2. 1 I.The Nature of Solutions Review Book Unit 7 Solutions HW P 120 QUESTIONS 1 TO 12

3. 2 DEFINITIONS – SOLUTION – SOLUTE – SOLVENT – HOMOGENEOUS MIXTURE CONCENTRATION – DILUTE VS CONCENTRATED THE NATURE OF SOLUTIONS SOLVATION

4. 3 Some Definitions A solution is a HOMOGENEOUS mixture of 2 or more substances in a single phase. One constituent is usually regarded as the SOLVENT and the others as SOLUTES.

5. 4 A. Definitions Solution -Solution - homogeneous mixture Solvent Solvent - present in greater amount Solute Solute - substance being dissolved

6. 5 Parts of a Solution SOLUTE – the part of a solution that is being dissolved (usually the lesser amount). Uniformly spread in the solvent SOLVENT – the part of a solution that dissolves the solute (usually the greater amount) Solute + Solvent = Solution

7. 6 What happens when a solute dissolves in a solvent? Solvation –Solvation – the process of dissolving solute particles are separated and pulled into solution solute particles are surrounded by solvent particles

8. 7 How Does a Solution Form? As a solution forms, the solvent pulls solute particles apart and surrounds, or solvates, them.

9. 8 B. Solvation DissociationDissociation –separation of an ionic solid into aqueous ions –Attractions between H 2 O and ions : molecule ion attractions NaCl(s)  Na + (aq) + Cl – (aq)

10. 9 CHARACTERISTICS OF A LIQUID SOLUTION 1.- Homogeneous mixtures, particles are evenly spread. 2.- Dissolved particles are too small to be seen, therefore solutions are clear and do not disperse light. 3.- Can not be separated by filtration. Dissolved particles are too small and will pass trough any filter. 4.- Stable. Dissolved particles will not come out of the solution and will not settle.

11. 10 Solutions Homogeneous mixtures. Solvation is the process by which the solution forms.

12. 11 April 7 In your notebook answer the following question Is air a solution? Explain your reasoning

13. 12 Solutions are not always liquids Solutions are homogeneous mixtures of two or more pure substances. In a solution, the solute (present in smaller amount) is dispersed uniformly throughout the solvent (present in largest amount).

14. 13 SOLUBILITY FACTORS THAT AFFECT SOLUBILITY NATURE OF SUBSTANCES TEMPERATURE PRESSURE

15. 14 Solubility A measure of how much solute can be dissolved in an amount of solvent at a given temperature.

16. 15 A substance can be… Soluble in a solvent. Example: sugar is soluble in water. Miscible is the term used when the two components are liquids and they dissolve in one another. Example: alcohol and water are miscible Insoluble in a solvent Example: sand is insoluble in water. Immiscible is the term used when the two components are liquids and they do not mix. Example: oil and water are immiscible

17. 16 What affects Solubility? 1. Nature of Solute Temperature 2. Temperature 3. Pressure * graph

18. 17 Nature of Solute A polar solute molecule (alcohol) dissolves in a polar solvent (water). A nonpolar solute (oil paint) dissolves in a nonpolar solvent (turpentine) “ Like Dissolves Like”

19. 18 Solubility for ionic compounds Table F This table is used to predict if a double replacement reaction will occur. If it the reaction produces an insoluble compound it occurs. If the products of the reaction are filtered the insoluble compound will remain in the filter paper

20. 19 Table F

21. 20 Pb(NO 3 ) 2 + 2KI  PbI 2 + 2KNO 3 NaCl + AgNO 3  AgCl + NaNO 3 CuSO 4 + Na 2 CO 3  Na 2 SO 4 + CuCO 3

22. 21 Questions Is NaCl soluble? Yes! Is AgBr soluble? No!

23. 22 TABLE F SOLUBILITY GUIDELINES FOR AQ SOL ANSWERS 1.4 2.1 3.3 4.4 5.3 6.4 7.4 8.3 9.4 10. 1 11. Hydroxide ion B) NaOH 12) a Yes B) Ba(OH)2 or Sr(OH)2 13 soluble

24. 23 When temperature increases… Solubility of a gas decreases Solubility of a solid increases

25. 24 3 Pressure Makes gas more soluble  ex. Soda can Has almost no effect on liquids and solids -High pressure forces carbon dioxide into water to make soda. - When you open the cap, there is less pressure on the soda b/c the soda fizzes and gas escapes.

26. 25 Gases are more soluble at high pressures EX: nitrogen narcosis, the “bends,” soda

27. 26

28. 27 Solubility Curves SolubilitySolubility –maximum grams of solute that will dissolve in 100 g of solvent at a given temperature –varies with temp –based on a saturated solution

29. 28 C. Solubility Solubility CurveSolubility Curve –shows the dependence of solubility on temperature

30. 29 C. Solubility SATURATED SOLUTION no more solute dissolves UNSATURATED SOLUTION more solute dissolves SUPERSATURATED SOLUTION becomes unstable, crystals form concentration

31. 30 Definitions Solutions can be classified as saturated or unsaturated. A saturated solution contains the maximum quantity of solute that dissolves at that temperature. An unsaturated solution contains less than the maximum amount of solute that can dissolve at a particular temperature

32. 31 Supersaturated Sodium Acetate Supersaturated Sodium Acetate One application of a supersaturated solution is the sodium acetate “heat pack.”One application of a supersaturated solution is the sodium acetate “heat pack.”

33. 32 C. Solubility Solids are more soluble at...Solids are more soluble at... –high temperatures. Gases are more soluble at...Gases are more soluble at... –low temperatures & –high pressures (Henry’s Law). –EX: nitrogen narcosis, the “bends,” soda

34. 33 Determining Electrical Conductivity When a solution is soluble, it has ions that can conduct electricity (electrolytes) Ex. NaCl When a solution is insoluble, it cannot conduct electricity (non-electrolytes or poor electrolytes) When a solution is insoluble, it cannot conduct electricity (non-electrolytes or poor electrolytes) Ex. AgBr

35. 34 Solubility curves worksheet answers 1 KI 2 KClO3 3 SO2 4 134 g 5 SO2, NH3 and HCl 6 Sol at 50 C 115 g at 10 C 80 g difference 35 g 7 ~ 47 C 8 KNO3 and NaNO3 9 NaCl 10 KNO3 11 ~ 57.5 g 12 ~ 46 g 13 38 C 14 114g will precipitate 15 60 g 16 12 g

36. 35 Set 1 Solubility curves 1.2 2.1 3.1(turn page around) 4.2 5.4(turn page again) 6.4 7.2 8.3 9.1 10. 6 to 8 g 11.a- As P decreases, solubility decreases too. b- As T increases, solubility decreases 12 a KNO3

37. 36 Ways of expressing concentration MOLARITY % BY MASS, BY VOLUME PPM

38. 37 CONCENTRATION The amount of solute in the solution. Relative terms Diluted: Small amount of solute in relation to the amount of solvent Concentrated: Large amount of solute in relation with the solvent.

39. 38 Concentration of Solute The amount of solute in a solution is given by its concentration The amount of solute in a solution is given by its concentration. Molarity (M) = moles solute liters of solution

40. 39 Example If you have 50 moles (mol) of solute (salt) in 25 liters (L) of solution and you want to know the molarity (concentration), look at Table T for the molarity formula. Molarity = moles of solute / liters of solution 50 moles / 25 liters = 2M

41. 40 1.0 L of water was used to make 1.0 L of solution. Notice the water left over.

42. 41 Preparing Solutions Weigh out a solid solute and dissolve in a given quantity of solvent.Weigh out a solid solute and dissolve in a given quantity of solvent. Dilute a concentrated solution to give one that is less concentrated.Dilute a concentrated solution to give one that is less concentrated.

43. 42 PROBLEM: Dissolve 5.00 g of NiCl 2 6 H 2 O in enough water to make 250 mL of solution. Calculate the Molarity. Step 1: Calculate moles of NiCl 2 6H 2 O Step 2: Calculate Molarity NiCl 2 6 H 2 O [NiCl 2 6 H 2 O ] = 0.0841 M

44. 43 Step 1: Change mL to L. 250 mL * 1L/1000mL = 0.250 L Step 2: Calculate. Moles = (0.0500 mol/L) (0.250 L) = 0.0125 moles Step 3: Convert moles to grams. (0.0125 mol)(90.00 g/mol) = 1.13 g USING MOLARITY moles = MV What mass of oxalic acid, H 2 C 2 O 4, is required to make 250. mL of a 0.0500 M solution?

45. 44 Learning Check How many grams of NaOH are required to prepare 400. mL of 3.0 M NaOH solution? 1)12 g 2)48 g 3) 300 g

46. 45 Calculating Concentrations Dissolve 62.1 g (1.00 mol) of ethylene glycol in 250. g of H 2 O. Calculate molality and % by mass of ethylene glycol.

47. 46 Try this molality problem 25.0 g of NaCl is dissolved in 5000. mL of water. Find the molality (m) of the resulting solution. m = mol solute / kg solvent 25 g NaCl 1 mol NaCl 58.5 g NaCl = 0.427 mol NaCl Since the density of water is 1 g/mL, 5000 mL = 5000 g, which is 5 kg 0.427 mol NaCl 5 kg water = 0.0854 m salt water

48. 47 DO NOW – REVIEW MOLARITY % BY MASS, % BY VOLUME, ppm COLLIGATIVE PROPERTIES HW SOLUTIONS TAKE HOME TEST!!!! DUE MONDAY APRIL 25

49. 48 Two Other Concentration Units grams solute grams solution % by mass = % by mass X 100

50. 49 Two Other Concentration Units grams solute grams solution Ppm = Ppm = parts per million X 1000,000

51. 50 Calculating Concentrations Dissolve 62.1 g (1.00 mol) of ethylene glycol in 250. g of H 2 O. Calculate m & % of ethylene glycol (by mass). Calculate weight %

52. 51 Learning Check A solution contains 15 g Na 2 CO 3 and 235 g of H 2 O? What is the mass % of the solution? 1) 15% Na 2 CO 3 2) 6.4% Na 2 CO 3 3) 6.0% Na 2 CO 3

53. 52 Using mass % How many grams of NaCl are needed to prepare 250 g of a 10.0% (by mass) NaCl solution?

54. 53 April 30 Why is salt added to the roads in a snow day?

55. 54 Colligative Properties On adding a solute to a solvent, the properties of the solvent are modified. Melting point decreasesMelting point decreases Boiling point increasesBoiling point increases These changes are called COLLIGATIVE PROPERTIES. They depend only on the NUMBER of solute particles relative to solvent particles, not on the KIND of solute particles.

56. 55 Freezing point depression and boiling point elevation One mole of particles dissolved in a 1000g of water lowers the freezing point of the water by 1.86 C and increases the boiling point of water by 0.52 C. Note that electrolytes (ionic substances) produce a greater effect than non electrolytes because they produce more particles in solution.

57. 56 Change in Freezing Point The freezing point of a solution is LOWER than that of the pure solvent Pure water Ethylene glycol/water solution

58. 57 Change in Freezing Point Common Applications of Freezing Point Depression Propylene glycol Ethylene glycol – deadly to small animals

59. 58 Common Applications of Freezing Point Depression Which would you use for the streets of New York to lower the freezing point of ice and why? Would the temperature make any difference in your decision? a)sand, SiO 2 b)Rock salt, NaCl c)Ice Melt, CaCl 2 Change in Freezing Point

60. 59 Change in Boiling Point Common Applications of Boiling Point Elevation

61. 60 Calculate the Freezing Point of a solution containing 4.00 mol of glycol in a 1000 g of water K f = 1.86 o C/mol Solution Change in FP= (1.86 o C/mol)(4.00 mol) ∆T FP = 7.44 FP = 0 – 7.44 = -7.44 o C (because water normally freezes at 0) If NaCl is used instead the fp will be lower because the number of particles double. Freezing Point Depression

62. 61 At what temperature will a solution with 4 mol of NaCl in a 1000 g H2O freeze? Solution NaCl  Na + Cl ∆T FP = (1.86 o C/molal) 8 mol ∆T FP = (1.86 o C/molal) 8 mol ∆T FP = 14.88 o C ∆T FP = 14.88 o C FP = 0 – 14.88 = -14.88 o C FP = 0 – 14.88 = -14.88 o C Freezing Point Depression

63. 62 ANSWERS TO P 123 13.2 14.4 15.3 16.1 17.1 18.4 19.4 20.3 21.2 22.1 23.2