2. Intermolecular Forces, Liquids, and Solids Chapter 11

3. What are the 3 states of matter? What physical properties do we know about them? Liquids and solids are similar compared to gases… –Incompressible –Density doesn’t change with temperature Because solids and liquids are molecules that are so close together while gases are so far apart. What holds them so close together? States of Matter

4. Intermolecular Forces Physical Properties of solids and liquids is due largely to Intermolecular Forces (IMFs). IMFs are forces that exist between molecules within a solid or a liquid There are many different types of IMFs. By understanding the nature & strength of IMFs we can relate composition & structure to physical properties.

5. Types of IMFs Strong Intermolecular Forces: –Covalent Bonding –Ionic Bonding Weak Intermolecular Forces: –Dipole-Dipole –Hydrogen “bonding” –London Dispersion Forces During phase changes, molecules stay intact… NRG used to overcome forces

6. What is a dipole? Dipole – Dipole Forces: Attraction between molecules with dipole moments. Maximizes + to – interactions and minimizes + to + or – to – interactions. 1% the strength of ionic bonds Force strength increases as polarity increases Not important for gases… why? Ion – Dipole Force: Force between Ion and dipole of a polar molecule. Weaker than Dipole – Dipole forces.

7. + - + - + - + - + - + - + - + - + - + -

8. Hydrogen Bonding Special dipole-dipole attraction –Hydrogen covalently bonded to highly electronegative elements –Especially strong with: N, O, F (because they have high electronegativity and they are small enough to get close) Stronger than other dipole-dipole attractions Important for bonding of Water and DNA Effects Boiling Points

9. CH 4 SiH 4 GeH 4 SnH 4 PH 3 NH 3 SbH 3 AsH 3 H2OH2O H2SH2S H 2 Se H 2 Te HF HI HBr HCl Boiling Points 0ºC 100 -100 200

10. Water ++ -- ++

11. London Dispersion Forces Instantaneous dipoles Electrons move randomly, and creates a momentary nonsymmetrical distribution of charge (even in nonpolar molecules) Induces short-lived dipole with neighboring molecules Exist between ALL molecules, but are the weakest forces of attraction

12. Example HH HH HH HH ++  HH HH ++ -- ++ 

13. London Dispersion Forces Weak and short lived Last longer at Low Temp. Polarizabliity – ability to distort the electron cloud; “squishiness” of the molecule. –increases with the number of electrons in a molecule… CCl 4 experiences greater LDFs than CH 4. Larger molecules = more LDF = higher melting and boiling points

14. IMF Summary London Dispersion Forces Dipole-Dipole Forces Hydrogen Bonding ALL 3 are sometimes referred to as “Van der Waals Forces” See flow chart on pg 417 to review what types of forces are involved with different types of bonds. Increasing strength

15. The Liquid State Many of the properties of liquids are due to the internal attraction of atoms: –Surface Tension –Beading –Capillary Action –Viscosity Stronger IMFs cause each of these to increase.

16. Surface tension Molecules in the middle are attracted in all directions. u Molecules at the top are only pulled inside. u Minimizes surface area. High IMF means Higher surface tension

17. Capillary Action Liquids spontaneously rise in a narrow tube. IMFs are cohesive (connect like things) Adhesive forces form between two separate “like” objects (polar molecules attract to polar bonds in material making up container)

19. Beading If a polar substance is placed on a non- polar surface. –There are cohesive, –But no adhesive forces. And Visa Versa H2OH2O Wax Paper

20. Viscosity Measure of a liquid’s resistance to flow. Increases with IMFs Increases with molecular size.

21. Liquids… a structural model Strong IMFs (like solids) Considerable molecular motion (like gases)

22. Phase Changes Phase of a substance determined by… 1.Temperature 2.Pressure 3.Strength of IMFs Stronger Forces = Bigger Effect –Hydrogen Bonding –Polar Molecule (dipole – dipole) –London Dispersion Forces Decreasing Strength

23. Melting / Freezing Vaporization / Condensation Sublimation / Deposition We’ll go more into the energy changes, heating curves, and phase diagrams later in this unit… Types of Phase Changes:

24. Solids Two major types. Amorphous- those with much disorder in their structure. Crystalline- have a regular arrangement of components in their structure.

25. Crystals Lattice- a three dimensional grid that describes the locations of the pieces in a crystalline solid. Unit Cell-The smallest repeating unit in of the lattice. Three common shapes: –Cubic –Body Centered Cubic –Face Centered Cubic

26. Cubic

27. Body-Centered Cubic

28. Face-Centered Cubic

29. Solids Amorphous Solids: Components “frozen in place” and lacking orderly arrangement. Ex: Glass and plastic We’ll focus more crystalline solids…

30. Crystalline Solids: Ionic Solids have ions at lattice points (Salt) Molecular Solids have molecules (Sugar) Metallic Solids (Cu, Fe, Al, Pt…) Covalent Network Solids (Diamonds, Quartz)

31. The book drones on about Using diffraction patterns to identify crystal structures. Talks about metals and the closest packing model. It is interesting, but trivial. We need to focus on metallic bonding. Why do metal atoms stay together. How their bonding effects their properties.

32. Metallic bonding 1s 2s 2p 3s 3p Filled Molecular Orbitals Empty Molecular Orbitals Magnesium Atoms

33. 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 Magnesium Atoms

34. Filled Molecular Orbitals Empty Molecular Orbitals 1s 2s 2p 3s 3p Magnesium Atoms The 3s and 3p orbitals overlap and form molecular orbitals.

35. Filled Molecular Orbitals Empty Molecular Orbitals 1s 2s 2p 3s 3p Magnesium Atoms Electrons in these energy level can travel freely throughout the crystal.

36. Filled Molecular Orbitals Empty Molecular Orbitals 1s 2s 2p 3s 3p Magnesium Atoms This makes metals: -conductors -Malleable because the bonds are flexible.

37. Carbon- A Special Atomic Solid There are three types of solid carbon. Amorphous- coal uninteresting. Diamond- hardest natural substance on earth, insulates both heat and electricity. Graphite- slippery, conducts electricity. How the atoms in these network solids are connected explains why.

38. Diamond- each Carbon is sp3 hybridized, connected to four other carbons. Carbon atoms are locked into tetrahedral shape. Strong  bonds give the huge molecule its hardness.

39. Why is it an insulator? Empty MOsFilled MOs E The space between orbitals make it impossible for electrons to move around

40. Each carbon is connected to three other carbons and sp 2 hybridized. The molecule is flat with 120º angles in fused 6 member rings. The  bonds extend above and below the plane. Graphite is different.

41. This  bond overlap forms a huge  bonding network. Electrons are free to move through out these delocalized orbitals. The layers slide by each other.

42. Molecular solids. Molecules occupy the corners of the lattices. Different molecules have different forces between them. These forces depend on the size of the molecule. They also depend on the strength and nature of dipole moments.

43. Those without dipoles or H bonding. Most are gases at 25ºC. The only forces are London Dispersion Forces. These depend on size of atom. Large molecules (such as I 2 ) can be solids even without dipoles. Example: the Halogens

44. The Halogens F Cl Br I Boiling points 85 K 239 K 333 K 458 K State at 20 o C Gas Liquid Solid

45. Those with dipoles. Dipole-dipole forces are generally stronger than L.D.F. Hydrogen bonding is stronger than Dipole-dipole forces. No matter how strong the intermolecular force, it is always much, much weaker than the forces in bonds. Stronger forces lead to higher melting and freezing points.

46. Water is special Each molecule has two polar O-H bonds. Each molecule has two lone pair on its oxygen. Each oxygen can interact with 4 hydrogen atoms. H O H  

47. Water is special H O H   H O H   H O H   This gives water an especially high melting and boiling point.

48. Ionic Solids The extremes in dipole dipole forces- atoms are actually held together by opposite charges = IONIC BONDS Huge melting and boiling points. Atoms are locked in lattice so hard and brittle. Every electron is accounted for so they are poor conductors-good insulators.

49. Vapor Pressure Vaporization - change from liquid to gas at boiling point. Evaporation - change from liquid to gas below boiling point Heat of Vaporization (  H vap ) - the energy required to vaporize 1 mole at 1 atm.

50. Vaporization is an endothermic process - it requires heat. Energy is required to overcome intermolecular forces. Responsible for cool earth. Why we sweat.

51. Condensation Change from gas to liquid. Achieves a dynamic equilibrium with vaporization in a closed system. What is a closed system? A closed system means matter can’t go in or out. What is a “dynamic equilibrium?”

52. Dynamic Equilibrium 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. 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 condensation equals rate of vaporization

53. Rate of Vaporization = Rate of Condensation Molecules are constantly changing phase “Dynamic” The total amount of liquid and vapor remains constant “Equilibrium” Dynamic equilibrium

54. Vapor pressure The pressure above the liquid at equilibrium. Liquids with high vapor pressures evaporate easily. They are called volatile. Decreases with increasing intermolecular forces. –Bigger molecules (bigger LDF) –More polar molecules (dipole-dipole)

55. Vapor pressure Increases with increasing temperature. Easily measured in a barometer.

56. Dish of Hg P atm = 760 torr A barometer will hold a column of mercury 760 mm high at one atm. If we inject a volatile liquid in the barometer it will rise to the top of the mercury. There it will vaporize and push the column of mercury down. Water

57. Dish of Hg 736 mm Hg Water Vapor The mercury is pushed down by the vapor pressure. P atm = P Hg + P vap P atm - P Hg = P vap 760 - 736 = 24 torr

58. Temperature Effect Kinetic energy # of molecules T1T1 Energy needed to overcome intermolecular forces

59. Kinetic energy # of molecules T1T1 Energy needed to overcome intermolecular forces T1T1 T2T2 At higher temperature more molecules have enough energy - higher vapor pressure. Energy needed to overcome intermolecular forces

60. Changes of state The graph of temperature versus heat applied is called a heating curve. Heat of Fusion represents the transition from Solid to Liquid Heat of Vaporization represents the transition from Liquid to Gas

61. Heating Curve for Water (axis not to scale) Ice Water and Ice Water Water and Steam Steam

62. Heating Curve for Water (axis not to scale) Heat of Fusion Heat of Vaporization Slope is Heat Capacity

63. Copy picture of Heating Curve for Water 1.In which region(s) of the graph is temperature constant? 2.Describe how the amount of energy required to melt a given mass of ice compares to the energy required to vaporize the same mass of water? Explain. 3.Which region of the graph represents the coexistence of solid and liquid? Liquid and Vapor? HEAT IN CHANGES OF STATE

64. 4.When a material condenses or freezes is energy released or absorbed? Explain your reasoning. 5.When a material evaporates or melts, is energy released or absorbed? Explain your reasoning.

65. Energy (as heat) during Changes of State Compounds absorb energy (heat) when breaking bonds. Solids absorb energy to melt into liquids. (ice cubes melting) Liquids absorb energy to evaporate into gases. Energy goes to changing state (breaking bonds), there is no temperature change!

66. Compounds release energy when forming bonds. Gases release energy to condense into a liquid. Liquids release energy to freeze into a solid. Energy goes to changing state (making bonds), there is no temperature change!

67. Equations for Heating Curve No Phase Change: Q = m(▲T)C p Latent Heat of Fusion: (SolidLiquid) Q = m ▲H fus Latent Heat of Vaporization: (LiquidGas) Q= m ▲H vap

68. Melting Point Melting point is determined by the vapor pressure of the solid and the liquid. At the melting point… the vapor pressure of the solid = vapor pressure of the liquid

69. Solid Water Liquid Water Water Vapor Vapor

70. Solid Water Liquid Water Water Vapor Vapor If the vapor pressure of the solid is higher than that of the liquid the solid will release molecules to achieve equilibrium.

71. Solid Water Liquid Water Water Vapor Vapor While the molecules of water condense to a liquid.

72. This can only happen if the temperature is above the freezing point since solid is turning to liquid. Solid Water Liquid Water Water Vapor Vapor

73. If the vapor pressure of the liquid is higher than that of the solid, the liquid will release molecules to achieve equilibrium. Solid Water Liquid Water Water Vapor Vapor

74. Solid Water Liquid Water Water Vapor Vapor While the molecules condense to a solid.

75. The temperature must be below the freezing point since the liquid is turning to a solid. Solid Water Liquid Water Water Vapor Vapor

76. 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. Solid Water Liquid Water Water Vapor Vapor

77. Boiling Point Reached when the vapor pressure equals the external pressure. Normal boiling point is the boiling point at 1 atm pressure. Super heating - Heating above the boiling point. Supercooling - Cooling below the freezing point.

78. Using Vapor Pressure Data Clausius-Clapeyron equation is used to determine the vapor pressure of a liquid at a different temp. or to find the heat of vaporization of a substance. ln(P vap,T1 /P vapT2 ) = (  H vap /R)[(1/T 2 ) – (1/T 1 )] The graph of ln P vs 1/T is linear with a negative slope.

79. Determine the boiling point of water in Breckenridge, Colorado, where the atmospheric pressure is 520 torr. For water,  H vap = 40.7 kJ/mol. P 1 = 520 Torr; P 2 = 760 Torr R = 8.3145 J/K x mol T 2 = 373 K; T 1 = ? T1 = 362 K or 89 o C

80. Phase Diagrams. A plot of temperature versus pressure for a closed system, with lines to indicate where there is a phase change.

81. Temperature Solid Liquid Gas 1 Atm A A B B C C D D D Pressure D

82. Dissecting our Phase Diagram A = Sublimation / Deposition B = Vaporization / Condensation C = Melting / Freezing D = Liquid moving up past the critical point and transforming into a gas

83. Phase Diagram terminology Triple Point – all three phases exist in equilibrium in a “closed system” Critical Temperature – temp. above which the substance cannot exist as a liquid regardless of how great the pressure. Critical Pressure – pressure required to produce liquefaction at the critical temperature. Critical Point – Point defined by the critical temp. and pressure (For Water: 374 o C and 218 atm)

84. Solid Liquid Gas Triple Point Critical Point Temperature Pressure

85. Solid Liquid Gas This is the phase diagram for water. The density of liquid water is higher than solid water. Temperature Pressure

86. Solid Liquid Gas 1 Atm This is the phase diagram for CO 2 The solid is more dense than the liquid The solid sublimes at 1 atm. Temperature Pressure

87. How does the melting/freezing line differ for the phase diagram of Water and Carbon Dioxide? What may cause that difference?