phase Changes  Equilibrium trapped molecules reach a balance between evaporation & condensation

22. D. Phase Changes  Vapor Pressure pressure of vapor above a liquid at equilibrium IMFv.p.tempv.p. temp v.p.

23. D. Phase Changes  Boiling Point- temp at which vapor pressure of liquid equals external atmospheric pressure VAPOR PRESSURE ATMOSPHERIC PRESSURE IMFb.p.P atm b.p.

24.  Which has a higher m.p.? polar or nonpolar? D. Phase Changes  Melting Point equal to freezing point polar IMFm.p.

25. E. Phase Changes  Sublimation solid  gas ATMOSPHERIC PRESSURE VAPOR PRESSURE  EX: dry ice, mothballs, solid air fresheners  Opposite is deposition: H 2 0 vapor snow

26. E. Phase Diagrams  Show the phases of a substance at different temps and pressures.

27. D. Types of Solids  Crystalline - repeating geometric pattern called a unit cell (ie. salt, sugar, snow, metals etc.) Amorphous - no geometric pattern (for example: glass) covalent network metallic ionic covalent molecular increasing melting point

28. D. Types of Solids Ionic (NaCl) Metallic

29. D. Types of Solids Covalent Molecular (H 2 O) Covalent Network (SiO 2 - quartz) Amorphous (SiO 2 - glass)

31. © 2014 Pearson Education, Inc. Crystal Lattice When allowed to cool slowly, the particles in a liquid will arrange themselves to give the maximum attractive forces. –Therefore, minimize the energy. The result will generally be a crystalline solid. The arrangement of the particles in a crystalline solid is called the crystal lattice. The smallest unit that shows the pattern of arrangement for all the particles is called the unit cell.

32. © 2014 Pearson Education, Inc. *Unit Cells Unit cells are three-dimensional. –Usually containing two or three layers of particles Unit cells are repeated over and over to give the macroscopic crystal structure of the solid. Starting anywhere within the crystal results in the same unit cell. Each particle in the unit cell is called a lattice point. Lattice planes are planes connecting equivalent points in unit cells throughout the lattice.

33. © 2014 Pearson Education, Inc. Orthorhombic a ≠ b ≠ c all 90° Seven Unit Cells Hexagonal a = c < b 2 faces 90° 1 face 120° Cubic a = b = c all 90° Tetragonal a = c < b all 90° Monoclinic a ≠ b ≠ c 2 faces 90° Rhombohedral a = b = c no 90° Triclinic a ≠ b ≠ c no 90°

34. © 2014 Pearson Education, Inc. *Unit Cells The number of other particles each particle is in contact with is called its coordination number. –For ions, it is the number of oppositely charged ions an ion is in contact with Higher coordination number means more interaction; therefore, stronger attractive forces hold the crystal together. The packing efficiency is the percentage of volume in the unit cell occupied by particles. –The higher the coordination number, the more efficiently the particles are packing together.

35. © 2014 Pearson Education, Inc. *Cubic Unit Cells All 90° angles between corners of the unit cell The length of all the edges – equal If the unit cell is made of spherical particles –⅛ of each corner particle is within the cube. –½ of each particle on a face is within the cube. –¼ of each particle on an edge is within the cube.

36. © 2014 Pearson Education, Inc. *The Cubic Crystalline Lattices

37. © 2014 Pearson Education, Inc. *Cubic Unit Cells – Simple Cubic Eight particles, one at each corner of a cube ⅛ of each particle lies in the unit cell. –Each particle part of eight cells –Total = one particle in each unit cell 8 corners × ⅛ Edge of unit cell = twice the radius Coordination number of 6

38. © 2014 Pearson Education, Inc. *Simple Cubic

39. © 2014 Pearson Education, Inc. *Cubic Unit Cells – Body-Centered Cubic Nine particles, one at each corner of a cube + one in center ⅛ of each corner particle lies in the unit cell –Two particles in each unit cell 8 corners × 1/8 + 1 center Edge of unit cell = ( ) times the radius of the particle Coordination number of 8

40. © 2014 Pearson Education, Inc. *Body-Centered Cubic

41. © 2014 Pearson Education, Inc. *Cubic Unit Cells - Face-Centered Cubic 14 particles, one at each corner of a cube + one in center of each face ⅛ of each corner particle + ½ of face particle lies in the unit cell –4 particles in each unit cell 8 corners × ⅛ + 6 faces × ½ Edge of unit cell = 2 2 times the radius of the particle Coordination number of 12

42. © 2014 Pearson Education, Inc. *Face-Centered Cubic

43. © 2014 Pearson Education, Inc. *Closest-Packed Structures First Layer With spheres, it is more efficient to offset each row in the gaps of the previous row than to line up rows and columns.

44. © 2014 Pearson Education, Inc. *Closest-Packed Structures Second Layer The second layer atoms can sit directly over the atoms in the first layer—called an AA pattern. Or the second layer can sit over the holes in the first layer—called an AB pattern.

45. © 2014 Pearson Education, Inc. *Hexagonal Closest-Packed Structures

46. © 2014 Pearson Education, Inc. *Cubic Closest-Packed Structures

47. © 2014 Pearson Education, Inc. *Classifying Crystalline Solids Crystalline solids are classified by the kinds of particles found. Some of the categories are subclassified by the kinds of attractive forces holding the particles together.

48. © 2014 Pearson Education, Inc. *Classifying Crystalline Solids Molecular solids are solids whose composite particles are molecules. Ionic solids are solids whose composite particles are ions. Atomic solids are solids whose composite particles are atoms. –Nonbonding atomic solids are held together by dispersion forces. –Metallic atomic solids are held together by metallic bonds. –Network covalent atomic solids are held together by covalent bonds.

49. © 2014 Pearson Education, Inc. *Types of Crystalline Solids

50. © 2014 Pearson Education, Inc. *Molecular Solids The lattice sites are occupied by molecules. –CO 2, H 2 O, C 12 H 22 O 11 The molecules are held together by intermolecular attractive forces. –Dispersion forces, dipole–dipole attractions, and H-bonds Because the attractive forces are weak, they tend to have low melting points. –Generally < 300 °C

51. © 2014 Pearson Education, Inc. *Ionic Solids Lattice sites are occupied by ions. They are held together by attractions between oppositely charged ions. –Nondirectional –Therefore, every cation attracts all anions around it, and vice versa. The coordination number represents the number of close cation–anion interactions in the crystal. The higher the coordination number, the more stable the solid. –Lowers the potential energy of the solid The coordination number depends on the relative sizes of the cations and anions that maintain charge balance. –Generally, anions are larger than cations. –the number of anions that can surround the cation is limited by the size of the cation. –The closer in size the ions are, the higher the coordination number.

52. © 2014 Pearson Education, Inc. *Cesium Chloride Structures Coordination number = 8 ⅛ of each Cl ─ (184 pm) inside the unit cell Whole Cs + (167 pm) inside the unit cell –Cubic hole = hole in simple cubic arrangement of Cl ─ ions Cs:Cl = 1: (8 × ⅛); therefore the formula is CsCl.

53. © 2014 Pearson Education, Inc. *Rock Salt Structures Coordination number = 6 Cl ─ ions (181 pm) in a face- centered cubic arrangement. –⅛ of each corner Cl ─ inside the unit cell –½ of each face Cl ─ inside the unit cell Na + (97 pm) in holes between Cl ─ –Octahedral holes –1 in center of unit cell –1 whole particle in every octahedral hole –¼ of each edge Na + inside the unit cell Na:Cl = (¼ × 12) + 1: (⅛ × 8) + (½ × 6) = 4:4 = 1:1 Therefore, the formula is NaCl.

54. © 2014 Pearson Education, Inc. *Zinc Blende Structures Coordination number = 4 S 2─ ions (184 pm) in a face– centered cubic arrangement –⅛ of each corner S 2─ inside the unit cell –½ of each face S 2─ inside the unit cell Each Zn 2+ (74 pm) in holes between S 2─ –Tetrahedral holes –1 whole particle in ½ the holes Zn:S = (4 × 1) : (⅛ × 8) + (½ × 6) = 4:4 = 1:1 Therefore, the formula is ZnS.

55. © 2014 Pearson Education, Inc. *Fluorite Structures Coordination number = 4 Ca 2+ ions (99 pm) in a face-centered cubic arrangement –⅛ of each corner Ca 2+ inside the unit cell –½ of each face Ca 2+ inside the unit cell Each F ─ (133 pm) in holes between Ca 2+ –Tetrahedral holes –1 whole particle in all the holes Ca:F = (⅛ × 8) + (½ × 6): (8 × 1) = 4:8 = 1:2 Therefore, the formula is CaF 2. –Fluorite structure common for 1:2 ratio Usually get the antifluorite structure when the cation:anion ratio is 2:1 –The anions occupy the lattice sites and the cations occupy the tetrahedral holes.

56. © 2014 Pearson Education, Inc. *Lattice Holes

57. © 2014 Pearson Education, Inc. *Nonbonding Atomic Solids Noble gases in solid form Solid held together by weak dispersion forces –Very low melting Tend to arrange atoms in closest-packed structure –Either hexagonal cP or cubic cP –Maximizes attractive forces and minimizes energy

58. © 2014 Pearson Education, Inc. Metallic Atomic Solids Solid held together by metallic bonds –Strength varies with sizes and charges of cations Coulombic attractions Melting point varies Mostly closest-packed arrangements of the lattice points –Cations

59. © 2014 Pearson Education, Inc. Closest-Packed Crystal Structures in Metals

60. © 2014 Pearson Education, Inc. Metallic Bonding Metal atoms release their valence electrons. Metal cation “islands” fixed in a “sea” of mobile electrons

61. © 2014 Pearson Education, Inc. Network Covalent Solids Atoms attach to their nearest neighbors by covalent bonds. Because of the directionality of the covalent bonds, these do not tend to form closest– packed arrangements in the crystal. Because of the strength of the covalent bonds, these have very high melting points. –Generally > 1000 °C Dimensionality of the network affects other physical properties.

62. © 2014 Pearson Education, Inc. The Diamond Structure: A Three-Dimensional Network The carbon atoms in a diamond each have four covalent bonds to surrounding atoms. –sp 3 –Tetrahedral geometry This effectively makes each crystal one giant molecule held together by covalent bonds. –You can follow a path of covalent bonds from any atom to every other atom.

63. © 2014 Pearson Education, Inc. Properties of Diamond Very high melting point, ~3800 °C –Need to overcome some covalent bonds Very rigid –Due to the directionality of the covalent bonds Very hard –Due to the strong covalent bonds holding the atoms in position –Used as abrasives Electrical insulator Thermal conductor –Best known Chemically very nonreactive

64. © 2014 Pearson Education, Inc. The Graphite Structure: A Two-Dimensional Network In graphite, the carbon atoms in a sheet are covalently bonded together. –Forming six-member flat rings fused together Similar to benzene Bond length = 142 pm –sp 2 Each C has three sigma and one pi bond. –Trigonal-planar geometry –Each sheet a giant molecule The sheets are then stacked and held together by dispersion forces. –Sheets are 341 pm apart.

65. © 2014 Pearson Education, Inc. Properties of Graphite Hexagonal crystals High melting point, ~3800 °C –Need to overcome some covalent bonding Slippery feel –Because there are only dispersion forces holding the sheets together, they can slide past each other. Glide planes –Lubricants Electrical conductor –Parallel to sheets Thermal insulator Chemically very nonreactive

66. © 2014 Pearson Education, Inc. Silicates ~90% of Earth’s crust Extended arrays of Si  O –Sometimes with Al substituted for Si – aluminosilicates Glass is the amorphous form.

67. © 2014 Pearson Education, Inc. Quartz SiO 2 in pure form –Impurities add color. Three-dimensional array of Si covalently bonded to 4 O – Tetrahedral Melts at ~1600 °C Very hard

68. © 2014 Pearson Education, Inc. Micas There are various kinds of mica that have slightly different compositions, but are all of the general form X 2 Y 4–6 Z 8 O 20 (OH,F) 4. –X is K, Na, or Ca, or less commonly Ba, Rb, or Cs. –Y is Al, Mg, or Fe, or less commonly Mn, Cr, Ti, Li, etc. –Z is chiefly Si or Al, but also may include Fe 3+ or Ti. Minerals that are mainly two-dimensional arrays of Si bonded to O –Hexagonal arrangement of atoms Sheets Chemically stable Thermal and electrical insulator

69. © 2014 Pearson Education, Inc. *Band Theory The structures of metals and covalent network solids result in every atom’s orbitals being shared by the entire structure. For large numbers of atoms, this results in a large number of molecular orbitals that have approximately the same energy; we call this an energy band.

70. © 2014 Pearson Education, Inc. *Band Theory When two atomic orbitals combine they produce both a bonding and an antibonding molecular orbital. When many atomic orbitals combine they produce a band of bonding molecular orbitals and a band of antibonding molecular orbitals. The band of bonding molecular orbitals is called the valence band. The band of antibonding molecular orbitals is called the conduction band.

71. © 2014 Pearson Education, Inc. Molecular Orbitals of Polylithium

72. © 2014 Pearson Education, Inc. *Band Gap At absolute zero, all the electrons will occupy the valence band. As the temperature rises, some of the electrons may acquire enough energy to jump to the conduction band. The difference in energy between the valence band and conduction band is called the band gap. –The larger the band gap, the fewer electrons there are with enough energy to make the jump.

73. © 2014 Pearson Education, Inc. Types of Band Gaps and Conductivity

74. © 2014 Pearson Education, Inc. *Band Gap and Conductivity The more electrons at any one time that a substance has in the conduction band, the better conductor of electricity it is. If the band gap is ~0, then the electrons will be almost as likely to be in the conduction band as the valence band and the material will be a conductor. –Metals –The conductivity of a metal decreases with temperature. If the band gap is small, then a significant number of the electrons will be in the conduction band at normal temperatures and the material will be a semiconductor. –Graphite –The conductivity of a semiconductor increases with temperature. If the band gap is large, then effectively no electrons will be in the conduction band at normal temperatures and the material will be an insulator.

75. © 2014 Pearson Education, Inc. *Doping Semiconductors Doping is adding impurities to the semiconductor’s crystal to increase its conductivity. The goal is to increase the number of electrons in the conduction band. n-type semiconductors do not have enough electrons themselves to add to the conduction band, so they are doped by adding electron-rich impurities. p-type semiconductors are doped with an electron- deficient impurity, resulting in electron “holes” in the valence band. Electrons can jump between these holes in the valence band, allowing conduction of electricity.

76. © 2014 Pearson Education, Inc. Diodes When a p-type semiconductor adjoins an n-type semiconductor, the result is a p–n junction. Electricity can flow across the p–n junction in only one direction; this is called a diode. This also allows the for accumulation of electrical energy, called an amplifier.

77. F. Heating Curves Melting - PE  Solid - KE  Liquid - KE  Boiling - PE  Gas - KE 

79. F. Heating Curves  Temperature Change change in KE (molecular motion) depends on heat capacity  Heat Capacity energy required to raise the temp of 1 gram of a substance by 1°C water has a very high heat capacity

80. F. Heating Curves  Phase Change change in PE (molecular arrangement) temp remains constant  Heat of Fusion (  H fus ) energy required to melt 1 gram of a substance at its m.p.

81. F. Heating Curves  Heat of Vaporization H 2 0 (l) H 2 0 (g)  H=44 Kj  Heat of Fusion H 2 0 (s) H 2 0 (l)  H=6 Kj

82. B. Heating Curves  Heat of Vaporization (  H vap ) energy required to boil 1 gram of a substance at its b.p. usually larger than  H fus …why?  EX: sweating, steam burns