Solutions. NaCl and water form homogeneous mixture Mixture is even throughout

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1 Solutions 1

2 Types of Mixtures 2

3 Types of Mixtures Some materials it is easy to see it s a mixture, b/c you can see component parts Others, like milk, do not look like mixtures, but they are Can see round droplets Heterogeneous mixture Composition not equal 3

4 Solutions NaCl and water form homogeneous mixture Mixture is even throughout 4

5 Sugar is soluble in water Soluble - ability to be dissolved What happens as sugar dissolves in water? 5

6 Lump of sugar disappears as sugar molecules leave surface of their crystals and mix with water molecules Eventually all sugar molecules are evenly spread through water molecules All visible traces of sugar are gone It becomes solution - homogeneous mixture of two or more substances in a single phase 6

7 Components of Solutions Particles of one substance randomly mixed with particles of another Solvent - dissolving medium in a solution Solute - substance dissolved in solution 7

8 Solute usually less than solvent Dissolved solute particles cannot be seen Stay mixed with solvent forever, so long as conditions stay the same Solute-particles from nm in diameter 8

9 Types of Solutions Solutions can exist as gases, liquids, or solids 9

10 Many alloys, such as brass (zinc and copper) are sterling silver (silver and copper) are solid solutions Atoms of two or metals are evenly mixed By properly choosing percentages of each metal, can get many desirable properties 10

11 Ex. Alloys have higher strength and greater resistance to corrosion than pure metals Pure gold (24K) too soft for jewelry Alloy with silver (14K) increases strength and hardness but keeps appearance of gold 11

12 Suspension - if particles in solvent are so large they settle out unless mixture constantly stirred Think of sand and water Particles over 1000 nm in diameter Suspensions 12

13 Colloids Particles that are medium in size between those in solutions and suspensions from mixtures are known as colloids (colloidal dispersions) Particles between nm diameter 13

14 In muddy water, large soil particles settle Water still cloudy b/c colloidal particles stay dispersed in water Filter? Colloidal particles pass through and stays cloudy 14

15 Particles in colloid small enough to be suspended in solvent by constant movement of surrounding molecules Colloidal particles make up dispersed phase Water is dispersing medium 15

16 Emulsion and foam are colloids Mayonnaise is emulsion of oil droplets in water Egg yolk acts as emulsifying agent (keeps oil droplets dispersed) 16

17 17

18 Class of Colloid Phases Example Sol Solid dispersed in liquid Paints, mud Gel Solid network extending through liquid Gelatin Liquid emulsion Liquid dispersed in liquid Milk, mayonnaise Foam Gas dispersed in liquid Shaving cream, whipped cream Solid aerosol Solid dispersed in gas Smoke, auto exhaust Liquid aerosol Liquid dispersed in gas Fog, mist, clouds, aerosol spray 18 Solid emulsion Liquid dispersed in solid Cheese, butter

19 Tyndall Effect Many colloids appear homogeneous b/c individual particles cannot be seen Particles large enough to scatter light In fog, can see headlight beam Tyndall effect - occurs when light is scattered by colloidal particles dispersed in transparent medium 19

20 20

21 Solutions Colloids Suspensions Homogeneous Heterogeneous Heterogeneous Particle size: nm; can be atoms, ions, molecules Particle size: nm, dispersed; can be combined or large molecules Particle size: over 1000 nm, suspended; can be large particles or combined particles Do not separate on standing Do not separate on standing Particles settle out Cannot be separated by filtration Cannot be separated by filtration Can be separated by filtration Do not scatter light Scatter light (Tyndall effect) May scatter light, but are not transparent 21

22 Electrolytes vs. Nonelectrolytes Substances that dissolve in water are classified according to whether they produce molecules or ions in solution 22

23 When ionic compound dissolves, cations and anions separate and are surrounded by water molecules Solute ions free to move, making possible to carry electric current 23

24 Electrolyte - substance that dissolves in water to give a solution that conducts electric current NaCl electrolyte Usually highly polar compounds become electrolytes 24

25 Solution containing neutral solute molecules does not conduct current b/c it doesn t have mobile charged particles Nonelectrolyte - substance that dissolves in water to give solution that does not conduct electric current Ex. Sugar 25

26 26

27 the solution process 27

28 Factors Affecting Rate of Dissolution Ever tried to dissolve sugar in iced tea or coffee? You notice temperature has something to do with how quickly solute dissolves What other factors affect how quickly solutes dissolve? 28

29 1. Increasing Surface Area of Solute b/c dissolution process occurs at surface of solute Can increase dissolution by increasing surface area 29

30 2. Agitating (Stirring) Solution Close to surface of solute, concentration of dissolved solute is high Stirring/shaking helps disperse solute particles 30

31 Brings fresh solvent into contact with solute surface Contact between solvent and solute surface increased

32 3. Heating a Solvent Solutes dissolve faster in warmer solvent As temp increases, solvent molecules move faster Average kinetic energy increases Collisions between solvent and solute more frequent Helps separate solute molecules and spread them out 32

33 Solubility If you add sugar to tea, eventually no more sugar dissolves For every combination of solvent with solute at given temp, there is limit to amount of solute that can be dissolved 33

34 When sugar first dropped into water, sugar molecules leave solid surface and move about at random in solvent Some dissolved molecules may collide with crystal and stay there As more solid dissolves and concentration of dissolved molecules increases, collisions happen more often 34

35 Eventually, molecules are returning to crystal at same rate they are going into solution Dynamic equilibrium established between dissolution and crystallization Solution equilibrium - physical state in which the opposing processes of dissolution and crystallization of a solute occur at equal rates 35

36 Saturated vs. Unsaturated Solutions Saturated solution - solution that contains the maximum amount of dissolved solute How do you know if solution is saturated? If more NaCl added, it falls to bottom and doesn t dissolve Equilibrium already established 36

37 If more water added to saturated solution more NaCl will dissolve At 20, 35.9 g NaCl is max that will dissolve in 100 g water Unsaturated solution - solution that contains less solute than a saturated solution under the existing conditions 37

38 Supersaturated Solutions When saturated solution where solubility increases as temp increases is cooled, excess solute usually comes out of solution Sometimes, if solution is left to cool the excess solute doesn t separate 38

39 Supersaturated solution - solution that contains more dissolved solute than a saturated solution contains under same conditions 39

40 40

41 Supersaturated solution can remain unchanged over long time if it isn t disturbed Once crystals begin to form, continues until equilibrium is recreated at lower temp 41

42 Example Sodium thiosulfate, Na2S2O3 Solute added to hot water until solution saturated Hot solution filtered Drop small crystal in seeding crystallization Continues until equilibrium created 42

43 Solubility Values Solubility - amount of substance required to form a saturated solution with specific amount of solvent at specified temp Ex. Solubility of sugar = 204 g per 100 g water at 20 Must specify temp b/c solubility changes with temp 43

44 For gases, pressure must also be specified Rate at which substance dissolves is unrelated to solubility Max amount of solute that dissolves and reaches equilibrium is always same under same conditions 44

45 Solute-Solvent Interactions like dissolves like is generally useful rule for predicting whether one substance will dissolve in another Depends on type of bonding, polarity/ nonpolarity, and intermolecular forces between solute and solvent 45

46 Dissolving Ionic Compounds in Aqueous Solution Polarity of water molecules plays important role in formation of solution of ionic compounds Charged ends of water molecules attract ions and surround them to keep separated from other ions 46

47 Hydration - solution process with water as solvent Ions said to be hydrated As hydrated ions diffuse into solution, other ions exposed and drawn away from crystal surface Entire crystal gradually dissolves 47

48 When crystallized from aqueous solutions, some ionic compounds form crystals that include water molecules The crystalline compounds, known as hydrates, have specific ratios of water molecules Represented by formulas like CuSO4 5H2O 48

49 When crystalline hydrate dissolves in water, water of hydration returns to solvent Behavior of solute in hydrated form no different than anhydrous form 49

50 Nonpolar Solvents Nonpolar solvents do not attract ions of ionic compound strongly enough to overcome forces holding crystal together 50

51 Liquid Solutes and Solvents When you shake bottle of salad dressing, oil droplets disperse in water Stop shaking strong attraction between water molecules squeeze out oil and form separate layers Liquid solutes and solvents that are not soluble in each other are immiscible 51

52 Nonpolar substances (fat, oil, grease) are generally quite soluble in nonpolar liquids (carbon tetrachloride, toluene, gasoline) Only attractions between nonpolar molecules are weak London forces 52

53 London Dispersion Forces b/c e- are constantly moving, at any time, there may be uneven distribution of them around a nonpolar atom Temporarily creates positive end and negative end 53

54 London dispersion forces - intermolecular attractions resulting from constant motion of e- and creation of instantaneous dipoles 54

55 55

56 Liquids that dissolve freely in one another in any ratio are said to be completely miscible Benzene and carbon tetrachloride (both nonpolar) are completely miscible Nonpolar molecules of these two apply no strong forces of attraction or repulsion, and molecules mix freely 56

57 Ethanol and water are also completely miscible -OH group on ethanol is slightly polar Can form H-bonds with water as well as other ethanol molecules Intermolecular forces in mixture are very similar to those in pure liquids so they are mutually soluble in any amount 57

58 Gasoline contains mainly nonpolar hydrocarbons Excellent solvent for fats, oils, greases Major intermolecular forces are weak London forces 58

59 Ethanol less polar than water, more than carbon tetrachloride Better solvent for less-polar substances b/c of nonpolar region 59

60 Effects of Pressure on Solubility Little effect on solids or liquids Increase pressure - increase solubilities of gases in liquids 60

61 When gas in contact with surface of liquid, gas molecules can enter liquid As amount of dissolve gas increases, some molecules start to escape and reenter gas phase Eq eventually established between rates of entering/leaving gas phase 61

62 As long as eq undisturbed, solubility of gas in liquid unchanged at given pressure Gas + solvent solution 62

63 Increasing pressure of solute gas above solution puts stress on eq Molecules collide with surface more often Increase in pressure partially offset by increase in rate of gas molecules entering solution In turn, increase in amount of dissolved gas causes increase in rate molecules escape liquid and becomes gas 63

64 Eventually, eq restored at higher gas solubility Expected from Le Chatelier s principle Increase in gas pressure causes eq to shift so fewer molecules are in gas phase 64

65 Henry s Law Henry s law - solubility of a gas in liquid is directly proportional to the partial pressure of that gas on the surface of the liquid Applies to gas-liquid solutions at constant temp 65

66 When mixture of ideal gases is limited in constant volume at constant temp, each gas applies same pressure it would apply if it occupied the space alone Assuming gases do not react, each gas dissolves to the extent it would if no other gases were there 66

67 In carbonated drinks (Coke, Pepsi, etc.) solubility of CO2 increased by increasing pressure At bottling factory, CO2 gas forced into solution of flavored water at pressure of 5-10 atm Gas-in-liquid solution then sealed in bottles/ cans 67

68 When cap removed, pressure reduced to 1 atm, and some CO2 escapes as gas bubbles Effervescence - rapid escape of a gas from a liquid in which it is dissolved 68

69 Effects of Temperature on Solubility Let s consider gas solubility Increasing temperature usually decreases gas solubility As temp increases, average kinetic energy of molecules in solution increases Greater number of solute molecules escape from attraction of solvent and return to 69

70 At higher temps, eq is reached with fewer gas molecules in solution Gases generally less soluble 70

71 Effect of temp on solubility of solids in liquids more difficult to predict Often, increasing temp increases solubility However, the same temp increase can result in large increase in solubility in one case, and only slight increase in the next 71

72 72

73 enthalpies of Solution Formation of solution comes with energy change Dissolve KI in water, container feels cold Dissolve LiCl in water, feels hot Formation of solid-liquid solution can absorb or release heat 73

74 During formation of solution, solvent and solute particles experience changes in forces attracting them to other particles Before dissolving begins, solvent molecules held by intermolecular forces (solvent-solvent attraction) In solute, molecules held by Ifs (solute-solute attraction) 74

75 Energy is required to separate solute molecules and solvent molecules from their neighbors A solute particle that is surrounded by solvent molecules is said to be solvated 75

76 Step 1: solute particles become separated from solid (energy absorbed) Step 2: solvent particles move apart to allow solute particles to enter liquid (energy absorbed) 76

77 Step 3: solvent particles attracted to and solvate solute particles (energy released) 77

78 Heat of solution - net amount of heat energy absorbed or released when a specific amount of solute dissolves in a solvent Heat of solution negative (heat is released) if sums of energy in steps 1 and 2 is less than step 3 If 1+2 > 3, HOS is positive (heat absorbed) 78

79 In gaseous state, molecules are so far apart there are nearly no forces between them Solute-solute interaction has little effect on heat of solution of a gas Energy released when gas dissolved in liquid b/ c attraction between solute gas and solvent molecules greater than energy needed to separate solvent molecules 79

80 Concentration of Solutions 80

81 Concentration - measure of amount of solute in given amount of solvent or solution dilute just means small amount of solute concentrated just means relatively large amount of solute Unrelated to degree to which solution is saturated Ex. Saturated solution of substance not very soluble might be dilute 81

82 Molarity Molarity - number of moles of solute in one liter of solution To find molarity, must know molar mass A one molar solution of NaOH is 1 mole NaOH per liter of solution (1 M) 82

83 1 mol NaOH = 40.0 g If this quantity dissolve in enough water to make EXACTLY 1.00 L solution, solution is 1 M If 20.0 g NaOH dissolve in enough to make 1.00 L solution, what is molarity? M 83

84 1 M solution is not made by adding 1 mol solute to 1 L solvent That would be more than 1 L Instead, 1 mol solute dissolved in less than 1 L Then diluted with more solvent to bring TOTAL VOLUME to 1 L 84

85 Practice Problem You have 3.40 L of solution that contains 90.0 g sodium chloride. What is the molarity of that solution? M NaCl You have 0.8 L of a 0.5 M HCl solution. How many moles of HCl does this solution contain? 0.4 mol HCl 85

86 What is the molarity of a solution composed of 5.85 g of potassium iodide, dissolved in enough water to make a L of solution? M KI 86

87 How many moles of H2SO4 are present in L of a M H2SO4 solution? mol What volume of 3.00 M NaCl is needed for reaction that requires g of NaCl? L 87

88 Molality Molality - concentration of solution expressed in moles of solute per kilogram of solvent Solution that contains 1 mol solute (NaOH) dissolved in exactly 1 kg of solvent is one molal solution (1 m NaOH) 88

89 1 mol NaOH = 40.0 g 40.0 g NaOH dissolved in 1 kg water is 1 m NaOH 20.0 g NaOH in 1 kg water = m NaOH 89

90 Practice Problem A solution was prepared by dissolving 17.1 g sucrose (C12H22O11) in 125 g of water. Find the molal concentration of this solution m C12H22O11 90

91 A solution of iodine in carbon tetrachloride is used when iodine is needed for certain chemical tests. How much iodine must be added to prepare a m solution of iodine in CCl4 if g of CCl4 is used? 12.2 g I2 91

92 What is the molality of a solution composed of 255 g acetone, (CH3)2CO, dissolved in 200. g of water? 22 m acetone What quantity, in grams, of methanol, CH3OH, is required to prepare a m solution in 400. g water? 3.12 g methanol 92

93 How many grams of AgNO3 are needed to prepare a m solution in 250. ml of water? 5.31 g AgNO3 What is the molality of a solution containing 18.2 g HCl and 250. g water? 2.00 m 93

94 Ions in Aqueous Solutions and Colligative Properties 94

95 compounds in aqueous solutions 95

96 Dissociation When compound made from ions dissolves in water, ions separate Dissociation - separation of ions that occurs when an ionic compound dissolves 96

97 Notice number of ions made per formula unit in equations 1 formula unit of NaCl gives 2 units of ions in solution 1 formula unit of CaCl2 gives 3 units of ions 97

98 Assuming 100% dissociation, solution that contains 1 mol NaCl contains 1 mol Na+ and 1 mol Cl- Can assume 100% dissociation of all soluble ionic compounds 98

99 Practice Problem Write the equation for the dissolution of aluminum sulfate in water. How many moles of aluminum ions are made by dissolving 1 mol aluminum sulfate? What is the total number of moles of ions made by dissolving 1 mol of aluminum sulfate? 99

100 Given: 1. Analyze Amount of solute = 1 mol Al2(SO4)3 Solvent identity = water Unknown a. moles of aluminum ions and sulfate ions Total number of moles of solute ions produced 100

101 2. Plan The coefficients in the balanced dissociation equation will tell mole relationships You can use the equation to find out the number of moles of solute ions produced 101

102 3. Compute a. b. 102

103 Write the equation for the dissolution of each of the following in water, and then determine the number of moles of each ions made as well as the total number of ions made. a. 1 mol ammonium chloride b. 1 mol sodium sulfide c. 0.5 mol barium nitrate 103

104 a. 1 mol ammonium chloride 1 mol NH4+ 1 mol Cl- 2 mol total ions 104

105 b. 1 mol sodium sulfide 2 mol Na+ 1 mol S-2 3 mol total ions 105

106 c. 0.5 mol barium nitrate 0.5 mol Ba+2 1 mol NO3-1.5 mol total ions 106

107 Precipitation Reactions Even though no compound is completely insoluble, compounds of very low solubility can be considered insoluble for practical purposes There are general guidelines to help predict whether a compound made of certain combination of ions is soluble 107

108 General Solubility Guidelines Sodium, potassium, and ammonium compounds are soluble in water. 108

109 Nitrates, acetates, and chlorates are soluble. 109

110 Most chlorides are soluble, except those of silver, mercury (I), and lead. Lead (II) chloride is soluble in hot water. 110

111 Most sulfates are soluble, except those of barium, strontium, lead, calcium, and mercury. 111

112 Most carbonates, phosphates, and silicates are insoluble, except those of sodium, potassium, and ammonium. 112

113 Most sulfides are insoluble, except those of calcium, strontium, sodium, potassium, and ammonium. 113

114 Is calcium phosphate, Ca3(PO4)2 soluble or insoluble? Not one of the exceptions, so it is insoluble Dissociation equations are not written for insoluble compounds 114

115 Guidelines also useful in predicting what will happen if 2 solutions of 2 different soluble compounds are mixed If mixing results in combination of ions that forms insoluble compound, a double-replacement and precipitation reaction will happen 115

116 Will precipitate form when solutions of ammonium sulfide and cadmium nitrate are combined? You can tell calcium nitrate is soluble from the guidelines (Most nitrates, acetates, and chlorates are soluble) You can also tell ammonium sulfide is soluble (Most sulfides are insoluble, except those of calcium, strontium, sodium, potassium, and ammonium) 116

117 117

118 To decide whether precipitate can form, must know solubilities of two possible products Ammonium nitrate soluble (Most nitrates, acetates, and chlorates are soluble) Cadmium sulfide insoluble (Most sulfides are insoluble, except those of calcium, strontium, sodium, potassium, and ammonium) 118

119 b/c one product is insoluble, doublereplacement and precipitation reaction will happen (NH4)2S (aq) + Cd(NO3)2 (aq) NH4NO3 (aq) + CdS (s) 119

120 Net Ionic Equations Reactions of ions in aqueous solution usually represented by net ionic equations instead of formula equations Net ionic equation - includes only those compounds and ions that undergo a chemical change in a reaction in an aqueous solution 120

121 121

122 To write NIE (net ionic equation) first convert chemical equation into overall ionic equation All soluble ionic compounds shown as dissociated ions in solution Precipitates shown as solids 122

123 Notice ammonium and nitrate ions appear on both sides They haven t gone through any chemical change Spectator ions - ions that do not take part in a chemical reaction and are found in solution both before and after the reaction 123

124 To change ionic equation to NIE, eliminate spectator ions Cd +2 (aq) + S -2 (aq) CdS(s) This NIE applies to any reaction in which precipitate of CdS forms 124

125 Practice Problem Identify the precipitate that forms when aqueous solutions of zinc nitrate and ammonium sulfide are combined. Write the equation for the possible double-replacement reaction. Then write the formula equation, overall ionic equation, and the net ionic equation. 125

126 1. Analyze Given: Identity of reactants: zinc nitrate and ammonium sulfide Reaction medium: aqueous solution Unknown: a. equation for the possible double-replacement reaction b. identity of the precipitate c. formula equation d. overall ionic equation e. net ionic equation 126

127 2. Plan Write the possible double-replacement reaction between Zn(NO3)2 and (NH4)2S Use the rules to determine if any of the products will precipitate Write a formula equation and overall net ionic equation Then cancel spectator ions to make net ionic equation 127

128 3. Compute a. possible double-replacement reaction: b. rules show that zinc sulfide is not soluble so will precipitate (ammonium nitrate is soluble) c. formula equation: 128

129 d. overall ionic equation: e. net ionic equation: 129

130 Will a precipitate form if solutions of potassium sulfate and barium nitrate are combined? If so, write the net ionic equation for the reaction. Yes Ba +2 (aq) + SO4-2 (aq) BaSO4(s) 130

131 Will a precipitate form if solutions of potassium nitrate and magnesium sulfate are combined? If so, write the net ionic equation for the reaction. No 131

132 Will a precipitate form if solutions of barium chloride and sodium sulfate are combined? If so, identify the spectator ions and write the net ionic equation for the reaction. Yes; Na+ and Cl- Ba +2 (aq) + SO4-2 (aq) BaSO4(s) 132

133 Write the net ionic equation for the precipitation of nickel(ii) sulfide. Ni +2 (aq) + S -2 (aq) NiS(s) 133

134 Ionization Some molecular compounds can also form ions in solution Usually polar Ionization - ions formed from solute molecules by action of the solvent (creation of ions where there were none) Ionization different from dissociation 134

135 Dissociation: ionic compounds dissolve and ions already present separate Ionization: ions formed where non existed before Ions formed are hydrated Heat released during hydration of ions gives enough energy to break covalent bonds 135

136 Extent to which solute ionizes depends on strength of bonds within molecules of solute and strength of attraction between solute and solvent molecules 136

137 HCl molecular compound that ionizes in aqueous solution Contains highly polar bond Attraction between polar HCl molecule and polar water strong enough to break HCl bond 137

138 The Hydronium Ion H+ ions from HCl attracts other molecules or ions so strongly that it doesn t normally exist alone Ionization of HCl better described as direct transfer of proton from HCl to H2O, forming H3O+ 138

139 139

140 Hydration of H+ to form hydronium is highly exothermic Energy released gives a lot of energy needed to ionize a molecular solute Many molecular compounds that ionize in aqueous solution contain hydrogen and form hydronium ion 140

141 Strong and Weak Electrolytes Substances that make ions and conduct current in solution are electrolytes HCl is one of series of compounds that has hydrogen and halogen Hydrogen halides are all molecular compounds with single polar-covalent bonds 141

142 All are gases Very soluble in water All are electrolytes HCl, HBr, HI strongly conduct current in solution HF only weakly conducts current at same concentration Strength of conduction related to ability to ionize 142

143 143

144 Strong Electrolytes HCl, HBr, HI are 100% ionized in solution Strong electrolyte - any compound whose dilute aqueous solutions conduct electricity well; this is due to the presence of all or almost all of the dissolve compound in the form of ions 144

145 Unique characteristic of strong electrolyte is that they yield only ions Ex. Ionic compound may be highly soluble in water and dissociate into ions (NaCl) Others may not dissolve much, but amount that does dissolve exists only as ions 145

146 Weak Electrolytes HF dissolves in water to give acid solution called hydrofluoric acid HF bond stronger than bonds between hydrogen and other halogens When HF dissolves, some molecule ionize Reverse reaction also happens HF(aq) + H2O(l) H3O + (aq) + F - (aq) 146

147 HF(aq) + H2O(l) H3O+ (aq) + F-(aq) Concentration of dissolved unionized HF stays high and concentration of ions stays low Weak electrolyte - any compound whose aqueous solutions conduct electricity poorly; this is due to the presence of small amount of the dissolved compound in the form of ions 147

148 Different from nonelectrolyte NO ions at all Description of electrolyte as strong or weak not related to concentration of solution Electrolytes differ in degree of ionization, not amount of solute dissolved 148

149 Colligative properties of solutions 149

150 Presence of solute affects properties of solutions Some properties not dependent on nature of dissolved substance but on how may dissolve particles are present Colligative properties - properties that depend on the concentration of solute particles but not on their identity 150

151 Vapor-Pressure Lowering Boiling and freezing point of solution different from pure solvent Nonvolatile solute raises boiling point and lowers freezing point Nonvolatile substance - one that has little tendency to become gas under existing conditions 151

152 152

153 To understand why nonvolatile solute changes boiling and freezing point, consider equilibrium vapor pressure Vapor pressure: pressure caused by molecules that have escape liquid phase to gas phase Can be thought of a measure of tendency of molecules to escape from a liquid 153

154 Aqueous solution of Pure water nonvolatile solute Addition of sucrose (nonvolatile solute) lowers concentration of water molecules at surface of liquid This lowers tendency of water molecules to leave solution and enter gas phase Vapor pressure of solution is lower than vapor pressure of pure water 154

155 Nonelectrolyte solutions of same molality have same concentration of particles Dilute solutions of same solvent and equal molality of any nonelectrolyte solute lower vapor pressure equally 155

156 Example 1 m aqueous solution of nonelectrolyte glucose lowers vapor pressure of water 5.5 x 10-4 atm at 25 1 m aqueous solution of sucrose also lowers water vapor pressure to 5.5 x 10-4 atm at

157 b/c vapor-pressure lowering depends on concentration of nonelectrolyte solute and doesn t depend on the type of solute, it is a colligative proprerty 157

158 b/c vapor pressure has been lowered, solution remains liquid over larger temperature range Lowers freezing point and raises boiling point 158

159 Can assume changes in boiling and freezing point also depend on concentration of solute They are colligative properties 159

160 Freezing-Point Depression Freezing point of 1-molal solution of any nonelectrolyte solute in water is found (by experiment) to be 1.86 lower than freezing point of water When 1 mol of nonelectrolyte solute dissolved in 1 kg water, freezing point is -1.86, not

161 When 2 mol nonelectrolyte solute dissolved in 1 kg water, freezing point is For any concentration of nonelectrolyte solute in water, decrease in freezing point can be determined using value of /m 161

162 This value, called the molal freezing-point contstant (Kf) is the freezing-point depression of the solvent in a 1-molal solution of a nonvolatile, nonelectrolyte solute 162

163 Each solvent has own characteristic molal freezing-point constant Freezing-point depression (Δtf) the difference between the freezing points of the pure solvent and a solute of nonelectrolyte in that solvent, and it is directly proportional to the molal concentration of the solution Δtf = Kfm 163

164 Kf expressed as /m Δtf = Kfm m expressed in mol solute/kg solvent (molality) 164

165 Sample Problem What is the freezing point depression of water in a solution of 17.1 g of sucrose, C12H22O11 and 200. g of water? What is the actual freezing point of the solution? 165

166 Given: 1. Analyze Solute mass and chemical formula 17.1 g C12H22O11 Solvent mass and identity g water Unknown: a. freezing-point depression b. freezing point of the solution 166

167 2. Plan 167

168 Calculate 168

169 Unknown: a. freezing-point depression b. freezing point of the solution 169

170 Practice Problem A water solution containing an unknown quantity of a nonelectrolyte solute is found to have a freezing point of 0.23 C.What is the molal concentration of the solution? 0.12 m 170

171 A solution consists of 10.3 g of the nonelectrolyte glucose, C6H12O6, dissolved in 250. g of water. What is the freezing-point depression of the solution? C 171

172 In a laboratory experiment, the freezing point of an aqueous solution of glucose is found to be C.What is the molal concentration of this solution? m 172

173 If mol of a nonelectrolyte solute are dissolved in g of ether, what is the freezing point of the solution? C 173

174 The freezing point of an aqueous solution that contains a nonelectrolyte is 9.0 C. a.what is the freezing-point depression of the solution? b.what is the molal concentration of the solution? a. 9.0 C b. 4.8 m 174

175 Boiling-Point Elevation Boiling point is temp at which vapor pressure is equal to atmospheric pressure Change in vapor pressure will cause corresponding change in boiling point 175

176 Vapor pressure of nonvolatile solution is lower than vapor pressure of pure solvent More heat will be required to raise vapor pressure of solution Boiling point of solution is higher than boiling point of pure solvent 176

177 Molal boiling-point constant (Kb) - boilingpoint elevation of the solvent in a 1-molal solution of a nonvolatile, nonelectrolyte solute b-p elevation of any nonelectrolyte in water found by experiment to be 0.51 So, molal b-p constant for water is 0.51 /m 177

178 For different solvents, b-p elevations of 1- molal solutions have different values (Table 14-2) Like freezing-point constants, values most accurate for dilute solutions 178

179 Boiling-point elevation, Δtb, difference between the boiling points of the pure solvent and a nonelectrolyte solution of that solvent, and it is directly proportional to the molal concentration of the solution B-p elevation can be calculated using Δtb = Kbm Δtb (boiling point elevation) expressed in /m and m expressed in mol solute/kg solvent

180 Sample Problem What is the boiling-point elevation of a solution made from 20.0 g of a nonelectrolyte solute and g of water? The molar mass of the solute is 62.0 g/mol. 180

181 Given: 1. Analyze solute mass = 20.0 g solute molar mass = 62.0 g/mol solvent mass and identity = g of water Unknown: boiling-point elevation 181

182 2. Plan 182

183 3. Compute 183

184 Practice Problem A solution contains 50.0 g of sucrose, C12H22O11, a nonelectrolyte, dissolved in g of water. What is the boiling-point elevation? 0.15 C 184

185 A solution contains g of sucrose, C12H22O11, a nonelectrolyte, dissolved in 250 g of water.what is the boiling point of the solution? C 185

186 If the boiling point elevation of an aqueous solution containing a nonvolatile electrolyte is 1.02 C, what is the molality of the solution? 2.0 m 186

187 The boiling point of an aqueous solution containing a nonvolatile electrolyte is C. a.what is the boiling-point elevation? b.what is the molality of the solution? a C b. 1.5 m 187

188 Electrolytes and Colligative Properties Early investigators confused by experiments where certain substances depressed freezing point or elevated boiling point more than expected 188

189 Ex. 0.1 m solution of NaCl lowers freezing point of solvent almost twice as much as 0.1 m solution of sucrose 0.1 m CaCl2 lowers f.p. almost 3 times as much as 0.1 m solution of sucrose 189

190 To understand why this happens, contrast behavior or sucrose with NaCl in aqueous solution Each sucrose dissolves to make only 1 particle in solution So 1 mol sucrose dissolves to 1 mol particles 1 mol NaCl dissolves to make 2 moles of particles (1 mol Na+ and 1 mol Cl-) 190

191 191

192 Calculated Values for Electrolyte Solutions Colligative properties depend on total concentration of solute particles regardless of their identity Electrolytes cause changes in colligative properties proportional to total molality in terms of all dissolves particles instead of formula units 192

193 What about barium nitrate, Ba(NO3)2? Each mole of barium nitrate yields 3 mol of ions in solution Ba(NO3)2(s) Ba 2+ (aq) + 2NO3 - (aq) You would expect solution of given molality to lower f.p. of solvent 3 times as much as nonelectrolytic solution of same molality 193

194 Sample Problem What is the expected change in the freezing point of water in a solution of 62.5 g of barium nitrate, Ba(NO3)2, in 1.00 kg of water? 194

195 Given: 1. Analyze solute mass and formula = 62.5 g Ba(NO3)2 solvent mass and identity = 1.00 kg water Δtf = Kfm Unknown: expected freezing-point depression 195

196 2. Plan The molality can be calculated by converting the solute mass to moles and then dividing by the number of kilograms of solvent That molality is in terms of formula units of Ba(NO3)2 and must be converted to molality in terms of dissociated ions in solution It must be multiplied by the number of moles of ions produced per mole of formula unit This adjusted molality can then be used to calculate the freezing-point depression 196

197 197

198 3. Compute 198

199 Practice Problem What is the expected freezing-point depression for a solution that contains 2.0 mol of magnesium sulfate dissolved in 1.0 kg of water? 7.4 C 199

200 What is the expected boiling-point elevation of water for a solution that contains 150 g of sodium chloride dissolved in 1.0 kg of water? 2.7 C The freezing point of an aqueous sodium chloride solution is 20.0 C. What is the molality of the solution? 5.4 m NaCl 200

201 Actual Values for Electrolyte Solutions Remember values calculated are only expected values Actual values of colligative properties for all strong electrolytes are almost what would be expected Difference between expected and actual values caused by attractive forces that exist between dissociated ions in aqueous solution 201

202 Attraction between hydrated ions in solution is small compared with those in crystalline solid However, forces of attraction don t interfere with movements of aqueous ions More concentrated solution means ions closer together and attraction is greater Also depends on charge of ions higher charge, greater attraction 202