1 Chapter 17a Ionic Equilibria: Part II Buffers and Titration Curves

2 Chapter Goals 1. The Common Ion Effect and Buffer Solutions 2. Buffering Action 3. Preparation of Buffer Solutions 4. Acid-Base Indicators Titration Curves 5. Strong Acid/Strong Base Titration Curves 6. Weak Acid/Strong Base Titration Curves 7. Weak Acid/Weak Base Titration Curves 8. 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: 1 Solutions made from a weak acid plus a soluble ionic salt of the weak acid. 2 Solutions made from a weak base plus a soluble ionic salt of the weak base

5 The Common Ion Effect and Buffer Solutions 1. 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 CH 3 COOH The soluble ionic salt - sodium acetate NaCH 3 COO

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 Substitute the quantities determined in the previous relationship into the ionization expression.

8 The Common Ion Effect and Buffer Solutions Apply the simplifying assumption to both the numerator and denominator.

9 The Common Ion Effect and Buffer Solutions This is a comparison of the acidity of a pure acetic acid solution and the buffer described in Example 19-1.

10 The Common Ion Effect and Buffer Solutions Compare the acidity of a pure acetic acid solution and the buffer described in Example Solution[H + ]pH 0.15 M CH 3 COOH1.6 x M CH 3 COOH & 0.15 M NaCH 3 COO buffer 1.8 x [H + ] is 89 times greater in pure acetic acid than in buffer solution.

11 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:

12 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:

13 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:

14 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.

15 Weak Bases plus Salts of Weak Bases 2. Buffers that contain a weak base plus the salt of a weak base One example of this buffer system is ammonia plus ammonium nitrate.

16 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, NH 3, and 0.30 M in ammonium nitrate, NH 4 NO 3.

17 Weak Bases plus Salts of Weak Bases Substitute the quantities determined in the previous relationship into the ionization expression for ammonia.

18 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 NH x M M NH 3 & 0.15 M NH 4 NO 3 buffer 9.0 x M8.95 The [OH-] in aqueous ammonia is 180 times greater than in the buffer.

19 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:

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

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

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

23 Buffering Action These movies show that buffer solutions resist changes in pH.

24 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. 1 Calculate the pH of the original buffer solution.

25 Buffering Action

26 Buffering Action 2 Next, calculate the concentration of all species after the addition of the gaseous HCl. The HCl will react with some of the ammonia and change the concentrations of the species. This is another limiting reactant problem.

27 Buffering Action

28 Buffering Action 3 Using the concentrations of the salt and base and the Henderson-Hassselbach equation, the pH can be calculated.

29 Buffering Action

30 Buffering Action 4 Finally, calculate the change in pH.

31 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!

32 Buffering Action pH of the original buffer solution is 8.95, from above. 1. First, calculate the concentration of all species after the addition of NaoH. NaOH will react with some of the ammonium chloride. The limiting reactant is the NaOH.

33 Buffering Action

34 Buffering Action 2 Calculate the pH using the concentrations of the salt and base and the Henderson-Hasselbach equation.

35 Buffering Action 3 Calculate the change in pH.

36 Buffering Action This table is a summary of examples 19-3 and Notice that the pH changes only slightly in each case. Original Solution Original pH Acid or base added New pH  pH 1.00 L of solution containing M NH 3 and M NH 4 Cl mol NaOH mol HCl

37 Preparation of Buffer Solutions This move shows how to prepare a buffer.

38 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. Determine the amounts of acetic acid and sodium hydroxide prior to the acid-base reaction.

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

40 Preparation of Buffer Solutions After the two solutions are mixed, the total volume of the solution is 300 mL (100 mL of NaOH mL of acetic acid). The concentrations of the acid and base are:

41 Preparation of Buffer Solutions Substitution of these values into the ionization constant expression (or the Henderson-Hasselbach equation) permits calculation of the pH.

42 Preparation of Buffer Solutions For biochemical situations, it is sometimes important to prepare a buffer solution of a given pH. Example 19-6:Calculate the number of moles of solid ammonium chloride, NH 4 Cl, 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 Because pH = 9.15, the pOH can be determined.

43 Preparation of Buffer Solutions The appropriate equilibria representations are:

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

45 Preparation of Buffer Solutions

46 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:

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

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

49 Acid-Base Indicators Color change ranges of some acid-base indicators Indicator Color in acidic rangepH range Color in basic range Methyl violetYellow0 - 2Purple Methyl orangePink3.1 – 4.4Yellow LitmusRed4.7 – 8.2Blue PhenolphthaleinColorless8.3 – 10.0Red

50 Titration Curves 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.

51 Strong Acid/Strong Base Titration Curves Before any KOH is added the pH of the HClO 4 solution is Remember perchloric acid is a strong acid that ionizes essentially 100%.

52 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 ___?

53 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 ___?

54 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 ____?

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

56 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.

57 Weak Acid/Strong Base Titration Curves As an example, consider the titration of mL of M acetic acid, CH 3 COOH, (a weak acid) with M KOH (a strong base). The acid and base react in a 1:1 mole ratio.

58 Weak Acid/Strong Base Titration Curves Before the equivalence point is reached, both CH 3 COOH and KCH 3 COO are present in solution forming a buffer. The KOH reacts with CH 3 COOH to form KCH 3 COO. A weak acid plus the salt of a weak acid form a buffer. Hypothesize how the buffer production will effect the titration curve.

59 Weak Acid/Strong Base Titration Curves 1. Determine the pH of the acetic acid solution before the titration is begun. Same technique as used in Chapter 18.

60 Weak Acid/Strong Base Titration Curves

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

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

63 Weak Acid/Strong Base Titration Curves At the equivalence point, the solution is M in KCH 3 COO, 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.

64 Weak Acid/Strong Base Titration Curves 1. Set up the equilibrium reaction:

65 Weak Acid/Strong Base Titration Curves 2. Determine the concentration of the salt in solution.

66 Weak Acid/Strong Base Titration Curves 3. Perform a hydrolysis calculation for the potassium acetate in solution.

67 Weak Acid/Strong Base Titration Curves 4. After the equivalence point is reached, the pH is determined by the excess KOH just as in the strong acid/strong base example.

68 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.

69 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.

70 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.

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

72 Synthesis Question Aspirin is acetyl salicylic acid. So to buffer it all that would have to be added is the salt of acetyl salicylic acid.

73 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?

74 Group Question

75 End of Chapter 19 We have examined : 1 Gas phase equilibria in Chapter 17 2 Hydrolysis equilibria in Chapter 18 3 Acid/base equilibria in Chapter 19 Chapter 20 is the last equilibrium chapter. It involves solid/solution equilibria.