1. Chapter 13 Properties of Solutions
Chemistry, The Central Science, 11th edition
Theodore L. Brown, H. Eugene LeMay, Jr., and Bruce E. Bursten
Chapter 13 Properties of Solutions
MC 9 out of 75
FRQ almost every year
Extension of Chapter 4
Thanks to
John D. Bookstaver
© 2009, Prentice-Hall, Inc.

2. 13.1 The Solution Process
Solutions are homogeneous mixtures of two or more pure substances.
Formed when 1 substance (solute) disperses uniformly throughout another (solvent).
Solute – present in lesser quantity
Solvent – present in greater quantity

3. Solutions can be mixes involving any state.
13.1 The Solution Process
Solutions can be mixes involving any state.
This mix always forms a solution!

4. How Does a Solution Form? Solvation
As a solution forms, the solvent pulls solute particles apart and surrounds, or solvates, them.
If the solvent is water, this process is termed hydration.

5. How Does a Solution Form? Dispersion
In the absence of a strong solvent-solute attraction, the particles randomly spread out or Disperse.
This is typical for nonpolar solutions!

6. Why do Solutions Form?
The ability of substances to form solutions depends on 2 factors:
Type of intermolecular interactions involved
Natural tendency of substances to spread into larger volumes when not restrained

7. The Effect of Intermolecular Forces
There are 3 forces at work during solvation.
Break Solvent-Solvent attraction.
Break solute-solute attraction.
Make solvent-solute attraction
If #3 is comparable to, or greater, than #1 and #2, than a solutions forms!

8. Break Solvent – Solvent Forces
H to N,O, or F
Hydrogen
bonding
Polar
Solvent
Dipole-dipole
Nonpolar
solvent
London or
Dispersion
(induced dipoles)

9. Break Solute – Solute Forces
H to N,O, or F
Hydrogen
bonding
Dipole-dipole
Polar Molecules
London or
Dispersion
(induced dipoles)
Nonpolar
Molecules

10. Make Solute – Solvent Forces
Strong
Weak

11. NaCl and Water
NaCl is soluble in water, it is because the ion-dipole interactions are strong enough to overcome the lattice energy of the salt crystal.

12. The solute-solvent force is too weak to overcome the ionic bonding!
NaCl and C6H14 +
>
Solute-solvent
Induced dipole
(weak)
Solute-solute
ionic
(strong)
Solvent-solvent
London dispersion
(weak)
Insoluble!
The solute-solvent force is too weak to overcome the ionic bonding!

13. Energy Changes and Solution Formation
Simply put, three processes affect the energetics of solution:
separation of solute particles,
endothermic
separation of solvent particles,
new interactions between solute and solvent.
exothermic

14. Energy Changes in Solution
The enthalpy change of the overall process depends on H for each of these steps.

15. Energy Changes in Solution
Things do not tend to occur spontaneously (i.e., without outside intervention) unless the energy of the system is lowered.
(overall exothermic)

16. Why Do Endothermic Processes Occur?
Yet we know that in some processes, like the dissolution of NH4NO3 in water, heat is absorbed, not released.

17. Solution Formation, Spontaneity, and Entropy
This is where the second factor effecting solutions is involves!
Natural tendency of substances to spread into larger volumes when not restrained

18. Solution Formation, Spontaneity, and Entropy
Dissolving, when it occurs, is a spontaneous process!
If you remember from chemistry, reactions that increase the entropy (disorder and randomness) tend to be spontaneous.

19. Enthalpy Is Only Part of the Picture
The reason is that increasing the disorder or randomness (known as entropy) of a system tends to lower the energy of the system.
Diffusion is therefore a spontaneous reaction

20. Enthalpy Is Only Part of the Picture
So even though enthalpy may increase (endothermic), the overall energy of the system can still decrease if the system becomes more disordered.

21. Silver chloride is essentially insoluble in water.
Would you expect a significant change in entropy of the system when 10 g of AgCl is added to 100 g of water?
NO!

22. Solution Formation and Chemical Reactions
Just because a substance disappears when it comes in contact with a solvent, it doesn’t mean the substance dissolved.

23. Solution Formation and Chemical Reactions
Dissolution is a physical change — you can get back the original solute by evaporating the solvent.
If you can’t, the substance didn’t dissolve, it reacted.

24. 13.2 Types of Solutions Saturated
In a saturated solution, the solvent holds as much solute as is possible at that temperature.
Dissolved solute is in dynamic equilibrium with solid solute particles.

25. 13.2 Types of Solutions Unsaturated
If a solution is unsaturated, less solute than can dissolve in the solvent at that temperature is dissolved in the solvent.

26. 13.2 Types of Solutions Supersaturated
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.

27. 13.3 Factors Affecting Solubility
Chemists use the axiom “like dissolves like."
Polar substances tend to dissolve in polar solvents.
Nonpolar substances tend to dissolve in nonpolar solvents.

28. Solute-Solvent Interactions
The more similar the intermolecular attractions, the more likely one substance is to be soluble in another.

29. Solute-Solvent Interactions
Glucose (which has hydrogen bonding) is very soluble in water, while cyclohexane (which only has dispersion forces) is not.

30. Decrease, -CH3 is nonpolar
What would happen to the solubility of glucose if the hydroxyl groups (-OH) were replaced with methyl groups (-CH3)?
Decrease, -CH3 is nonpolar

31. Solute-Solvent Interactions
Vitamin A is soluble in nonpolar compounds (like fats).
Vitamin C is soluble in water.

32. Pressure Effects
In general, the solubility of gases in water increases with increasing mass.
Larger molecules have stronger dispersion forces.

33. Pressure Effects
The solubility of liquids and solids does not change appreciably with pressure.
The solubility of a gas in a liquid is directly proportional to its pressure.

34. Henry’s Law Sg = kPg where Sg is the solubility of the gas,
k is the Henry’s Law constant for that gas in that solvent, and
Pg is the partial pressure of the gas above the liquid.

35. Henry’s Law
How can this relationship be expressed?

36. Henry’s Law Problem x = (4.0 atm)(0.030 M)/(1.0 atm) x = 0.12 M
Calculate the concentration of CO2 in a soft drink that is bottled with a partial pressure of CO2 at atm over the liquid at 25 ºC if the solubility at 1.0 atm is M?
x = (4.0 atm)(0.030 M)/(1.0 atm)
x = 0.12 M

37. Temperature Effects
Generally, the solubility of solid solutes in liquid solvents increases with increasing temperature.

38. Why do bubbles form on the inside wall of a cooking pot when the water is heated on the stove, even though the temperature is well below the boiling point of water?
Increasing the temperature of the water decreases the solubility of the dissolved gasses. These released gasses deposit on the side of the container.

39. Temperature Effects The opposite is true of gases.
Carbonated soft drinks are more “bubbly” if stored in the refrigerator.
Warm lakes have less O2 dissolved in them than cool lakes.

40. 13.4 Ways of Expressing Concentrations
Qualitative Concentrations
Dilute
Relatively small concentration
Concentrated
Relatively large concentration
Quantitative Concentration
Mass percent
ppm and ppb
Mole fraction
Molarity
Molality

41. % mass is also known as “pph” (parts per hundred)
Mass Percentage
mass of A in solution
total mass of solution
 100
What does a 4.0 % NaCl solution mean?
4.0 g NaCl for every 96.0 g of water!
% mass is also known as “pph” (parts per hundred)

42. Parts per Million and Parts per Billion
Parts per Million (ppm)
mass of A in solution
total mass of solution
ppm =
 106
Parts per Billion (ppb)
mass of A in solution
total mass of solution
ppb =
 109

43. total moles in solution
Mole Fraction (χ)
moles of A
total moles in solution
XA =
In some applications, one needs the mole fraction of solvent, not solute — make sure you find the quantity you need!
Xsolute + Xsolvent = 1

44. Molarity (M) mol of solute M = L of solution
You will recall this concentration measure from Chapter 4.
Since volume is temperature-dependent, molarity can change with temperature.

45. Molality (m) mol of solute m = kg of solvent
Since both moles and mass do not change with temperature, molality (unlike molarity) is not temperature-dependent.

46. Changing Molarity to Molality
If we know the density of the solution, we can calculate the molality from the molarity and vice versa. (p545)

47. 13.5 Colligative Properties
Changes in colligative properties depend only on the number of solute particles present, not on the identity of the solute particles.
Among colligative properties are
Vapor pressure lowering
Boiling point elevation
Melting point depression
Osmotic pressure

48. Vapor Pressure
Because of solute-solvent intermolecular attraction, higher concentrations of nonvolatile solutes make it harder for solvent to escape to the vapor phase.

49. Vapor Pressure
Therefore, the vapor pressure of a solution is lower than that of the pure solvent.

50. Raoult’s Law PA = APA where
A is the mole fraction of compound A, and
PA is the normal vapor pressure of A at that temperature.
NOTE: This is one of those times when you want to make sure you have the vapor pressure of the solvent.

51. Boiling Point Elevation and Freezing Point Depression
Nonvolatile solute-solvent interactions also cause solutions to have higher boiling points and lower freezing points than the pure solvent.

52. Boiling Point Elevation
The change in boiling point is proportional to the molality of the solution:
Tb = Kb  m
where Kb is the molal boiling point elevation constant, a property of the solvent.
Tb is added to the boiling point of the pure solvent.

53. Freezing Point Depression
The change in freezing point can be found similarly:
Tf = Kf  m
Here Kf is the molal freezing point depression constant of the solvent.
Tf is subtracted from the boiling point of the pure solvent.

54. Boiling Point Elevation and Freezing Point Depression
Note that in both equations, T does not depend on what the solute is, but only on how many particles are dissolved.
Tb = Kb  m
Tf = Kf  m

55. Colligative Properties of Electrolytes
Since these properties depend on the number of particles dissolved, solutions of electrolytes (which dissociate in solution) should show greater changes than those of nonelectrolytes.

56. van’t Hoff Factor
We use the van’t Hoff factor, i, to represent the number of particles we get when a substance dissolves.
Nonelectrolytes (molecules)
i = 1
Electrolytes (ionic compounds)
i = # of ions
NaCl → i = 2
K2SO4 → i = 3
MgSO4 → i = 2

57. van’t Hoff Factor Tf = Kf  m  i Tb = Kb  m  i
We modify the previous equations by multiplying by the van’t Hoff factor, i.
Tf = Kf  m  i
Tb = Kb  m  i

58. van’t Hoff Factor
However, a 1M solution of NaCl does not show twice the change in freezing point that a 1M solution of methanol does.

59. van’t Hoff Factor
One mole of NaCl in water does not really give rise to two moles of ions.

60. van’t Hoff Factor
Some Na+ and Cl- reassociate for a short time, so the true concentration of particles is somewhat less than two times the concentration of NaCl.

61. van’t Hoff Factor
Reassociation is more likely at higher concentration and this effectively lowers the i value.
Therefore, the number of particles present is concentration-dependent.
© 2009, Prentice-Hall, Inc.

62. van’t Hoff Factor
In addition, higher charge values have greater attractive forces (coulomb's law!) and cause greater reassociation and this also lowers the actual value of i.
© 2009, Prentice-Hall, Inc. 62

63. Osmosis
Some substances form semipermeable membranes, allowing some smaller particles to pass through, but blocking other larger particles.
In biological systems, most semipermeable membranes allow water to pass through, but solutes are not free to do so.

64. Osmosis
In osmosis, there is net movement of solvent from the area of higher solvent concentration (lower solute concentration) to the are of lower solvent concentration (higher solute concentration).

65. Osmotic Pressure n  = ( )RT = MRT V
The pressure required to stop osmosis, known as osmotic pressure, , is n V
 = ( )RT = MRT
where M is the molarity of the solution.
If the osmotic pressure is the same on both sides of a membrane (i.e., the concentrations are the same), the solutions are isotonic.

66. Osmosis in Blood Cells
If the solute concentration outside the cell is greater than that inside the cell, the solution is hypertonic.
Water will flow out of the cell, and crenation results.

67. Osmosis in Cells
If the solute concentration outside the cell is less than that inside the cell, the solution is hypotonic.
Water will flow into the cell, and hemolysis results.

68. 13.6 Colloids
Suspensions of particles larger than individual ions or molecules, but too small to be settled out by gravity are called colloids.

69. Tyndall Effect Colloidal suspensions can scatter rays of light.
This phenomenon is known as the Tyndall effect.

70. Mixtures Review Settle out? Particle size Tyndall Effect
Type of mixture
Suspension
Colloid
Solution
yes
yes
large
heterogeneous no
intermediate
yes
heterogeneous no
small no
homogeneous

71. Colloids in Biological Systems
Some molecules have a polar, hydrophilic (water-loving) end and a non-polar, hydrophobic (water-hating) end.

72. Hydrophilic Colloid
In this molecules the hydrophilic parts are on the surface end and the hydrophobic parts are folded in to make a hydrophilic colloid.
Hydrophilic parts become attracted to water and partially hydrated. 72

73. Hydrophilic Colloid
In this molecules the hydrophilic parts are on the surface end and the hydrophobic parts are folded in to make a hydrophilic colloid.
Hydrophilic parts become attracted to water and partially hydrated. 73

74. How do hydrophilic colloid form?
If ions get adsorbed (stick to surface), it causes them to be attracted by water.
Ions stabilize the colloid and allow limited hydration and prevents clumping nonpolar colloid by ion to ion repulsion.

75. How do hydrophilic colloid form?
They can also be stabilized as a colloid by emulsifying agents (molecule with polar and nonpolar portions like soap)