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 11.2-11.3)

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)

AP Learning Objectives, Margin Notes and References 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”)

Solutions are homogeneous mixtures of two or more pure substances. In a solution, the solute is dispersed uniformly throughout the solvent.

Various Types of Solutions Copyright © Cengage Learning. All rights reserved

Solution Composition Copyright © Cengage Learning. All rights reserved

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. Copyright © Cengage Learning. All rights reserved

Copyright © Cengage Learning. All rights reserved

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 Copyright © Cengage Learning. All rights reserved

Copyright © Cengage Learning. All rights reserved

AP Learning Objectives, Margin Notes and References 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?”)

The ability of substances to form solutions depends on intermolecular forces natural tendency toward mixing

Intermolecular Forces of Attraction Any intermolecular force of attraction (Chapter 10) can be the attraction between solute and solvent molecules.

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

Steps in the Dissolving Process Gas Mixing of gases Gas Copyright © Cengage Learning. All rights reserved

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

Enthalpy (Heat) of Solution Enthalpy change associated with the formation of the solution is the sum of the ΔH values for the steps: ΔHsoln = ΔH1 + ΔH2 + ΔH3 ΔHsoln may have a positive sign (energy absorbed) or a negative sign (energy released). For a reaction to occur, ΔHsoln must be close to the sum of ΔHsolute and ΔHsolvent. Copyright © Cengage Learning. All rights reserved

Enthalpy (Heat) of Solution Copyright © Cengage Learning. All rights reserved

Enthalpy (Heat) of Hydration (Hhyd) ΔHhyd combines the term ΔH2 (for expanding the solvent) and ΔH3 (for solute-solvent interaction). Enthalpy change associated with the dispersal of gaseous solute in water. Copyright © Cengage Learning. All rights reserved

CONCEPT CHECK! Explain why water and oil (a long chain hydrocarbon) do not mix. In your explanation, be sure to address how ΔH plays a role. Oil is a mixture of nonpolar molecules that interact through London dispersion forces, which depend on molecule size. ΔH1 will be relatively large for the large oil molecules. The term ΔH3 will be small, since interactions between the nonpolar solute molecules and the polar water molecules will be negligible. However, ΔH2 will be large and positive because it takes considerable energy to overcome the hydrogen bonding forces among the water molecules to expand the solvent. Thus ΔHsoln will be large and positive because of the ΔH1 and ΔH2 terms. Since a large amount of energy would have to be expended to form an oil-water solution, this process does not occur to any appreciable extent. Copyright © Cengage Learning. All rights reserved

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.

The Energy Terms for Various Types of Solutes and Solvents ΔHsoln Outcome Polar solute, polar solvent Large Large, negative Small Solution forms Nonpolar solute, polar solvent Large, positive No solution forms Nonpolar solute, nonpolar solvent Polar solute, nonpolar solvent Copyright © Cengage Learning. All rights reserved

AP Learning Objectives, Margin Notes and References 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.

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.

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.

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.

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.

Affecting aqueous solutions Structure Effects: Polarity Pressure Effects: Henry’s law Temperature Effects: Affecting aqueous solutions Copyright © Cengage Learning. All rights reserved

Structure Effects Hydrophobic (water fearing) Non-polar substances Hydrophilic (water loving) Polar substances Copyright © Cengage Learning. All rights reserved

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.

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.

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.

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. Nonpolar (A,D,E,K) Polar (C,B)

Pressure Effects C = concentration of dissolved gas k = constant 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

A Gaseous Solute Figure 11.5 | (a) A gaseous solute in equilibrium with a solution. (b) The piston is pushed in, which increases the pressure of the gas and the number of gas molecules per unit volume. This causes an increase in the rate at which the gas enters the solution, so the concentration of dissolved gas increases. (c) The greater gas concentration in the solution causes an increase in the rate of escape. A new equilibrium is reached. Copyright © Cengage Learning. All rights reserved

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. Delta H solution refers to formation of 1M ideal solution and is not relevant to the process of dissolving a solid I a saturated solution. Delta H solution of limited use when predicting solubility. Copyright © Cengage Learning. All rights reserved

The Solubilities of Several Solids as a Function of Temperature Copyright © Cengage Learning. All rights reserved

The Solubilities of Several Gases in Water Opposite behavior is observed for most nonaqueous solvents. Copyright © Cengage Learning. All rights reserved

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

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.

An Aqueous Solution and Pure Water in a Closed Environment Figure 11.9 | An aqueous solution and pure water in a closed environment. (a) Initial stage. (b) After a period of time, the water is transferred to the solution. Copyright © Cengage Learning. All rights reserved

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

A Solution Obeying Raoult’s Law Linear relationship. Y=mxx+b Copyright © Cengage Learning. All rights reserved

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. Linear relationship. Y=mxx+b

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

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 .

Vapor Pressure for a Solution of Two Volatile Liquids Copyright © Cengage Learning. All rights reserved

Summary of the Behavior of Various Types of Solutions Interactive Forces Between Solute (A) and Solvent (B) Particles ΔHsoln 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) Positive Negative Acetone-water A  A, B  B > A  B Positive (endothermic) Ethanol-hexane Copyright © Cengage Learning. All rights reserved

Hexane (C6H14) and chloroform (CHCl3) Acetone (C3H6O) and water CONCEPT CHECK! 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? Hexane (C6H14) and chloroform (CHCl3) Acetone (C3H6O) and water Hexane (C6H14) and octane (C8H18) a) Positive deviation; Hexane is non-polar, chloroform is polar. b) Negative deviation; Both are polar, and the acetone molecules can not form stronger hydrogen bonding with itself, but it can with the water molecules. c) Ideal; Both are non-polar with similar molar masses. Copyright © Cengage Learning. All rights reserved

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 = Kbmsolute ΔT = boiling-point elevation Kb = molal boiling-point elevation constant (solvent) msolute = molality of solute Boiling point elevation is directly proportional to the concentration of the solute. Copyright © Cengage Learning. All rights reserved

EXERCISE! A solution was prepared by dissolving 25.00 g of glucose in 200.0 g water. The molar mass of glucose is 180.16 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. 100.35 °C The change in temperature is ΔT = Kbmsolute. Kb is 0.51 °C·kg/mol. To solve formsolute, use the equation m = moles of solute/kg of solvent. Moles of solute = (25.00 g glucose)(1 mol / 180.16 g glucose) = 0.1388 mol glucose Kg of solvent = (200.0 g)(1 kg / 1000 g) = 0.2000 kg water msolute = (0.1388 mol glucose) / (0.2000 kg water) = 0.6938 mol/kg ΔT = (0.51 °C·kg/mol)(0.6938 mol/kg) = 0.35 °C. The boiling point of the resulting solution is 100.00 °C + 0.35 °C = 100.35 °C. Note: Use the red box animation to assist in explaining how to solve the problem. Copyright © Cengage Learning. All rights reserved

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 = Kfmsolute ΔT = freezing-point depression Kf = molal freezing-point depression constant (solvent) msolute = molality of solute Copyright © Cengage Learning. All rights reserved

The change in temperature is directly proportional to molality (using the van’t Hoff factor).

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.

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). NaCl i = 2 KNO3 i = 2 Na3PO4 i = 4 Copyright © Cengage Learning. All rights reserved

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

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

Changes in Boiling Point and Freezing Point of Water Copyright © Cengage Learning. All rights reserved

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. Copyright © Cengage Learning. All rights reserved

Osmotic pressure is a colligative property.  = atm R = 0.08206 L atm/K mole T = Kelvin If two solutions separated by a semipermeable membrane have the same osmotic pressure, no osmosis will occur.

Osmosis 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 HERE Copyright © Cengage Learning. All rights reserved

EXERCISE! 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 The molar mass is 111 g/mol. Note: Use the red box animation to assist in explaining how to solve the problem. Copyright © Cengage Learning. All rights reserved

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

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.

EXERCISE! 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? The cell will explode (or at least expand). Thus, the cell has a higher concentration, and water will enter the cell. Note: Use the red box animation to assist in explaining how to solve the problem. Copyright © Cengage Learning. All rights reserved

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

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

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

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

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.

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