1. Chapter 8 Solutions A.Definition A solution is a homogeneous mixture of two or more different substances only one phase composition is variable almost always clear (transparent) A solution cannot be separated into its components by filtration Solute – dissolved substance (smaller amount) Solvent – dissolving medium (larger amount) Solutions can be in gas, liquid or solid phases 1

2. B.Types of solutions a) Solid – liquid 1. Ionic compound NaCl in water NaCl Na + + Cl - breaking of ionic bond H2OH2O Hydration Ion-dipole attraction 2

3. 2. Covalent compounds sugar in water 3

4. b. Liquid - liquid C O H and H 2 O H H H (CH 3 OH) Polar polar H O CH 3 H O H CH 3 Polar dissolves in polar Nonpolar dissolves in nopolar 4

5. c. Gas – liquid All gases are slightly soluble in water d. Gas-gas All mixtures of gases are solutions e.Solid – solid Cu (copper) in Au (gold) Alloy Brass (Cu/Zn) 5

6. C.Solubility Soluble – dissolves a large amount Insoluble – dissolves an negligible amount Miscible – both components are liquid and can dissolve (mix) in any proportion example: H 2 O and CH 3 OH a)Solubility - # of grams that can be dissolved in 100 g of solvent at saturation ( equilibrium) 6

7. Figure 8.3 In a saturated solution, the dissolved solute is in equilibrium with the undissolved solute. Example: 7

8. Table 8.1 Solubilities of Various Compounds in Water at 0 o C, 50 o C, and 100 o C. 8

9. Unsaturated solution – dissolved amount less than solubility Supersaturated solution - dissolved amount more than solubility Aqueous solution – solution in water 9

10. D. Factors affecting solubility 1. Nature of solute and solvent In general, for non-ionic solutes, “like dissolves like’. polar solutes dissolve in polar solvents. nonpolar solutes dissolve in nonpolar solvents. Ionic compounds do not dissolve in nonpolar solvents CH 4 in water? NaCl in octane (nonploar)? CH 4 OH in water? HCl in water? Examples 10

11. Solubility Guidelines for common Ionic Compounds in Water. 1. All compounds containing group IA (Li +, Na +, K +, etc.) and NH 4 + are soluble in water. 2.All nitrates (NO 3 -) and acetates (CH 3 COO - or C 2 H 3 O - ) are soluble in water. 3.All chlorides (Cl - ), bromides (Br - ) and iodides (I - ) are soluble in water except those of Ag +, Pb 2+ and Hg 2 2+. 4.All sulfates (SO 4 2- ) are soluble except PbSO 4, BaSO 4, SrSO 4 and CaSO 4 5.All hydroxides (OH - ) are insoluble except those of IA & Ca(OH) 2, Sr(OH) 2, Ba(OH) 2 6.Most other ionic compounds are insoluble in water 11

12. Li +, Na +, K + and NH 4 + soluble NO 3 -, CH 3 COO - soluble Cl -, Br -, I - soluble AgX, PbX 2, Hg 2 X 2. SO 4 2- soluble PbSO 4, BaSO 4,CaSO 4, SrSO 4 OH - insoluble IA, Ca 2+, Sr 2+, Ba 2+ Solubility except K 2 SO 4 NaCl PbCl 2 MgSO 4 Soluble in water? CaCO 3 Fe(NO 3 ) 2 NH 4 MnO 4 Soluble in water? BaSO 4 CuCr 2 O 7 12

13. Which of the following would be expected to be the most soluble in water? 13

14. 2. Temperature Generally, solubility increases as T increases, except for most gases in liquid where solubility decreases as T increases. Example: O 2 in water 14

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16. 3. Pressure (above the surface of the solution) Only affects solubility of gases in liquid. Solubility increase as P increases. 16

17. E. Concentrations of solutions - amount of solute in certain amount of solvent or solution. a) % by mass (mass-mass %) b) % by volume (V-V %) 17

18. Figure 8.7 When volumes of two different liquids are combined, the volumes are not additive. 18

19. Figure 8.8 Identical volumetric flasks are filled to the 50.0-mL mark with ethanol and with water. 19

20. c) mass-volume % d) parts per million (ppm) e) parts per billion (ppb) 20

21. If an aqueous solution is 2.5 % w/v in aluminum sulfate, Al 2 (SO 4 ) 3, how many grams of aluminum sulfate are there in a liter of solution? Examples How many grams of zinc fluoride, ZnF 2, are required to make a 5.00 % w/v aqueous solution in a 250 mL volumetric flask? 21

22. e) Molarity, M (molar concentration) # of moles of solute / liters of solution Example 1 0.50 mol KOH is dissolved in 2.0L of solution. Find the molarity. 1L NaCl 1 mole NaCl 1 M NaCl solution 22

23. Example 2 How many grams of NaCl is needed to make 250 mL of a 0.50 M NaCl solution? 23

24. Example 3. Make a 500mL of 0.250 M K 2 Cr 2 O 7 solution MW of K 2 Cr 2 O 7 = 294.2 24

25. F. Dilution or MV = n In general, C 1 V 1 = C 2 V 2, where C stands for concentration M1V1M1V1 = M 2 V 2 M2V2M2V2 M1V1M1V1 add water 25

26. M1V1M1V1 = M 2 V 2 Example 1. A 250 mL 2.0M NaOH solution is diluted to 1.0 L. What is the molarity of the final solution? Example 2. What volume of a 5.0M HCl solution would be needed to prepare 2.0 L of a 0.25M HCl solution? 26

27. G. Colligative properties of solutions A colligative property is a physical property of a solution that depends only on the concentration. a) Lowering of vapor pressure water Next day sugar in water sugar in water 27

28. water sugar in water b) Elevation of the Boiling point 28

29. c) Freezing point (depression) lowering Example: Antifreeze; salt-water. 29

30. H2OH2Osolution Semi-permeable membrane d) Osmosis and Osmotic pressure Osmosis: the flow of solvent through a semipermeable membrane from a dilute to a more concentrated solution. At equilibrium, the molecules move back and forth at equal rates. 30

31. H2OH2Osolution This measures the osmotic pressure of the solution The process is called osmosis Pressure applied to prevent osmosis = osmotic pressure 31

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33. V.P. lowering, B.P. elevation and F.P depression and osmotic pressure depend on the concentration of the solute particles but not on the type of solute. Examples: 33

34. Example: 1.5 mol of K 2 CrO 4 is dissolved in 1000 g of water. What is the freezing point of the solution? Freezing point (depression) lowering mol of particle = 34

35. Osmolarity NaCl Na + + Cl - 1M NaCl 2 mol ions osmolarity = 2M osmolarity = M x i i = # of particles produced from the dissociation of one formula unit of solute H2OH2O Osmolarity in cell = 0.31M (osmol) Osmolarity = 0.31M isotonic solution 5.0% m/v glucose 0.92% m/v NaCl (physiological saline solution) 35

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37. Osmolarity = 0.31M isotonic solution Osmolarity > 0.31M hypertonic solution Osmolarity < 0.31M hypotonic solution cell crenation Burst (hemolysis) 37

38. Red blood cell in a) isotonic solution b) Hypertonic solution c) Hypotonic solution crenation hemolysis 38

39. Example Consider the following solutions 1M sugar 2M NaCl 1.5M Na 2 SO 4 1M Ca(NO 3 ) 2 solution 39

40. H. Colloidal dispersions A colloidal dispersion is a mixture in which a material is dispersed rather than dissolved.................................. dispersed phase Dispersing medium 10 -7 – 10 -5 cm Cannot be filtered 40

41. Some proteins have size colloidal dispersion solution colloidal dispersion 41

42. I. Dialysis Dialyzing membrane - has pores large enough to allow some ions and small molecules to pass along with water and gases. Kidney cells, blood capillaries, intestinal walls, etc. function as dialyzing membranes. Dialysis - the movement of ions and small molecules (urea), including water (solvent), across a dialyzing membrane. Large molecules such as proteins cannot pass through a dialyzing membrane. 42

43. Figure 8.18 In dialysis, there is a net movement of ions from a region of higher concentration to a region of lower concentration. (a) Before dialysis. (b) After dialysis 43

44. Impurities (ions) can be removed from a solution by using a dialysis procedure. 44

45. Artificial kidney: a Hemodialysis Machine 45