1. Water and Aqueous Systems 1

2. The Water Molecule Bent Two lone electron pairs Polar molecule 2

3. Hydrogen Bonding: The intermolecular forces in which hydrogen that is covalently bonded to a very electronegative atom is also weakly bonded to an unshared electron pair of another atom (N, O, F). 3

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5. Hydrogen bonding is responsible for many of the unique properties of water such as: high surface tension having a low vapor pressure having a lower density in the solid form than in the liquid form. 5

6. Surface tension causes water to form nearly spherical droplets. 6

7. http://www.chemistryland.com/CHM107/Water/WaterBeadsOnPlantPlusWaterCharge.jpg 7

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9. Surface Tension – Surface tension is an inward force that causes the surface to behave as a skin. – The higher the intermolecular forces the greater the surface tension. 9

10. Surface Tension of Water http://quest.nasa.gov/space/teachers/microgravity/image/66.gif 10

11. Water – Mercury 11

12. Surfactants: Substances that interfere with the hydrogen bonding between molecules and reduce the surface tension. Cause spreading or wetting. Examples of surfactants are soaps. 12

13. Water’s Low Vapor Pressure Because of the hydrogen bonds holding the water molecules together, the molecules have a low tendency to break free from the surface into the vapor phase. 13

14. Water has a relatively high boiling point. Why? 14

15. Density of Water Ice has a lower density than liquid water. Maximum density of water happens at 4 O C. 15

16. Density of Water 16

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19. Ice- (Honeycomb Shape) 19

20. The Solution Process 20

21. Solution A homogeneous mixture. One phase. Stainless Steel (Fe, Cr, Ni) 21

22. Solute, Solvent Solute—the substance being dissolved. Example: When you dissolve CuCl 2 in water, CuCl 2 is the solute. Solvent- the substance that dissolves the solute. Example: water 22

23. Aqueous Solution A solution that has water as the solvent. Possible substances that can dissolve in water: – Ionic compounds – Polar covalent compounds 23

24. Solvation The surrounding of solute particles by solvent particles. 24

25. Dissociation of Ionic Compounds the process by which an ionic compound separates into its ions as it dissolves. 25

26. Dissociation of NaCl in Water 26

27. Dissociation of NaCl 27

28. Movie Clip- Dissociation of Salt in Water http://www.youtube.com/watch?v=EBfGcTAJF4o 28

29. “Like dissolves like” SolventSoluteIs Solution Likely? Polar Yes PolarNonpolarNo NonpolarPolarNo Nonpolar Yes 29

30. Polar ethanol molecule is dissolved by the polar water molecule. Ethanol remains intact. 30

31. Will s ugar dissolve in water? 31

32. Will petroleum dissolve in water? 32

33. Oil on water- 2 phases 33

34. Will ionic compounds conduct electric current when dissolved in water? Yes Why? 34

35. Electrolytes and Non-electrolytes Electrolytes: conduct an electric current when in the molten state or in aqueous solution. Ionic compounds Non-electrolytes do not conduct a current. Usually molecular compounds 35

36. Is sugar C 6 H 12 O 6 electrolyte? 36

37. Do all electrolytes conduct electricity to the same degree? Weak electrolytes: partially ionize in water and conduct electricity in solution poorly (ex. Ammonia) Strong electrolytes: fully ionize in water and conduct electricity in solution strongly(ex. NaCl). 37

38. Hydrate: A crystalline compound in which the ions are attached to one or more water molecules. 38

39. Example: CuSO 4 5H 2 O copper(II) sulfate pentahydrate 39

40. Prefixes for naming Hydrates mono-1 di-2 tri-3 tetra-4 penta-5 hexa-6 hepta-7 octa-8 nona-9 deca-10 40

41. Analyzing Hydrates Simulation click on the link below: http://www.chem.iastate.edu/group/Greenbo we/sections/projectfolder/flashfiles/stoichiom etry/empirical.html http://www.chem.iastate.edu/group/Greenbo we/sections/projectfolder/flashfiles/stoichiom etry/empirical.html 41

42. Problem Calculate the percent by mass of water in washing soda (Na 2 CO 3. 10 H 2 O) % mass of H 2 O = MM water x100 % MM Hydrate Answer: 62.9% 42

43. Efflorescent Hydrates Hydrates that have high vapor pressures compared to water. When the vapor pressure of the surrounding is lower than the vapor pressure of the hydrate, the hydrate will lose its water; it effloresces. 43

44. Hygroscopic Hydrates and Dessicants Hydrates that have a low vapor pressure compared to water. These hydrates can absorb water from the air. These can be used as dessicants (ex. CaSO 4 ). 44

45. Deliquescent Materials that absorb so much water that they will become wet (form solutions). Ex. NaOH. 45

46. Part II Heterogeneous Aqueous Systems 46

47. Colloids and Suspensions Heterogeneous Mixtures 47

48. Suspension A mixture whose particles are temporarily suspended in a medium, but eventually settle down. Particle size>100nm Ex: dust in air. 48

49. Colloid A mixture whose particles (of size ~1 to ~100nm) are dispersed through a continuous medium. (The word colloid means “glue-like”) Heterogeneous because there are distinct phases. Tyndall Effect: Scattering of light. 49

50. Tyndall Effect 50

51. Types of Colloids Aerosol: liquid or solid dispersed in gases (fog, smoke). Foam: gas in liquid (whipped cream). Emulsion: both substances are liquids (mayonnaise). Sol: solid in liquid (jelly) 51

52. SOLUTIONS Ch 16.1-16.4 Solubility, Concentration, Colligative Properties….

53. Colligative Properties Properties of SOLUTIONS that depend only on the number of solute particles and not on their identity. 53

54. Some Colligative Properties are: 1.Vapor pressure lowering 2.Boiling point elevation *Remember? High-altitude cooking directions? 3.Freezing/ Melting Point depression *Chemists use this to determine purity of synthesized drugs! *We use this to make ice cream & salt our roads during snow/ ice storms… 54

55. Solubility and Concentration 55

56. Part I Solubility: The ability of a solute to dissolve in a solvent. Expressed as g solute/100g H 2 O or g solute/100mL H 2 O 56

57. Soluble & Insoluble Soluble: a solute that has appreciable solubility. (ex. Sugar in water) Insoluble: a solute that has a low or negligible solubility (ex. Flour in water) 57

58. Miscible & Immiscible TERM used for 2 LIQUIDS: Miscible – 2 liquids dissolve in each other. (ex. Alcohol & Water) Immiscible – 2 liquids do NOT dissolve in each other (ex. Oil & water) 58

59. Solubility Curves 59

60. Solubility Curves 60

61. Types of Solutions Saturated contains the maximum amount of solute that can be dissolved at the given conditions of T and P (precipitate or excess, undissolved solute is often found at bottom). Unsaturated contains less than the saturated amount of solute dissolved. Supersaturated contains more than the saturated amount of solute dissolved by dissolving at higher temp and then cooling. (extremely unstable – a crystal or movement can precipitate out excess solute) 61

62. Precipitate: solute that comes out of solution. 62

63. A supersaturated solution crystallizes after a seed crystal is introduced. http://www.chem.ufl.edu/~itl/2045/change/C12F11.GIF 63

64. Rates of Solution Does every candy you eat take the same time to dissolve? 64

65. Factors affecting the rate of dissolving: Surface Area Stirring Temperature (average kinetic energy), (affects solid, liquid and gaseous solutes) Pressure (affects gaseous solutes) 65

66. Effect of temperature on solubility of gases. Example: Compare the amount of oxygen dissolved in the waters of the arctic ocean to the amount of oxygen dissolved in warm tropical waters. 66

67. Effect of temperature on solubility of MOST solids. Example: Can you dissolve more sugar in warm water or in cold water? 67

68. Solubility Curves 68

69. Effect of Pressure on solubility of solids. Negligible. 69

70. Effect of Pressure on solubility of gases. Henry’s Law: the amount of gas dissolved in a solution is directly proportional to the pressure of the gas above the solution. The higher the pressure the higher the solubility of the gas. P 1= P 2 P=pressure above liquid S 1 S 2 S=solubility in liquid 70

71. Effect of Pressure on solubility of gases. http://www.chem.ufl.edu/~itl/2045/lectures/lec_i.html 71

72. Example An unopened bottle of soda has a pressure of 5atm above the liquid, so the concentration of CO 2 in the soda is high. Compare that to a pressure of 1 atm above the liquid when the bottle has been opened. 72

73. Part II Concentration: A measure of the amount of solute dissolved in the solution. – Molarity (M) – Molality (m) – Mole Fraction (x) – Percent by Mass (%) 73

74. Part II Concentration 74

75. Molarity A unit of concentration of a solution expressed in moles solute per liter of solution. (Note: 1L = 1 dm 3 ) Molarity (M) = Moles of Solute Liters of Solution 75

76. Why Molarity ? http://www.chem.ucla.edu/~gchemlab/volumetric_soln_conc.jpg 76

77. Steps involved in the preparation of a standard aqueous solution 77

78. Process of making 500 mL of a 1.00 M acetic acid solution 78

79. Ex. 1 Molarity What is the molarity of a solution of 8g NaOH in 100mL of solution? Answer: 2M NaOH 79

80. Ex. 2 Molarity How many grams of NaOH are contained in 2L of a 3M NaOH solution? Answer: 240g NaOH 80

81. Making Dilutions Moles Solute=M 1 V 1 =M 2 V 2 M molarity V volume 81

82. Ex. 3 Dilutions How many milliliters of 2.00M MgSO 4 solution must be diluted with water to prepare 100.00 mL of 0.400M MgSO 4 ? Answer: 20.0mL 82

83. Percent by Mass % by mass= Mass of solute x 100 % Mass of solution 83

84. Percent by Volume % by volume= Volume of solute x 100 % Volume of solution 84

85. Part III Molality and Mole Fraction 85

86. Molality Moles of Solute per kilogram of Solvent Molality (m) = moles solute kg solvent 86

87. Mole Fraction x solute = Moles of solute Moles of solution x solvent = Moles of solvent Moles of solution x solute + x solvent = 1 87

88. Colligative Properties of Solutions 88

89. How do you get from this… 89

90. …to this? 90

91. Add an ionic compound! 91

92. Colligative Properties Properties of SOLUTIONS that depend only on the number of solute particles and not on their identity. 92

93. Some Colligative Properties are: Vapor pressure lowering Boiling point elevation *Remember? High-altitude cooking directions? Freezing/ Melting Point depression *Chemists use this to determine purity of synthesized drugs! *We use this to make ice cream & salt our roads during snow/ ice storms… 93

94. Vapor Pressure 94

95. Vapor Pressure Lowering The particles of solute are surrounded by and attracted to particles of solvent. Now the solvent particles have less kinetic energy and tend less to escape into the space above the liquid. So the vapor pressure is less. 95

96. Ionic vs Molecular Solutes Ionic solutes produce two or more ion particles in solution. They affect the colligative properties proportionately more than molecular solutes (that do not ionize). The effect is proportional to the number of particles in the solution. 96

97. How many particles do each of the following give upon solvation? NaCl CaCl 2 Glucose 97

98. Freezing Point Depression 98

99. Example Salt is added to melt ice by reducing the freezing point of water. 99

100. Boiling Point Elevation 100

101. Example Addition of ethylene glycol C 2 H 6 O 2 (antifreeze) to car radiators. 101

102. Freezing Point Depression and Boiling Point Elevation Boiling Point Elevation ∆T b =mk b (for water k b =0.51 o C/m) Freezing Point Depression ∆T f =mk f (for water k f =1.86 o C/m) Note: m is the molality of the particles, so if the solute is ionic, multiply by the #of particles it dissociates to. 102

103. Which is more effective for lowering the freezing point of water? NaCl or CaCl 2 103

104. Example 1: Find the new freezing point of 3m NaCl in water. 104

105. Example 2: Find the new boiling point of 3m NaCl in water. 105