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Chapter 11 Properties of Solutions AP*. AP Learning Objectives  LO 1.16 The student can design and/or interpret the results of an experiment regarding.

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Presentation on theme: "Chapter 11 Properties of Solutions AP*. AP Learning Objectives  LO 1.16 The student can design and/or interpret the results of an experiment regarding."— Presentation transcript:

1 Chapter 11 Properties of Solutions AP*

2 AP Learning Objectives  LO 1.16 The student can design and/or interpret the results of an experiment regarding the absorption of light to determine the concentration of an absorbing species in a solution. (Sec 11.1)  LO 2.8 The student can draw and/or interpret representations of solutions that show the interactions between the solute and solvent. (Sec 11.1)  LO 2.9 The student is able to create or interpret representations that link the concept of molarity with particle views of solutions. (Sec 11.2)  LO 2.14 The student is able to apply Coulomb’s Law qualitatively (including using representations) to describe the interactions of ions, and the attractions between ions and solvents to explain the factors that contribute to the solubility of ionic compounds. (Sec 11.2)  LO 2.15 The student is able to explain observations regarding the solubility of ionic solids and molecules in water and other solvents on the basis of particle views that include intermolecular interactions and entropic effects. (Sec )

3 AP Learning Objectives  LO 5.10 The student can support the claim about whether a process is a chemical or physical change (or may be classified as both) based on whether the process involves changes in intramolecular versus intermolecular interactions. (Sec 11.2)  LO 6.24 The student can analyze the enthalpic and entropic changes associated with the dissolution of a salt, using particulate level interactions and representations. (Sec 11.2)

4 Section 11.1 Solution Composition AP Learning Objectives, Margin Notes and References  Learning Objectives  LO 1.16 The student can design and/or interpret the results of an experiment regarding the absorption of light to determine the concentration of an absorbing species in a solution.  LO 2.8 The student can draw and/or interpret representations of solutions that show the interactions between the solute and solvent.  AP Margin Notes  Spectral analysis is a common method for analyzing the composition of a solution. See Appendix 3 “Spectral Analysis” for a discussion of the Beer-Lambert law.  Additional AP References  LO 1.16 (see APEC #2, “The Percentage of Copper in Brass”)

5 Section 11.1 Solution Composition  Solutions are homogeneous mixtures of two or more pure substances.  In a solution, the solute is dispersed uniformly throughout the solvent.

6 Section 11.1 Solution Composition Copyright © Cengage Learning. All rights reserved 6 Various Types of Solutions

7 Section 11.1 Solution Composition Copyright © Cengage Learning. All rights reserved 7 Solution Composition

8 Section 11.1 Solution Composition Copyright © Cengage Learning. All rights reserved 8  Molarity (M) depends n the volume of solution, so it will change slightly with temperature.  Molality (m) is independent of temperature because it depends only on mass.

9 Section 11.1 Solution Composition Copyright © Cengage Learning. All rights reserved 9

10 Section 11.1 Solution Composition Copyright © Cengage Learning. All rights reserved 10  Normality (N) = number of equivalents liters of solution  The definition of equivalent depends on the reaction type.  Acid/ Base – the mass of acid or base that can furnish or accept exactly 1 mole of protons  Redox – the quantity of oxidizing or reducing agent that can furnish or accept exactly 1 mole of electrons

11 Section 11.1 Solution Composition Copyright © Cengage Learning. All rights reserved 11

12 Section 11.2 The Energies of Solution Formation AP Learning Objectives, Margin Notes and References  Learning Objectives  LO 2.9 The student is able to create or interpret representations that link the concept of molarity with particle views of solutions.  LO 2.14 The student is able to apply Coulomb’s Law qualitatively (including using representations) to describe the interactions of ions, and the attractions between ions and solvents to explain the factors that contribute to the solubility of ionic compounds.  LO 2.15 The student is able to explain observations regarding the solubility of ionic solids and molecules in water and other solvents on the basis of particle views that include intermolecular interactions and entropic effects.  LO 5.10 The student can support the claim about whether a process is a chemical or physical change (or may be classified as both) based on whether the process involves changes in intramolecular versus intermolecular interactions.  LO 6.24 The student can analyze the enthalpic and entropic changes associated with the dissolution of a salt, using particulate level interactions and representations.  Additional AP References  LO 5.10 (see Appendix 7.6, “Distinguishing between Chemical and Physical Changes at the Molecular Level”)  LO 6.24 (see Appendix 7.7, “Intermolecular Forces and Thermodynamics: Why Aren’t All Ionic Solids Soluble in Water?”)

13 Section 11.2 The Energies of Solution Formation The ability of substances to form solutions depends on  intermolecular forces  natural tendency toward mixing

14 Section 11.2 The Energies of Solution Formation Intermolecular Forces of Attraction Any intermolecular force of attraction (Chapter 10) can be the attraction between solute and solvent molecules.

15 Section 11.2 The Energies of Solution Formation Formation of a Liquid Solution 1.Separating the solute into its individual components (expanding the solute). 2.Overcoming intermolecular forces in the solvent to make room for the solute (expanding the solvent). 3.Allowing the solute and solvent to interact to form the solution. Copyright © Cengage Learning. All rights reserved 15

16 Section 11.2 The Energies of Solution Formation Steps in the Dissolving Process Copyright © Cengage Learning. All rights reserved 16 Gas

17 Section 11.2 The Energies of Solution Formation Steps in the Dissolving Process  Steps 1 and 2 require energy, since forces must be overcome to expand the solute and solvent.  Step 3 usually releases energy.  Steps 1 and 2 are endothermic, and step 3 is often exothermic. Copyright © Cengage Learning. All rights reserved 17

18 Section 11.2 The Energies of Solution Formation Enthalpy (Heat) of Solution  Enthalpy change associated with the formation of the solution is the sum of the ΔH values for the steps: ΔH soln = ΔH 1 + ΔH 2 + ΔH 3  ΔH soln may have a positive sign (energy absorbed) or a negative sign (energy released).  For a reaction to occur, ΔH soln must be close to the sum of ΔH solute and ΔH solvent. Copyright © Cengage Learning. All rights reserved 18

19 Section 11.2 The Energies of Solution Formation Enthalpy (Heat) of Solution Copyright © Cengage Learning. All rights reserved 19

20 Section 11.2 The Energies of Solution Formation Enthalpy (Heat) of Hydration (  H hyd ) Copyright © Cengage Learning. All rights reserved 20  ΔH hyd combines the term ΔH 2 (for expanding the solvent) and ΔH 3 (for solute-solvent interaction).  Enthalpy change associated with the dispersal of gaseous solute in water.

21 Section 11.2 The Energies of Solution Formation Explain why water and oil (a long chain hydrocarbon) do not mix. In your explanation, be sure to address how ΔH plays a role. Copyright © Cengage Learning. All rights reserved 21 CONCEPT CHECK!

22 Section 11.2 The Energies of Solution Formation Natural Tendency toward Mixing  Mixing of gases is a spontaneous process.  Each gas acts as if it is alone to fill the container.  Mixing causes more randomness in the position of the molecules, increasing a thermodynamic quantity called entropy.  The formation of solutions is favored by the increase in entropy that accompanies mixing.

23 Section 11.2 The Energies of Solution Formation The Energy Terms for Various Types of Solutes and Solvents Copyright © Cengage Learning. All rights reserved 23 ΔH1ΔH1 ΔH2ΔH2 ΔH3ΔH3 ΔH soln Outcome Polar solute, polar solventLarge Large, negativeSmallSolution forms Nonpolar solute, polar solventSmallLargeSmallLarge, positiveNo solution forms Nonpolar solute, nonpolar solvent Small Solution forms Polar solute, nonpolar solventLargeSmall Large, positiveNo solution forms

24 Section 11.2 The Energies of Solution Formation

25 Section 11.3 Factors Affecting Solubility AP Learning Objectives, Margin Notes and References  Learning Objectives  LO 2.15 The student is able to explain observations regarding the solubility of ionic solids and molecules in water and other solvents on the basis of particle views that include intermolecular interactions and entropic effects.

26 Section 11.3 Factors Affecting Solubility Aqueous Solution vs. Chemical Reaction Just because a substance disappears when it comes in contact with a solvent, it does not mean the substance dissolved. It may have reacted, like nickel with hydrochloric acid.

27 Section 11.3 Factors Affecting Solubility Opposing Processes  The solution-making process and crystallization are opposing processes.   When the rate of the opposing processes is equal, additional solute will not dissolve unless some crystallizes from solution. This is a saturated solution.  If we have not yet reached the amount that will result in crystallization, we have an unsaturated solution.

28 Section 11.3 Factors Affecting Solubility  Solubility is the maximum amount of solute that can dissolve in a given amount of solvent at a given temperature.  Saturated solutions have that amount of solute dissolved.  Unsaturated solutions have any amount of solute less than the maximum amount dissolved in solution.  Surprisingly, there is one more type of solution.

29 Section 11.3 Factors Affecting Solubility Supersaturated Solutions  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.  These are uncommon solutions.

30 Section 11.3 Factors Affecting Solubility  Structure Effects:  Polarity  Pressure Effects:  Henry’s law  Temperature Effects:  Affecting aqueous solutions Copyright © Cengage Learning. All rights reserved 30

31 Section 11.3 Factors Affecting Solubility Structure Effects  Hydrophobic (water fearing)  Non-polar substances  Hydrophilic (water loving)  Polar substances Copyright © Cengage Learning. All rights reserved 31

32 Section 11.3 Factors Affecting Solubility Solute–Solvent Interactions  Simply put: “Like dissolves like.”  That does not explain everything!  The stronger the solute–solvent interaction, the greater the solubility of a solute in that solvent.  The gases in the table only exhibit dispersion force. The larger the gas, the more soluble it will be in water.

33 Section 11.3 Factors Affecting Solubility Organic Molecules in Water  Polar organic molecules dissolve in water better than nonpolar organic molecules.  Hydrogen bonding increases solubility, since C–C and C–H bonds are not very polar.

34 Section 11.3 Factors Affecting Solubility Liquid/Liquid Solubility  Liquids that mix in all proportions are miscible.  Liquids that do not mix in one another are immiscible.  Because hexane is nonpolar and water is polar, they are immiscible.

35 Section 11.3 Factors Affecting Solubility Solubility and Biological Importance  Fat-soluble vitamins (like vitamin A) are nonpolar; they are readily stored in fatty tissue in the body.  Water-soluble vitamins (like vitamin C) need to be included in the daily diet.

36 Section 11.3 Factors Affecting Solubility Pressure Effects  Little effect on solubility of solids or liquids  Henry’s law:C = kP C = concentration of dissolved gas k = constant P =partial pressure of gas solute above the solution  Amount of gas dissolved in a solution is directly proportional to the pressure of the gas above the solution.  Obeyed most accurately dilute solutions of gases that do not dissociate or react with the solvent. (ex. HCl does not obey) Copyright © Cengage Learning. All rights reserved 36

37 Section 11.3 Factors Affecting Solubility A Gaseous Solute Copyright © Cengage Learning. All rights reserved 37

38 Section 11.3 Factors Affecting Solubility Temperature Effects (for Aqueous Solutions)  Although the solubility of most solids in water increases with temperature, the solubilities of some substances decrease with increasing temperature.  Predicting temperature dependence of solubility is very difficult (best to determine experimentally).  Solubility of a gas in solvent typically decreases with increasing temperature. Copyright © Cengage Learning. All rights reserved 38

39 Section 11.3 Factors Affecting Solubility The Solubilities of Several Solids as a Function of Temperature Copyright © Cengage Learning. All rights reserved 39

40 Section 11.3 Factors Affecting Solubility The Solubilities of Several Gases in Water Copyright © Cengage Learning. All rights reserved 40

41 Section 11.4 The Vapor Pressures of Solutions Colligative Properties  Colligative properties depend only on the quantity, not on the identity of the solute particles.  Among colligative properties are:  Vapor-pressure lowering  Boiling-point elevation  Freezing-point depression  Osmotic pressure

42 Section 11.4 The Vapor Pressures of Solutions Because of solute–solvent intermolecular attraction, higher concentrations of nonvolatile solutes make it harder for solvent to escape to the vapor phase. Therefore, the vapor pressure of a solution is lower than that of the pure solvent.

43 Section 11.4 The Vapor Pressures of Solutions An Aqueous Solution and Pure Water in a Closed Environment Copyright © Cengage Learning. All rights reserved 43

44 Section 11.4 The Vapor Pressures of Solutions Vapor Pressures of Solutions  Nonvolatile solute lowers the vapor pressure of a solvent.  Raoult’s Law: P soln =observed vapor pressure of solution solv =mole fraction of solvent =vapor pressure of pure solvent Copyright © Cengage Learning. All rights reserved 44

45 Section 11.4 The Vapor Pressures of Solutions A Solution Obeying Raoult’s Law Copyright © Cengage Learning. All rights reserved 45

46 Section 11.4 The Vapor Pressures of Solutions The phenomenon of lowering vapor pressure gives an experimental way to determine molar mass. Vapor pressure depression can also be used to characterize solutes. Ex. NaCl lowers the VP almost 2x as expected due to 2 ions per formula unit.

47 Section 11.4 The Vapor Pressures of Solutions

48 Section 11.4 The Vapor Pressures of Solutions Nonideal Solutions  Liquid-liquid solutions where both components are volatile.  Modified Raoult’s Law:  Nonideal solutions behave ideally as the mole fractions approach 0 and 1 or if the solute and solvent have similar interactions. Copyright © Cengage Learning. All rights reserved 48

49 Section 11.4 The Vapor Pressures of Solutions When a solution contains two volatile components, both contribute to the total vapor pressure. Note that in this case the solution contains equal numbers of the components and, but the vapor contains more than. This means that component is more volatile (has a higher vapor pressure as a pure liquid) than component.

50 Section 11.4 The Vapor Pressures of Solutions Vapor Pressure for a Solution of Two Volatile Liquids Copyright © Cengage Learning. All rights reserved 50

51 Section 11.4 The Vapor Pressures of Solutions Summary of the Behavior of Various Types of Solutions Copyright © Cengage Learning. All rights reserved 51 Interactive Forces Between Solute (A) and Solvent (B) Particles ΔH soln ΔT for Solution Formation Deviation from Raoult’s Law Example A  A, B  B  A  B Zero None (ideal solution) Benzene- toluene A  A, B  B < A  B Negative (exothermic) PositiveNegative Acetone- water A  A, B  B > A  B Positive (endothermic) NegativePositive Ethanol- hexane

52 Section 11.4 The Vapor Pressures of Solutions For each of the following solutions, would you expect it to be relatively ideal (with respect to Raoult’s Law), show a positive deviation, or show a negative deviation? a)Hexane (C 6 H 14 ) and chloroform (CHCl 3 ) b)Acetone (C 3 H 6 O) and water c)Hexane (C 6 H 14 ) and octane (C 8 H 18 ) Copyright © Cengage Learning. All rights reserved 52 CONCEPT CHECK!

53 Section 11.5 Boiling-Point Elevation and Freezing-Point Depression Boiling-Point Elevation  Since vapor pressures are lowered for solutions, it requires a higher temperature to reach atmospheric pressure.  Hence, nonvolatile solute elevates the boiling point of the solvent.  ΔT = K b m solute ΔT= boiling-point elevation K b = molal boiling-point elevation constant (solvent) m solute = molality of solute Copyright © Cengage Learning. All rights reserved 53

54 Section 11.5 Boiling-Point Elevation and Freezing-Point Depression A solution was prepared by dissolving g of glucose in g water. The molar mass of glucose is g/mol. What is the boiling point of the resulting solution (in °C)? Glucose is a molecular solid that is present as individual molecules in solution °C Copyright © Cengage Learning. All rights reserved 54 EXERCISE!

55 Section 11.5 Boiling-Point Elevation and Freezing-Point Depression Freezing-Point Depression  When a solute is dissolved in a solvent, the freezing point of the solution is lower than that of the pure solvent.  ΔT = K f m solute ΔT = freezing-point depression K f = molal freezing-point depression constant (solvent) m solute = molality of solute Copyright © Cengage Learning. All rights reserved 55

56 Section 11.5 Boiling-Point Elevation and Freezing-Point Depression

57 Section 11.5 Boiling-Point Elevation and Freezing-Point Depression  The change in temperature is directly proportional to molality (using the van’t Hoff factor).

58 Section 11.5 Boiling-Point Elevation and Freezing-Point Depression The van’t Hoff Factor (i)  What is the van’t Hoff factor?  It takes into account dissociation in solution!  Theoretically, we get 2 particles when NaCl dissociates. So, i = 2.  In fact, the amount that particles remain together is dependent on the concentration of the solution.

59 Section 11.7 Colligative Properties of Electrolyte Solutions Examples  The expected value for i can be determined for a salt by noting the number of ions per formula unit (assuming complete dissociation and that ion pairing does not occur).  NaCli = 2  KNO 3 i = 2  Na 3 PO 4 i = 4 Copyright © Cengage Learning. All rights reserved 59

60 Section 11.7 Colligative Properties of Electrolyte Solutions Ion Pairing  At a given instant a small percentage of the sodium and chloride ions are paired and thus count as a single particle. Copyright © Cengage Learning. All rights reserved 60

61 Section 11.7 Colligative Properties of Electrolyte Solutions Ion Pairing  Ion pairing is most important in concentrated solutions.  As the solution becomes more dilute, the ions are farther apart and less ion pairing occurs.  Ion pairing occurs to some extent in all electrolyte solutions.  Ion pairing is most important for highly charged ions. Copyright © Cengage Learning. All rights reserved 61

62 Section 11.5 Boiling-Point Elevation and Freezing-Point Depression Changes in Boiling Point and Freezing Point of Water Copyright © Cengage Learning. All rights reserved 62

63 Section 11.6 Osmotic Pressure Copyright © Cengage Learning. All rights reserved 63  Some substances form semipermeable membranes, allowing some smaller particles to pass through, but blocking larger particles.  The net movement of solvent molecules from solution of low to high concentration across a semipermeable membrane is osmosis. The applied pressure to stop it is osmotic pressure.

64 Section 11.6 Osmotic Pressure  Osmotic pressure is a colligative property.  = atm R = L atm/K mole T = Kelvin  If two solutions separated by a semipermeable membrane have the same osmotic pressure, no osmosis will occur.

65 Section 11.6 Osmotic Pressure Osmosis Copyright © Cengage Learning. All rights reserved 65 To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERECLICK HERE

66 Section 11.6 Osmotic Pressure When 33.4 mg of a compound is dissolved in 10.0 mL of water at 25°C, the solution has an osmotic pressure of 558 torr. Calculate the molar mass of this compound. 111 g/mol Copyright © Cengage Learning. All rights reserved 66 EXERCISE!

67 Section 11.6 Osmotic Pressure Types of Solutions & Osmosis 1)Isotonic solutions: Same osmotic pressure; solvent passes the membrane at the same rate both ways. 2)Hypotonic solution: Lower osmotic pressure; solvent will leave this solution at a higher rate than it enters with. 3)Hypertonic solution: Higher osmotic pressure; solvent will enter this solution at a higher rate than it leaves with.

68 Section 11.6 Osmotic Pressure Osmosis and Blood Cells  Red blood cells have semipermeable membranes.  If stored in a hypertonic solution, they will shrivel as water leaves the cell; this is called crenation.  If stored in a hypertonic solution, they will grow until they burst; this is called hemolysis.

69 Section 11.5 Boiling-Point Elevation and Freezing-Point Depression A plant cell has a natural concentration of 0.25 m. You immerse it in an aqueous solution with a freezing point of –0.246°C. Will the cell explode, shrivel, or do nothing? Copyright © Cengage Learning. All rights reserved 69 EXERCISE!

70 Section 11.8 Colloids Suspensions of particles larger than individual ions or molecules, but too small to be settled out by gravity, are called colloids.

71 Section 11.8 Colloids Tyndall Effect  Colloidal suspensions can scatter rays of light. (Solutions do not.)  This phenomenon is known as the Tyndall effect.

72 Section 11.8 Colloids Colloids and Biomolecules Some molecules have a polar, hydrophilic (water- loving) end and a nonpolar, hydrophobic (water- fearing) end.

73 Section 11.8 Colloids Stabilizing Colloids by Adsorption  Ions can adhere to the surface of an otherwise hydrophobic colloid.  This allows it to interact with aqueous solution.

74 Section 11.8 Colloids Colloids in Biological Systems  Colloids can aid in the emulsification of fats and oils in aqueous solutions.  An emulsifier causes something that normally does not dissolve in a solvent to do so.

75 Section 11.8 Colloids Coagulation  Destruction of a colloid.  Usually accomplished either by heating (increase velocity of molecules causing them to collide with enough energy to break the in barrier) or by adding an electrolyte which neutralizes the adsorbed ion barriers (formation of deltas). Copyright © Cengage Learning. All rights reserved 75


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