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Buffers and Titrations

Buffers and Titrations. Chapter 19. The Common Ion Effect & 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

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Buffers and Titrations

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  1. Buffers and Titrations Chapter 19

  2. The Common Ion Effect & 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. Weak Acids plus Salts of Weak Acids For example ~ acetic acid CH3COOH and sodium acetate NaCH3COO

  4. Ex. 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. Ka = 1.8 x 10-5 • (note: sodium acetate completely dissociates) • R CH3COOH + H2O  CH3COO- + H3O+ • 0.15 0.15 0 • -x +x +x • E. 0.15 – x 0.15 + x x

  5. Compare the acidity of a pure acetic acid solution and the buffer we just described. Notice that [H+] is 89 times greater in pure acetic acid than in buffer solution.

  6. Weak Bases plus Salts of Weak Bases Ex.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. Kb = 1.8 x 10-5 R NH3 + H2O  NH4+ + OH- I 0.15 0.30 0 C -x + x + x E 0.15 –x 0.30 + x x

  7. Substitute these values into the ionization expression for ammonia and solve algebraically.

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

  9. Henderson-Hasselbach equation For acids: For bases: Remember: pX = -log X

  10. Buffering Action • Buffer solutions resist changes in pH. Ex. 3) 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.

  11. 1st ~ Calculate the pH of the original buffer solution

  12. 2nd ~ Calculate the concentration of all species after the addition of HCl. • HCl will react with some of the ammonia

  13. 3rd ~ Now that you have the concentrations of our salt and base, you can calculate the new pH.

  14. 4th ~ Calculate the change in pH.

  15. Ex. 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.

  16. Preparation of Buffer Solutions Ex. 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 (before reaction)

  17. Preparation of Buffer Solutions • For biochemical situations, it is sometimes important to prepare a buffer solution of a given pH. Ex. 6) A) Find 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 B) What mass is needed?

  18. 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

  19. Common Acid-Base Indicators

  20. Strong Acid/Strong Base Titration Curves • Titration curves are graphs that show the pH at various amounts of titrate added. Allows you to find the equivalence point. • For Titration curves, Plot pH vs. Volume of acid or base added in titration.

  21. Ex. 7) Consider the titration of 100.0 mL of 0.100 Mperchloric acid with 0.100 M potassium hydroxide. Find the equivalence point of this rxn. • Plot pH vs. mL of KOH added • 1:1 mole ratio

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

  23. After 20.0 mL of 0.100 M KOH has been added the new pH is 1.17.

  24. After 50.0 mL of 0.100 M KOH has been added the pH is 1.48.

  25. After 90.0 mL of 0.100 M KOH has been added the pH is 2.28.

  26. After 100.0 mL of 0.100 M KOH has been added the pH is 7.00.

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

  28. Weak Acid/Strong Base Titration Curves Salts of weak acids and strong bases hydrolyze to give basic solns so the soln is basic at the equivalence point and the soln is buffered before the equivalence point.

  29. Strong Acid/Weak BaseTitration Curves • Titration curves for Strong Acid/Weak Bases look similar to Strong Base/Weak Acid but they are inverted. The soln is buffered before the equivalence point and is acidic at the equivalence point.

  30. Weak Acid/Weak BaseTitration Curves • Titration curves have very short vertical sections. • Solution is buffered both before and after the equivalence point. • Visual indicators cannot be used. Instead you can measure the conductivity in order to find the end point. • The math is complex, we will not worry about it in AP Chem. 

  31. Fun Chem problem for you • Blood is slightly basic, having a pH of 7.35 to 7.45. What chemical species causes our blood to be basic? How does our body regulate the pH of blood?

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