1. 1 Solutions There are TWO types of Mixtures: 1) Heterogeneous mixtures (have parts with different makeup – different phases usually) Many are SUSPENSIONS - mixtures of larger insoluble “specks” suspended in another substance Remember the yellow PbI 2 ? Lead Iodide precipitate at the molecular level. The yellow stuff is solid The water is liquid

2. 2 Characteristics of suspensions: -specks are “visible” to the naked eye (appears cloudy – called the “Tyndall effect”) -Specks will settle out -suspensions can be filtered Other examples : Fog, Dust in air, muddy water Tyndall effect: Can you see a beam of light in the left beaker?

3. 3 The droplets of water suspended in the air scatter the light Fog is a liquid suspended in the gaseous air!

4. 4 2. Homogeneous mixtures (same makeup throughout) Are known as Solutions  mixtures of individual particles (molecules, ions, etc.) of one substance dispersed in another substance Solutions are made of Two parts: the SOLUTE: The dissolved substance - usually a solid the SOLVENT: The dissolving substance -Usually a liquid; usually water – “aqueous” (aq) Examples: in Salt-water salt = solute, water = the solvent CuSO 4(aq) NaCl (aq)

5. 5 Notice how the Na + and Cl - ions are dispersed throughout the water? Salt water

6. 6 Homogeneous Characteristics: -the dissolved particles won't settle out -are clear and transparent (may be colored) -solutions can’t be separated by a filter The Solution runs through the filter suspension

7. 7 Homogeneous Characteristics: -dissolved particles won't settle out -Clear and transparent (may be colored) -Can’t be filtered Examples: Salt-water -NaCl (aq) seltzer -CO 2 gas dissolved in water Metal alloys- ex: brass (Cu and Zn mixture) Actually once the bubbles appear in soda it becomes a suspension.

8. 8 Try these: HDYK

9. 9

10. 10 HDYK

11. 11 HDYK

12. 12 The Dissolving process: The solvent (H 2 O) strikes solute particle and pulls it into the solution Each ion becomes surrounded by water molecules (“hydrated”) - - - - H--- O Na + O---H O---H Cl - H---O | | | | +H H+ H + + H Notice: The positive aligns to the negative and vice versa

13. 13 Recall that due to their shape, water molecules are polar?

14. 14 Try this:

15. 15 Electrolytes Electrolytes are ionic compounds (salts) that break into ions when dissolved in water – conduct electricity Ex: CaCl 2 (s)  Ca 2+ (aq) + 2Cl - (aq) Salt crystal breaks up into charged ions

16. 16 Electrolytes Non- electrolytes – molecules stay together when dissolved in water – don’t conduct Ex:sucrose: C 12 H 22 O 11 (s)  C 12 H 22 O 11 (aq) Molecules separate but remain neutral

17. 17 Electrolytes Weak electrolytes – few molecules break apart, most stay together- fair electrical conductors Ex: weak acids: acetic acid HC 2 H 3 O 2  H + + C 2 H 3 O - Some stay as neutral molecules while others split into ions

18. 18 Try these: HDYK

19. 19 Try this: Illustrate:

20. 20 SOLUBILITY Solubility is the ability of one substance to dissolve in another substance ex: “sugar is soluble in water; chalk is not” Sugar is soluble in water Lead iodide is not soluble (that’s why its precipitating) Oil and water are not miscible

21. 21 We say… "LIKES DISSOLVE LIKES“ -Polar solvents dissolve polar solutes and ionic compounds Polar: NH 3, HCl, NaCl (salts); dissolve easily in water Nonpolar: O 2 CO 2 (nonpolar) do not dissolve easily in water Solubility Factors

22. 22 We say… "LIKES DISSOLVE LIKES“ Solubility Factors -nonpolar solutes dissolve in nonpolar solvents (oily/organic) - most molecular compounds (organic, waxy) Ex. fat-soluble vitamins (dangerous!) - Miscible vs. immiscible liquids (ex. oil and water).

23. 23 Factors affecting solubility We say… "LIKES DISSOLVE LIKES“ Purple Iodine crystals were stirred into a beaker of oil and water. Which dissolved the iodine? Why? Oil is nonpolar Water is polar I 2 is nonpolar

24. 24 Why does sugar (C 12 H 22 O 11 ) dissolve so easily in water? Remember O bonded to H? It Forms Hydrogen bonds!

25. 25 Try these: HDYK

26. 26 Solubility factors (cont.) Temperature - Solids: greater temperature, greater solubility (example: washing clothes) - Gases: greater temperature, lesser solubility Example: air bubbles escaping from water as its heated

27. 27 Pressure - Gases only - greater pressure = greater solubility Example: Opening a soda bottle What happens to the gas pressure when the bottle is opened? What happens to the solubility of CO 2 gas in the water, when the bottle is opened?

28. 28 Be Careful! RATE OF SOLUTION speed of dissolving (is not the same as solubility) Factors: Size of Particles (surface area) / Stirring / Temperature / Amount already dissolved Dissolving a bullion cube

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31. 31 Try these solubility factor questions:

32. 32 16 23 31 42

33. 33 SOLUBILITY TABLES: (Table F): Table showing solubility of ionic compounds in water Soluble vs. Insoluble* ions and exceptions *Note: Ions, which are insoluble, may dissolve to a very small extent Used to predict products and extent of double replacement reactions These ions are soluble /except when bonded to these These ions are not soluble /except when bonded to these

34. 34 Which chlorides are least soluble? Write their formulas: Write the formula for any soluble carbonate compound:

35. 35 Try these: HDYK

36. 36 SOLUBILITY CURVES: graphs showing solubility of a substance over a range of temperatures. Points on the line represent greatest amount of solute which can be dissolved in each 100 mL of water at a particular temperature At 30 o C one can dissolve about 50 grams of KNO 3 in 100 grams of water At 100 o C one can dissolve about 50 grams 100 grams 150 grams 200 grams 260 grams of KNO 3

37. 37 SOLUBILITY CURVES: graphs showing solubility of a substance over a range of temperatures. Points on the line represent greatest amount of solute which can be dissolved in each 100 mL of water at a particular temperature Notice as solution temperature increases, Solubility increases (More solute can be dissolved)

38. 38 SOLUBILITY CURVES: graphs showing solubility of a substance over a range of temperatures. Points on the line represent greatest amount of solute which can be dissolved in each 100 mL of water at a particular temperature At 100 0 C 250 grams of KNO 3 Can dissolve in 100 g Of water What mass of KNO 3 will dissolve in 100 grams of water at 100 0 C? In 200 grams of water? Recall that 100 mL or water = 100 grams? 250 g KNO 3 = X g KNO 3 100 g H 2 O 200 g H 2 O Twice as much?

39. 39 Types of Solution Saturated Solution: Solution holding the maximum quantity of solute at a given temperature. At 100 0 C 250 grams of KNO 3 Is required to saturate 100 mL Of water At 100 0 C 100 grams of KNO 3 Dissolved is an Unsaturated solution Unsaturated solution: holding less than the Maximum quantity of solute at a given temperature (Points on the curve) (Points below the curve) It can still hold 150 grams more!

40. 40 Types of solutions Saturated unsaturated Super–saturated: Holding more than the maximum at a given temperature Supersaturated (above the curve)

41. 41 1)A Hot saturated solution cools (but nothing precipitates) 2) Adding a small crystal starts the precipitation (the excess comes out) 3) Solution is back to saturated Supersaturated solution

42. 42 1. How many grams of KNO 3 will dissolve in 100 grams of water at 30 0 C? 2. Is a solution with 100 grams of NaNO 3 Dissolved In 100 grams of water at 50 0 C Saturated? OR Unsaturated? 3. List one solute which is most likely a gas. How do you know? 4. 30 grams KCl is dissolved In 100 grams of water at 80 0 C. How much additional solute must be added to saturate the solution? 5. 100 grams of water saturated with NH 4 Cl at 90 0 C is cooled to 70 0 C. How much solute will precipitate out? Do these:

43. 43 Relative to how concentrated Concentrated -relatively large amount of solute in a given volume of water (high on table) Types of solutions Unsaturated solution of 130 grams Of KI in 100 g of water at 30 0 C More solute than water! Orange juice “concentrate A lot of orange in a little water!

44. 44 Relative to how concentrated Concentrated -relatively large amount of solute in a given volume of water (high on table) Vs. Dilute -relatively small amount of solute in a given volume of water (low on table) Types of solutions Saturated solution of 12 grams of KClO 3 In 100 g of water at 30 0 C

45. 45 What must be done to the concentrate to prepare the juice? Dilute it with water!

46. 46 Try these:

47. 47 Try these:

48. 48 In saturated solutions rate of dissolving = rate of recrystallization Solution Equilibrium They are at “Equilibrium”

49. 49 Expressing Concentration: Molarity – moles of solute in every 1 liter of solution Molarity = #moles solute # liters solution Moles per liter mol/L Ex: What is the molarity of a solution with 4 moles of solute dissolved in 2 liters of solution? Ex: How many moles of solute are dissolved in 0.5 liters of a 1 molar solution? 4 moles / 2 liters = 2 moles per liter or 2 M 0.5 liters x 1 mole/liter = 0.5 moles Or moles = Molarity X Liters

50. 50 Expressing Concentration

51. 51 1 mole dissolved in 500 mL of solution 2 moles dissolved in 1000 mL of solution 0.5 mol in 250 mL Which of the three is most concentrated? 1 0.500 2 1.00 0.5 0.250 In each case the ratio of moles of chemical to volume of solution is the same Each has 2 moles of chemical per 1000 mL (liter) of solution We say that each is “2 molar” concentration - abbreviated “2 M” This means 2 moles of solute dissolved per liter of solution.

52. 52 2 mole dissolved in 400 mL of solution 3 mole dissolved in 900 mL of solution 1 mole dissolved in 100 mL of solution 2 mol = 5 M 0.4 L 3 mol = 3.3 M 0.9 L 1 mol = 10 M 0.1 L Has the least chemical, but even less volume! Which of the three is most concentrated?

53. 53 1. What is the concentration of a solution with 0.25 moles of HCl dissolved in 250 mL of solution? 2. How many moles of HCl are dissolved in 250 mL of a 0.1 M solution? M = moles liters 0.25 moles= 0.250 L 1.0 Molar M = moles liters moles = M x liters 0.1 M = mol 0.250 L mol = 0.1 M x 0.250 L mol = 0.0250

54. 54 3. What mass of NaCl is dissolved in 0.3 liters of a 3 M solution? 4. What is the molarity of 29 grams of NaCl dissolved in 500 mL of solution? (Find moles of solute first from formula, then convert to mass) (Find moles of solute first, then calculate molarity) Try these on your own:

55. 55 Try these Regents problems: Setup

56. 56 setup

57. 57 HDYK

58. 58 Parts Per Million (PPM) - grams of solute per 1,000,000 grams of solution - for substances which have very low solubility or very toxic substances ex: oxygen 8 PPM at 25 0 C – 8 milligrams dissolved in 1 liter of water (8 grams per 1,000,000 grams or 0.008 grams per 1000 g ) 1. What is the concentration in PPM of a solution with 0.36 grams of solute dissolved in 2000 grams of solution? PPM = mass of (part) solute X 1,000,000 mass of (whole) solution PPM = 0.36g solute x 10 6 2000. g solution = 0.000 18 x 10 6 = 180 PPM

59. 59 Try these: setup

60. 60 Colligative Properties Depend ONLY on the concentration of dissolved particles: (not on the identity) Adding solute raises the Boiling point but lowers the Freezing point (Adding solutes to water changes water’s properties) Salting roads lowers water’s freezing point causing ice to melt (Water’s freezing point decreases until it starts to melt)

61. 61 Colligative Properties Depend ONLY on the concentration of dissolved particles: (not on the identity) Adding solute raises the Boiling point but lowers the Freezing point (Adding solutes to water changes water’s properties) Why does pasta cook faster in boiling salt water? (The boiling point increased to the water is now hotter than normal)

62. 62 Colligative Properties Depend ONLY on the concentration of dissolved particles: (not on the identity) Adding solute raises the Boiling point but lowers the Freezing point (Adding solutes to water changes water’s properties) How does anti-”freeze” keep the water in your car’s radiator from boiling away (“over heating”) in hot weather? Antifreeze is also anti-”boil” Called “ethylene glycol”

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64. 64 Solute takes up space in the solution making it more difficult for water to escape as vapor. Vapor pressure decreases. Water has to get even hotter to get back up to atmospheric pressure!

65. 65 Colligative Properties Also: Electrolytes (ionic) have greater effect than non-electrolytes (molecular) ex: 1 mole of sucrose dissolves  1 mole molecules 1 mole of NaCl dissolves  1 mole Na + 1 mole Cl = 2 moles total ions (Adding solutes to water changes water’s properties)

66. 66 1 CaCl 2  1Ca 2+ + 2Cl - 1 NaCl  1Na + + 1Cl - Which salt is more effective?

67. 67 Which has a higher boiling point? Circle one: a. 10 g NaCl in 1000 grams of water or 20 g in 1000 grams of water? b. 1 mole of NaCl or 1 mole of C 12 H 22 O 11 ? c. 10 grams of NaCl in 1000 g water or 10 grams in 500 grams of water?

68. 68 Try these: HDYK

69. 69 HDYK