1. PROPERTIES OF SOLUTIONS

2. Solution A homogeneous mixture of two or more substances in a single phase. Does not have to involve liquids -- air is a solution of nitrogen, oxygen, carbon dioxide etc.; solder is a solution of lead, tin etc.

3. Solute Component in lesser concentration. “Dissolvee”

4. Solvent Component in greater concentration. “Dissolver”

5. Solubility Maximum amount of material that will dissolve in a given amount of solvent at a given temp. to produce a stable solution.

6. Saturated Solution A solution containing the maximum amount of solute that will dissolve under a given set of conditions.

7. Saturated solutions are at dynamic equilibrium with any excess undissolved solute present. Solute particles dissolve and recrystallize at equal rates.  This point is the same as solubility for that substance.

8. Unsaturated Solution A solution containing less than the maximum amount of solute that will dissolve under a given set of conditions. (more solute can dissolve)

9. Supersaturated Solution A solution that has been prepared at an elevated temperature and then slowly cooled. It contains more than the usual maximum amount of solution dissolved.

10. Miscible When two or more liquids mix. (example: Water and food coloring)

11. Immiscible When two or more liquids DON’T mix.--they usually layer if allowed to set for a while. (example: Water and oil)

12. UNITS OF SOLUTION CONCENTRATION

13. Molarity (M) # of moles of solute per liter of solution. IS temperature dependent.

14. The liquid solvent can expand and contract with changes in temperature. Thus, not a constant ratio of solute:solvent particles.

15. Most M solutions are made at 25  C so this point is subtle and picky!!

16. Mass Percent (weight percent) Percent by mass of the solute in the solution.

17. Mole Fraction (  ) Ratio of the number of moles of a given component to the total number of moles of solution. Mole fraction a =  a

18. Molality (m) # of moles of solute per kilogram of solvent. NOT temperature dependent.

19. Represents a ratio of solute:solvent molecules at all times.

20. Normality (N) Number of equivalents per liter of solution.

21. Equivalent (for an acid-base reaction) The mass of acid or base that can furnish or accept exactly one mole of protons (H + ). M (# H + or OH - ) = N

22. Equivalent (for a redox reaction) The mass of oxidizing or reducing agent that can accept or furnish one mole of electrons.

24. THE SOLUTION PROCESS

25. Energies Involved in Solution Formation When a solute is dissolved in a solvent, the attractive forces between solute and solvent particles are great enough to overcome the attractive forces within the pure solvent and solute.

26. The solute becomes solvated (usually by dipole-dipole or ion- dipole forces). When the solvent is water, the solute is hydrated.

27. Heat of Solution The Heat of Solution is the amount of heat energy absorbed (endothermic) or released (exothermic) when a specific amount of solute dissolves in a solvent.

28. Steps in Solution Formation  H 1 Step 1 -  H 1 Step 1 - Expanding the solute Separating the solute into individual components  H solute = -  H lattice energy

29. Steps in Solution Formation  H 2 Step 2 -  H 2 Step 2 - Expanding the solvent Overcoming intermolecular forces of the solvent molecules

30. Steps in Solution Formation  H 3 Step 3 -  H 3 Step 3 - Interaction of solute and solvent to form the solution  H soln =  H 1 +  H 2 +  H 3  H 2 +  H 3 =  H hyd  H hyd

31. Predicting Solution Formation

32. Solubility Chart

33. Some heats of solution are positive (endothermic). The reason that the solute dissolves is that the solution process greatly increases the entropy (disorder). This makes the process spontaneous.

34. The solution process involves two factors, the change in heat and the change in entropy, and their relative sizes determine whether a solute dissolves in a solvent.

35. Hot and Cold Packs These often consist of a heavy outer pouch containing water and a thin inner pouch containing a salt. A squeeze on the outer pouch breaks the inner pouch and the salt dissolves.

36. Some hot packs use anhydrous CaCl 2 (  H soln = -82.8 kJ/mol), whereas, many cold packs use NH 4 NO 3 (  H soln = 25.7 kJ/mol).

37. Other hot packs function on the principle of liquid to solid which is exothermic, while others contain iron filings and the process of rusting is sped up, thus, producing energy.

38. Factors Affecting Solubility Molecular Structure

39. Fat Soluble Vitamins A, D, E, & K - Nonpolar Can be stored in the body tissue such as fat.

40. Water Soluble Vitamins B & C - Polar Are not stored, must be consumed regularly.

41. Pressure Effects The solubility of a gas is higher with increased pressure. Pressure has very little effect on the solubility of liquids and solids.

42. Vapor Pressure Francois Marie Raoult William Henry

43. Carbonated beverages must be bottled at high pressures to ensure a high concentration of carbon dioxide in the liquid.

44. Henry’s Law The concentration of a dissolved gas in a solution is directly proportional to the pressure of the gas above the solution Applies most accurately for dilute solutions of gases that do not dissociate or react with the solvent Yes  CO 2, N 2, O 2 No  HCl, HI

45. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 45 Temperature Effects (for aqueous solutions) Dissolving of a solid usually occurs more rapidly at higher T, but the amt. of solid that can be dissolved may increase or decrease w/increasing T. Solubility of a gas decreases w/increasing T Causes thermal pollution

46. Solubility generally increases with temperature if the solution process is endothermic (  H soln > 0). Solubility generally decreases with temperature if the solution process is exothermic (  H soln < 0). What compounds will decrease in solubility when heated?

47. Figure 11.6: The solubilities of several solids as a function of temperature. Na 2 sO 4 and Ce 2 (SO 4 ) 3

48. Figure 11.7: The solubilities of several gases in water as a function of temperature at a constant pressure of 1 atm of gas above the solution.

49. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 49 Lake Nyos

50. Raoult’s Law The presence of a nonvolatile solute lowers the vapor pressure of the solvent. P solution P solution = Observed Vapor pressure of the solution P 0 solvent P 0 solvent = Vapor pressure of the pure solvent  solvent  solvent = Mole fraction of the solvent A nonvolatile solute has no tendency to escape from solution into the vapor phase.

51. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 51 Figure 11.9: An aqueous solution (H 2 SO 4 ) and pure water in a closed environment. (a) Initial stage. (b) After a period of time, the water is transferred to the solution. VP pure solvent > solution

52. Figure 11.10: The presence of a nonvolatile solute inhibits the escape of solvent molecules from the liquid and so lowers the vapor pressure of the solvent.

53. Non Ideal Solutions Liquid-liquid solutions in which both components are volatile Modified Raoult's Law: P 0 P 0 is the vapor pressure of the pure solvent P A P B P A and P B are the partial pressures

54. Raoult’s Law – Ideal Solution A solution that obeys Raoult’s Law is called an ideal solution

55. Negative Deviations from Raoult’s Law Strong solute-solvent interaction results in a vapor pressure lower than predicted Exothermic mixing = Negative deviation

56. Positive Deviations from Raoult’s Law Weak solute-solvent interaction results in a vapor pressure higher than predicted Endothermic mixing = Positive deviation

57. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 57

58. Colligative Properties of Solutions Jacobus Henricus van 't Hoff (1852-1911)

59. Colligative Properties Colligative properties are those that depend on the concentration of particles in a solution, not upon the identity of those properties.  Boiling Point Elevation  Freezing Point Depression  Osmotic Pressure

60. Freezing Point Depression Each mole of solute particles lowers the freezing point of 1 kilogram of water by 1.86 degrees Celsius. K f K f = 1.86  C  kilogram/mol m m = molality of the solution i van’t Hoff i = van’t Hoff factor

61. Boiling Point Elevation Each mole of solute particles raises the boiling point of 1 kilogram of water by 0.51 degrees Celsius. K b K b = 0.51  C  kilogram/mol m m = molality of the solution i van’t Hoff i = van’t Hoff factor

62. Freezing Point Depression and Boiling Point Elevation Constants,  C/ m

63. The van’t Hoff Factor, i Electrolytes may have two, three or more times the effect on boiling point, freezing point, and osmotic pressure, depending on its dissociation.

64. Dissociation Equations and the Determination of i NaCl(s)  AgNO 3 (s)  MgCl 2 (s)  Na 2 SO 4 (s)  AlCl 3 (s)  Na + (aq) + Cl - (aq) Ag + (aq) + NO 3 - (aq) Mg 2+ (aq) + 2 Cl - (aq) 2 Na + (aq) + SO 4 2- (aq) Al 3+ (aq) + 3 Cl - (aq) i = 2 i = 3 i = 4

65. Ideal vs. Real van’t Hoff Factor The ideal van’t Hoff Factor is only achieved in VERY DILUTE solution.

66. Osmotic Pressure The minimum pressure that stops the osmosis is equal to the osmotic pressure of the solution

67. Osmotic Pressure Calculations   = Osmotic pressure M M = Molarity of the solution R R = Gas Constant = 0.08206 L  atm/mol  K i i = van’t Hoff Factor

68. Suspensions and Colloids Suspensions and colloids are NOT solutions. Suspensions: The particles are so large that they settle out of the solvent if not constantly stirred. Colloids: The particles intermediate in size between those of a suspension and those of a solution.

69. Types of Colloids

70. The Tyndall Effect Colloids scatter light, making a beam visible. Solutions do not scatter light. Which glass contains a colloid? solution colloid