1. Chapter 13 Properties of Solutions
CHEMISTRY The Central Science 9th Edition
Chapter 13 Properties of Solutions

2. Text, P. 417, review (Chapter 11)

3. 13.1: The Solution Process Solutions homogeneous mixtures
Solution formation is affected by
strength and type of intermolecular forces
forces are between and among the solute and solvent particles

4. Text, P. 486

5. Hydration of solute
Attractive forces between solute & solvent particles are comparable in magnitude with those between the solute or solvent particles themselves
Note attraction of charges
What has to happen to:
Water’s H-bonds?
NaCl?
What intermolecular
force is at work in
solvation?
Text, P. 486

7. Energy Changes and Solution Formation
There are three energy steps in forming a solution:
the enthalpy change in the solution process is
Hsoln = H1 + H2 + H3
Hsoln can either be + or - depending on the intermolecular forces
Text, P. 487

8. Text, P. 488
MgSO4 Hot Pack
NH4NO3 Cold Pack

9. H1 and H2 are both positive
Breaking attractive intermolecular forces is always endothermic
Forming attractive intermolecular forces is always exothermic
To determine whether Hsoln is positive or negative, consider the strengths of all solute-solute and solute-solvent interactions:
H1 and H2 are both positive
H3 is always negative

10. Rule: Polar solvents dissolve polar solutes
Non-polar solvents dissolve non-polar solutes
(like dissolves like)
WHY?
If Hsoln is too endothermic a solution will not form
NaCl in gasoline: weak ion-dipole forces (gasoline is non-polar)
The ion-dipole forces do not compensate for the separation of ions

11. Solution Formation, Spontaneity, and Disorder
A spontaneous process occurs without outside intervention
When energy of the system decreases, the process is spontaneous
Some spontaneous processes do not involve the system moving to a lower energy state (e.g. an endothermic reaction)
If the process leads to a greater state of disorder, then the process is spontaneous
Entropy

12. Example: a mixture of CCl4 and C6H14 is less ordered than the two separate liquids
Therefore, they spontaneously mix even though Hsoln is very close to zero
Text, P. 489

13. Ni(s) + 2HCl(aq)  NiCl2(aq) + H2(g)
Solution Formation and Chemical Reactions
Example:
Ni(s) + 2HCl(aq)  NiCl2(aq) + H2(g)
When all the water is removed from the NiCl2 solution, no Ni is found only NiCl2·6H2O (a chemical reaction that results in the formation of a solution)
Water molecules fit into the crystal lattice in places not specifically occupied by a cation or an anion
Hydrates
Water of hydration
Think about it: What happens when NaCl is dissolved in water and then heated to dryness?

14. NaCl(s) + H2O (l)  Na+(aq) + Cl-(aq)
When the water is removed from the solution, NaCl is found
NaCl dissolution is a physical process

15. Sample problem # 3

16. 13.2: Saturated Solutions and Solubility
Dissolve: solute + solvent  solution
Crystallization: solution  solute + solvent
Saturation: crystallization and dissolution are in equilibrium
Solubility: amount of solute required to form a saturated solution
Supersaturated: a solution formed when more solute is dissolved than in a saturated solution

17. 13.3: Factors Affecting Solubility
1. Solute-Solvent Interaction
“Like dissolves like”
Miscible liquids: mix in any proportions
Immiscible liquids: do not mix

18. Generalizations:
Intermolecular forces are important:
Water and ethanol are miscible
broken hydrogen bonds in both pure liquids are
re-established in the mixture
The number of carbon atoms in a chain affects solubility: the more C atoms in the chain, the less soluble the substance is in water

19. Generalizations, continued:
The number of -OH groups within a molecule increases solubility in water
The more polar bonds in the molecule, the better it dissolves in a polar solvent (like dissolves like)
Network solids do not dissolve
the strong IMFs in the solid are not re-established in any solution

20. Text, P. 493

21. Read “Chemistry & Life”, P. 494
Fat soluble vitamin
Water soluble vitamin

22. 2. Pressure Effects
Solubility of a gas in a liquid is a function of the pressure of the gas

23. High pressure means
More molecules of gas are close to the solvent
Greater solution/gas interactions
Greater solubility
If Sg is the solubility of a gas
k is a constant
Pg is the partial pressure of a gas
then Henry’s Law gives:
Carbonated Beverages!

24. As temperature increases Solubility of solids generally increases
3. Temperature Effects
As temperature increases
Solubility of solids generally increases
Solubility of gases decreases
Thermal pollution
Text, P. 497

25. Figure 13.17, P. 497

26. Sample problem # 17

27. 13.4: Ways of Expressing Concentration
All methods involve quantifying amount of solute per amount of solvent (or solution)
Amounts or measures are masses, moles or liters
Qualitatively solutions are dilute or concentrated

28. Definitions: 1.

29. 2. 3.
Recall mass can be converted to moles using the molar mass

30. 4.
Converting between molarity (M) and molality (m) requires density
Molality doesn’t vary with temperature
Mass is constant
Molarity changes with temperature
Expansion/contraction of solution changes volume

31. Text, P. 501

32. Sample Problems #31, 33, 37, 39, 41

33. 13.5: Colligative Properties
Colligative properties depend on quantity of solute particles, not their identity
Electrolytes vs. nonelectrolytes
0.15m NaCl  0.15m in Na+ & 0.15m in Cl-  0.30m in particles
0.050m CaCl2  0.050m in Ca+2 & 0.1m in Cl-  0.15m in particles
0.10m HCl  0.10m in H+ & 0.10m in Cl-  0.20m in particles
0.050m HC2H3O2  between 0.050m & 0.10m in particles
0.10m C12H22O11  0.10m in particles
Compare physical properties of the solution with those of the pure solvent

34. 1. Lowering Vapor Pressure
Non-volatile solutes reduce the ability of the surface solvent molecules to escape the liquid
Vapor pressure is lowered
Raoult’s Law:
PA is the vapor pressure with solute
PA is the vapor pressure without solute
A is the mole fraction of solvent in solution A
Increase X of solute, decrease vapor pressure above the solution

35. Ideal solution: one that obeys Raoult’s law
Raoult’s law breaks down (Real solutions)
Real solutions approximate ideal behavior when
solute concentration is low
solute and solvent have similar IMFs
Assume ideal solutions for problem solving
2. Boiling-Point Elevation
The triple point - critical point curve is lowered

36. At 1 atm (normal BP of pure liquid) there is a lower vapor pressure of the solution
A higher temperature is required to reach a vapor pressure of 1 atm for the solution (Tb)
Molal boiling-point-elevation constant, Kb, expresses how much Tb changes with molality, m:

37. Text, P. 505

38. 3. Freezing Point Depression
The solution freezes at a lower temperature (Tf) than the pure solvent
lower vapor pressure for the solution
Decrease in FP (Tf) is directly proportional to molality (Kf is the molal freezing-point-depression constant):

39. Applications: Antifreeze!
Text, P. 505
Applications: Antifreeze!

41. Examples: # 45, 47, 49, 51 & 53
A neat link

42. 4. Osmosis
Semipermeable membrane: permits passage of some components of a solution
Example: cell membranes and cellophane
Osmosis: the movement of a solvent from low solute concentration to high solute concentration
There is movement in both directions across a semipermeable membrane
“Where ions go, water will flow” ~ Mrs. Moss

43. Eventually the pressure difference between the arms stops osmosis
Text, P. 507

44. Osmotic pressure, , is the pressure required to stop osmosis:
It is colligative because it depends on the concentration of the solute in the solvent

45. Isotonic solutions: two solutions with the same  separated by a semipermeable membrane
Hypertonic solution: a solution that is more concentrated than a comparable solution
Hypotonic solution: a solution of lower  than a hypertonic solution
Osmosis is spontaneous
Read text, P. 508 – 509 for practical examples

46. Examples: #57, 59 & 61

47. There are differences between expected and observed changes due to colligative properties of strong electrolytes
Electrostatic attractions between ions
“ion pair” formation temporarily reduces the number of particles in solution
van’t Hoff factor (i): measure of the extent of ion dissociation

48. Ratio of the actual value of a colligative property to the calculated value (assuming it to be a nonelectrolyte)
Ideal value for a salt is the # of ions per formula unit
Factors that affect i:
Dilution
Magnitude of charge on ions
lower charges, less deviation

49. Sample Problem, # 63, 82

50. 11.6: Colloids Read Text, Section 13.6, P. 511 – 515 Terms/Processes:
Tyndall effect
Hydrophilic
Hydrophobic
Adsorption
Coagulation

51. 11.6: Colloids Read Text, Section 13.6, P. 511 – 515
Suspensions in which the suspended particles are larger than molecules
too small to drop out of the suspension due to gravity
Tyndall effect: ability of a colloid to scatter light
The beam of light can be seen through the colloid

52. Text, P. 512

53. Hydrophilic and Hydrophobic Colloids
“Water loving” colloids: hydrophilic
“Water hating” colloids: hydrophobic
Molecules arrange themselves so that hydrophobic portions are oriented towards each other

54. Adsorption: when something sticks to a surface we say that it is adsorbed
Ions stick to a colloid (colloids appears hydrophilic)
Oil drop and soap (sodium stearate)
Sodium stearate has a long hydrophobic tail (Carbons) and a small hydrophilic head (-CO2-Na+)

55. Text, P. 514

56. Removal of Colloidal Particles
Coagulation (enlarged) until they can be removed by filtration
Methods of coagulation:
heating (colloid particles are attracted to each other when they collide)
adding an electrolyte (neutralize the surface charges on the colloid particles)

57. End of Chapter 13 Properties of Solutions