1. PROPERTIES OF SOLUTIONS
Chapter Eleven:
PROPERTIES OF SOLUTIONS

2. Mixtures and Aqueous Solutions
We use solutions all the time
What are they?
Where do we find them?
How do we describe them?

3. Soluble versus insoluble
Some solids are soluble in water, ie: table salt, NaCl. Soluble means: able to be dissolved.
Soluble ionic solids (made of cation and anion) dissociate into their ions in water.
Soluble covalent solids (like sugar) dissolve because they are relatively polar.
In a solution, the dissolved particles cannot be easily seen or separated from the solution.
Alloys are solutions of metals!

4. Parts of a solution
The dissolving medium is the solvent (what does the dissolving…the dissolver)
The dissolved substance is the solute (what gets dissolved…the dissolvey)
The solute and solvent together form the solution.
Solvents and solutes can be any phase.
solution

5. Special types of mixtures - Suspensions
mixtures where the solutes particles are very large, so they don’t completely dissolve into their solvent.
Solute particles will settle out of the solution if left undisturbed. – this creates two phases.
Muddy water and Italian salad dressing are good examples of suspensions.

6. Special types of mixtures - Colloids
mixtures where the solute particle is smaller than particles in a suspension, but not small enough to dissolve.
Colloids have two phases:
Dispersed phase – the solute
Dispersing medium – the solvent.
Mayonnaise and hair gel are good examples of colloids.
There are 7 types of colloids.

7. 7 Types of Colloids Two groups of colloids:
Colloid Type
Phases
Example
Sol
solid in liquid, liquid substance
Paint
Gel
solid in liquid, solid substance
Gelatin
Foam
gas in liquid
Whipped cream
Liquid Emulsion
liquid in liquid
Milk, mayonnaise
Solid Emulsion
liquid in solid
Cheese, butter
Solid Aerosol
solid in gas
Smoke
Liquid Aerosol
liquid in gas
Clouds, fog
Two groups of colloids:
Heterogeneous colloids – two phases are clearly seen
Homogeneous colloids – appears to be one phase

8. The Tyndall Effect
John Tyndall, Brittish, c1860
The Tyndall effect allows us to distinguish between solutions, colloids, and suspensions.
It works by shining a beam of light into the mixture. If…
Light doesn’t pass through
the mixture is a suspension or a heterogeneous colloid.
Light passes through unobstructed
the mixture is a solution.
Light passes, but the beam can be seen in the mixture
the mixture is a homogeneous colloid

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

10. Electrolytes Electrolytes: Solutions that conduct electricity.
Ionic solutions are electrolytes.
Covalent solutions are nonelectrolytes.
Is saltwater (NaCl in water) an electrolyte?
Is sugar water (C6H12O6 in water) an electrolyte?
Conductivity tester
can tell us if a solution is an electrolyte, and sometimes, how strong an electrolyte is.

11. Supersaturated Solution
solution has as much solute as it will allow (equal to solubility)
Unsaturated Solution
more solute can dissolve into solution (less than solubility)
Supersaturated Solution
too much solute in solution-some will fall out (more than solubility)
We express the quantitative amount of solute in a solution with concentration …

12. Solution Composition

13. Expressing Concentration of Solutions
A.    Molarity (M) – Moles of solute per liter of solution. Unit - mol/l
B.    Molality (m) – Moles of solute per kilogram of solvent. Unit – mol/kg
C.    Mole Fraction (X) – Moles of solute divided by total moles of all components in solution. Unit - none
D.    Percent Composition (by mass) – mass of solute divided by total mass of solution times 100. Unit - %
E.     Normality – Moles of equivalents (ions) per liter of solution. Unit – moleq/l
F.     Parts per million (ppm) / parts per billion (ppb)

14. Solution Composition 1. Molarity (M) = 2. Mass (weight) percent =
3. Mole fraction (A) =
4. Molality (m) =

15. Concentration - Molarity
the quantitative amount of solute present in a solution
Molarity (M) – moles/liter
number of moles solute in liters of solution

16. Try these Molarity questions
What is the concentration [in Molarity] when 3.0 moles of NaCl are dissolved in 2.0 Liters of solution?
How much (in liters) of a 0.1 M solution do you need to get 2 moles of solute?
How many moles of NaOH are present in 300mL of a 1M solution?
How many grams of HCl are found in 100. mL of a 2.0 M solution?
1.5 M “molar”
20 L
.3 moles
7.3 grams

17. Concentration - Molality
Molality (m) – moles/kilogram – number of moles solute in kilogram of solvent.
molality is used less often, but is important when we discuss colligative properties
Remember, where Molarity is “per liter solution”, molality is “per kilogram of solvent”

18. Try these molality questions
What is the concentration in molality when 2 moles of NaCl are dissolved in 4kg of water?
How many moles of solute are present in 1 kg of a 3 m solution?
What mass of water do you need to add to 4 moles of NaCl to make a 2 m solution?
What is the molality of a solution created by dissolving 3.5 moles methanol in 340g of CCl4.
.5 m “molal”
3 moles
2 kilograms
10.3 m

19. React 1
Consider two aqueous solutions, A and B (which contain different solutes). Is it possible for solution A to have a greater concentration than solution B in terms of mass percent, but a smaller concentration in terms of molarity? If no, explain why not. If yes, provide an example. What if the solutions contained the same solute?

20. Solution Preparation By solid dissolving:
1. calculate how many grams are needed to create our volume of our desired molarity solution
2. weigh out that mass, and add it to a volumetric flask
3. add some water and allow to dissolve
4. add water to the desired volume
By dilution of a standard solution:
1. use the relationship M1V1=M2V2
2. calculate volume of standard molarity solution to use to get desired volume of desired molarity solution.

21. 11.2 The Energies of Solution Formation

22. Formation of Solutions
A.    Spontaneous Process- Solution forms without input of outside energy
a.  Changes in energy – solute-solute and solvent-
solvent attractions to solvent- solute attractions
       i.   Ideal solutions – no energy change because the attractions
between solvent molecules (ions) are the same as the
attractions between solute and solvent molecules (ions).
       ii.   Processes in which the energy content of the system
tends to decrease occur spontaneously.
b. Changes in disorder of the components.
         i.   Changes in which the disorder of the system (entropy)
increases occur spontaneously.

23. The Steps in the Dissolving Process

24. The Energy Terms for Various Types of Solutes and Solvents
Hsoln
Outcome
Polar solvent, polar solute
Large
Large, negative
Small
Solution forms
Polar solvent, nonpolar solute
Large, positive
No solution forms
Nonpolar solvent, nonpolar solute
Nonpolar solvent, polar solute

25. React 2
Refer to Figure 11.1 in your text when answering the following. In your explanations, estimate the H values of the various processes for solution formation (large or small, positive or negative), order these according to magnitude (absolute values), and explain.
a) Explain why water and oil (a long chain hydrocarbon) do not mix.
b) Explain why two different oils (oils are long chain hydrocarbons) will mix.
c) Explain why water and another polar liquid will mix.

26. Think About It What does the term “hydrophobic” mean?
Oil is sometimes termed “hydrophobic”. Why? Also, why is the term “hydrophobic” not really correct?

27. 11.3 Factors Affecting Solubility

28. A. Nature of Solute and Solvent – Likes dissolve likes
11-3 Solubility factors
  A.  Nature of Solute and Solvent – Likes dissolve likes
a.  Dissolution of ionic compounds – Occurs due to ion-dipole attraction
b. Dissolution of molecular electrolytes – Ionize in water and have ion-dipole attractions.
C.  Dissolution of non-polar solutes – Size matters – small non-polar solute molecules are soluble in water. Large non-polar molecules can only dissolve in non-polar solvents.
B.  Temperature increase makes solid solutes more soluble (in general) and gas solutes less soluble.
C.    Pressure increases do not affect solubility of liquid and solid solutes but increases the solubility of gases (Why?). Henry’s Law C=kP
D.     Rate of solution for solid solutes
a.  Temperature
b. Stirring
c.  Surface area

29. Like Dissolves Like!
Some solvents are polar, like magnets, having partial negative and partial positive ends. (H2O)
Other solvents are nonpolar, having no “+” “-” poles
Polar solutes tend to dissolve well in polar solvents…
Nonpolar solutes tend to dissolve well into nonpolar solvents.
“Like dissolves like”
Water is very polar. Does it dissolve polar substances or non polar substance?

30. Factors Affecting Solubility
Structural Effects
Pressure Effects
Temperature Effects

31. Solid-Liquid and Gas-Liquid solubility with temperature

32. Gasses in liquids
In addition to cold temperatures, high pressures increase solubility of gasses in liquids.
Henry’s Law:
solubility of a gas in a liquid increases with increasing pressure of that gas above the liquid.

33. A Gaseous Solute

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

35. The Solubilities of Several Gases in Water

36. Colligative Properties
“Colligative” means depends on amount.
A colligative property of a solution depends on the amount of solute dissolved in solution.
Physical Properties of a solution change because solute particles act like impurities, getting in the way of solvent particles.
We add impurities to lower freezing points, increase boiling points, or reduce vapor pressure.
The more impurities, the greater the change
Ethyl glycol is added to water in your car’s radiator to increase water’s boiling point.

37. Colligative Properties
Depend only on the number, not on the identity, of the solute particles in an ideal solution.
Boiling point elevation
Freezing point depression
Osmotic pressure

38. Colligative Properties
I. Properties which depend only on the concentration
of solute species present (not their identity).
A.   Vapor Pressure Reduction- The greater the number
of solute molecules the lower the vapor pressure
of the solution.
a.  Raoult’s Law – The vapor pressure of a solvent in
an ideal solution, Psolv, is equal to the mole
fraction of the solvent, Xsolv, times the vapor
pressure of the pure solvent, P0solv.
                           i.      Psolv= Xsolv P0solv
or restated
                         ii.      ΔP = Xsolute P0solv
as the change in vapor pressure as a function of solute concentration. Remember vapor pressure decreases w/ addition of solute.

39. 11.4 The Vapor Pressure of Solutions

40. A Solution Obeying Raoult’s Law

41. Vapor Pressure Lowering: Liquid/Vapor Equilibrium

42. Vapor Pressure Lowering: Addition of a Solute

43. Vapor Pressure Lowering: Solution/Vapor Equilibrium

44. React 3
Consider an ideal mixture of two volatile liquids, A and B. Answer the following questions. Use pictures in your explanations. 
Suppose you have an equimolar mixture of A and B, and the vapor pressure of pure liquid A is twice that of pure liquid B. Determine the mole fraction of A in the vapor above the mixture. 
Suppose you make a mixture with twice as much A as B (in moles), and the vapor pressure of pure liquid A is twice that of pure liquid B. Determine the mole fraction of A in the vapor above the mixture.

45. React 4
Benzene and toluene form ideal solutions. Imagine mixing benzene and toluene at a temperature where the vapor pressure of pure benzene is torr and the vapor pressure of toluene is torr. Consider the following:
You make a solution in which the mole fraction of the benzene is You then place this solution in a closed container, and wait for the vapor to come into equilibrium with the solution. Then, you condense the vapor. Determine the composition (mole percent) of the vapor. Explain why the relative numbers make sense.

46. React 4
Benzene and toluene form ideal solutions. Imagine mixing benzene and toluene at a temperature where the vapor pressure of pure benzene is torr and the vapor pressure of toluene is torr. Consider the following:
You make a solution by pouring some toluene into some benzene. You then place this solution in a closed container, and wait for the vapor to come to equilibrium with the solution. Then, you condense the vapor and determine the mole fraction of the toluene to be Determine the composition (mole percent) of the original solution. Explain why the relative numbers make sense.

47. Summary of the Behavior of Various Types of Solutions
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

48. React 5
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)

49. React 6
Would a solution with a positive deviation from Raoult’s Law have a higher or lower boiling point than ideal? Explain.

50. Boiling-Point Elevation and Freezing-Point Depression

51. 11-4 Colligative Properties cont…
 C.     Elevation of the boiling point – Reduction in vapor pressure of solvent due to addition of solute results in elevation of the boiling point of the solvent.
a.  ΔT = Kbm where
                i.    ΔT is the change in boiling point.
                ii.   Kb is the molal boiling point elevation constant for that
solvent. See handout for values of Kb for various solvents
                 iii.   m is the molality of the solution in moles of solute per
kilogram of solvent.
D.    Depression of the freezing point – Reduction in vapor pressure of the solvent due to addition of solute results in depression of the freezing point. This occurs because the solid and liquid have the same vapor pressure and are in equilibrium at the freezing (melting) point.
a.  ΔT = Kfm where
              i.    ΔT is the change in freezing point.
              ii.   Kf is the molal freezing point depression constant for that
solvent. See handout for values of Kf for various solvents.
               iii.  m is the molality of the solution in moles of solute per

52. Boiling Point Elevation: Liquid/Vapor Equilibrium

53. Boiling Point Elevation: Addition of a Solute

54. Boiling Point Elevation: Solution/Vapor Equilibrium

55. Freezing Point Depression: Solid/Liquid Equilibrium

56. Freezing Point Depression: Addition of a Solute

57. Freezing Point Depression: Solid/Solution Equilibrium

58. Changes in Boiling Point and Freezing Point of Water

59. React 7
Consider two solutions. In solution A, you add 1.0 mole of table sugar to 1.0 L of water. In solution B, you add 1.0 mole of table salt to 1.0 L of water.
How do the freezing points of each solution compare to water?
How do the freezing points of each solution compare to each other?

60. React 8
Why is molality used (not molarity) for colligative properties?

61. React 9
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 faction of sucrose in the original mixture.

62. React 10 The freezing point of an aqueous solution is –2.79°C.
Determine the boiling point of this solution.
Determine the vapor pressure (in mm Hg) of this solution at 25°C (the vapor pressure of pure water at 25°C is mm Hg).
Explain any assumptions you make in solving a and b.

63. React 11
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?

64. 11-4 Colligative Properties cont…
D.     Osmosis and Osmotic Pressure
a.  Osmosis – Occurs when a semipermeable membrane separates a
solution from the pure solvent. The solvent molecules pass through
and dilute the solution.
b. Osmotic Pressure (π)– The pressure required to stop the
osmosis from a pure solvent into a solution.
                      i.      π = MRT where
1. π – Osmotic Pressure
2. M – Molarity of solution in moles per liter
3. R – Gas constant = l atm/ K mol
4. T – Kelvin Temperature.
c.  Reverse Osmosis – Occurs when pressure is applied to solution which
is greater than the osmotic pressure. This increases the volume of
pure solvent.
d. Isotonic solutions have the same osmotic pressure as blood.
Hypotonic solutions are more dilute than blood and cause the cells to
swell or burst. Hypertonic solutions are more concentrated and
cause cells to shrivel as water diffuses out of the cell
E. Determining Molar mass
a.  Any of the colligative properties can be used to determine the molar
mass of solutes. Lab #12 addresses this. Do the practice problems.

65. Osmosis and Osmotic Pressure
Osmosis is the travel of a solvent from an area of low concentration (high purity) to high concentration (low purity).
Examples of osmosis:
A Cucumber placed in a conc. NaCl solution (brine) loses water, shrivels up, and becomes a pickle.
Limp carrots and celery, placed in water, become firm because water enters via osmosis.
From pure to impure

66. Osmotic Pressure

67. Osmosis

68. React 12
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.

69. React 13
An aqueous solution is 1.00% NaCl by mass. Its density is g/mL at 25°C. The observed osmotic pressure of this solution is 7.83 atm at 25°C.
What fraction of the moles in NaCl in this solution are ion pairs?
Calculate the freezing point that would be observed for this solution.

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

71. Will it dissolve? (Solubility Rules)
Not all ionic solids (salts) will dissolve.
We use solubility rules to decide if the substance will dissolve.
Salts containing…
Alkali metal cations(+) are soluble.
NH4+, NO3-, ClO3-, SO42- are soluble.
Pb+, Ag+, Hg2+ are insoluble.
CO3-, PO43-, S2- are insoluble.
Which of the following salts are soluble?
NaCl, HgCO3, Ca(NO3)2, AgF, PbI2, FeSO4
HUGE ON THE AP TEST!
BaSO4, SrSO4, and PbSO4 are insoluble

72. Dissociation and Ions Present
Dissociation = a salt dissolving into its ions:
How many moles of ions are in a solution of 1 mole of NaCl?
How many moles of ions are in solutions of 1 mole of each of the following?:

73. Net Ionic Equations
HUGE ON THE AP TEST!
When we write a balanced chemical equation, we show all species present (all reactants and all products):
In a net ionic equation, we show only precipitates formed, and the reactants that form them:
The chemicals that stay ions are called spectator ions, And are left out (Na+, NO3-)
Remember to Balance

74. Net Ionic Equation Practice
Write the net ionic equations for the following:

75. Strong/Weak Electrolytes
Recall that a solid compound made up of a cation and anion is called a salt.
Salts that dissolve completely into their ions when put in water dissociate completely.
Salts that dissociate completely form strong electrolytes – solutions that conduct electricity well.
Some salts only partially dissociate, forming weak electrolytes – solutions that conduct electricity, but do so poorly.

76. H+ / OH- Ions – (Acids and Bases)
When a H+ ion is released into solution, a H3O+ ion is produced, called Hydronium ion.
When a OH- ion is produced, we call this a Hydroxide ion.
Hydronium (H3O+) and Hydroxide (OH-) are the fundamental ions involved in acid/base chem.
Acids that dissociate completely, releasing H+ ions form strong electrolytes.
Bases that dissociate completely releasing OH- ions form strong electrolytes.

77. Increasing temperature increases the solubility of solids in liquids
The extent to which a solute will dissolve in a solvent. (how much solute will dissolve)
High solubility
large amounts of solute will dissolve in a solvent
Low solubility
only small amounts of solute will dissolve
Increasing temperature increases the solubility of solids in liquids
Increasing temperature decreases the solubility of gasses in liquids! …