1. AP Chapter 17 Ionic Equilibria of Weak Electrolytes

2. Review Test 20 AP Multiple Choice Questions

3. pH and pOH scales basicacidic pH + pOH = 14.00

4. H + = H 3 O +

5. pH = -log[H + ] The pH scale is a logarithmic scale. This means that in order to change the pH by one unit there must be a tenfold change in the [H + ]. Find the pH of a solution with [H + ] = 0.010M Find the pH of a solution with [H + ] = 0.10M

6. Example 17.1 Page 521 What is the pH of a solution of HCl with a concentration of 1.2 x 10 -3 M? pH = 2.92

7. Example 17.2 Page 521 Calculate the pH of: –0.10 M solution of HNO 3 –0.10 M solution of CH 3 CO 2 H (1.3% ionized) pH = 2.89 pH = 1.00

8. Example 17.4 Page 522 Calculate the [H 3 O + ] of a solution with a pH of 9.0 [H 3 O + ] = 1.0 x 10 -9

9. Example 17.6 Page 523 Calculate the pH and pOH of a 0.0125 M KOH solution. pOH = 1.903pH = 12.097

10. Calculate the [OH - ] of a solution with a pH of 5.56. [OH - ] = 3.63 x 10 -9

11. What are the ion concentrations in a 0.10M HCl solution? 0.10M HCl 0.10M H + 0.10M Cl - Strong electrolytes dissociate completely

12. What are the ion concentrations in a 0.15 M K 2 SO 4 solution? 0.15M K 2 SO 4 0.30M K + 0.15M SO 4 2-

13. At equilibrium, a solution of acetic acid, [ CH 3 CO 2 H ] = 0.0788M and [H 3 O + ] = [CH 3 CO 2 - ] = 0.0012M. What is the K a of acetic acid? Example 17.8 page 527

14. Example 17.9 page 527 The pH of a 0.0516 M solution of nitrous acid, HNO 2, is 2.34. What is the K a ?

16. Additional K a values Table contains more Ka values than are pictured here

17. Calculate the [H 3 O + ], [CH 3 CO 2 - ], and [ CH 3 CO 2 H ] in a 0.100 M solution of acetic acid. What is the K a = 1.8 x 10 -5. Example 17.10 page 528

19. What else can you determine in the previous example? pH, pOH, [OH - ] percent ionization Calculate this value. Percent ionization = 1.3%

20. What is the percent ionization of a 0.25 M solution of trimethylamine, (CH 3 ) 3 N, a weak base with a K b = 7.4 x 10 -5 1.7%

21. What is the pH of the trimethylamine solution? [OH-] = 4.3 x 10 -3 pOH = -log[4.3 x 10 -3 ] = 2.37 pH = 14 – 2.37 = 11.63 [OH-] = 4.3 x 10 -3 [H+] = 1 x 10 -14 ÷ 4.3 x 10 -3 = 2.3 x 10 -12 pH = -log[2.3 x 10 -12 ] = 11.63

22. Appendix G has additional values

23. Diprotic and Triprotic Acids A diprotic acid ionizes in two steps because it has two ionizable hydrogens. A triprotic acid ionizes in three steps because it has three ionizable hydrogens.

24. The Stepwise Dissociation of Phosphoric Acid. A triprotic acid. H 3 PO 4 (aq) + H 2 O (l) H 2 PO 4 - (aq) + H 3 O + (aq) H 2 PO 4 - (aq) + H 2 O (l) HPO 4 2- (aq) + H 3 O + (aq) HPO 4 2- (aq) + H 2 O (l) PO 4 3- (aq) + H 3 O + (aq) H 3 PO 4 (aq) + 3 H 2 O (l) PO 4 3- (aq) + 3 H 3 O + (aq)

25. Each ionization step occurs to a lesser extent than one preceding it. H 3 PO 4 (aq) + H 2 O (l) H 2 PO 4 - (aq) + H 3 O + (aq) H 2 PO 4 - (aq) + H 2 O (l) HPO 4 2- (aq) + H 3 O + (aq) HPO 4 2- (aq) + H 2 O (l) PO 4 3- (aq) + H 3 O + (aq) H 3 PO 4 (aq) + 3 H 2 O (l) PO 4 3- (aq) + 3 H 3 O + (aq) K a = 7.5 x 10 -3 K a = 6.3 x 10 -8 K a = 3.6 x 10 -13 K a = ?

26. 0.10 mol CH 3 CO 2 H and 0.50 mol NaCH 3 CO 2 are dissolved in 1.0L of solution. Calculate [H 3 O + ]. Example 17.19 page 548

28. 0.10 mol CH 3 CO 2 H and 0.50 mol NaCH 3 CO 2 are dissolved in 1.0L of solution. Calculate [H 3 O + ]. Example 17.19 page 548 We can work the problem using either K a or K b

29. Compare this answer to [H 3 O + ] from problem 17.10 pp. 528. How does this illustrates LeChatlier’s Principle.

30. Example 17.10 page 528 CH 3 CO 2 H + H 2 O ↔ CH 3 CO 2 - + H 3 O + [H 3 O + ] = 1.3 x 10 -3 0.10 mol CH 3 CO 2 H and 0.50 mol NaCH 3 CO 2 are dissolved in 1.0L of solution. Calculate [H 3 O + ]. [H 3 O + ] = 3.6 x 10 -6 Calculate the [H 3 O + ], [CH 3 CO 2 - ], and [ CH 3 CO 2 H ] in a 0.100 M solution of acetic acid. What is the Ka = 1.8 x 10 -5.

31. 10.0ml of 0.10M HCl and 25.0ml of 0.10M NH 3 are mixed. Calculate the [OH-].

32. Example 17.20 page 549

34. Common Ion Problems. What is the “common ion” in each example? Example 17.20 page 549 10.0ml of 0.10M HCl and 25.0ml of 0.10M NH 3 are mixed. Calculate the [OH - ]. Example 17.19 page 548 0.10 mol CH 3 CO 2 H and 0.50 mol NaCH 3 CO 2 are dissolved in 1.0L of solution. Calculate [H 3 O + ]. Buffers are also examples of the common ion effect since they are mixtures where both substances produce the same ion. Both of the solution mixtures above are buffers.

35. Buffers A buffer solution is an aqueous solution consisting of a mixture of a weak acid and its conjugate base or a weak base and its conjugate acid. Buffer solutions have the property that the pH of the solution changes very little when a small amount of acid or base is added to it.

36. Demonstration: Buffered vs. Non-buffered solutions

37. Why are these solutions buffers? Example 17.20 page 549 10.0ml of 0.10M HCl and 25.0ml of 0.10M NH 3 are mixed. Calculate the [OH - ]. Example 17.19 page 548 0.10 mol CH 3 CO 2 H and 0.50 mol NaCH 3 CO 2 are dissolved in 1.0L of solution. Calculate [H 3 O + ].

38. How buffers work Weak acids and weak bases tend to remain in high concentrations when added to water because by definition they do not ionize much in water since they are weak. However, they are very likely to react with any added strong base or strong acid.

39. Why are weak acids/bases used to create buffers?

40. Remember this problem Calculate the pH of: –0.10 M solution of HNO 3 –0.10 M solution of CH 3 CO 2 H (1.3% ionized) What would the [CH 3 CO 2 H] have to be for it to have the same pH as the HNO 3 assuming the 1.3% ionization factor does not change? 0.10M = [ CH 3 CO 2 H](0.013) [ CH 3 CO 2 H] = 7.7 M pH = 2.89 pH = 1.00

41. Why are weak acids/bases used to create buffers? That’s right 7.7 M [ CH 3 CO 2 H] vs. 0.10M have the same pH. In other words the [ CH 3 CO 2 H] is 77 times greater [HNO 3 ]. It takes much more base to change the pH of a weak acid solution as compared to a strong acid solution of the same pH because there are many more moles of acid available to neutralize base in the weak acid as compared to the strong acid.

42. Weak Acid (HA) and its conjugate base (A - ) buffer

43. Adding a strong base to the buffer If a strong base is added to a buffer, the weak acid will give up its H + in order to transform the base (OH - ) into water (H 2 O) and the conjugate base: HA + OH - → A - + H 2 O. Since the added OH - is consumed by this reaction, the pH will change only slightly.

44. Adding a strong acid to the buffer If a strong acid is added to a buffer, the weak base will react with the H + from the strong acid to form the weak acid, HA: H + + A - → HA The H + gets absorbed by the A - instead of reacting with water to form H 3 O + (H + ), so the pH changes only slightly.

45. An effective buffer requires relatively equal amounts of weak acid and conjugate base.

46. Basic Buffers Note that the same ideas hold true for weak bases, (B), and their conjugate acids (BH + ).

47. For the most effective buffers K a = [H 3 O + ] HA + H 2 O ↔ A - + H 3 O + Consider acetic acid

48. You want to make the most “effective” buffer you can using acetic acid. What would the [H 3 O + ] be? Calculate the pH of this buffer.

49. I want to make a buffer with a pH of 3.14. Which acid should I use?

50. Did you choose well? Or will you be flung into the “Gorge of Eternal Peril”.

51. The Bridge of Death ≈ 3:10

52. I want to make a buffer with a pH of 3.14. Which acid should I use? K a = [H 3 O + ] [H 3 O + ] = antilog (-pH) Ka = antilog (-3.14) = 7.2 x 10 -4 HF

53. I want to make a buffer with a pH of 3.75. Which acid should I use?

54. Formic Acid (HCO 2 H) K a = [H 3 O + ] [H 3 O + ] = antilog (-pH) Ka = antilog (-3.14) = 7.2 x 10 -4

55. For the most effective buffers K b = [OH - ] Consider ammonia

56. For the most effective buffers K b = [OH - ] B + H 2 O ↔ BH + + OH -

57. I want to make a buffer with a pH of 10.64. Which base should I use?

58. K b = [OH - ] K b = antilog (-pOH) K b = antilog (-3.36) = 4.4 x 10 -4 Methylamine (CH 3 NH 2 )

59. Effective Buffers The previous formulas only apply when the concentrations of weak acid and conjugate base or weak base and conjugate acid are equal. However it is important to note that you do not have to have equal concentrations to have a buffer.

60. Henderson – Hasselbach Equation

61. 0.10 mol CH 3 CO 2 H and 0.50 mol NaCH 3 CO 2 are dissolved in 1.0L of solution. Calculate [H 3 O + ]. Example 17.19 page 548

62. Henderson – Hasselbach Equation

63. 10.0ml of 0.10M HCl and 25.0ml of 0.10M NH 3 are mixed. Calculate the [OH - ]. Example 17.20 page 549

64. I want to make a the most effective buffer possible using acetic acid. What chemicals should I get from the stock room? Describe how you would make the buffer? What would the pH of the buffer be?

65. I have no acetate salts. Can I still make the buffer using acetic acid? Explain.

66. I want to make a the most effective buffer possible using ammonia. What chemicals should I get from the stock room? Describe how you would make the buffer? What would the pH of the buffer be?

67. I have no ammonium salts. Can I still make a buffer using ammonia? Explain.

68. Making Effective Buffers 1.Equal moles of a weak acid and a salt of its conjugate base. 2.Equal moles of a weak base and a salt of its conjugate acid. 3.A weak acid and half as many moles of strong base. 4.A weak base and half as many moles of strong acid. 5.“Use a pH meter”

69. pH meter

70. Calculate the pH of a buffer that is 0.10 M acetic acid and 0.10 M sodium acetate. pH = 4.74 Example 17.21 page 551

71. Buffer Capacity Buffer capacity: the amount of an acid or base that can be added to a volume of a buffer solution before its pH changes significantly. Buffer capacity depends on the amount (moles) of the conjugate pair used to make the buffer. Buffer solutions have essentially lost their buffering capabilities when one component of the conjugate pair is about 10% or less of the other.

72. Page 555

73. Buffers Which buffer has the greater capacity? Which buffer is more effective? Buffer A: –100mL of 0.1M CH 3 CO 2 H with 100 mL of 0.1M NaCH 3 CO 2. Buffer B: –100mL of 1.0M CH 3 CO 2 H with 100 mL of 1.0M NaCH 3 CO 2.

74. Calculate the [OH - ] of a 0.050 M solution of NaCH 3 CO 2 [OH - ] = 5.3 x 10 -6

75. Calculate the pH of a 0.050 M solution of NaCH 3 CO 2 pH = 8.72

76. Calculate the percent reaction of a 0.050 M solution of NaCH 3 CO 2 0.011%

77. Calculate the pH of a 0.10 M solution of AlCl 3 (K a = 1.4 x 10 -5 ) pH = 2.93

78. Titration titrant analyte

79. Titration Curves A titration curve is a plot of the pH against the volume of acid or base added in a titration. The equivalence point or endpoint for a titration is the point at which exactly enough of the titrant has been added to completely react with the analyte. In other words, at the equivalence point, the number of moles of titrant added corresponds exactly to the number of moles of substance being titrated, the analyte, according to the reaction stoichiometry.

80. Table 17.8 Page 568

83. How do we choose an appropriate indicator for a titration?

84. Indicators An acid – base indicator is a substance that indicates how acidic or basic a solution is using certain color changes. Indicators are normally weak acids that have a different color than their conjugate base. HA + H 2 O ↔ H 3 O + + A - Indicators can also be weak bases that have a different color than their conjugate acid.

85. Indicators Remember that the equivalence point of a titration is where you have mixed the two substances in exact stoichiometric proportions. You obviously need to choose an indicator which changes color as close as possible to that equivalence point. The indicator should have a pKa value near the pH of the titration's endpoint. That varies from titration to titration.

86. pK a and pH range of Acid-Base Indicators

87. Which indicator(s) would be appropriate for the titrations below?

88. Indicators

90. Choosing an Indicator Remember that an indicator is normally a weak acid. HA + H 2 O ↔ H 3 O + + A - Write the equilibrium expression for the reaction of the indicator.

91. Choosing an Indicator HA + H 2 O ↔ A - + H 3 O + How would the [HA] compare to the [A-] halfway to the equivalence point of the reaction of HA with water?

92. Choosing an Indicator HA + H 2 O ↔ A - + H 3 O + [HA] = [A-] halfway to the equivalence point.

93. Choosing an Indicator HA + H 2 O ↔ A - + H 3 O + Therefore halfway to the equivalence point for the reaction of any indicator is an excellent point to pick the indicator for our acid-base titration. Why?

94. Choosing an Indicator HA + H 2 O ↔ A - + H 3 O + [HA] = [A-] halfway to the equivalence point. This is the point that the indicator changes color.

95. Choosing an Indicator HA + H 2 O ↔ A - + H 3 O + [HA] = [A-] halfway to the equivalence point. Less than halfway to the equivalence point of this reaction [HA] > [A-] and therefore the color is still red. More than halfway to the equivalence point of this reaction [A-] > [HA] and therefore the color is still blue.

96. Choosing an Indicator [HA] = [A-] halfway to the equivalence point. At half way to the equivalence point K a = [H 3 O + ] At half way to the equivalence point pK a = pH

97. Choosing an Indicator At half way to the equivalence point K a = [H 3 O + ] At half way to the equivalence point pK a = pH Therefore we want to choose an indicator that has a pK a = pH at the equivalence point of the titration we are performing. This is because this is the point at which the indicator will go through its most significant color change.

98. Now that we have chosen an appropriate indicator for a titration. Let’s consider in more detail what happens during the titration of different types of acids.

99. What substances are involved in each of the titrations? (a) (b) (c) (d)

100. What species are present at half way to the equivalence point? NaOH + HCl → NaCl + H 2 O At half way to the equivalence point the main species present are H + (hydronium), Cl - and Na +. The solution is still acidic.

101. What species are present at half way to the equivalence point? NaOH + CH 3 CO 2 H → NaCH 3 CO 2 + H 2 O At half way to the equivalence point the main species present are CH 3 CO 2 H, Na +, and CH 3 CO 2 - and to a lesser extent H + (hydronium). The solution is still acidic.

102. The titration curve is produced when a 10.0 ml sample of HCl is titrated with 0.100M NaOH. What is the concentration of the HCl solution. 0.25M HCl

103. The titration curve is produced when a 10.0 ml sample of CH 3 CO 2 H is titrated with 0.100M NaOH. What is the concentration of the CH 3 CO 2 H solution. 0.25M CH 3 CO 2 H

104. Determine the K a of the weak acid using the titration curve. See next slide for a clue.

105. a

106. Determine the Ka of the weak acid using the titration curve.

107. Diprotic Acid When titrating a diprotic acid with a strong base it is essentially like doing two titrations at once.

108. Titration curve of a diprotic acid The titration curve shown above is for a diprotic acid such as H 2 SO 4. This proves that polyprotic acids lose their protons in a stepwise manner. H 2 SO 4 → H + + HSO 4 - HSO 4 - → H + + SO 4 2-

109. How would you determine the Ka for each acid (H 2 SO 4 and HSO 4 - )?

111. Why do you have to learn all this very challenging information?

112. This stuff is hard. Why do you have to learn it?

113. Why? Because I had to suffer through this crap and now its payback time!