1. Physical Properties of Solutions Honors Unit 10

2. Solutions in the World Around Us

3. Review from Unit 4: Solution: homogeneous mixture of two or more substances.  Solute: the substance being dissolved.  Solvent: the substance doing the dissolving The solute gets dissolved in the solvent!!

4. Colloid – solution of solid particles small enough to be suspended in solution  Colloids may look clear when dilute enough.  Examples: milk, paint, smoke

5. Tyndall effect – scattering of light by particles in a colloid or suspension ▫Causes the beam of light to become visible ▫Why you can see rays from the sun (particles in the air scatter the light)

6. Suspension – a mixture from which some of the particles settle out slowly upon standing ▫Particles are too big to be dissolved ▫Suspensions can be filtered! ▫Examples: sand in water, dust in air

7. Concentration Concentration of a solution = the quantity of a solute in a given quantity of solution (or solvent). ▫A concentrated solution contains a relatively large amount of solute vs. the solvent ▫A dilute solution contains a relatively small concentration of solute vs. the solvent “Concentrate” and “dilute” aren’t very quantitative!!

8. Need different concentration units to describe different properties of solutions:  Molarity (review)  Mass Percent  Molality Solution Concentration

9. Molarity (M) = moles of solute liters of solution ▫ Moles of solute dissolved in 1 liter of solution ▫ “M” is pronounced “molar” Example: 0.23 M NaCl solution = 0.23 moles of NaCl dissolved in 1 L of solution (water)

10. Example #1: What is the molarity of a solution made by dissolving 12.5 g of oxalic acid in 456 mL of solution? Example #1

11. How many grams of sodium carbonate are needed to prepare.250 L of an aqueous 0.300 M soln.? Example #2

12. Example #3 What is the % by mass of the solute in a solution made by adding 1.20 g of methyl alcohol to 16.8 g of H 2 O?

13. Why use Molality?? Molality (m) is the number of moles of solute per one kilogram of solvent (not solution!) Molality does NOT change with temperature! Molarity (M) varies with temperature due to the expansion or contraction in the volume of the solution 

14. Example #4 A solution contains 15.5 g of urea, NH 2 CONH 2, in 74.3 g of water. Calculate the molality of the urea.

15. Dilutions are used to decrease the concentration (or molarity) of a solution M 1 V 1 =M 2 V 2

16. Example #5 How would you prepare 0.250 L of 0.300 M Na 2 CO 3 starting with 1.33 M solution?

17. Example #6

18. Solubility Vocabulary  Miscible - two liquids that are soluble in each other (mix in all proportions)  Immiscible - liquids that are insoluble in each other (do not mix)

19. Solubility  Saturated Solution - contains the maximum amount of dissolved solute  Unsaturated Solution - contains less than the maximum amount of dissolved solute  Supersaturated Solution – contains more solute than can theoretically be dissolved at a given temperature  How is this possible???

20. Supersaturation A supersaturated solution is created when a warm, saturated solution is allowed to cool without the precipitation of the excess solute Testing for saturation: add crystal of a solid and watch for crystallization.

21. Supersaturated Sodium Acetate One application of a supersaturated solution is the sodium acetate “heat pack.” Sodium acetate has an ENDOthermic heat of solution.

22. Supersaturated Sodium Acetate Sodium acetate has an ENDOTHERMIC heat of solution. NaCH 3 CO 2 (s) + heat  Na + (aq) + CH 3 CO 2 - (aq) Therefore, formation of solid sodium acetate from its ions is EXOTHERMIC. Na + (aq) + CH 3 CO 2 - (aq)  NaCH 3 CO 2 (s) + heat

23. Solubility – the maximum amount of solute that will dissolve in a given quantity of solvent at a specific temperature and pressure to produce a saturated solution Units for solubility : grams of solute per 100 g solvent Example: At 20˚C, NaNO 3 has a solubility of 74 g/100 g H 2 O Solubility

24. Factors Which Influence Solubility Temperature Pressure (for gases) Surface Area (Impacts how QUICKLY a substance will dissolve)

25. “Like Dissolves Like” Solvents of a specific polarity or type will dissolve solute of similar polarities or types! Polar substances dissolve easily in water  Alcohols, CH 3 OH  Solubility of alcohols decreases as the molar mass of the alcohol increases Nonpolar substances have poor affinity for water  Petroleum  Hydrocarbons (pentane, C 5 H 12 )

26. Solubility Graphs or Curves The concentration of the solute in a saturated solution (the solubility) can be shown on a graph or curve called a “solubility curve.”

27. Solubility Curves Example #7: What mass of solute will dissolve in 100 g of water at the following temperatures. a)KNO 3 at 70°C b)NaCl at 100°C

28. Solubility Curves Example #8: At 20°C, if 100 grams of NaNO 3 are dissolved in 100 grams of water, is this solution saturated, unsaturated, or supersaturated?

29. Solubility Curves Example #9: Which term - saturated, unsaturated, or supersaturated – best describes: A solution that contains 70g of NaNO 3 per 50 g H 2 O at 30°C

30. Solubility Curves Example #10: What term - saturated, unsaturated, or supersaturated – best describes: A solution that contains 70 g of dissolved KCl per 200 g H 2 O at 80°C.

31. Solubility Curves Example #11: Determine the molality of a saturated NaCl solution at 25°C.

32. The Solubilities of Gases Most gases become less soluble in liquids as the temperature increases. Example: can of soda going flat on a hot day

33. Effect of Pressure  Henry’s Law - At low to moderate pressure, the concentration of a gas increases with the pressure  Solubility increases with increasing pressure  Pressure has a major effect on the solubility of a gas in a liquid, but little effect on other systems

34. Colligative Properties When adding a solute to a solvent, the properties of the solvent are modified.  Vapor pressure decreases  Melting point decreases  Boiling point increases These changes are called COLLIGATIVE PROPERTIES.

35. Colligative Properties  Colligative means “depending on the collection.”  Depends only on the number of dissolved particles, not on the identity of dissolved particles.  Examples of colligative properties:  Vapor pressure lowering, boiling point elevation, freezing point depression, and osmotic pressure

36.  Boiling occurs when vapor pressure equals atmospheric pressure.  The boiling point of a solution is higher than the boiling point of the pure solvent.  Dissolving substances increases (elevates) the boiling point of a solvent. Boiling Point Elevation  Ex.) Adding salt to water allows the water temp. to exceed 100°C, thereby cooking food faster

37. Elevation of Boiling Point Boiling Point Elevation Formula: ∆T b = k b mi k b = constant; depends on the solvent i = van’t Hoff factor = # of ions formed in soln. (1 for nonelectrolytes) m = molality ∆T b = change in temperature (number of degrees the boiling pt. goes UP)

38. Freezing Point Depression  The freezing point of a solution is lower than the freezing point of the pure solvent.  Dissolving substances lowers (depresses) the freezing point of a solvent.  Ex: Icy pavement - throw down CaCl 2 or NaCl, and the water will then freeze at a lower temperature

39. Antifreeze: Ethylene glycol/water soln. Uses: 1.Prevents car’s radiator from freezing in the winter. 2.Prevents car’s radiator from boiling over in the summer The more ethylene glycol in the water, the lower the freezing point, and the higher the boiling point.

40. Freezing Point Depression Formula: ∆T f = k f mi k f = constant; depends on the solvent i = van’t Hoff factor = # of ions formed in soln. (1 for nonelectrolytes) m = molality ∆T f = change in temperature (number of degrees the FP goes down) Depression (lowering) of Freezing Point

41. Colligative Properties of Electrolytes  Electrolytes = Soluble ionic compounds. When they dissolve in solution, they dissociate into their component ions and conduct electricity.  Ex.) NaCl (aq)  Na + (aq) + Cl - (aq)  Covalent molecules in aqueous solution:  Covalent particles do not dissociate when in solution, so the # of molecules = the # of particles.

42. Colligative Properties of Electrolytes Nonelectrolytes vs. electrolytes Nonelectrolytes produce only molecules in solution; electrolytes produce ions. NaCl  Na + + Cl – The greater the product of molality and number of ions, the greater the boiling point elevation or freezing point depression!

43. Example #12

44. Example #13 Rank the following aqueous solutions in order of lowest to highest melting point: (1) 0.010 m C 6 H 12 O 6 (3) 0.0080 m HCl (2) 0.0050 m MgCl 2 (4) 0.0040 m Al 2 (SO 4 ) 3