1. Chapter 8 & 9 Acids, Bases & Buffers

2. Chapter 8 Introducing Acids & Bases Water pH (Acid rain) in the USA in 2001

3. Conjugate acids & bases

4. Relation between [H + ], [OH - ], and pH?

5. 8-3 Strengths of acids & bases Strong Acids and Bases Common strong acids and bases are listed in Table 8-1. A strong acid or strong base is completely dissociated in aqueous solution. (8-4) (8-5) P.175

6. Carboxylic Acids Are Weak Acids and Amines Are Weak Bases (8-6) P.175

7. 8-3 Strengths of acids & bases Ask yourself at p.178 Carboxylic Acids are Weak Acids and Amines are Weak Bases

8. Metal Ions with Charge ≧ 2 Weak Acids Metal Ions with Charge ≧ 2 Are Weak Acids A proton can dissociate from M(H 2 O) w n+ to reduce the positive charge on the metal complex. P.177 Relation Between Ka and Kb

9. 8-4 pH of strong Acids & Bases Example at p.180 The pH of 4.2 x 10 -3 M HClO 4 ? The pH of 4.2 x 10 -3 M KOH? Can we dissolve base in water and obtain an acidic pH (<7)?

10. 8-5 Tools for Dealing with Weak Acids and Bases pK the negative logarithm of an equilibrium constant Weak Is Conjugate to Weak The conjugate base of a weak acid is a weak base. The conjugate acid of a weak base is a weak acid. P.181

11. P.182 Using Appendix B Acid dissociation constants appear in Appendix B. Each compound is shown in its fully protonated form. Pyridoxal phosphate is given in its fully protonated form as follows:

12. P.182

13. 8-6 Weak-Acid Equilibrium P.182

14. Fraction of Dissociation Figure 8-4 compares the fraction of dissociation of two weak acids as a function of formal concentration. acid increase as it is diluted. P.184

15. Chapter 9 Buffers Buffered solution resists changes in pH when small amounts of acids or base are added or when dilution occurs. pH dependence of the rate of a particular enzyme-catalyzed reaction. The rate near pH 8 is twice as the rate at pH 7 or 9

16. 9.2 The Henderson-Hasselbalch eqn

17. If pH = pKa, [HA] = [A - ] If pH [A - ] If pH > pKa, [HA] < [A - ]

18. 9-3 A Buffer in Action Example: find the pH of a buffer solution at p. 198 Effect of adding acid to a buffer

19. 9-4 Preparing Buffers Example at p. 202 In the real life p. 203

20. Preparing a Buffer in Real Life Suppose you wish to prepare 1.00 L of buffer containing 0.100 M tris at pH 7.60. When we say 0.100 M tris, we mean that the total concentration of tris plus tris H+ will be 0.100M. Procedure: 1. Weigh out 0.100 mol tris hydrochloride and dissolve it in a beaker containing about 800 mL water and a stirring bar. 2. Place a pH electrode in the solution and monitor the pH. 3. Add NaOH solution until the pH is exactly 7.60. The electrode does not respond instantly. 4. Transfer the solution to a volumetric flask and wash the beaker and stirring bar a few times. Add the washings to the volumetric flask. 5. Dilute to the mark and mix. P.202

21. 9.5 Buffer capacity -1 The amount of H + or OH - that buffered solution can absorb without a significant change in pH Buffer capacity measures how well a solution resists changes in pH when acid or base is added. The greater the buffer capacity, the less the pH changes.

22. 9.5 Buffer capacity -2 2) Magnitudes of [HA] and [A - ]  the capacity of a buffered soln. Ex : soln A : 5.00 M HOAc + 5.00 M NaOAc soln B : 0.05 M HOAc + 0.05 M NaOAc pH change when 0.01 mol of HCl(g) is added

23. 9.5 Buffer capacity -3 3) [A - ] / [HA] ratio  the pH of a buffered soln.

24. Table 9-2 Structures and pK a values for common buffers P.205

25. 9.6 How indicators work -1 1) Usually a weak organic acid or base that has distinctly different colors in its nonionized & ionized forms. HIn (aq)  H + (aq) + In - (aq) pK HIn nonionized ionized form form

26. 9.6 How indicators work -2 1) 2) 3)

27. 9.6 How indicators work -3 2) The useful pH range for indicator is pK HIn ± 1  (Fig 10.3) encompass the pH at equivalence point (titration curve) 3)Not all indicators change color at the same pH. (Table 9.3)

28. Table 9-3

29. Two different sets of colors