Chapter Fourteen

1. Chapter Fourteen • Aqueous Equilibria

2. The Common Ion Effect and Buffer Solutions • Common ion effect - solutions in which the same ion is produced by two different compounds • Buffer solutions - resist changes in pH when acids or bases are added to them • due to common ion effect • Two common kinds of buffer solutions • solutions of a weak acid plus a soluble ionic salt of the weak acid • solutions of a weak base plus a soluble ionic salt of the weak base

3. The Common Ion Effect and Buffer Solutions • Weak Acids plus Salts of Weak Acids • acetic acid CH3COOH • sodium acetate NaCH3COO

4. The Common Ion Effect and Buffer Solutions • Example: Calculate the concentration of H+and the pH of a solution that is 0.15 M in acetic acid and 0.15 M in sodium acetate.

5. The Common Ion Effect and Buffer Solutions • Example: Calculate the concentration of H+and the pH of a solution that is 0.15 M in acetic acid and 0.15 M in sodium acetate.

6. The Common Ion Effect and Buffer Solutions • Substitute these quantities into the ionization expression.

7. The Common Ion Effect and Buffer Solutions • Apply the simplifying assumption

8. The Common Ion Effect and Buffer Solutions • Compare the acidity of a pure acetic acid solution and the buffer we just described.

9. The Common Ion Effect and Buffer Solutions • Compare the acidity of a pure acetic acid solution and the buffer we just described. • [H+] is 89 times greater in pure acetic acid than in buffer solution.

10. The Common Ion Effect and Buffer Solutions • General expression for the ionization of a weak monoprotic acid.

11. The Common Ion Effect and Buffer Solutions • Its ionization constant expression is

12. The Common Ion Effect and Buffer Solutions • Solve the expression for [H+]

13. The Common Ion Effect and Buffer Solutions • Making the assumption that the concentrations of the weak acid and the salt are reasonable. • The expression reduces to

14. The Common Ion Effect and Buffer Solutions • The above relationship is valid for buffers containing a weak monoprotic acid and a soluble, ionic salt. • The relationship changes if the salt’s cation is not univalent to

15. The Common Ion Effect and Buffer Solutions • Simple rearrangement of this equation and application of algebra yields the Henderson-Hasselbach equation

16. Weak Bases plus Salts of Weak Bases • Buffers that contain a weak base plus the salt of a weak base - for example - ammonia plus ammonium nitrate.

17. Weak Bases plus Salts of Weak Bases • Buffers that contain a weak base plus the salt of a weak base - for example - ammonia plus ammonium nitrate.

18. Weak Bases plus Salts of Weak Bases • Example: Calculate the concentration of OH- and the pH of the solution that is 0.15 M in aqueous ammonia, NH3, and 0.30 M in ammonium nitrate, NH4NO3.

19. Weak Bases plus Salts of Weak Bases • Substitute these values into the ionization expression for ammonia and solve algebraically.

20. Weak Bases plus Salts of Weak Bases • Let’s compare the aqueous ammonia concentration to that of the buffer described above.

21. Weak Bases plus Salts of Weak Bases • Let’s compare the aqueous ammonia concentration to that of the buffer described above. • The [OH-] in aqueous ammonia is 180times greater than in the buffer.

22. Weak Bases plus Salts of Weak Bases • Derive a general relationship for buffer solutions that contain a weak base plus a salt of a weak base. • Ionization equation

23. Weak Bases plus Salts of Weak Bases • Ionization expression • general form

24. Weak Bases plus Salts of Weak Bases • For salts that have univalent ions: • For salts that have divalent or trivalent ions

25. Weak Bases plus Salts of Weak Bases • Simple rearrangement of this equation and application of algebra yields the Henderson-Hasselbach equation

26. Buffering Action • Buffer solutions resist changes in pH. • Example: If 0.020 mole of HCl is added to 1.00 liter of solution that is 0.100 M in aqueous ammonia and 0.200 M in ammonium chloride, how much does the pH change? Assume no volume change due to addition of the gaseous HCl. • Calculate the pH of the original buffer solution

27. Buffering Action

28. Buffering Action • Now we calculate the concentration of all species after the addition of HCl. • HCl will react with some of the ammonia

29. Buffering Action • Now that we have the concentrations of our salt and base, we can calculate the pH.

30. Buffering Action • Calculate the change in pH.

31. Preparation of Buffer Solutions • Example: Calculate the concentration of H+ and the pH of the solution prepared by mixing 200 mL of 0.150 M acetic acid and 100 mL of 0.100 M sodium hydroxide solutions. • Determine the amounts of acetic acid and sodium hydroxide (before reaction)

32. Preparation of Buffer Solutions • Example: Calculate the concentration of H+ and the pH of the solution prepared by mixing 200 mL of 0.150 M acetic acid and 100 mL of 0.100 M sodium hydroxide solutions. • Determine the amounts of acetic acid and sodium hydroxide (before reaction)

33. Preparation of Buffer Solutions • Sodium hydroxide and acetic acid react in a 1:1 mole ratio.

34. Preparation of Buffer Solutions • After the two solutions are mixed, the total volume is 300 mL (100 + 200), and the concentrations are:

35. Preparation of Buffer Solutions • Substitution into the ionization constant expression (or Henderson-Hasselbach equation) gives

36. Preparation of Buffer Solutions • For biochemical situations, it is sometimes important to prepare a buffer solution of a given pH. • Example: Calculate the number of moles of solid ammonium chloride, NH4Cl, that must be used to prepare 1.00 L of a buffer solution that is 0.10 M in aqueous ammonia, and that has a pH of 9.15. • Because pH=9.15

37. Preparation of Buffer Solutions • For biochemical situations, it is sometimes important to prepare a buffer solution of a given pH. • Example:Calculate the number of moles of solid ammonium chloride, NH4Cl, that must be used to prepare 1.00 L of a buffer solution that is 0.10 M in aqueous ammonia, and that has a pH of 9.15. • Because pH=9.15

38. Preparation of Buffer Solutions • Appropriate equations and equilibria representations are:

39. Preparation of Buffer Solutions • Substitute into the ionization constant expression (or Henderson-Hasselbach equation) for aqueous ammonia

40. Acid-Base Indicators • Equivalence point - point at which chemically equivalent amounts of acid and base have reacted • End point - point at which chemical indicator changes color

41. Acid-Base Indicators • Equilibrium constant expression for an indicator would be

42. Acid-Base Indicators • Rearrange this expression to get a feeling for range over which indicator changes color.

43. Acid-Base Indicators Some Acid-Base Indicators

44. Strong Acid/Strong Base Titration Curves • Plot pH vs. Volume of acid or base added in titration. • Consider the titration of 100.0 mL of 0.100 M perchloric acid with 0.100 M potassium hydroxide. • Plot pH vs. mL of KOH added • 1:1 mole ratio

45. Strong Acid/Strong Base Titration Curves • Before titration starts the pH of the HClO4 solution is 1.00. • Remember perchloric acid is a strong acid

46. Strong Acid/Strong Base Titration Curves • After 20.0 mL of 0.100 M KOH has been added the pH is 1.17.

47. Strong Acid/Strong Base Titration Curves • After 50.0 mL of 0.100 M KOH has been added the pH is 1.48.

48. Strong Acid/Strong Base Titration Curves • After 90.0 mL of 0.100 M KOH has been added the pH is 2.28.

49. Strong Acid/Strong Base Titration Curves • After 100.0 mL of 0.100 M KOH has been added the pH is 7.00.

50. Strong Acid/Strong Base Titration Curves • We have calculated only a few points on the titration curve. Similar calculations for remainder of titration show clearly the shape of the titration curve.