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Chapter 17a

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

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Chapter 17a

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  1. Chapter 17a • 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 • 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 19-1. • [H+] is 89 times greater in pure acetic acid than in buffer solution.

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

  12. 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: The Common Ion Effect and Buffer Solutions

  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 • 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, NH3, and 0.30 M in ammonium nitrate, NH4NO3.

  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. • The [OH-] in aqueous ammonia is 180times 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. The general form of the ionization expression is: Solve for the [OH-] Weak Bases plus Salts of Weak Bases

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

  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 0.020 mole of gaseous HCl is added to 1.00 liter of a buffer 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 HCl. • Calculate the pH of the original buffer solution.

  25. Buffering Action

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

  29. Buffering Action

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

  31. Buffering Action • Example 19-4: If 0.020 mole of NaOH 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 solid NaOH. You do it!

  32. Buffering Action • pH of the original buffer solution is 8.95, from above. • 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 • Calculate the pH using the concentrations of the salt and base and the Henderson-Hasselbach equation.

  35. Buffering Action • Calculate the change in pH.

  36. This table is a summary of examples 19-3 and 19-4. Notice that the pH changes only slightly in each case. Buffering Action

  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 0.150 M acetic acid and 100 mL of 0.100 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 + 200 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, 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, 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

  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 100.0 mL of 0.100 M perchloric acid with 0.100 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.

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