1. COMMON ION
EFFECT

2. COMMON ION
an ion common with one in a system at equilibrium which places a stress on the equilibrium
Common Ion

3. Uses of Common Ion Effect
1. control pH of a weak acid or base
2. control formation of a precipitate

4. BUFFER Example Non-example
A solution which resists a change in pH when an acid or base is added
consists of a weak acid or base and a salt containing a common ion of its conjugate

5. LeChatelier’s Principle explain the
How does
LeChatelier’s Principle explain the
operation of a buffer?

6. Example of a buffer system
CH3COOH + HOH  CH3COO- + H3O+
NaCH3COO(aq)  Na+ + CH3COO-

7. Characteristics of a Good Buffer

8. 1. operates over a narrow pH range (< 1 pH unit)
2. no reactions between buffers in a multiple buffer system
3. range can be extended using more than one buffer

9. Henderson-Hasselbalch
Equation

10. Maximum buffering will occur when ratio is close to 1, or when
pH = pKa

11. 1. What is the pH of a 0.20 M acetic acid solution?

12. Add 10. 0 mL of 0. 20 M NaOH to 50. 0 mL of the preceding solution
Add 10.0 mL of 0.20 M NaOH to 50.0 mL of the preceding solution. What is the pH?

13. Add 5. 0g sodium acetate (MM 82. 05) to 500. mL of the 0
Add 5.0g sodium acetate (MM 82.05) to 500. mL of the 0.20 M acetic acid solution. What is the pH?

14. Add 10. 0 mL of 0. 20 M NaOH to 50. 0 mL of the preceding solution
Add 10.0 mL of 0.20 M NaOH to 50.0 mL of the preceding solution. What is the pH?

15. 2. Calculate the mass of ammonium chloride (MM 43
2. Calculate the mass of ammonium chloride (MM 43.6) needed to buffer 250. mL of 2.0 M ammonia to a pH of 10.

17. TITRATION
CURVES

18. A graphical history of a titration
Titration Curve
A graphical history of a titration
typically a plot of the pH (dependent variable) and volume titrant (independent variable)

19. Uses of Titration Curves
1. determine equivalence point
2. determine number of ionization reactions
3. determine optimum buffer region
4. determine possible indicators

20. Shape of Titration Curve
Strong acid - strong base
Weak acid - strong base

21. Shape of Titration Curve

22. Shape of Titration Curve

23. Equivalence Point 1. Midpoint between points of inflection
2. Plot of the slope of each point of the curve against volume titrant (DpH/DV vs Vavg)

24. Number of Ionization Reactions
CH3COOH - NaOH
H2C2O4 - NaOH

27. Optimum Buffer Region
Area where the concentration of molecules and their conjugate ions are relatively high

28. Indicators Need to choose for each titration system
Dependent on pH at equivalence point

29. ACID-BASE
INDICATORS

30. Acid-base indicators are weak Bronsted-Lowry compounds that are different colors in acid and base form.

31. Acid-base indicators are all large organic molecules.
HIn <===> H+ + In-
Color Color 2

32. Phenolphthalein
Colorless acid form, HIn

33. Phenolphthalein
Pink base form, In-

34. The color change occurs at a different pH for different indicators.
The pH at which the indicator changes color is dependent on the Ka of the indicator as a weak acid.

35. HIn <===> H+ + In-

36. Experiments have shown that the minimum amount of change of
HIn <==> In-
that can be detected visually is

38. Thus, from the Henderson-Hasselbalch equation, one can select an appropriate indicator for a titration based upon the Ka of the indicator and the pH at the equivalence point.

39. What is the pH at the equivalence point of a titration of 25
What is the pH at the equivalence point of a titration of 25.0 mL each of 0.10 M HCl and 0.10 M NaOH?

40. What is the pH at the equivalence point of a titration of 25
What is the pH at the equivalence point of a titration of 25.0 mL each of 0.10 M CH3COOH and 0.10 M NaOH?

41. Phenolphthalein
Ka = 1 x 10-9
pH of perceptible
color change?

42. SOLUBILITY
EQUILIBRIA

43. Saturated Solution
Maximum amount of solute dissolved in a specific volume of solvent at a specific temperature

44. Saturated Solution
Equilibrium is established between a solid solute and ions from the solute

45. Super-Saturated Solution
More than the normal maximum amount of solute is dissolved in a solution.

46. Question
at a constant temperature, what is the difference in concentration of a saturated solution:
(1 mL vs 1 ML solution)
(1 mg vs 1 kg solid)

47. The concentration of a saturated solution remains the same, no matter how much solid is present, as long as the temperature remains constant.

48. The “concentration” of a solid remains the same at a constant temperature.

49. solid <===> solution AgCl(s) <===> Ag+ + Cl-
By convention, equations for the formation of saturated solutions are written in the format
solid <===> solution
AgCl(s) <===> Ag+ + Cl-

50. AgCl(s) <===> Ag+ + Cl-

51. AgCl(s) <===> Ag+ + Cl-

52. 3. What is the solubility of silver chloride in water at 25oC?
(Ksp = 1.6 x 10-10)

53. 4. What is the solubility of lead(II) bromide at 25oC. (Ksp = 4
4. What is the solubility of lead(II) bromide at 25oC? (Ksp = 4.6 x 10-6)

54. 6. What mass of nickel is dissolved in 100
6. What mass of nickel is dissolved in 100. mL of saturated nickel(II) hydroxide? (Ksp = 1.6 x 10-16)
What is the pH of this solution?

55. Which is more soluble?
Ag2CO3 [Ksp = 8.5 x 10-13] or
CaCO3 [Ksp = 3.4 x 10-9]

56. SOLUBILITY
----
ACIDITY
PRECIPITATION

57. 8. If 0. 581 gram of magnesium hydroxide (MM 58. 3) is added to 1
8. If gram of magnesium hydroxide (MM 58.3) is added to 1.00L of water, will it all dissolve? (Ksp = 8.9 x 10-12)
Below what pH would the solution be buffered so that it does all dissolve?

58. 9. Calculate the concentration of NH4+ from ammonium chloride required to prevent the precipitation of Ca(OH)2 in a liter of solution that contains 0.10 mole of ammonia and 0.10 mole of calcium ion.

59. 10. If 50. mL of 0. 012M barium chloride are mixed with 25 mL of 1
10. If 50. mL of 0.012M barium chloride are mixed with 25 mL of 1.0 x 10-6M sulfuric acid, will a precipitate form?
HINT: use the concentration quotient “Q” as we used it before

60. You have a aqueous solution of Zn2+ and Pb2+ both 0. 0010 M
You have a aqueous solution of Zn2+ and Pb2+ both M. Both form insoluble sulfides. Approximately what pH will allow maximum precipitation of one ion and leave the other in solution?
[Ksp ZnS = 2.5 x 10-22]
[Ksp PbS = 7 x 10-29]

61. SOLUBILITY
----
COMMON IONS
COMPLEX IONS

62. 12. Calculate the molar solubility of silver thiocyanate, AgSCN, in pure water and in 0.010M NaSCN.

63. A charged species consisting of a metal ion surrounded by ligands
Complex Ion
A charged species consisting of a metal ion surrounded by ligands

64. LIGAND
An ion or molecule, acting as a Lewis base, attached to the central metal ion using the d-orbitals of the metal

65. The number of ligands attached to the central metal ion.
Coordination Number
The number of ligands attached to the central metal ion.
2, 4, or 6 are most common CN

66. Metal ions add ligands one step at a time.
Ag+ + NH3 <==> Ag(NH3)+
Kf1 = 2.1 x 103
Ag(NH3)+ + NH3 <==> Ag(NH3)2+
Kf2 = 8.2 x 103
where Kf = formation constant

68. You need to familiarize yourself with “typical” complex ions, Appendix K

69. Note that a formation constant reflects the stability of the complex.

70. 13. Calculate the equilibrium constant for
AgI(s) + 2NH3(aq) <===>
[Ag(NH3)2]+(aq) + I-(aq)

71. 14. Will 5.0 mL of 2.5 M NH3 dissolve 0.0001 mole AgCl?

72. 15. A solution is prepared by adding 0. 10 mole Ni(NH3)6Cl2 to 0
15. A solution is prepared by adding 0.10 mole Ni(NH3)6Cl2 to 0.50 L of 3.0 M NH3. Calculate the [Ni(NH3)62+] and [Ni2+] in the solution.