1. Ch 10 Particle Forces

2. States of Matter Solid- Particles moving about a fixed point Liquid-Particles moving about a moving point Gas-Particles filling the volume of the container with complete random motions.

3. Particle Forces Affect Solubility Vapor Pressures Freezing Points Boiling Points

4. Particle Forces Intramolecular forces (Relative strength = 100)  Ionic bonding  Covalent bonding Interparticle forces  Ion-dipole forces  Dipole-dipole (Polar molecules)  (relative Strength = 1)  London Forces (Dispersion forces)( Nonpolar molecules)  (relative strength = 1)  Hydrogen Bonding (Relative strength = 10)

5. Ion-Ion Interactions Coulomb’s law states that the energy (E) of the interaction between two ions is directly proportional to the product of the charges of the two ions (Q 1 and Q 2 ) and inversely proportional to the distance (d) between them.

6. Predicting Forces of Attraction Coulombs Law indicates the increases in the charges of ions will cause an increase in the force of attraction between a cation and an anion. Increases in the distance between ions will decrease the force of attraction between them.

7. Size of Ions

8. Lattice Energy The lattice energy (U) of an ionic compound is the energy released when one mole of the ionic compound forms from its free ions in the gas phase. M + (g) + X - (g) ---> MX(s)

9. Comparing Lattice Energies Lattice Energies of Common Ionic Compounds CompoundU(kJ/mol) LiF-1047 LiCl-864 NaCl-790 KCl-720 KBr-691 MgCl 2 -2540 MgO-3791

10. Practice Determine which salt has the greater lattice energy. A. MgO and NaF B. MgO and MgS

11. Lattice Energy Using Hess’s Law

12. Electron Affinity Electron affinity is the energy change occurring when one mole of electrons combines with one mole of atoms or ion in the gas phase. Step 4 in diagram on the last slide. Cl(g) + e - (g) ---> Cl - (g) ΔH Ea = -349 kj/mole

13. Calculating U Na + (g) + e - (g) ---> Na(g) -  H IE1 Na(g) ---> Na(s)-  H sub Cl - (g) ---> Cl(g) + e - (g)-  H EA Cl(g) ---> 1/2Cl 2 (g) -1/2  H BE Na(s) + 1/2Cl 2 (g) ---> NaCl(s)  H f Na + (g) + Cl - (g) ---> NaCl(s) U U =  H f - 1/2  H BE -  H EA -  H sub -  H IE1

14. Lattice energy for NaCl.

15. Interactions Involving Polar Molecules An ion-dipole interaction occurs between an ion and the partial charge of a molecule with a permanent dipole. The cluster of water molecules that surround an ion in aqueous medium is a sphere of hydration.

16. Illustrates of Ion-Dipole Interaction

17. The Solution Process Bond Breaking Processes Break solute particle forces (expanding the solute), endothermic Break solvent particle forces (expanding the solvent), endothermic

18. The Solution Process Attractive Forces Energy released when solute solvent are attracted, exothermic Energy is released due to new attractions  Ion dipole if the solute is ionic and the solvent polar.  London-Dipole for nonpolar solute and polar solvent  Dipole-dipole for polar solute and polar solvent

19. The Solution Process Theromodynamics Enthalpy Entropy ΔS (Perfect crystal, assumed to be zero) Gibbs free energy ΔG ΔG = ΔH - T ΔS ΔG < 0, spontaneous change ΔG = 0, equilibrium ΔG > 0, nonspontaneous

20. The Solution Process Oil dissolving in water London forces holding the oil molecules together are large do to the large surface area of the oil The hydrogen bonds holding water molecules together are large The forces of attraction of between nonpolar oil and polar water are weak at best Thus the overall process is highly endothermic and not allowed thermo chemically

21. The Solution Process Oil dissolving in water Entropy should be greater than zero Free energy should be greater than zero, since the process is highly endothermic Thus the overall process is nonspontaneous

22. The Solution Process Sodium chloride dissolving in water Large amount of energy is required to break the ionic lattice of the sodium chloride (expand solute) Large amount of energy is required to separate the water molecules to expand the solvent breaking hydrogen bonds Formation of the ion dipole forces releases a large amount of energy, strong forces (why?) The sum of the enthalpies is about +6 kJ (slightly endothermic), which is easily overcome by the entropy of the solution formation.

23. Water as a Solvent Water most important solvent, important to understand its solvent properties Most of the unusual solvent properties of water stem from it hydrogen bonding nature Consider the following ∆S of solution KCl →75j/K-mole LiF→-36j/K-mole CaS→-138 j/K-mole

24. Water as a Solvent We would expect ∆S>0 for all solutions, right? But two are negative, why? Obviously, something must be happening for the increased order. Ion-dipole forces are ordering the water molecules around the ions, thus causing more order in water i.e. less positions for water than in the pure liquid state

25. Water as a Solvent Smaller ions, have stronger ion dipole forces, thus pulling water closer, therefore less positions Also, ions with a charge greater than one will attract to water stronger than a one plus charge, thus more order due to less space between particles

26. Dipole-Dipole Interactions Dipole-dipole interactions are attractive forces between polar molecules. An example is the interaction between water molecules. The hydrogen bond is a special class of dipole-dipole interactions due to its strength.

27. Dipole-Dipole Forces Dipole-dipole (Polar molecules)  Alignment of polar molecules to two electrodes charged + and δ–  Forces compared to ionic/covalent are about 1 in strength compared to a scale of 100, thus 1% H Cl H Cl H Cl δ–δ– δ–δ– δ–δ– δ+δ+δ+δ+ δ+δ+

28. Slide 28 of 35 Dipole Dipole Interactions

29. Hydrogen Bonding Hydrogen bonding a stronger intermolecular force involving hydrogen and usually N, O, F, and sometimes Cl –Stronger that dipole-dipole, about 10 out of 100, or 10 –Hydrogen needs to be directly bonded to the heteroatom –Since hydrogen is small it can get close to the heteroatom –Also, the second factor is the great polarity of the bond.

30. Slide 30 Hydrogen Bonding in HF(g)

31. Slide 31 Hydrogen Bonding in Water around a molecule in the solid in the liquid

32. Boiling Points of Binary Hydrides

33. Interacting Nonpolar Molecules Dispersion forces (London forces) are intermolecular forces caused by the presence of temporary dipoles in molecules. A instantaneous dipole (or induced dipole) is a separation of charge produced in an atom or molecule by a momentary uneven distribution of electrons.

34. Illustrations

35. Strength of Dispersion Forces The strength of dispersion forces depends on the polarizability of the atoms or molecules involved. Poarizability is a term that describes the relative ease with which an electron cloud is distorted by an external charge. Larger atoms or molecules are generally more polarizable than small atoms or molecules.

36. London Forces (Dispersion) Induced dipoles (Instantaneous ) Strength is surface area dependent More significant in larger molecules All molecules show dispersion forces Larger molecules are more polarizable

37. Slide 37 Instantaneous and Induced Dipoles

38. Molar Mass and Boiling Points of Common Species. HalogenM(g/mol)Bp(K)Noble GasM(g/mol)Bp(K) He24 F2F2 3885Ne2027 Cl 2 71239Ar4087 Br 2 160332Kr84120 I2I2 254457Xe131165 Rn211 Molar Mass and Boiling Point

39. HydrocarbonAlcohol Molecular Formula Molar Mass Bp ( o C) Molecular Formula Molar Mass Bp ( o C) CH 4 16.04-161.5 CH 3 30.07-88CH 3 OH32.0464.5 CH 3 CH 2 CH 3 44.09-42CH 3 CH 2 OH46.0778.5 CH 3 CH(CH)CH 3 58.12-11.7CH 3 CH(OH)CH 3 60.0982 CH 3 CH 2 CH 2 CH 3 58.12-0.5CH 3 CH 2 CH 2 OH60.0997 London vs Hydrogen Bonding

40. The Effect of Shape on Forces

41. Practice Rank the following compound in order of increasing boiling point. CH 3 OH, CH 3 CH 2 CH 2 CH 3, and CH 3 CH 2 OCH 3

42. Practice Rank the following compound in order of increasing boiling point. CH 3 OH, CH 3 CH 2 CH 2 CH 3, and CH 3 CH 2 OCH 3 CH 3 OH CH 3 CH 2 CH 2 CH 3 CH 3 CH 2 OCH 3 MM 32.0 58.0 60.0 IM Forces London and H-bonding London, only London and Dipole-dipole

43. Practice Rank the following compound in order of increasing boiling point. CH 3 OH, CH 3 CH 2 CH 2 CH 3, and CH 3 CH 2 OCH 3 CH 3 OH CH 3 CH 2 CH 2 CH 3 CH 3 CH 2 OCH 3 MM 32.0 58.0 IM Forces London and H-bonding London, only London and Dipole-dipole The order is: CH 3 CH 2 CH 2 CH 3 <CH 3 CH 2 OCH 3 <CH 3 OH

44. Polarity and Solubility If two or more liquids are miscible, they form a homogeneous solution when mixed in any proportion. Ionic materials are more soluble in polar solvents then in nonpolar solvents. Nonpolar materials are soluble in nonpolar solvents. Like dissolves like

45. Polarity and Solubility If two or more liquids are miscible, they form a homogeneous solution when mixed in any proportion. Ionic materials are more soluble in polar solvents then in nonpolar solvents. Nonpolar materials are soluble in nonpolar solvents.

46. Polarity and Solubility How does polarity effect solubility? The thermodynamic argument, is that the lower the potential energy, the more stable the system. If subtracting the potential energy of the solute from the potential energy of the original solute and solvent is negative (exothermic) then solution is thermodynamically favored.

47. Polarity and Solubility How does polarity effect solubility? Non polar solute and solvent: The forces holding these particles together are London Dispersion forces, the weakest of all of the inter-particle forces. The strength of these forces are relative to the surface area if solute and solvent are of similar size, then about the same amount of energy is required to separate solute and solvent particles from each other. And about the same amount of energy is released when solute and solvent are attracted to each other forming a solution. Thus we predict non polar solutes and solvents should dissolve

48. Polarity and Solubility How does polarity effect solubility? Non polar solute and polar solvent: Considering solutes and solvents of similar surface area it should be noted that more energy is required to separate the polar solvent molecules from each other, since dipole- dipole interactions are stronger. The only interaction between a nonpolar solute and polar solvent would be London Dispersion forces, so the energy released is much less than required for separating the solvent and solute. Subtracting the potential energy of the products from reactants would give a positive (endothermic) result and the solution would be less stable than the dissolution.

49. Practice Rank the following compound in order of increasing boiling point. CH 3 OH, CH 3 CH 2 CH 2 CH 3, and CH 3 CH 2 OCH 3

50. Solubility of Gases in Water Henry’s Law states that the solubility of a sparingly soluble chemically unreactive gas in a liquid is proportional to the partial pressure of the gas. C gas = k H P gas where C is the concentration of the gas, k H is Henry’s Law constant for the gas.

51. Henry’s Law Constants Gask H [mol/(Latm)]k H [mol/(kgmmHg)] He3.5 x 10 -4 5.1 x 10 -7 O2O2 1.3 x 10 -3 1.9 x 10 -6 N2N2 6.7 x 10 -4 9.7 x 10 -7 CO 2 3.5 x 10 -2 5.1 x 10 -5

52. Terms A hydrophobic (“water-fearing) interaction repels water and diminishes water solubility. A hydrophilic (“water-loving”) interaction attracts water and promotes water solubility.

53. Affects of Intermolecular Forces Solubility Vapor Pressures Freezing Points Boiling Points Surface tension

54. Vapor Pressure Vaporization or evaporation is the transformation of molecules in the liquid phase to the gas phase. Vapor pressure is the force exerted at a given temperature by a vapor in equilibrium with its liquid phase.

55. Vapor Pressure

56. The normal boiling point of a liquid is the temperature at which its vapor pressure equals 1 atmosphere.

57. Vapor Pressure of Solutions What evaporates faster, sugar water or pure water? s w s w s w w w w

58. Vapor Pressure of Solutions What evaporates faster, sugar water or pure water? s w s w s w w w w w(g)

59. Vapor Pressure of Solutions What evaporates faster, sugar water or pure water? s w s w s w w w w w(g) Pure water evaporates faster, since there are more water particles on the surface, thus lowering the average kinetic energy. Evaporation of a solution is inversely proportional to concentration.

60. July 2009General Chemistry: Chapter 11 Slide 60 of 46 Vapor Pressures of Solutions Roault, 1880s.  Dissolved solute lowers vapor pressure of solvent.  The partial pressure exerted by solvent vapor above an ideal solution is the product of the mole fraction of solvent in the solution and the vapor pressure of the pure solvent at a given temperature. P A =  A P°

61. Vapor Pressure of Solutions Raoult’s Law  P solution = X solvent (P solvent ) P - vapor pressure X - mole fraction X solute + X solvent = 1

62. Vapor pressure of solutions What evaporates faster pure water of sugar water? What is vapor pressure Roualts Law P sol = X solvent P° solvent  Equation of a straight line  The slope is the P° solvent  The y-intercept is zero July 2009General Chemistry: Chapter 11 62

63. For a Solution that Obeys Raoult's Law, a Plot fo P soln Versus X solvent, Give a Straight Line

64. Vapor Pressure of Solvent and Solution

65. An Aqueous Solution and Pure Water in a Closed Environment

66. July 2009General Chemistry: Chapter 11 Slide 66 of 46 Liquid-Vapor Equilibrium

67. Two Volatile Liquids Positive deviation Ideal Solution Negative deviation Positive deviation exists when experimental value is larger than calculated value, weaker solute solvent attraction; more evaporation. Negative deviation exists when experimental value is smaller than calculated value; stronger solvent solute attraction; less evaporation

68. July 2009General Chemistry: Chapter 11 Slide 68 of 46 Fractional Distillation

69. July 2009General Chemistry: Chapter 11 Slide 69 of 46 Fractional Distillation

70. Phase Diagram of Water Note: Negative slope

71. Phase Diagram of Water

72. Phase Diagram CO 2 Note: Posititve Slope

73. Phase Diagram CO 2

74. Special Properties of Water Hydrogen bonding found in water gives it special properties listed below: Surface tension Capillary Rise Viscosity

75. Surface Tension

76. Surface Tension units J/m 2

77. Surface Tension Adhesive forces stronger in red aqueous solution Cohesive forces stronger in mercury liquid

78. Capillary Rise

79. Viscosity

80. Terms Capillary action is the rise of a liquid up a narrow tube as a result of adhesive forces between the liquid and the tube and cohesive forces within the liquid. Viscosity is a measure of the resistance to flow of a fluid.

81. THE END

82. ChemTour: Lattice Energy Click to launch animation PCPC | MacMac Students learn to apply Coulomb’s law to calculate the exact lattice energies of ionic solids. Includes Practice Exercises.

83. ChemTour: Intermolecular Forces Click to launch animation PCPC | MacMac This ChemTour explores the different types of intermolecular forces and explains how these affect the boiling point, melting point, solubility, and miscibility of a substance. Includes Practice Exercises.

84. ChemTour: Henry’s Law Click to launch animation PCPC | MacMac Students learn to apply Henry’s law and calculate the concentration of a gas in solution under varying conditions of temperature and pressure. Includes interactive practice exercises.

85. ChemTour: Molecular Motion Click to launch animation PCPC | MacMac Students use an interactive graph to explore the relationship between kinetic energy and temperature. Includes Practice Exercises.

86. ChemTour: Raoult’s Law Click to launch animation PCPC | MacMac Students explore the connection between the vapor pressure of a solution and its concentration as a gas above the solution. Includes Practice Exercises.

87. ChemTour: Phase Diagrams Click to launch animation PCPC | MacMac Students use an interactive phase diagram and animated heating curve to explore how changes in temperature and pressure affect the physical state of a substance.

88. ChemTour: Capillary Action Click to launch animation PCPC | MacMac In this ChemTour, students learn that certain liquids will be drawn up a surface if the adhesive forces between the liquid on the surface of the tube exceed the cohesive forces between the liquid molecules.

89. ChemTour: Boiling and Freezing Points Click to launch animation PCPC | MacMac Students learn about colligative properties by exploring the relationship between solute concentration and the temperature at which a solution will undergo phase changes. Interactive exercises invite students to practice calculating the boiling and freezing points of different solutions.

90. ChemTour: Osmotic Pressure Click to launch animation PCPC | MacMac Students discover how a solute can build up pressure behind a semipermeable membrane. This tutorial also discusses the osmotic pressure equation and the van’t Hoff factor.

91. Solubility of CH4, CH2Cl2, and CCl4 Which of the following three compounds is most soluble in water? A) CH 4 (g) B) CH 2 Cl 2 (λ) C) CCl 4 (λ)

92. Solubility of CH4, CH2Cl2, and CCl4 Consider the following arguments for each answer and vote again: A.A gas is inherently easier to dissolve in a liquid than is another liquid, since its density is much lower. B.The polar molecule CH 2 Cl 2 can form stabilizing dipole-dipole interactions with the water molecules, corresponding to a decrease in ΔH° soln. C.The nonpolar molecule CCl 4 has the largest molecular mass, and so is most likely to partially disperse into the water, corresponding to an increase in ΔS° soln.

93. The End