1. Intermolecular Forces, Liquids and Solids. Goals: 1.Describe intermolecular forces and their effects. 2.List the effects of hydrogen bonding. 3.List the properties of liquids. 4.Draw basic cubic unit cells. 5.Relate unit cells for ionic compounds to formulas. 6.List the properties of solids. 7.Use phase diagrams to predict physical state of materials.

2. Intermolecular Forces Have studied INTRA molecular forces— the forces holding atoms together to form molecules. Now turn to forces between molecules — INTER molecular forces. Forces between molecules, between ions, or between molecules and ions.

3. Liquids and Solids Why at atmospheric pressure and room temperature, O 2 is a gas H 2 O is a liquid C 6 H 12 O 6 is a solid - little interaction between particles: no specific volume or shape - - particles held together with sufficient force that give a volume - - particles are rigidly held in a volume and shape

4. Intermolecular Forces Intramolecular forces – chemical bonds – bind atoms to one another within molecules. Intermolecular forces – attractive forces between molecules. – Determine physical properties of liquids and solids – without intermolecular forces there would not be liquids or solids. – Intermolecular forces are relatively unimportant in gases because molecules are far apart (the ideal gas law assumes there are non).

5. Intermolecular Forces Determine the physical properties of substances. Molecules have a tendency to remain apart from each other. Intermolecular forces of attraction overcome this tendency more effectively at low temperature (molecules have low energies) and at high pressures (molecules are close together). A substance is likely to exist as a gas at ________ temperatures (energetic molecules) and at _______ pressures (molecules far apart). A substance is likely to exist as a solid at ______ temperatures and ________________pressures (closely packed molecules). The liquid – in-between state – exists at intermediate temperatures and moderate to high pressures.

6. For comparison: Ion-Ion Forces Na + —Cl - in salt These are the strongest forces. Lead to solids with high melting temperatures. NaCl, mp = 800 o C MgO, mp = 2800 o C

7. For comparison: Covalent Bond Forces C–H, 413 kJ/mol C=C, 610 kJ/mol C–C, 346 kJ/mol CN, 887 kJ/mol

8. Summary of Intermolecular Forces p. 604

9. Attraction between Ions and Permanent Dipoles Water is highly polar and can interact with positive ions to give hydrated ions in water.

10. Attraction between Ions and Permanent Dipoles Many metal ions are hydrated. This is the reason metal salts dissolve in water.

11. Attraction between Ions and Permanent Dipoles Attraction between ions and dipole depends on ion _______ and ion-dipole ________. Measured by ∆H for M n+ + H 2 O --> [M(H 2 O) x ] n+ -1922 kJ/mol -405 kJ/mol -263 kJ/mol

12. Attraction between Dipoles Such forces bind molecules having permanent dipoles to one another.

13. Dipole- Dipole Forces Influence of dipole-dipole forces is seen in the boiling points of simple molecules. CompdMol. Wt.Boil Point N 2 28-196 o C CO28-192 o C Br 2 16059 o C ICl16297 o C

14. Dipole- Dipole Forces Which has the higher boiling point? If a molecule is polar will have dipole-dipole interactions, and then the boiling point needed to break the interactions will be higher than that needed for a nonpolar molecule. –HCl or F 2 –SiH 4 or PH 3

15. Hydrogen Bonding A special form of dipole-dipole attraction, which enhances dipole-dipole attractions. H-bonding is strongest when X and Y are N, O, or F

16. Hydrogen Bonding A hydrogen bond between molecules is an intermolecular force in which a hydrogen atom covalently bonded to a nonmetal atom in a molecule is simultaneously attracted to a nonmetal atom of a neighboring molecule. The strongest hydrogen bonds are formed if the nonmetal atoms are small and highly electronegative (N, O, and F). H 2 O

17. Hydrogen Bonding H-Bonding Between Methanol and WaterH-bondH-bond

18. Hydrogen Bonding H-Bonding Between Two Methanol Molecules H-bondH-bond ---- ++++ ----

19. Hydrogen Bonding in Water H-bonding is especially strong in water because the O—H bond is very polar there are 2 lone pairs on the O atom Accounts for many of water’s unique properties.

20. Hydrogen Bonding in Water Ice has open lattice-like structure. Ice density is _______ than liquid. And so solid floats on water.

21. Hydrogen Bonding in Water Ice has open lattice-like structure. Ice density is < liquid and so solid floats on water. One of the VERY few substances where solid is LESS DENSE than the liquid.

22. Hydrogen Bonding in Water A consequence of H-bonding

23. Hydrogen Bonding in Water H bonds ---> abnormally high specific heat capacity of water (4.184 J/gK) This is the reason water is used to put out fires, it is the reason lakes/oceans control climate, and is the reason thunderstorms release huge energy.

24. Hydrogen Bonding in Water H bonds leads to abnormally high boiling point of water.

25. Summary of Intermolecular Forces p. 604

26. Boiling Points of Simple H- containing Compounds

27. Hydrogen Bonding in Biology H-bonding is especially strong in biological systems — such as DNA. DNA — helical chains of phosphate groups and sugar molecules. Chains are helical because of tetrahedral geometry of P, C, and O. Chains bind to one another by specific hydrogen bonding between pairs of Lewis bases. —adenine with thymine —guanine with cytosine

28. Hydrogen Bonding and DNA Hydrogen bonding is also the force that binds the two chains of a DNA molecule together, so hydrogen bonding is a key to understanding the replication of organisms. Adenine and Thymine form 2 hydrogen bonds Guanine and Cytosine form 3 hydrogen bonds

29. Base Pairing through Hydrogen Bonds

30. Hydrogen Bonds Which has a higher boiling point? H 3 C-CH 3 (ethane) or H 3 C-OH (methanol) H 2 O (water) or HCl (hydrogen chloride)

31. Hydrogen Bonds In which of the following compounds would hydrogen bonding be an important intermolecular force?

32. Forces Involving Induced Dipoles How can non-polar molecules such as O 2 and I 2 dissolve in water? Dipole-induced dipole

33. Forces Involving Induced Dipoles Solubility increases with mass the gas

34. Forces Involving Induced Dipoles Process of inducing a dipole is ________. Degree to which electron cloud of an atom or molecule can be distorted in its _____________. The larger the molecule, the more polarizable.

35. Induced Dipoles Consider I 2 dissolving in ethanol, CH 3 CH 2 OH O H -  +  I-I R -  +  O H +  -  R The alcohol temporarily creates or INDUCES a dipole in I 2.

36. Induced Dipoles Formation of a dipole in two nonpolar I 2 molecules. Induced dipole- induced dipole

37. Induced Dipoles The induced forces between I 2 molecules are very weak, so solid I 2 sublimes (goes from a solid to gaseous molecules).

38. Induced Dipoles The magnitude of the induced dipole depends on the tendency to be distorted. Higher molec. weight ---> larger induced dipoles. MoleculeBoiling Point ( o C) CH 4 (methane) - 161.5 C 2 H 6 (ethane)- 88.6 C 3 H 8 (propane) - 42.1 C 4 H 10 (butane) - 0.5 CH 4 C2H6C2H6C2H6C2H6 C3H8C3H8C3H8C3H8 C 4 H 10

39. Induced Dipoles Which has the higher boiling point? Ne or Xe C 2 H 6 (ethane) or C 4 H 10 (butane) C 2 H 5 OH (ethanol) or C 3 H 7 OH (propanol)

40. What intermolecular force(s) must be overcome to: Melt ice? Sublime solid I 2 ? Convert liquid NH 3 to NH 3 vapor? H H H N

41. Intermolecular Forces Students should be familiar with identifying the types of intermolecular forces and predicting physical properties of substances.

42. In Liquids: Molecules are in constant motion, but the movement is restricted by neighboring molecules (___ diffusion, compared to gases). Molecules of a liquid are much closer together than those of a gas (_____ compressibility). There are appreciable intermolec. Forces. Liquids do not fill the container.

43. Liquids The two key properties we need to describe are EVAPORATION and its opposite—CONDENSATION break IM bonds make IM bonds Add energy Remove energy LIQUID VAPOR<---condensation evaporation--->

44. Liquids - Evaporation To evaporate, molecules must have sufficient energy to break IM forces. Breaking IM forces requires energy. The process of evaporation is endothermic.

45. Liquids – Distribution of Energies 0 Number of molecules Molecular energy higher Tlower T See Figure 13.14 Minimum energy req’d to break IM forces and evaporate Distribution of molecular energies in a liquid. KE is proportional to T.

46. Liquids – Distribution of Energies At higher T a much larger number of molecules has high enough energy to break IM forces and move from liquid to vapor state. High E molecules carry away E. You cool down when sweating or after swimming.

47. Equilibrium Vapor Pressure Liquid in flask evaporates and exerts pressure on manometer.

48. Summary of Intermolecular Forces

49. Equilibrium Vapor Pressure When molecules of liquid are in the vapor state, they exert a VAPOR PRESSURE. EQUILIBRIUM VAPOR PRESSURE is the pressure exerted by a vapor over a liquid in a closed container when the rate of evaporation = the rate of condensation. System is in a state of Dynamic Equilibrium Higher Intermolecular Forces -> Lower Vapor Pressure

50. Boiling Point Boiling point – of a liquid is the temperature at which its vapor pressure becomes equal to atmospheric pressure. –Increases when external pressure is increased. –Normal boiling point – temperature at which a liquid boils under standard pressure (1 atm).

51. Equilibrium Vapor Pressure Vapor Pressure as a function of T: 1. The curves show all conditions of P conditions of P and T where LIQ and VAP and T where LIQ and VAP are in EQUILIBRIUM are in EQUILIBRIUM 2. The Vapor Pressure rises with T. with T. 3. When VP = external P, the liquid boils. This means that BP’s of liquids change with altitude. This means that BP’s of liquids change with altitude.

52. Boiling at lower Pressure When pressure is lowered, the vapor pressure can equal the external pressure at a lower temperature.

53. Consequences of Vapor Pressure Changes When can cools, vapor pressure of water drops. Pressure in the can is less than that of atmosphere, so can is crushed.

54. Liquids 4. If external P = 760 mm Hg, T of boiling is the NORMAL BOILING POINT 5.VP of a given molecule at a given T depends on IM forces. Here the VP’s are in the order C 2 H 5 H 5 C 2 H H 5 C 2 H H wateralcoholether increasing strength of IM interactions extensive H-bonds dipole- dipole O O O

55. Liquids HEAT OF VAPORIZATION is the heat req’d (at constant P) to vaporize the liquid. LIQ + heat ---> VAP Compd.∆H vap (kJ/mol) IM Force H 2 O40.7 (100 o C)H-bonds SO 2 26.8 (-47 o C)dipole Xe12.6 (-107 o C)induced dipole

56. Heat of Vaporization Heat is required for the conversion of a liquid to a vapor. Some volatile liquids are used to cool the skin. When a liquid is condensed it releases heat. The heat of vaporization is characteristic of a given liquid. Molar heat of vaporization – quantity of heat required to vaporize 1 mol of a liquid at a constant pressure (units: cal/mol). Given the molar heat of vaporization, we can calculate the heat of vaporization in cal/g or kJ/g using the molar mass.

57. Clasius-Clapeyron Equation Clausius-Clapeyron equation — used to find ∆H˚ vap.Clausius-Clapeyron equation — used to find ∆H˚ vap. The logarithm of the vapor pressure P is proportional to ∆H vaporization and to 1/T.The logarithm of the vapor pressure P is proportional to ∆H vaporization and to 1/T. ln P = –(∆H˚ vap /RT) + C ln P = –(∆H˚ vap /RT) + C

58. Calculate the enthalpy of vaporization (  H o vap ) of ethylene glycol. This compound has a vapor pressure of 14.9 mmHg at 373 K and a vapor pressure of 49.1 mmHg at 398 K.

59. Liquids Molecules at surface behave differently than those in the interior. Molecules at surface experience net INWARD force of attraction.Molecules at surface experience net INWARD force of attraction. This leads to SURFACE TENSION — the energy req’d to break the surface.This leads to SURFACE TENSION — the energy req’d to break the surface.

60. Surface Tension Surface tension also leads to spherical liquid droplets.

61. Surface Tension Surface tension – force or tension that resists disruption (horizontal needle over water, insects over water). –A molecule at the surface of the liquid is attracted only by molecules at its sides and below it; forces tend to pull inward and cause the liquid to contract – Small amount of liquid will “bead” to minimize its surface area. –Soaps and other detergents act in part by lowering surface tension, enabling water to spread out and wet a solid surface. Leaf surface without surfactant with surfactant Surface

62. Properties of Liquids: Surface tension – force or tension that resists disruption (horizontal needle over water, insects over water). Viscosity – resistance to flow.

63. Properties of Liquids Viscosity – resistance to flow. – Liquids with low viscosity (“thin liquids”) generally consist of small, symmetrical molecules with weak intermolecular forces. Viscous liquids are generally made up or large or unsymmetrical molecules with fairly strong intermolecular forces. – Viscosity generally decreases with increasing temperature. HexaneOctane Iso-octane

64. Properties of Liquids Intermolecular forces also lead to CAPILLARY action and to the existence of a concave meniscus for a water column. concave meniscus H 2 O in glass tube ADHESIVE FORCES between water and glass COHESIVE FORCES between water molecules Movement of water up a piece of paper depends on H- bonds between H 2 O and the -OH groups of the cellulose in the paper.

65. Properties of Liquids Answer with increases, decreases, or does not change: If the intermolecular forces in a liquid increase, the normal boiling point of the liquid. If the intermolecular forces in a liquid decrease, the vapor pressure of the liquid. If the temperature of a liquid increases, the vapor pressure. increases

66. Metallic and Ionic Solids

67. Solids Particles (atoms, molecules, or ions) are close together (incompressible). In solids there is little motion other than vibration about a fixed point (particles in liquid are in constant, restricted motion). An increase in temperature will increase the vigor of the vibrations in a solid. If vibrations become strong enough, the solid will melt. Solid crystals – stacking unit cells – portion of a crystal which represents the regular, repeating manner extending in three dimensions that atoms, ions, or molecules are organized. –Crystal lattice – framework on which a pattern is outline. There are 14 different lattices to describe all crystalline solids. Amorphous solids – are solids that lack this ordered arrangement: glasses.

68. From Solid to Liquid: Melting (Fusion) Melting point – temperature at which the solid becomes a liquid. The heat energy is absorbed by the particles of the solid; the particles vibrate with more and more vigor until, the forces holding the particles in a particular arrangement are overcome. A high melting point is one indication of the forces holding a solid together are very strong.

69. Heat of Fusion Molar heat of fusion – quantity of heat required to convert 1 mol of a solid to liquid. Generally, for 1 mol of a substance, it takes more heat energy to vaporize it than to melt it. It takes more energy to vaporize the liquid because the attraction between particles must be almost completely overcome (to get to gas).

70. Types of Solids TYPEEXAMPLEFORCE Ionic NaCl, CaF 2, ZnSIon-ion MetallicNa, FeMetallic MolecularIce, I 2 Dipole Ind. dipole NetworkDiamondExtended GraphiteCovalent

71. Network Solids Diamond Graphite A comparison of diamond (pure carbon) with silicon.

72. Network Solids SiC also is important for tooling in the semiconductor industry, for laser mirrors, as a substrate for wear-resistant diamond coatings, as an abrasive and grinding wheel, as heating elements and igniters, as an additive for reinforcement of metals, and for numerous refractories applications.

73. Crystal Lattices Regular 3-D arrangements of equivalent LATTICE POINTS in space.Regular 3-D arrangements of equivalent LATTICE POINTS in space. Lattice points define UNIT CELLSLattice points define UNIT CELLS –smallest repeating internal unit that has the symmetry characteristic of the solid.

74. Cubic Unit Cells There are 7 basic crystal systems, but we are only concerned with CUBIC. All angles are 90 degrees All sides equal length

75. Cubic Unit Cells Very rare: Po Fe, K, Na, Cr, Mo, W Al, Cu, Pb, Ag

76. Cubic Unit Cells Each atom is at a corner of a unit cell and is shared among 8 unit cells. Each edge is shared with 4 cells Each face is part of two cells.

77. Atom Packing in Units Cells Assume atoms are hard spheres and that crystals are built by PACKING of these spheres as efficiently as possible. FCC is more efficient than either BC or SC. Leads to layers of atoms.

78. Solids can be classified by the types of intermolecular forces holding the particles together: – –Ionic solids Have_________ at definite points in the lattice. Ionic forces are very strong. Ionic solids have ______ melting points and _____vapor pressures, and are quite hard. – –Molecular crystals Have discrete _______________ at the lattice points Held together by rather _______ dispersion forces (nonpolar), dipolar forces (polar), or hydrogen bonds (hydrogen-bonded). Molecular solids often have ________ melting points than ionic solids. – –Covalent network crystals Have _________ at the lattice points. _________ are joined into extensive networks by ______________; each crystal is in essence one large molecule. They are extremely hard and nonvolatile, melt with decomposition at _______________temperatures. – –Metallic solids Considered as _______________ at the lattice sites. The lost valance electrons are released and distributed throughout the lattice. Electrons can move freely about the lattice; metals are ________ conductors of heat and electricity.

79. Classify the Solids Sodium Chloride (NaCl) Solid Nitrogen (N 2 (s) ) Copper (Cu) Quartz (SiO 2 ) Potassium (K)Sulfur (S 8 ) Iron (II) Sulfate (FeSO 4 )Brass (Copper and Zinc alloy) Ice (H 2 O (s) ) Carbon (C) Glucose IonicCovalent NetworkMetallic Molecular (ionic compounds) (atoms in the lattice)(metals) (molecules in the lattice)

80. Atom Packing in Units Cells Unit Cell Type Net Number Atoms Unit Cell Type Net Number Atoms SC SCBCCFCC 1 2 4

81. Phase Diagrams

82. Transitions between Phases Lines connect all conditions of T and P where EQUILIBRIUM exists between the phases on either side of the line. (At equilibrium particles move from liquid to gas as fast as they move from gas to liquid, for example.)

83. Phase Diagram of Water Solid phase Liquid phase Gas phase

84. Phase Equilibria for Water Solid- liquid Gas- Liquid Gas- Solid

85. Triple Point for Water At the TRIPLE POINT all three phases are in equilibrium.

86. Critical T and P Above critical T no pure liquid exists no matter how high the pressure. As P and T increase, you finally reach the CRITICAL T and P

87. Critical T and P COMPDT c ( o C)P c (atm) H 2 O374218 CO 2 3173 CH 4 -8246 Freon-1211241 (CCl 2 F 2 ) Notice that T c and P c depend on intermolecular forces.

88. A fluorocarbon, CF4, has a critical temperature of -45.7 o C, a critical pressure of 37 atm, and a normal bp of -128 o C. Are there any conditions under which this compound can be a liquid at room temperature?

89. CO 2 Phase Diagram

90. CO 2 Phases Separate phases Increasing pressure More pressure Supercritcal CO 2 Densities are the same For engineering purposes, supercritical fluids can be regarded as “hybrid solvents” with properties between those of gases and liquids, i.e. a solvent with a low viscosity, high diffusion rates and no surface tensionviscositydiffusion ratessurface tension

91. Solid-Liquid Equilibria In any system, if you increase Pressure the DENSITY will go up. Therefore — as P goes up, equilibrium favors phase with the larger density (or SMALLER volume/gram). Liquid H 2 OSolid H 2 O Density1 g/cm 3 0.917 g/cm 3 cm 3 /gram11.09

92. Solid-Liquid Equilibria Raising the pressure at constant T causes water to melt. The NEGATIVE SLOPE of the line is unique to H 2 O. Almost everything else has positive slope.

93. Solid-Liquid Equilibria The behavior of water under pressure is an example of LE CHATELIER’S PRINCIPLE At Solid/Liquid equilibrium, raising P squeezes the solid. It responds by going to phase with greater density, i.e., the liquid phase.

94. Solid-Liquid Equilibria At P < 4.58 mmHg and T < 0.0098 ˚C solid H 2 O can go directly to vapor. This process is called SUBLIMATION

95. Sublimation Sublimation – passage of molecules directly from the solid to the vapor state. Deposition – the reverse process of sublimation – condensation of a vapor to a solid. The vapor pressure of ice at 0 o C is 4.58 mmHg. The vapor pressure of iodine at 39.4 o C is 760 mmHg (1 atm).

96. Water: An unusual liquid At room temperature it is the only liquid compound with a molar mass as low as 18 g/mol. Unlike most substances, the solid form of water (ice) is less dense than the liquid (ice remains on top layer of frozen lakes). In liquid state, water molecules are close together but randomly arranged. In ice water molecules are ordered with large hexagonal holes. Liquid water has a higher density than most other familiar liquids (those insoluble in water –oil - float on its surface). 2 Boiling point ( o C) Period 345 100 0 -100 -200 H2OH2O H2SH2S H 2 Se H 2 Te

97. Water: An unusual liquid Structure of ice at normal atmospheric pressure. It is a hydrogen-bonded network of water molecules.

98. Water: An unusual liquid Water has a very high specific heat (Table 7.6). –It takes 1 cal of heat to raise the temperature of 1 g of water 1 o C (10 times larger than that of iron). Water acts as a giant thermostat to moderate daily temperatures due to the large amount of heat water gives off for a drop in temperature. Water has a high heat of vaporization. –Large amounts of body heat, produced as a by-product of metabolic processes, can be dissipated by the evaporation of small amounts of water (perspiration) from the skin. The heat of vaporization is obtained from the body, and the body is cooled. The water molecule is polar; in the liquid state water molecules are strongly associated by hydrogen bonding. A input of a large amount of energy is needed if vaporization is to take place.

99. Remember Go over all the contents of your textbook. Practice with examples and with problems at the end of the chapter. Practice with OWL tutor. Work on your OWL assignment for Chapter 13.