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Chapter 14

Chapter 14. Applications of Equilibrium Principles (acid/base systems). Chapter 14. Buffers - applies stoichiometry and equilibria Acid-Base Titrations. Chapter 14. Two system types: a weak acid, HB, and its conjugate base, B - (buffer system). Chapter 14.

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Chapter 14

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  1. Chapter 14 • Applications of Equilibrium Principles (acid/base systems)

  2. Chapter 14 • Buffers - applies stoichiometry and equilibria • Acid-Base Titrations

  3. Chapter 14 • Two system types: • a weak acid, HB, and its conjugate base, B- (buffer system)

  4. Chapter 14 • An acid and a base used in an acid-base titration (especially a polyprotic acid)

  5. Buffer System • Buffers contain roughly equal amounts of a weak acid and its conjugate base.

  6. Buffer System Highly resistant to changes in pH brought about by the addition of a strong acid or strong base. (why?)

  7. Buffer System • The buffer can react with either H3O+ (H+) ions or OH- ions. • HB(aq) + OH-(aq) --> B-(aq) + H2O • B-(aq) + H+(aq) --> HB(aq)

  8. Buffer System • The reactions have large Keq values. They go almost to completion, so the H+ ions or OH- ions are consumed. • The pH is not directly affected.

  9. Buffer System Has a pH close to the pKa of the weak acid. Why?

  10. Buffer System • HB(aq) <--> H+(aq) + B-(aq) • Ka = [H+][B-]/[HB] • Solve for [H+]

  11. Buffer System • 1.) Assume equilibrium is established without appreciably changing the concentrations of either HB or B-.

  12. Buffer System • 1.) [HB] = [HB]o • [B-] = [B-]o

  13. Buffer System • 2.) Since HB and B-are present in the same solution, the ratio of their concentrations equals their mole ratio. • [HB]/[B-] = nHB/nB-

  14. Buffer System • 2.) [H+] = Ka(nHB/nB-) • Or… Henderson-Hasselbach equation: pH=pKa+log([B-]/[HB])

  15. Buffer System Calculate [H+] and pH in a solution prepared by adding 0.200 mol of acetic acid and 0.200 mol of sodium acetate to one liter. [H+] = (1.8 x 10-5) x (.200/.200) = 1.8 x 10-5 M, pH = 4.74

  16. Choosing a Buffer • To make a buffer in the lab at a required pH: note that HB and B- are present in roughly equal amounts and that pH is roughly equal to pKa of the weak acid.

  17. Choosing a Buffer • Most important factor: choose a buffer system with a Ka close to the pH desired (see table 14.1, p. 382 - weak acids and their salts).

  18. Choosing a Buffer • Ex. To establish a buffer of pH 7, choose a system such as H2PO4-, HPO42-, where Ka = 6.2 x 10-8.

  19. Choosing a Buffer • Ex. At pH 7, [H+] = 1.0 x 10-7 M. • So…[H2PO4-] / [HPO42-] = [H+] / Ka = (1.0 x 10-7)/(6.2 x 10-8)=1.6

  20. Choosing a Buffer • Ex. So… add 160 ml of 0.100 M H2PO4- to 100 ml of 0.100 M HPO42-.

  21. Choosing a Buffer • Another method: partial neutralization of a weak acid and weak base gives a buffer.

  22. Choosing a Buffer • Ex. Reacting 0.10 mol of HCl with 0.20 mol of NH3, consumes the H+, producing 0.10 mol of NH4+ and 0.10 mol NH3

  23. Distribution Curves • See figure 14.3, p. 385. • Buffers are only effective over small ranges: there must be appreciable amounts of both HB and B- for the system to work.

  24. Distribution Curves • Ex. P. 385, 389, 390 • At a pH of 6.4, [HB] = [B-], pH = pKa. • At pH of 5.4 - 7.4 there is enough of either species to consume the addition of a strong acid or base.

  25. Added H+ or OH- • The pH of a buffer does change slightly as a strong acid or base is added. • Addition of H+ converts B- to HB. Addition of OH- converts HB to B- and H2O.

  26. Added H+ or OH- • Use stoichiometry to determine how the [HB] to [B-] ratio has changed. Then use equilibrium principles to ultimately calculate the new pH.

  27. Added H+ or OH- Calculate the pH of a buffer containing 0.200 mol of acetic acid and 0.200 mol of sodium acetate after the addition of 0.020 mol NaOH. [H+] = (1.8 x 10-5) x (.180/.220) = 1.5 x 10-5 M, pH = 4.82

  28. Indicators • Useful in determining the equivalence point of a titration.

  29. Indicators • It is a weak acid and its conjugate base. When one dominates, it is a different color then when the other dominates.

  30. Indicators • HIn (aq) <--> H+ (aq) + In- (aq). Ka = [H+][In-]/[HIn] • 1. If [HIn]/[In-] is >10, then the acid color dominates. • 2. If < 0.1, then the base color dominates.

  31. Indicators • 3. If [HIn]/[In-] approx. 1, then the color is an intermediate of the two.

  32. Indicators • Ka = [H+][In-]/[HIn], so…. • [HIn]/[In-] = [H+]/Ka • What two factors, then, effect [HIn]/[In-], and hence, the color of the indicator?

  33. Indicators • 1. [H+] or the pH of the solution. • High pH vs. low pH? • 2. Ka of the indicator. Indicators change color at different pHs (when pH = pKa).

  34. Distribution Curves • See table 14.2, p. 387. • Indicators are only effective over small ranges. There must be appreciable amounts of both the HB and B- for the system to work.

  35. Titrations • Titration: an analytical method in which a standard solution is used to determine the concentration of another solution.

  36. Titrations • A neutralization reaction occurs. Dependent on the indicator used, we can determine the volume at which the mixture reaches a certain pH.

  37. Titrations • Solving steps: 1.) balance the neutralization equation. • 2.) from the given determine the amount of moles used.

  38. Titrations • Solving steps: 3.) use the equation to determine the mole ratios. • 4.) Determine the molarity or the volume of the unknown.

  39. Titrations • If 26.4 cm3 of LiOH solution are required to neutralize 21.7 cm3 of 0.500 M HBr, what is the concentration of the basic solution?

  40. Acid/Base Titrations • Strong Acid-Strong Base • H+ + OH- --> H2O, K = 1/Kw = 1.0 x 1014. • pH = 7 at the end point, changes rapidly near end point.

  41. Acid/Base Titrations • Weak Acid-Strong Base • HA + OH- --> H2O + A-, K = 1/Kb. • Ex. If Ka = 1.0 x 10-5, Kb = 1.0 x 10-9, K = 1.0 x 109.

  42. Acid Base Titrations 0.10 M HA (Ka = 1 x 10-5) is titrated with a strong base. Determine the initial pH, the pH when the titration is half finished, and the pH at the end point. pH = 3.0, 5.0, and 9.0

  43. Acid/Base Titrations • In general, pH > 7 at the end point, changes slowly near the end point (buffer system).

  44. Acid/Base Titrations • Strong Acid-Weak Base • H+ + A- --> HA, K = 1/(Ka of HA). • pH < 7 at the end point, changes slowly near the end point.

  45. Choice of Indicator • 1. Weak acid - strong base: use an indicator such as (table 14.2 p. 387) phenolphthalein, which changes color above pH 7.

  46. Choice of Indicator • 2. Strong acid - weak base: use an indicator such as methyl red, which changes color below pH 7.

  47. Choice of Indicator • 3. Strong acid - strong base: use any indicator because the pH changes rapidly near the end point.

  48. Polyprotic Acids • Virtually all the H+ comes from the first dissociation. • Ex. H2CO3 <--> H+ + HCO3-, K1 = 4x10-7. • HCO3- <--> H+ + CO32-, K2 = 4 x 10-11

  49. Polyprotic Acids • In 0.1 M H2CO3, [H+] = [HCO3-] = 2 x 10-4 M, [CO32-] = 4 x 10-11 M

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