1. Solutions Solutions are homogeneous mixtures of two or more pure substances. In a solution, the solute is dispersed uniformly throughout the solvent. © 2012 Pearson Education, Inc.

2. 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. © 2012 Pearson Education, Inc.

3. Solutions Click Here to play movie:How Does a Solution Form? As a solution forms, the solvent pulls solute particles apart and surrounds, or solvates, them. © 2012 Pearson Education, Inc.

4. Solutions 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. © 2012 Pearson Education, Inc.

5. Solutions Energy Changes in Solution Simply put, three processes affect the energetics of solution: –Separation of solute particles, –Separation of solvent particles, –New interactions between solute and solvent. © 2012 Pearson Education, Inc.

6. Solutions Energy Changes in Solution The enthalpy change of the overall process depends on  H for each of these steps. © 2012 Pearson Education, Inc.

7. Solutions 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. © 2012 Pearson Education, Inc.

8. Solutions Why Do Endothermic Processes Occur? Yet we know that in some processes, like the dissolution of NH 4 NO 3 in water, heat is absorbed, not released. © 2012 Pearson Education, Inc.

9. Solutions Enthalpy Is Only Part of the Picture The reason is that increasing the disorder or randomness (known as entropy) of a system tends to lower the energy of the system. © 2012 Pearson Education, Inc.

10. Solutions Enthalpy Is Only Part of the Picture So even though enthalpy may increase, the overall energy of the system can still decrease if the system becomes more disordered. © 2012 Pearson Education, Inc.

11. Solutions 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. © 2012 Pearson Education, Inc.

12. Solutions 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. © 2012 Pearson Education, Inc.

13. Solutions 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. © 2012 Pearson Education, Inc.

14. Solutions 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. © 2012 Pearson Education, Inc.

15. Solutions Exercise 1 Decide whether liquid hexane (C6H14) or liquid methanol (CH3OH) is the more appropriate solvent for the substances grease (C20H42) and potassium iodide (KI).

16. Solutions Factors Affecting Solubility The more similar the intermolecular attractions, the more likely one substance is to be soluble in another. © 2012 Pearson Education, Inc.

17. Solutions Factors Affecting Solubility Glucose (which has hydrogen bonding) is very soluble in water, while cyclohexane (which only has dispersion forces) is not. © 2012 Pearson Education, Inc.

18. Solutions Factors Affecting Solubility Vitamin A is soluble in nonpolar compounds (like fats). Vitamin C is soluble in water. © 2012 Pearson Education, Inc.

19. Solutions Gases in Solution In general, the solubility of gases in water increases with increasing mass. Larger molecules have stronger dispersion forces. © 2012 Pearson Education, Inc.

20. Solutions Gases in Solution The solubility of liquids and solids does not change appreciably with pressure. Click here to play movie:But the solubility of a gas in a liquid is directly proportional to its pressure.Click here to play movie:But the solubility of a gas in a liquid is directly proportional to its pressure. © 2012 Pearson Education, Inc.

21. Solutions 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. © 2012 Pearson Education, Inc.

22. Solutions Exercise 2 A certain soft drink is bottled so that a bottle at 25°C contains CO2 gas at a pressure of 5.0 atm over the liquid. Assuming that the partial pressure of CO2 in the atmosphere is 4.0 × 10−4 atm, calculate the equilibrium concentrations of CO2 in the soda both before and after the bottle is opened. The Henry’s law constant for CO2 in aqueous solution is 0.031 mol/L atm at 25°C. before = 0.16 mol/L after = 1.2 × 10—5 mol/L

23. Solutions Temperature Generally, the solubility of solid solutes in liquid solvents increases with increasing temperature. Most solution formation is endothermic— LeChatelier’s Principle ). Solubility generally decreases with temperature if the solution process is exothermic © 2012 Pearson Education, Inc.

24. Solutions 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. © 2012 Pearson Education, Inc.

25. Solutions Temperature/Solubility of gases The opposite is true of gases. –Carbonated soft drinks are more “bubbly” if stored in the refrigerator. The solubility of a gas in water always decreases with increasing temperature. © 2012 Pearson Education, Inc.

26. Solutions Ways of Expressing Concentrations of Solutions © 2012 Pearson Education, Inc.

27. Solutions Mass Percentage Mass % of A = mass of A in solution total mass of solution  100 © 2012 Pearson Education, Inc.

28. Solutions moles of A total moles of all components X A = Mole Fraction (X) In some applications, one needs the mole fraction of solvent, not solute— make sure you find the quantity you need! © 2012 Pearson Education, Inc.

29. Solutions moles of solute kilograms of solvent m = Molality (m) Since both moles and mass do not change with temperature, molality (unlike molarity) is not temperature- dependent. © 2012 Pearson Education, Inc.

30. Solutions Changing Molarity to Molality If we know the density of the solution, we can calculate the molality from the molarity, and vice versa. © 2012 Pearson Education, Inc.

31. Solutions Colligative Properties Changes in colligative properties depend only on the number of solute particles present, not on the identity of the solute particles. Among colligative properties are –Vapor-pressure lowering –Boiling-point elevation –Melting-point depression –Osmotic pressure –Click Here for the Chemguy’s discussion of colligative propertiesClick Here for the Chemguy’s discussion of colligative properties –Stop here © 2012 Pearson Education, Inc.

32. Solutions

33. How many moles of ions in 250.0 ml of.500 M Al 2 (SO 4 ) 3

34. Solutions 1. Vapor Pressure Lowering Because of solute–solvent intermolecular attraction, higher concentrations of nonvolatile solutes make it harder for solvent to escape to the vapor phase. © 2012 Pearson Education, Inc.

35. Solutions 1. Vapor Pressure Lowering OR A more complicated (more favored), entropy based explanation. When a solute is added to a pure solvent, the entropy (disorder) dramatically increases. This is the preferred situation (since positive  S helps  G be negative). Prior to the solute being added, the solvent could only increase the disorder by evaporating (becoming vapor). With the solute added the entropy has already been dramatically increased, and there is less incentive for the vaporization process.

36. Solutions Vapor Pressure Therefore, the vapor pressure of © 2012 Pearson Education, Inc.

37. The vapor pressure of a solution is directly proportional to the mole fraction of solvent present. If the solute ionizes, the number of ions further affects (lowers) the vapor pressure.

38. Solutions P solution = observed vapor pressure of the solvent in the solution χ solvent = mole fraction of solvent Po solvent - vapor pressure of pure solvent i = van’t Hoff factor (moles of electrolyte must be multiplied by this) number of moles particles in solution/number of moles particles dissolved

39. Solutions Changes in vapor pressure are governed by Raoult’s Law. Fall into 2 categories: 1. Those where the solute is non-volatile, 2.and those where the solute is volatile, i.e. those where the solution has two volatile components

40. Solutions Exercise 5 Calculating the Vapor Pressure of a Solution Calculate the expected vapor pressure at 25°C for a solution prepared by dissolving 158.0 g of common table sugar (sucrose, molar mass = 342.3 g/mol) in 643.5 cm3 of water. At 25°C, the density of water is 0.9971 g/cm3 and the vapor pressure is 23.76 torr. = 23.46 torr

41. Solutions Boiling-Point Elevation The boiling point of a solution is reached when the external atmospheric pressure is equal to its vapor pressure. Since addition of a solute lowers the vapor pressure of a solution, it must, by definition, raise the boiling point.

42. Solutions Boiling Point Elevation For non-electrolyte solutions (non- ionic), the change is summarized in the expression T b = K b m  Tb = change in boiling point of the solvent, K b = molal boiling point constant for the solvent and m = molality of the solute

43. Freezing point depression Freezing points are lowered by the addition of non-volatile solutes. Again, this can be explained in one of two ways; (i) As a physical phenomenon (less favored), where solute particles actually get in the way of the solid structure being formed. If the solidification (freezing) process is hindered, more heat must be removed from the solution, and hence the freezing point is lowered.

44. Freezing point depression

45. OR (ii) As an entropy based phenomenon (more favored), where the solution has more disorder than the pure solvent so needs more energy removed from it in order to create the more ordered solid.  T f = K f  m

47. Solutions 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. © 2012 Pearson Education, Inc.

48. Solutions EXAMPLE: Calculate the freezing point and boiling point of a solution of 100. g ethylene glycol (C2H6O2) in 900. g of water.

49. Solutions Osmosis Some substances form semipermeable membranes, allowing some smaller particles to pass through, but blocking other larger particles. In biological systems, most semipermeable membranes allow water to pass through, but solutes are not free to do so. © 2012 Pearson Education, Inc.

50. Solutions Osmosis In osmosis, there is net movement of solvent from the area of higher solvent concentration (lower solute concentration) to the area of lower solvent concentration (higher solute concentration). © 2012 Pearson Education, Inc.

51. Solutions Osmotic Pressure The pressure required to stop osmosis, known as osmotic pressure, , is  = ( ) RT = MRT nVnV 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. © 2012 Pearson Education, Inc.

52. Solutions Exercise 9 Determining Molar Mass from Osmotic Pressure To determine the molar mass of a certain protein, 1.00 × 10−3 g of it was dissolved in enough water to make 1.00 mL of solution. The osmotic pressure of this solution was found to be 1.12 torr at 25.0°C. Calculate the molar mass of the protein. = 1.66 × 10 4 g/mol

53. Solutions Electrolyte solutions and the van’t Hoff factor

54. Solutions Colligative Properties of Electrolytes Since the colligative properties of electrolytes depend on the number of particles dissolved, solutions of electrolytes (which dissociate in solution) should show greater changes than those of nonelectrolytes. © 2012 Pearson Education, Inc.

55. Solutions Colligative Properties of Electrolytes For example, one unit of MgCl2 splits into three particles when it is dissolved in solution (one Mg2+ ion and two Cl- ions). Since we are dealing with colligative properties (those that depend on the number of particles present), this should produce a freezing point depression three times greater than a non-electrolyte solute that does not produce a greater number of particles when dissolved, for example, molecular sucrose, C6H12O6. © 2012 Pearson Education, Inc.

56. Solutions Colligative Properties of Electrolytes In order to account for this, the van’t Hoff factor (i) is added to the expressions already discussed. T b = i K b m  T f = i K f  m  = ( n/v )RT = i MR T © 2012 Pearson Education, Inc.

57. Solutions

58. van’t Hoff Factor One mole of NaCl in water does not really give rise to two moles of ions. © 2012 Pearson Education, Inc.

59. Solutions van’t Hoff Factor Reassociation is more likely at higher concentration. Therefore, the number of particles present is concentration- dependent. © 2012 Pearson Education, Inc.

60. Solutions Colloids Suspensions of particles larger than individual ions or molecules, but too small to be settled out by gravity, are called colloids. © 2012 Pearson Education, Inc.

61. Solutions Tyndall Effect Colloidal suspensions can scatter rays of light. This phenomenon is known as the Tyndall effect. © 2012 Pearson Education, Inc.

62. Solutions Colloids in Biological Systems Some molecules have a polar, hydrophilic (water-loving) end and a nonpolar, hydrophobic (water- hating) end. © 2012 Pearson Education, Inc.

63. Solutions Colloids in Biological Systems These molecules can aid in the emulsification of fats and oils in aqueous solutions. © 2012 Pearson Education, Inc.