1. Prentice Hall ©2004 CHAPTER 11 SOLUTIONS AND THEIR PROPERTIES Chapter 11Slide 1

2. Prentice Hall ©2004 Chapter 11Slide 2 Saturated: Contains the maximum amount of solute that will dissolve in a given solvent. Unsaturated: Contains less solute than a solvent has the capacity to dissolve. Supersaturated: Contains more solute than would be present in a saturated solution. Crystallization: The process in which dissolved solute comes out of the solution and forms crystals. Solution Formation01 Saturated: Contains the maximum amount of solute that will dissolve in a given solvent. Unsaturated: Contains less solute than a solvent has the capacity to dissolve. Supersaturated: Contains more solute than would be present in a saturated solution. Crystallization: The process in which dissolved solute comes out of the solution and forms crystals.

3. Prentice Hall ©2004 Chapter 11Slide 3 Solution Formation02

4. Prentice Hall ©2004 Chapter 11Slide 4 Solution Formation02

5. Prentice Hall ©2004 Chapter 11Slide 5 Solution Formation03

6. Prentice Hall ©2004 Chapter 11Slide 6 Solution Formation04 Exothermic ∆H soln : The solute–solvent interactions are stronger than solute– solute or solvent– solvent. Favorable process.

7. Prentice Hall ©2004 Chapter 11Slide 7 Solution Formation05 Endothermic ∆H soln : The solute–solvent interactions are weaker than solute–solute or solvent–solvent. Unfavorable process.

8. Prentice Hall ©2004 Chapter 11Slide 8 Solution Formation06 Solubility: A measure of how much solute will dissolve in a solvent at a specific temperature. Miscible: Two (or more) liquids that are completely soluble in each other in all proportions. Solvation: The process in which an ion or a molecule is surrounded by solvent molecules arranged in a specific manner.

9. Prentice Hall ©2004 Chapter 11Slide 9 Solution Formation07 1. Predict the relative solubilities in the following cases: (a) Br 2 in benzene (C 6 H 6 ) and in water, (b) KCl in carbon tetrachloride and in liquid ammonia, (c) urea (NH 2 ) 2 CO in carbon disulfide and in water. 2. Is iodine (I 2 ) more soluble in water or in carbon disulfide (CS 2 )? 3. Which would have the largest (most negative) hydration energy and which should have the smallest? Al 3+, Mg 2+, Na +

10. Prentice Hall ©2004 Chapter 11Slide 10 Concentration Units01

11. Prentice Hall ©2004 Chapter 11Slide 11 Concentration: The amount of solute present in a given amount of solution. Percent by Mass (weight percent): The ratio of the mass of a solute to the mass of a solution, multiplied by 100%.  %bymassofsolute= mass ofsolute mass ofsolution  100% mass ofsolution=mass ofsolute+mass ofsolvent Concentration Units01

12. Prentice Hall ©2004 Chapter 11Slide 12 Concentration Units02 Parts per Million: Parts per million (ppm) = = % mass x 10 4 One ppm gives 1 gram of solute per 1,000,000 g or one mg per kg of solution. For dilute aqueous solutions this is about 1 mg per liter of solution.

13. Prentice Hall ©2004 Chapter 11Slide 13 Concentration Units03 A sample of 0.892 g of potassium chloride (KCl) is dissolved in 54.6 g of water. What is the percent by mass of KCl in this solution? An aqueous solution is 5.50% H 2 SO 4. How many moles of sulfuric acid (MM = 98.08 g/mol) are dissolved in 250.0 g of the solution?

14. Prentice Hall ©2004 Chapter 11Slide 14 Concentration Units04 Mole Fraction (X): Molarity (M): Molality (m):

15. Prentice Hall ©2004 Units of Concentration Mass Percent (mass %) Mass % = (mass of component/total mass of SOLUTION) x 100% Parts per million, ppm = (mass of component/total mass of SOLUTION) x 10 6 Parts per billion, ppb = (mass of component / total mass of SOLUTION) x 10 9

16. Prentice Hall ©2004 Chapter 11Slide 16 Concentration Units06 Molality from Mass: Calculate the molality of a sulfuric acid solution containing 24.4 g of sulfuric acid in 198 g of water. The molar mass of sulfuric acid is 98.08 g. Molality from Molarity: Calculate the molality of a 5.86 M ethanol (C 2 H 5 OH) solution whose density is 0.927 g/ml.

17. Prentice Hall ©2004 Chapter 11Slide 17 Concentration Units07 Molality from Mass %: Assuming that seawater is a 3.50 mass % aqueous solution of NaCl, what is the molality of seawater? Molarity from Molality: The density at 20°C of a 0.258 m solution of glucose in water is 1.0173 g/mL, and the molar mass of glucose is 180.2 g. What is the molarity of the solution?

18. Prentice Hall ©2004 Chapter 11Slide 18 Concentration Units08 Mole Fraction from Molality: An aqueous solution is 0.258 m in glucose (MM = 180.2 g/mol). What is the mole fraction of the glucose? Mass from Molality: What mass (in grams) of a 0.500 m aqueous solution of urea [(NH 2 ) 2 CO, MM = 60.1 g/mol] would you use to obtain 0.150 mole of urea?

19. Prentice Hall ©2004 Chapter 11Slide 19 Effect of Temperature on Solubility01 Solids:

20. Prentice Hall ©2004 Chapter 11Slide 20 Effect of Temperature on Solubility02 Gases:

21. Prentice Hall ©2004 Chapter 11Slide 21 Henry’s Law: The solubility of a gas is proportional to the pressure of the gas over the solution. c  P c = k·P The Effect of Pressure on the Solubility of Gases01

22. Prentice Hall ©2004 Chapter 11Slide 22 Flash Animation - Click to Continue The Effect of Pressure on the Solubility of Gases02

23. Prentice Hall ©2004 Chapter 11Slide 23 The Effect of Pressure on the Solubility of Gases02 Calculate the molar concentration of O 2 in water at 25°C for a partial pressure of 0.22 atm. The Henry’s law constant for O 2 is 3.5 x 10 –4 mol/(L·atm). The solubility of CO 2 in water is 3.2 x 10 –2 M at 25°C and 1 atm pressure. What is the Henry’s law constant for CO 2 in mol/(L·atm)?

24. Prentice Hall ©2004 Chapter 11Slide 24 Colligative Properties of Nonvolatile Solutes01 Colligative Properties: Depend only on the number of solute particles in solution. These affect properties of the solvent. There are four main colligative properties: 1. Vapor pressure lowering 2. Freezing point depression 3. Boiling point elevation 4. Osmotic pressure

25. Prentice Hall ©2004 Chapter 11Slide 25 Colligative Properties of Nonvolatile Solutes02 When solute molecules displace solvent molecules at the surface, the vapor pressure drops since fewer gas molecules are needed to equalize the escape rate and capture rates at the liquid surface.

26. Prentice Hall ©2004 Chapter 11Slide 26 Colligative Properties of Nonvolatile Solutes03 Raoult’s Law: P soln = P° solv X solv For a single solute solution, X solv = 1 – X solute, We can obtain an expression for the change in vapor pressure of the solvent (the vapor pressure lowering).  P soln = P ° solv – P soln = P ° solv – X solv P ° solv = P ° solv – (1 – X solute ) P ° solv ∆P = X solute P ° solv Where superscript o is for pure substance.

27. Prentice Hall ©2004 Van’t Hoff Factor For incompletely dissociating ionic solids Van’t Hoff Factor i = moles of particles in solution moles of solute dissolved Chapter 11Slide 27

28. Prentice Hall ©2004 Chapter 11Slide 28 Colligative Properties of Nonvolatile Solutes05 The vapor pressure of a glucose (C 6 H 12 O 6 ) solution is 17.01 mm Hg at 20°C, while that of pure water is 17.25 mm Hg at the same temperature. Estimate the molality of the solution. How many grams of NaBr must be added to 250 g of water to lower the vapor pressure by 1.30 mm Hg at 40°C? The vapor pressure of water at 40°C is 55.3 mm Hg.

29. Prentice Hall ©2004 Chapter 11Slide 29 Colligative Properties of a Mixture of Two Volatile Liquids01 What happens if both components are volatile (have measurable vapor pressures)? The vapor pressure has a value intermediate between the vapor pressures of the two liquids. P T = P A + P B = X A P ° A + X B P ° B = X A P ° A + (1 – X A )P ° B P T = P ° B + (P ° A – P ° B )X A

30. Prentice Hall ©2004 Chapter 11Slide 30 Boiling-Point Elevation and Freezing-Point Depression01 Boiling-Point Elevation (∆T b ): The boiling point of the solution (T b ) minus the boiling point of the pure solvent (T ° b ): ∆T b = T b – T ° b ∆T b is proportional to concentration: ∆T b = K b m K b = molal boiling-point elevation constant. Also for incompletely dissociating ionic solids ∆T b = K b m i

31. Prentice Hall ©2004 Chapter 11Slide 31 Boiling-Point Elevation and Freezing-Point Depression02 Freezing-Point Depression (∆T f ): The freezing point of the pure solvent (T ° f ) minus the freezing point of the solution (T f ). ∆T f = T ° f – T f ∆T f is proportional to concentration: ∆T f = K f m K f = molal freezing-point depression constant. ∆T b = K b m i

32. Prentice Hall ©2004 Chapter 11Slide 32 Boiling-Point Elevation and Freezing-Point Depression04

33. Prentice Hall ©2004 Chapter 11Slide 33 Boiling-Point Elevation and Freezing-Point Depression06 van’t Hoff Factor, i: This factor equals the number of ions produced from each molecule of a compound upon dissolving. i = 1 for CH 3 OHi = 3 for CaCl 2 i = 2 for NaCli = 5 for Ca 3 (PO 4 ) 2 For compounds that dissociate on dissolving, use: ∆T b = i  K b m ∆T f = i  K f m ∆P = i  x 2 P ° 1

34. Prentice Hall ©2004 Chapter 11Slide 34 Boiling-Point Elevation and Freezing-Point Depression07 How many grams of ethylene glycol antifreeze, CH 2 (OH)CH 2 (OH), must you dissolve in one liter of water to get a freezing point of –20.0°C. The molar mass of ethylene glycol is 62.01 g. For water, K f = 1.86 (°C·kg)/mol. What will be the boiling point?

35. Prentice Hall ©2004 Chapter 11Slide 35 Boiling-Point Elevation and Freezing-Point Depression08 What is the molality of an aqueous solution of KBr whose freezing point is –2.95°C? K f for water is 1.86 (°C·kg)/mol. What is the freezing point (in °C) of a solution prepared by dissolving 7.40 g of K 2 SO 4 in 110 g of water? The value of K f for water is 1.86 (°C·kg)/mol.

36. Prentice Hall ©2004 Chapter 11Slide 36 Osmosis and Osmotic Pressure01

37. Prentice Hall ©2004 Chapter 11Slide 37 Osmosis: The selective passage of solvent molecules through a porous membrane from a dilute solution to a more concentrated one. Osmotic pressure (π or ∏): The pressure required to stop osmosis. π = i  MRT R = 0.08206 (L  atm)/(mol  K) Osmosis and Osmotic Pressure01

38. Prentice Hall ©2004 Chapter 11Slide 38 Osmosis and Osmotic Pressure02

39. Prentice Hall ©2004 Chapter 11Slide 39 Osmosis and Osmotic Pressure03

40. Prentice Hall ©2004 Chapter 11Slide 40 Osmosis and Osmotic Pressure04 Isotonic: Solutions have equal concentration of solute, and so equal osmotic pressure. Hypertonic: Solution with higher concentration of solute. Hypotonic: Solution with lower concentration of solute.

41. Prentice Hall ©2004 Chapter 11Slide 41 Osmosis and Osmotic Pressure05 The average osmotic pressure of seawater is about 30.0 atm at 25°C. Calculate the molar concentration of an aqueous solution of urea [(NH 2 ) 2 CO] that is isotonic with seawater. What is the osmotic pressure (in atm) of a 0.884 M sucrose solution at 16°C?

42. Prentice Hall ©2004 Chapter 11Slide 42 Uses of Colligative Properties01 Desalination:

43. Prentice Hall ©2004 Chapter 11Slide 43 Uses of Colligative Properties02 A 7.85 g sample of a compound with the empirical formula C 5 H 4 is dissolved in 301 g of benzene. The freezing point of the solution is 1.05°C below that of pure benzene. What are the molar mass and molecular formula of this compound?

44. Prentice Hall ©2004 Chapter 11Slide 44 Uses of Colligative Properties03 A 202 ml benzene solution containing 2.47 g of an organic polymer has an osmotic pressure of 8.63 mm Hg at 21°C. Calculate the molar mass of the polymer. What is the molar mass of sucrose if a solution of 0.822 g of sucrose in 300.0 mL of water has an osmotic pressure of 149 mm Hg at 298 K?

45. Prentice Hall ©2004 Chapter 11Slide 45 Uses of Colligative Properties06 Two miscible liquids, A and B, have vapor pressures of 250 mm Hg and 450 mm Hg, respectively. They were mixed in equal molar amounts. What is the total vapor pressure of the mixture and what are their mole fractions in the vapor phase?