2. 1 I.The Nature of Solutions p 118 REVIEW BOOK HW P 120 Q 1 TO 7 Solutions

3. 2 MARCH 29 DEFINITIONS – SOLUTION – SOLUTE – SOLVENT – HOMOGENEOUS MIXTURE CONCENTRATION – DILUTE VS CONCENTRATED THE NATURE OF SOLUTIONS SOLVATION SOLUBILITY FACTORS THAT AFFECT SOLUBILITY

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 APRIL 2 Characteristics of a liquid solution Types of solution Solubility and factors that affect it

11. 10 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.

12. 11 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).

13. 12 Solutions Homogeneous mixtures. Solvation is the process by which the solution forms.

14. 13 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.

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

22. 21 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

23. 22 When temperature increases… Solubility of a gas decreases Solubility of a solid increases

24. 23 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.

25. 24 Gases are more soluble at high pressures EX: nitrogen narcosis, the “bends,” soda

26. 25

27. 26 C. Solubility 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

28. 27 C. Solubility Solubility CurveSolubility Curve –shows the dependence of solubility on temperature

29. 28 C. Solubility SATURATED SOLUTION no more solute dissolves UNSATURATED SOLUTION more solute dissolves SUPERSATURATED SOLUTION becomes unstable, crystals form concentration

30. 29 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

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 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

34. 33 1.0 L of water was used to make 1.0 L of solution. Notice the water left over.

35. 34 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

36. 35 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?

37. 36 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

38. 37 Two Other Concentration Units grams solute grams solution % by mass = % by mass

39. 38 Two Other Concentration Units grams solute grams solution Ppm = Ppm = parts per million

40. 39 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

41. 40 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.

42. 41 DO NOW – REVIEW CONCENTRATIONS COLLIGATIVE PROPERTIES Hw Review Book P 129-131 q38 to 50

43. 42 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 %

44. 43 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

45. 44 Using mass % How many grams of NaCl are needed to prepare 250 g of a 10.0% (by mass) NaCl solution?

46. 45 Colligative Properties On adding a solute to a solvent, the properties of the solvent are modified. Vapor pressure decreasesVapor pressure decreases 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.

47. 46 Change in Freezing Point The freezing point of a solution is LOWER than that of the pure solvent Pure water Ethylene glycol/water solution

48. 47 Change in Freezing Point Common Applications of Freezing Point Depression Propylene glycol Ethylene glycol – deadly to small animals

49. 48 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

50. 49 Change in Boiling Point Common Applications of Boiling Point Elevation

51. 50 Calculate the Freezing Point of a 4.00 molal glycol/water solution. K f = 1.86 o C/molal (See K f table) Solution ∆T FP = (1.86 o C/molal)(4.00 m)(1) ∆T FP = 7.44 FP = 0 – 7.44 = -7.44 o C (because water normally freezes at 0) Freezing Point Depression

52. 51 At what temperature will a 5.4 molal solution of NaCl freeze? Solution ∆T FP = K f m i ∆T FP = (1.86 o C/molal) 5.4 m 2 ∆T FP = (1.86 o C/molal) 5.4 m 2 ∆T FP = 20.1 o C ∆T FP = 20.1 o C FP = 0 – 20.1 = -20.1 o C FP = 0 – 20.1 = -20.1 o C Freezing Point Depression

53. 52 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.