1. Ionic Equilibria: Part II Buffers and Titration Curves
Chapter 19
Ionic Equilibria: Part II
Buffers and Titration Curves

2. Chapter Goals The Common Ion Effect and Buffer Solutions
Buffering Action
Preparation of Buffer Solutions
Acid-Base Indicators
Titration Curves
Strong Acid/Strong Base Titration Curves
Weak Acid/Strong Base Titration Curves
Weak Acid/Weak Base Titration Curves
Summary of Acid-Base Calculations

3. The Common Ion Effect and Buffer Solutions
If a solution is made in which the same ion is produced by two different compounds the common ion effect is exhibited.
Buffer solutions are solutions that resist changes in pH when acids or bases are added to them.
Buffering is due to the common ion effect.

4. The Common Ion Effect and Buffer Solutions
There are two common kinds of buffer solutions:
Solutions made from a weak acid plus a soluble ionic salt of the weak acid.
Solutions made from a weak base plus a soluble ionic salt of the weak base

5. The Common Ion Effect and Buffer Solutions
Solutions made of weak acids plus a soluble ionic salt of the weak acid
One example of this type of buffer system is:
The weak acid - acetic acid CH3COOH
The soluble ionic salt - sodium acetate NaCH3COO

6. The Common Ion Effect and Buffer Solutions
Example 19-1: 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.
This is another equilibrium problem with a starting concentration for both the acid and anion.

7. The Common Ion Effect and Buffer Solutions

8. The Common Ion Effect and Buffer Solutions
Compare the acidity of a pure acetic acid solution and the buffer described in Example 19-1.
Solution
[H+] pH
0.15 M CH3COOH
1.6 x 10-3
2.80
0.15 M CH3COOH &
0.15 M NaCH3COO buffer
1.8 x 10-5
4.74
[H+] is 89 times greater in pure acetic acid than in buffer solution.

9. The Common Ion Effect and Buffer Solutions
The general expression for the ionization of a weak monoprotic acid is:
The generalized ionization constant expression for a weak acid is:

10. The Common Ion Effect and Buffer Solutions
If we solve the expression for [H+], this relationship results:
By making the assumption that the concentrations of the weak acid and the salt are reasonable, the expression reduces to:

11. The Common Ion Effect and Buffer Solutions
The relationship developed in the previous slide is valid for buffers containing a weak monoprotic acid and a soluble, ionic salt.
If the salt’s cation is not univalent the relationship changes to:

12. The Common Ion Effect and Buffer Solutions
Simple rearrangement of this equation and application of algebra yields the
Henderson-Hasselbach equation.
The Henderson-Hasselbach equation is one method to calculate the pH
of a buffer given the concentrations of the salt and acid.

13. Weak Bases plus Salts of Weak Bases
Buffers that contain a weak base plus the salt of a weak base
One example of this buffer system is ammonia plus ammonium nitrate.

14. Weak Bases plus Salts of Weak Bases
Example 19-2: 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.

15. Weak Bases plus Salts of Weak Bases
A comparison of the aqueous ammonia concentration to that of the buffer described above shows the buffering effect.
Solution
[OH-] pH
0.15 M NH3
1.6 x 10-3 M
11.20
0.15 M NH3 &
0.15 M NH4NO3 buffer
9.0 x 10-6 M
8.95
The [OH-] in aqueous ammonia is 180 times greater than in the buffer.

16. Weak Bases plus Salts of Weak Bases
We can derive a general relationship for buffer solutions that contain a weak base plus a salt of a weak base similar to the acid buffer relationship.
The general ionization equation for weak bases is:

17. Weak Bases plus Salts of Weak Bases
The general form of the ionization expression is:
Solve for the [OH-]

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

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

20. Buffering Action

21. Buffering Action
Example 19-3: If mole of gaseous HCl is added to 1.00 liter of a buffer solution that is M in aqueous ammonia and M in ammonium chloride, how much does the pH change? Assume no volume change due to addition of the HCl.
Calculate the pH of the original buffer solution.

22. Buffering Action
Example 19-4: If mole of NaOH is added to 1.00 liter of solution that is M in aqueous ammonia and M in ammonium chloride, how much does the pH change? Assume no volume change due to addition of the solid NaOH.
You do it!

23. Buffering Action This table is a summary of examples 19-3 and 19-4.
Original Solution
Original pH
Acid or base added
New pH
pH
1.00 L of solution containing
0.100 M NH3 and M NH4Cl
8.95
0.020 mol NaOH
9.08
+0.13
0.020 mol HCl
8.81
-0.14
Notice that the pH changes only slightly in each case.

24. Preparation of Buffer Solutions

25. Preparation of Buffer Solutions
Example 19-5: Calculate the concentration of H+ and the pH of the solution prepared by mixing 200 mL of M acetic acid and 100 mL of M sodium hydroxide solutions.

26. Preparation of Buffer Solutions
For biochemical situations, it is sometimes important to prepare a buffer solution of a given pH.

27. Preparation of Buffer Solutions
Example 19-6: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.

28. Acid-Base Indicators
The point in a titration at which chemically equivalent amounts of acid and base have reacted is called the equivalence point.
The point in a titration at which a chemical indicator changes color is called the end point.
A symbolic representation of the indicator’s color change at the end point is:

29. Acid-Base Indicators
The equilibrium constant expression for an indicator would be expressed as:

30. Acid-Base Indicators
If the preceding expression is rearranged the range over which the indicator changes color can be discerned.

31. Color change ranges of some acid-base indicators
Color in acidic range
pH range
Color in basic range
Methyl violet
Yellow
0 - 2
Purple
Methyl orange
Pink
3.1 – 4.4
Litmus
Red
4.7 – 8.2
Blue
Phenolphthalein
Colorless
8.3 – 10.0

32. Strong Acid/Strong Base Titration Curves
These graphs are a plot of pH vs. volume of acid or base added in a titration.
As an example, consider the titration of mL of M perchloric acid with M potassium hydroxide.
In this case, we plot pH of the mixture vs. mL of KOH added.
Note that the reaction is a 1:1 mole ratio.

33. Strong Acid/Strong Base Titration Curves
Before any KOH is added the pH of the HClO4 solution is 1.00.
Remember perchloric acid is a strong acid that ionizes essentially 100%.

34. Strong Acid/Strong Base Titration Curves
After a total of 20.0 mL M KOH has been added the pH of the reaction mixture is ___?

35. Strong Acid/Strong Base Titration Curves
After a total of 50.0 mL of M KOH has been added the pH of the reaction mixture is ___?

36. Strong Acid/Strong Base Titration Curves
After a total of 90.0 mL of M KOH has been added the pH of the reaction mixture is ____?

37. Strong Acid/Strong Base Titration Curves
After a total of mL of M KOH has been added the pH of the reaction mixture is ___?

38. 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.

39. Weak Acid/Strong Base Titration Curves
As an example, consider the titration of mL of M acetic acid, CH3 COOH, (a weak acid) with M KOH (a strong base).
The acid and base react in a 1:1 mole ratio.

40. Weak Acid/Strong Base Titration Curves
Before the equivalence point is reached, both CH3COOH and KCH3COO are present in solution forming a buffer.
The KOH reacts with CH3COOH to form KCH3COO.
A weak acid plus the salt of a weak acid form a buffer.
Hypothesize how the buffer production will effect the titration curve.

41. Weak Acid/Strong Base Titration Curves
Determine the pH of the acetic acid solution before the titration is begun.
Same technique as used in Chapter 18.

42. Weak Acid/Strong Base Titration Curves

43. Weak Acid/Strong Base Titration Curves
After a total of 20.0 mL of KOH solution has been added, the pH is:

44. Weak Acid/Strong Base Titration Curves
Similarly for all other cases before the equivalence point is reached.

45. Weak Acid/Strong Base Titration Curves
At the equivalence point, the solution is M in KCH3COO, the salt of a strong base and a weak acid which hydrolyzes to give a basic solution.
This is a solvolysis process as discussed in Chapter 18.
Both processes make the solution basic.
The solution cannot have a pH=7.00 at equivalence point.
Let us calculate the pH at the equivalence point.

46. Weak Acid/Strong Base Titration Curves
Set up the equilibrium reaction:

47. Weak Acid/Strong Base Titration Curves
Determine the concentration of the salt in solution.

48. Weak Acid/Strong Base Titration Curves
Perform a hydrolysis calculation for the potassium acetate in solution.

49. Weak Acid/Strong Base Titration Curves
After the equivalence point is reached, the pH is determined by the excess KOH just as in the strong acid/strong base example.

50. Weak 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.

51. Strong Acid/Weak Base Titration Curves
Titration curves for Strong Acid/Weak Base Titration Curves look similar to Strong Base/Weak Acid Titration Curves but they are inverted.

52. Weak Acid/Weak Base Titration Curves
Weak Acid/Weak Base Titration curves have very short vertical sections.
The solution is buffered both before and after the equivalence point.
Visual indicators cannot be used.

53. Synthesis Question
Bufferin is a commercially prepared medicine that is literally a buffered aspirin. How could you buffer aspirin? Hint - what is aspirin?

54. Group Question
Blood is slightly basic, having a pH of 7.35 to What chemical species causes our blood to be basic? How does our body regulate the pH of blood?

55. Group Question