1. Solution Properties 11.1 Solution Composition
11.2 The Energies of Solution Formation
11.3 Factors Affecting Solubility
11.4 The Vapor Pressures of Solutions
11.5 Boiling-Point Elevation and Freezing-Point Depression
11.6 Osmotic Pressure
11.7 Colligative Properties of Electrolyte Solutions
11.8 Colloids

2. Various Types of Solutions
Example
State of Solution
State of Solute
State of Solvent
Air, natural gas
Gas
Vodka, antifreeze
Liquid
Brass
Solid
Carbonated water (soda)
Seawater, sugar solution
Hydrogen in platinum

3. Solution Composition

4. Molarity

5. Exercise #1
You have 1.00 mol of sugar in mL of solution. Calculate the concentration in units of molarity.
8.00 M
1.00 mol / (125.0 / 1000) = 8.00 M

6. Exercise #2
You have a 10.0 M sugar solution. What volume of this solution do you need to have 2.00 mol of sugar?
0.200 L
2.00 mol / 10.0 M = L

7. Exercise #3
Consider separate solutions of NaOH and KCl made by dissolving g of each solute in mL of solution. Calculate the concentration of each solution in units of molarity.
10.0 M NaOH
5.37 M KCl
[100.0 g NaOH / g/mol] / [250.0 / 1000] = 10.0 M NaOH
[100.0 g KCl / g/mol] / [250.0 / 1000] = 5.37 M KCl

8. Mass Percent

9. Exercise #4
What is the percent-by-mass concentration of glucose in a solution made my dissolving 5.5 g of glucose in 78.2 g of water?
6.6%
[5.5 g / (5.5 g g)] × 100 = 6.6%

10. Mole Fraction

11. Exercise #5
A solution of phosphoric acid was made by dissolving 8.00 g of H3PO4 in mL of water. Calculate the mole fraction of H3PO4. (Assume water has a density of 1.00 g/mL.)
0.0145
8.00 g H3PO4 × (1 mol / g H3PO4) = mol H3PO4
100.0 mL H2O × (1.00 g H2O / mL) × (1 mol / g H2O) = 5.55 mol H2O
Mole Fraction (H3PO4) = mol H3PO4 / [ mol H3PO mol H2O] =

12. Molality

13. Exercise #6
A solution of phosphoric acid was made by dissolving 8.00 g of H3PO4 in mL of water. Calculate the molality of the solution. (Assume water has a density of 1.00 g/mL.)
0.816 m
8.00 g H3PO4 × (1 mol / g H3PO4) = mol H3PO4
100.0 mL H2O × (1.00 g H2O / mL) × (1 kg / 1000 g) = kg H2O
Molality = mol H3PO4 / kg H2O] = m

14. 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.

15. Steps in the Dissolving Process

16. 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.

17. 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).

18. Enthalpy (Heat) of Solution

19. 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.

20. The Energy Terms for Various Types of Solutes and Solvents
Hsoln
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

21. In General
One factor that favors a process is an increase in probability of the state when the solute and solvent are mixed.
Processes that require large amounts of energy tend not to occur.
Overall, remember that “like dissolves like”.

22. Factors Affecting Solubility
Structural Effects:
Polarity – “like dissolves like”
Pressure Effects:
Henry’s law – for solubility of gases
Temperature Effects:
Affecting aqueous solutions

23. Pressure Effects 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 partial pressure of gas above the solution.

24. A Gaseous Solute

25. 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.
Solubility of a gas in solvent typically decreases with increasing temperature.

26. The Solubilities of Several Solids as a Function of Temperature

27. The Solubilities of Several Gases in Water

28. Ideal Solution: One that obeys Raoult’s Law

29. Ideal Solutions Consisting of Two Volatile Liquids
Two volatile Liquids form ideal solution if:
they are structurally very similar, and
molecular interactions between nonidentical molecules were relatively similar to identical molecules.
The vapor of each liquid obeys Raoult’s Law:
PA = XAPoA; PB = XBPoB
PT = PA + PB = XAPoA + XBPoB
(X : mole fraction; Po : vapor pressure of pure liquid)

30. Summary of the Behavior of Various Types of Solutions of Two Volatile Liquids
Interactive Forces Between Solute (A) and Solvent (B) Particles
Hsoln
T 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

31. Vapor Pressure for a Solution of Two Volatile Liquids

32. Laboratory Fractional Distillation Apparatus

33. Fractional Distillation Towers in Oil Refinaries

34. Refined Crude Oil Mixtures

35. 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)
Ethyl alcohol (C2H5OH) 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 ethyl alcohol molecules can form stronger hydrogen bonding with the water molecules than it can with other alcohol molecules.
c) Ideal; Both are non-polar with similar molar masses.

36. Exercise #7 (a) The mole fractions of benzene and toluene;
A solution of benzene (C6H6) and toluene (C7H8) contains 50.0% benzene by mass. The vapor pressures of benzene and pure toluene at 25oC are 94.2 torr and 28.4 torr, respectively. Assuming ideal behavior, calculate the following:
(a) The mole fractions of benzene and toluene;
(b) The vapor pressure of each component in the mixture, and the total vapor pressure above the solution.
(c) The composition of the vapor in mole percent.

37. Exercise #8
A solution composed of 24.3 g acetone (CH3COCH3) and 39.5 g of carbon disilfide (CS2) has a measured vapor pressure of 645 torr at 35oC.
(a) Is the solution ideal or nonideal?
(b) If not, does it deviate positively or negatively from Raoult’s law?
(c) What can you say about the relative strength of carbon disulfide-acetone interactions compared to the acetone-acetone and carbon disulfide-carbon disulfide interaction?
(Vapor pressures at 35oC of pure acetone and pure carbon disulfide are 332 torr and 515 torr, respectively.)

38. An Aqueous Solution and Pure Water in a Closed Environment

39. Liquid/Vapor Equilibrium

40. Vapor Pressure Lowering: Addition of a Solute

41. Vapor Pressures of Solutions of Nonvolatile Solutes
Nonvolatile solute lowers the vapor pressure of solvent.
Raoult’s Law:
Psoln = vapor pressure of solution
solv = mole fraction of solvent
= vapor pressure of pure solvent

42. Colligative Properties
Lowering of solvent vapor pressure
Freezing-point depression
Boiling-point elevation
Osmotic pressure
Colligative properties depend only on the number, not on the identity, of the solute particles in an ideal solution.

43. Lowering of Solvent Vapor Pressure
The presence of nonvolatile solute particles lowers the number of solvent molecules in the vapor that is in equilibrium with the solution.
The solvent vapor pressure is lowered;
Assuming ideal behavior, the lowering of vapor pressure is proportional to the mole fraction of solute:
DP = Xsolute.Posolvent (for nonelectrolytes)
= iXsolute.Posolvent (for electrolytes)
(i is the van’t Hoff’s factor, which approximately relates to the number of ions per formula unit of the compound)

44. Changes in Boiling Point and Freezing Point of Water

45. 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 (for nonelectrolytes)
= iKfmsolute (for electrolytes)
ΔT = freezing-point depression
Kf = freezing-point depression constant
msolute = molality of solute

46. Freezing Point Depression: Solid/Liquid Equilibrium

47. Freezing Point Depression: Addition of a Solute

48. Freezing Point Depression: Solid/Solution Equilibrium

49. Boiling-Point Elevation
Nonvolatile solute elevates the boiling point of the solvent.
ΔT = Kbmsolute
ΔT = boiling-point elevation
Kb = boiling-point elevation constant
msolute = molality of solute

50. Boiling Point Elevation: Liquid/Vapor Equilibrium

51. Boiling Point Elevation: Addition of a Solute

52. Boiling Point Elevation: Solution/Vapor Equilibrium

53. Exercise #9
What mass of ethylene glycol (C2H6O2), in grams, must be added to 1.50 kg of water to produce a solution that boils at 105oC?
(Boiling point elevation constant for water is Kb = 0.512oC/m)
At what temperature will the solution freeze?
(Freezing point depression constant for water is Kf = 1.86oC/m)

54. Osmotic Pressure = MRT = osmotic pressure (atm)
Osmosis – flow of solvent into the solution through a semipermeable membrane.
= MRT
= osmotic pressure (atm)
M = molarity of the solution
R = gas law constant
T = temperature (Kelvin)

56. Osmosis

58. van’t Hoff Factor, i
The relationship between the moles of solute dissolved and the moles of particles in solution is usually expressed as:

59. Modified Equations for the Colligative Properties of Electrolytes

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.

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.

62. Exercise #10
A solution was prepared by dissolving g 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.

63. Exercise #11
You take 20.0 g of a sucrose (C12H22O11) and NaCl mixture and dissolve it in 1.0 L of water. The freezing point of this solution is found to be °C. Assuming ideal behavior, calculate the mass percent composition of the original mixture, and the mole fraction of sucrose in the original mixture.
72.8% sucrose and 27.2% sodium chloride;
mole fraction of the sucrose is 0.313
The solution is 72.8% sucrose and 27.2% sodium chloride. The mole fraction of the sucrose is To solve this problem, the students must assume that i = 2 for NaCl.
Note: Use the red box animation to assist in explaining how to solve the problem.

64. Exercise #12
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/expand, shrivel, or do nothing?
The cell will explode (or at least expand). The concentration of the solution is m. 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.

65. Exercise #13
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.

66. 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

67. Colloidal Mixtures A suspension of tiny particles in some medium.
Tyndall effect – scattering of light by particles.
Suspended particles are single large molecules or aggregates of molecules or ions ranging in size from 1 to 1000 nm.

68. Scattering of Light by Colloid Particles

69. Tyndall Effect of Colloidal Mixture

70. Tyndall Effect of Morning Mist

71. Types of Colloids

72. Micelle – A Colloidal Suspension

73. Micelle in Soap Bubbles

74. Coagulation Destruction of a colloid.
Usually accomplished either by heating or by adding an electrolyte.