1. CHAPTER FOURTEEN: SOLUTIONS AND THEIR BEHAVIOR

2. SOLUTION CONCENTRATION Molarity(M): Moles solute/1L solution Molality (m): Moles solute/1kg solvent Mole fraction (X A ): Moles A* total moles solution Mass percent: Mass solute x 100 total mass of solution * In some applications, one needs the mole fraction of solvent, not solute. Make sure you find the quantity you need!

3. CHANGING MOLARITY TO MOLALITY If we know the density of the solution, we can calculate the molality from the molarity and vice versa. © 2009, Prentice-Hall, Inc.

4. CONVERTING MOLALITY TO MOLARITY Determine the molarity of a 0.273m aqueous solution of KCl? The density is 1.011g/mL

5. CONVERTING MOLALITY TO MOLARITY What is the molarity of a 0.273m aqueous solution of KCl? The density if 1.011g/mL Ans: 0.271M

6. CONVERTING MOLARITY TO MOLALITY DETERMINE THE MOLALITY OF A 0.907M SOLUTION OF PB(NO 3 ) 2. THE SOLUTION DENSITY IS 1.252 G/ML.

7. CONVERTING MOLARITY TO MOLALITY An aqueous solution is 0.907M Pb(NO 3 ) 2. What is the molaltiy of lead (II) nitrate in this solution? The density is 1.252 g/mL. Ans: 0.953m

8. CALCULATING MASS PERCENT How would you prepare 425g of an aqueous solution containing 2.40% by mass of sodium acetate? Ans: 10.2 g NaC 2 H 3 O 2 in 414.8g of water

9. MOLE FRACTION Calculate the mole fraction of NaCl in an aqueous 5% solution of NaCl Mole fraction NaCl = 0.016

10. SOLUTION CONCENTRATION Commercial concentrated aqueous ammonia is 27% NH 3 by mass and has a density of 0.90 g/mL. What is the molarity of this solution?

11. SOLUTIONS Solutions are homogeneous mixtures of two or more pure substances. In a solution, the solute is dispersed uniformly throughout the solvent. © 2009, Prentice-Hall, Inc.

12. METAL ALLOYS: Substitutional alloys are formed when the two metallic components have similar atomic radii and chemical-bonding characteristics. For example, silver and gold form such an alloy over the entire range of possible compositions.

13. METAL ALLOYS: Interstitial alloys are form when the component present in the interstitial positions between the solvent atoms must have a much smaller covalent radius than the solvent atoms. Typically, an interstitial element is a nonmetal that participates in bonding to neighboring atoms. The presence of the extra bonds provided by the interstitial component causes the metal lattice to become harder, stronger, and less ductile. For example, steel is an alloy of iron that contains up to 3 percent carbon. Steel is much harder and stronger than pure iron.

14. AQUEOUS SOLUTIONS The intermolecular forces between solute and solvent particles must be strong enough to compete with those between solute particles and those between solvent particles. © 2009, Prentice-Hall, Inc.

15. HOW DOES A SOLUTION FORM? As a solution forms, the solvent pulls solute particles apart and surrounds, or solvates, them. © 2009, Prentice-Hall, Inc. http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/molvie1.swf

16. HOW DOES A SOLUTION FORM If an ionic salt is soluble in water, it is because the ion- dipole interactions are strong enough to overcome the lattice energy of the salt crystal. © 2009, Prentice-Hall, Inc.

17. solvation.Solutes dissolve in solvents by a process called solvation. like dissolves likePolar solvent dissolve polar solutes, non-polar solvents dissolve non-polar solutes. (aka: “like dissolves like”. miscibleIf two liquids mix to an appreciable extent to form a solution, they are said to be miscible. immiscibleIn contrast, immiscible liquids do not mix to form a solution; they exist in contact with each other as separate layers. The Solution Process

18. cationnegative When a cation exists in solution, it is surrounded by the negative dipole ends of water molecules. anionpositive When as anion exists in solution, it is surrounded by the positive dipole ends of water molecules. + Solvation of Ions

19. FACTORS AFFECTING SOLUBILITY Chemists use the axiom “like dissolves like." Polar substances tend to dissolve in polar solvents. Nonpolar substances tend to dissolve in nonpolar solvents. © 2009, Prentice-Hall, Inc.

20. FACTORS AFFECTING SOLUBILITY The more similar the intermolecular attractions, the more likely one substance is to be soluble in another. © 2009, Prentice-Hall, Inc.

21. ENERGY CHANGES IN SOLUTION THE enthalpy change of the overall process depends on  H for each of these steps. © 2009, Prentice-Hall, Inc.

22. WHY DO ENDOTHERMIC PROCESSES OCCUR? Things do not tend to occur spontaneously (i.e., without outside intervention) unless the energy of the system is lowered. The reason is that increasing the disorder or randomness (known as entropy) of a system tends to lower the energy of the system. © 2009, Prentice-Hall, Inc.

23. TYPES OF SOLUTIONS Saturated In a saturated solution, the solvent holds as much solute as is possible at that temperature. Dissolved solute is in dynamic equilibrium with solid solute particles. © 2009, Prentice-Hall, Inc.

24. TYPES OF SOLUTIONS Unsaturated If a solution is unsaturated, less solute than can dissolve in the solvent at that temperature is dissolved in the solvent. © 2009, Prentice-Hall, Inc.

25. TYPES OF SOLUTIONS Supersaturated In supersaturated solutions, the solvent holds more solute than is normally possible at that temperature. These solutions are unstable; crystallization can usually be stimulated by adding a “seed crystal” or scratching the side of the flask. © 2009, Prentice-Hall, Inc.

26. GASES IN SOLUTION In general, the solubility of gases in water increases with increasing mass. Larger molecules have stronger dispersion forces. © 2009, Prentice-Hall, Inc.

27. GASES IN SOLUTION The solubility of liquids and solids does not change appreciably with pressure. The solubility of a gas in a liquid is directly proportional to its pressure. © 2009, Prentice-Hall, Inc.

28. HENRY’S LAW S G = KP G where S g is the solubility of the gas, k is the Henry’s Law constant for that gas in that solvent, and P g is the partial pressure of the gas above the liquid. © 2009, Prentice-Hall, Inc.

29. TEMPERATURE Generally, the solubility of solid solutes in liquid solvents increases with increasing temperature. © 2009, Prentice-Hall, Inc.

30. SUCROSE SOLUBILITY

31. TEMPERATURE The opposite is true of gases. Carbonated soft drinks are more “bubbly” if stored in the refrigerator. Warm lakes have less O 2 dissolved in them than cool lakes. © 2009, Prentice-Hall, Inc.

32. COLLIGATIVE PROPERTIES In chemistry, colligative properties are properties of solutions that depend upon the ratio of the number of solute particles to the number of solvent molecules in a solution, and not on the type of chemical species present.

33. COLLIGATIVE PROPERTIES: Vapor Pressure Freezing Point Depression Boiling Point Elevation Osmotic Pressure

34. © 2009, Prentice-Hall, Inc. VAPOR PRESSURE Because of solute- solvent intermolecular attraction, higher concentrations of nonvolatile solutes make it harder for solvent to escape to the vapor phase.

35. © 2009, Prenti ce- Hall, Inc. RAOULT’S LAW P A = X A P  A where X A is the mole fraction of compound A, and P  A is the normal vapor pressure of A at that temperature. NOTE: This is one of those times when you want to make sure you have the vapor pressure of the solvent.

36. VAPOR PRESSURE: Raoult’s Law: P solution = X solvent P o H2O 12g Sucrose (C 12 H 22 O 11 )is dissolved in 250.0g water at 90 °C. What is the vapor pressure of water over this solution? (P o H2O = 525.8 mmHg- from data table) Ans. 524 mmHg

37. © 2009, Prentice-Hall, Inc. BOILING POINT ELEVATION AND FREEZING POINT DEPRESSION Nonvolatile solute- solvent interactions also cause solutions to have higher boiling points and lower freezing points than the pure solvent.

38. © 2009, Prentice-Hall, Inc. BOILING POINT ELEVATION The change in boiling point is proportional to the molality of the solution:  T b = K b  m where K b is the molal boiling point elevation constant, a property of the solvent.  T b is added to the normal boiling point of the solvent.

39. © 2009, Prenti ce- Hall, Inc. BOILING POINT ELEVATION AND FREEZING POINT DEPRESSION Note that in both equations,  T does not depend on what the solute is, but only on how many particles are dissolved.  T B = K B  m  T F = K F  m

40. FREEZING POINT DEPRESSION: Freezing Point Depression: ΔT fp = K fp m solution A solution is prepared by adding 0.50g of caffeine (C 8 H 10 O 2 N 4 ) to 100g of benzene (C 6 H 6 ). Calculate the freezing point of this solution. The freezing point of pure benzene is 5.50°C K fp for benzene = 5.23 °C/m) ANS:  T F = 0.132 ° C T F = 5.37C

41. BOILING POINT ELEVATION: Boiling point elevation: ΔT bp = K bp m solute A glycerol solution (C 3 H 8 O 3 ) in water is prepared by dissolving glycerol is 500g water. The boiling point of the solution is 100.42°C at 760mmHg. What mass of glycerol was dissolved to make this solution? K bp = 0.5121 °C/m Ans: 38 g glycerol

42. © 2009, Prentice-Hall, Inc. COLLIGATIVE PROPERTIES OF ELECTROLYTES Since these properties depend on the number of particles dissolved, solutions of electrolytes (which dissociate in solution) should show greater changes than those of nonelectrolytes.

43. © 2009, Prentice-Hall, Inc. VAN’T HOFF FACTOR We modify the previous equations by multiplying by the van’t Hoff factor, i.  T f = K f  m  i

44. © 2009, Prentice-Hall, Inc. VAN’T HOFF FACTOR Re-association is more likely at higher concentration. Therefore, the number of particles present is concentration- dependent.

45. ROLE OF ELECTROLYTES ON COLLIGATIVE PROPERTIES: Van’t Hoff Factor: ΔT fp = K fp m solute i A 0.0711m aqueous solution of Sodium sulfate freezes at -0.32C. What is the actual value (i) of the van’t Hoff factor? K fp = 1.86 °C/m Ans: i = 2.42

46. RANK FOR FOLLOWING: Increasing boiling point (ΔTb) 0.25 m C 6 H 12 O 11 0.40 m NaCl 0.15 m MgCl 2 C 6 H 12 O 11  MgCl 2  NaCl

47. © 2009, Prentice-Hall, Inc. OSMOSIS In osmosis, there is net movement of solvent from the area of higher solvent concentration (lower solute concentration) to the are of lower solvent concentration (higher solute concentration).

48. © 2009, Prenti ce- Hall, Inc. OSMOTIC PRESSURE The pressure required to stop osmosis, known as osmotic pressure, (  ) nVnV  = ( ) RT = MRT where M is the molarity of the solution. If the osmotic pressure is the same on both sides of a membrane (i.e., the concentrations are the same), the solutions are isotonic.

49. © 2009, Prentice-Hall, Inc. OSMOSIS IN BLOOD CELLS If the solute concentration outside the cell is greater than that inside the cell, the solution is hypertonic. Water will flow out of the cell, and crenation results.

50. OSMOTIC PRESSURE Osmotic Pressure: π= MRT What is the osmotic pressure of a 0.1M solution of sucrose at 25C? (remember: R=.08206 L-atm/mol-K) Ans: 2.45 atm

51. ANOTHER OSMOTIC PRESSURE PROBLEM. 100. mg of a protein are dissolved in enough water to make 10.0 mL of a solution. If this solution has an osmotic pressure of 13.3 mmHg at 25°C, what is the molar mass of the protein? Ans: 1.4 x 10 4 g/mole

52. © 2009, Prentice-Hall, Inc. COLLOIDS Suspensions of particles larger than individual ions or molecules, but too small to be settled out by gravity are called colloids.

53. © 2009, Prentice-Hall, Inc. TYNDALL EFFECT Colloidal suspensions can scatter rays of light. This phenomenon is known as the Tyndall effect.

54. Soap molecules interact with water through the charged, hydrophilic end of the molecule. The long, hydrophobic end of the molecule binds through dispersion forces with nonpolar hydrocarbons and other non polar substances. How Soap Works: