2. 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. 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. 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”)

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

6. Various Types of Solutions
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7. Solution Composition
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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.
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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
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12. 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?”)

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

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

15. 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.
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16. Steps in the Dissolving Process
Gas
Mixing of gases
Gas
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17. 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.
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18. 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.
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19. Enthalpy (Heat) of Solution
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20. 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.
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21. 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.
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22. 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. 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
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25. 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.

26. 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. 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. 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. 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. Affecting aqueous solutions
Structure Effects:
Polarity
Pressure Effects:
Henry’s law
Temperature Effects:
Affecting aqueous solutions
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31. Structure Effects Hydrophobic (water fearing) Non-polar substances
Hydrophilic (water loving)
Polar substances
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32. 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. 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. 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. 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)

36. 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)
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37. 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.
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38. 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.
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39. The Solubilities of Several Solids as a Function of Temperature
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40. The Solubilities of Several Gases in Water
Opposite behavior is observed for most nonaqueous solvents.
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41. 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. 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. 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.
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44. 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
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45. A Solution Obeying Raoult’s Law
Linear relationship. Y=mxx+b
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46. 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

48. 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.
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49. 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. Vapor Pressure for a Solution of Two Volatile Liquids
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51. 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
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52. 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.
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53. 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.
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54. EXERCISE!
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
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 / g glucose) = mol glucose
Kg of solvent = (200.0 g)(1 kg / 1000 g) = kg water
msolute = ( mol glucose) / ( kg water) = mol/kg
ΔT = (0.51 °C·kg/mol)( mol/kg) = 0.35 °C.
The boiling point of the resulting solution is °C °C = °C.
Note: Use the red box animation to assist in explaining how to solve the problem.
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55. 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
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57. The change in temperature is directly proportional to molality (using the van’t Hoff factor).

58. 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. 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
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60. Ion Pairing
At a given instant a small percentage of the sodium and chloride ions are paired and thus count as a single particle.
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61. 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.
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62. Changes in Boiling Point and Freezing Point of Water
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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.
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64. 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.

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66. EXERCISE!
When 33.4 mg of a compound is dissolved in 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.
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67. 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.

68. 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. EXERCISE!
A plant cell has a natural concentration of 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.
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70. Suspensions of particles larger than individual ions or molecules, but too small to be settled out by gravity, are called colloids.

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

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

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

74. 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. 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).
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