1. Chapter 12 Solutions

2. Objectives Describe the components of a solution Explain why water is a good solvent Explain the expression “Likes Dissolve Likes” Define solubility Compare solubility of solids and gases in solution

3. Solutions Homogeneous mixture of two or more substances –Homogenous – Uniform Two components of solutions –Solute – Substance being dissolved Present in lower amount –Solvent – Substance that dissolves solute Present in larger amount

4. Types of Solutions Solid dissolved in a liquid –Sugar in water, sodium chloride in water Liquid dissolved in a liquid –Ethanol in water (wine) Liquids are Miscible Gas dissolved in a liquid –Oxygen in water (Lake water) Gas dissolved in a gas –Oxygen in nitrogen (Air)

5. What is the solute? What is the solvent?

6. Types of Solutions Solid dissolved in a solid - Alloy –Copper in Silver (Sterling Silver)

7. Aqueous Solutions Solutions where water is the solvent Very common type of solution Water is an excellent solvent –Dissolves a number of solutes –Often referred to as a “universal solvent”

8. Water as a Solvent Water is an excellent solvent because of its shape –Water is a bent molecule –Very Polar Oxygen end of molecule is very negative Hydrogen end is very positive

10. Water as a Solvent These ends of the molecule seek out opposite charges to rip apart solutes Process called hydration

14. Water as a Solvent This is how water dissolves all ionic compounds that are soluble –Creates ions in solution Process is slightly different, but similar, for polar covalent compounds –Water just dissolves groups of polar molecules into individual polar molecules

16. “Likes Dissolve Likes” This saying tells us what kind of solute a given solvent will dissolve Polar solvents dissolve –Polar and ionic solutes Alcohols, Sugars, Salts Nonpolar solvents dissolve –Nonpolar solutes Oils, Plastics

17. Polarity Demo

18. Solubility The maximum amount of solute that can dissolve in a given amount of solvent at a given temperature –Usually in units of grams solute / 100 grams solvent Solubility information exists for solids, liquids, and gases Solubility changes with temperature

19. Solubility of Sucrose vs. Temp

20. Think About This What are some solids that are very soluble in water?

22. Types of Solutions Unsaturated Solution –A solution that has less solute than the maximum solubility Saturated Solution –A solution that has reached solubility Supersaturated Solution –A solution that exceeds solubility

23. Solubility Demo

24. Solubility of Gases Gases can dissolve in liquids –Oxygen or Carbon Dioxide in water Depends on Pressure and Temperature –High pressure and low temp increases solubility of gases High pressure forces more gas into the liq. Low temp. slows down particles so fewer gas particles escape the surface.

28. Implications of Gas Solubility Habitat for specific fish species –Trout live in cold streams Water is cold so more oxygen dissolves –Bullheads live in prairie potholes Water is warm so less oxygen dissolves

29. Implications of Gas Solubility From: http://tellus.ssec.wisc.edu/outreach/teach/ideas/kotoski/Minifact_Sheets/Minifact4 _Dissolved_Oxygen.pdf

30. Implications of Gas Solubility Scuba Diving –Scuba Divers cant rise to the surface quickly –Pressure in deep water is high –More gas dissolves in their blood –A quick rise will cause gas to bubble out –Causes the Bends

31. Homework p.479 37,38,39,40,43,86,87,93

32. Objectives Explain molarity Use molarity in calculations Describe how to prepare a solution of a given concentration Describe the process of dilution Calculate dilution problems

33. Concentration Concentration refers to the amount of solute in a given amount of solvent –General Terms Dilute Concentrated –Not very specific

34. A Bad Time 4 General Concentrations “Let’s give Mr. Hermansen a dilute solution of potassium chloride and see how he does.” What kind of solution?!?!?!?

35. Molarity Moles of solute per liter of solution –Molarity = Moles/Liters Units of mol/L Also (M) Can be said “Molar” Ex. 3.0 mol/L = 3.0M = Three Molar

36. Molarity Most common way of expressing concentration in chemistry b/c involves moles

37. Simple Examples How many moles of NaCl are required to prepare 1.0L of a 2.0 M solution –2.0 mol How many moles of NaCl are required to prepare 500mL of a 2.0 M solution –1.0 mol What volume of a 2.0 M solution of NaCl contains.20 mol NaCl –0.10L

39. Example Calculate the molarity of a solution made by dissolving 11.5 g NaOH in 1.5L of solution.

40. Example Calculate the molarity when 1.56 g of HCl is dissolved in 26.8 mL of solution.

41. Example What mass of potassium dichromate is required to prepare 200.0 mL of a.200M solution?

42. Preparation of Solutions Determine the mass of solute necessary Mass solute into a weight boat Transfer solute to volumetric flask of proper volume –Rinse weigh boat Add water to flask and dissolve solute Fill to line Insert stopper Invert to stir

44. Lets Make a Solution 250.00mL of a 0.018M solution of copper (II) nitrate

45. Dilution Process of making a solution less concentrated by adding solvent Dilution equation C 1 V 1 = C 2 V 2 C is concentration, V is volume Concentration times volume = moles When a sample of solution is diluted the number of moles does not change

47. Example Explain how would you prepare 300.0 mL of 1.5 M HCl from 6.0 M HCl

48. Lets Dilute Our Solution Explain how to prepare 100.00mL of 0.009 M Copper (II) Nitrate from 0.018 M stock solution.

49. Better Concentration Information “Let’s give Mr. Hermansen a 0.10M solution of potassium chloride and see how he does.” Thanks Doc..10M

50. Homework p. 482 #’s 109,110,111,114,115,117

51. Solution Preparation Sodium Hydrogen Carbonate 1. 0.025M 2. 0.018M 3. 0.036M 4. 0.0095M 5. 0.013M 6. 0.0082M 7. 0.010M 8. 0.020M

52. Objectives Complete stoichiometry calculations involving solutions

53. Solution Stoichiometry Stoichiometry with solutions is VERY similar to other types of stoichiometry. Need to find moles based on volume and concentration –Concentration times Volume = Moles General Process Volume * Concentration * Mole Ratio * Molar Mass

54. Example What mass of Lead (II) Sulfate is produced when 2.00L of 0.025M of Sodium Sulfate is mixed with an excess of Lead Nitrate?

55. Example What volume (mL) of 0.100M HCl is required to neutralize 25.0 mL of 0.350M Sodium Hydroxide?

56. Example What mass of Silver Chloride is produced when 23.0 mL of 0.20M Sodium Chloride is mixed with 50.0 mL of 0.11M Silver Nitrate?

57. Homework p. 482 #’s 128,130ab,132

58. Demo Start water boiling!

59. Objectives Explain what is meant by freezing and boiling point Describe colligative properties Calculate changes in boiling and freezing points

60. Boiling Point Temperature at which the vapor pressure of the liquid is equal to the atmospheric pressure –Water boils at 100ºC At this temperature the vapor pressure of water is 760mmHg –Ethanol boils at 78.4ºC At this temperature the vapor pressure of ethanol is 760mmHg –Ethanol is more volatile than water

61. Demo Boiling water with a vacuum

63. When Liquids Boil Q)What happens to the temperature of a liquid when it is boiling? A)The temperature stays the same - The heat energy is used to overcome the intermolecular forces present in the liquid.

65. Freezing Point Temperature at which a solid and liquid have the same vapor pressure. –Water freezes at 0ºC –Ethanol freezes at -115ºC! Freezing Pt. and Melting Pt. are the same Q)Is freezing exothermic or endothermic A)Exothermic - Energy must be lost to freeze

66. Colligative Properties Properties that depend on the number of solute particles present but NOT their identity. –Only the number of dissolved particles present matters Properties of solvents change when a solute is added –Solutions have different properties than the solvent that is present

67. Which Bottle Contains More Dissolved Particles? 0.10M KI 0.10M CH 3 OH 0.10M Na 2 S

68. Colligative Properties When a solute is added to a solvent the vapor pressure decreases What will this do to the boiling point? –The boiling point increases Boiling Point Elevation –The freezing point decreases Freezing Point Depression –Can you think of any examples? Salt on Roads, Antifreeze

69. Boiling Point Elevation ΔT b = Change in the boiling point Kb = Boiling point elevation constant Units of (ºC*kg/mol) –Constant for a given substance Data found in table 12.2 on page 474 –If ΔT b for salt water is 2.2ºC then the new B.P. is 102.2

70. The van’t Hoff Factor i = The van’t Hoff factor –Tells you how many moles of particles a solute would form if 1 mole dissolves –All covalent compounds have an i of 1 They do not ionize –i for ionic compounds equals the # of ions NaCl = 2 MgCl 2 = 3 AlCl 3 = 4

71. Freezing Point Depression ΔT f = Change in the freezing point K f = Freezing point depression constant Units of (ºC*kg/mol) –If ΔT f for salt water is 2.2ºC then the new F.P. is -2.2ºC

72. Example What is the freezing point of water if 0.25 mol. of glucose is dissolved in 500mL of water? K f =1.86 ºC*kg/mol?

73. Example What is the new boiling point of water when 100. grams of NaCl is dissolved in 2.0L of water? K b =0.51 ºC*kg/mol

74. Example Ethylene Glycol (C 2 H 6 O 2 62.08 g/mol) is placed in radiators to prevent the water from boiling. How many moles of EG must be added to 10.0L of water to make the water boil at 110.00 ºC? K b =0.51 ºC*kg/mol

75. Homework p. 483 #’s 135,137,143,144,145,148