1. Reactions Between Acids & Bases Chapter 16

2. 16-1 & 16-2 Titrations of Strong Acids and Bases Objectives To define analyte and titrant To calculate the volume of titrants needed to reach the equivalence point To identify regions of the titration curve in which the analyte or titrant is in excess To express amounts of analyte and titrant in units of millimoles To calculate the concentration of all species present during the titration of a strong acid with a strong base To calculate and graph the titration curve

3. 16-1 & 16-2 Titrations of Strong Acids and Bases Titration – a procedure for the determination of the quantity of one substance by the addition of a measured amount of a second substance The analyte is the substance whose concentration is being determined The titrant is the substance added to react with the analyte Analyte + Titrant → Products

4. 16-1 & 16-2 Titrations of Strong Acids and Bases Basic procedure for acid-base titrations: – A known volume of an acid of unknown [ ] is placed in a flask – An indicator solution that changes color in a particular pH range is added to the acid Ex: phenolphthalein – colorless in acid/pink in base – A standard solution of base is added until a color change occurs (the endpoint)

5. 16-1 & 16-2 Titrations of Strong Acids and Bases Alternatively, the indicator can be left out and a pH meter can used to measure the pH as the titrant is added. A titration curve is a graph of the pH of a solution as a function of the amount of titrant that is added Equivalence Point = point in the titration when exactly one mole of base has been added for each mole of acid present in the sample Inflection point = point at which pH changes most rapidly

6. Strong Acid/Strong Base Titration A solution that is 0.10 M HCl is titrated with 0.10 M NaOH Equivalence point is at pH 7

7. Strong Acid/Strong Base Titration A solution that is 0.10 M NaOH is titrated with 0.10 M HCl Equivalence point is at pH 7 It is important to recognize that titration curves are not always increasing from left to right.

8. 16-1 & 16-2 Titrations of Strong Acids and Bases Calculate the pH in the titration of 20.0 mL of 0.125 M HCl with 0.250 M NaOH after adding 0.0, 5.0, 9.60, 10.40, and 12.00 mL of NaOH. pH = 0.9, 1.30, 2.47, 11.52, 12.19

9. Fig. 16-6, p. 693

10. 16-3 Buffers Objectives To define a pH buffer and identify buffer solutions To calculate the pH of a buffer solution containing a weak acid and its conjugate base To identify the circumstances under which the common approximation equations fail to describe the system

11. 16-3 Buffers Buffers are solutions that resist changes in pH when either OH - or H + ions are added Strong acids and bases cannot be used as buffers! They form salt and water. HCl + NaOH → NaCl + H 2 O Rather, buffered solutions contain either: – a weak acid and its salt or – a weak base and its salt

12. 16-3 Buffers Consider a buffer made from a weak acid and its conjugate base: If hydrogen ions are added, they react with the conjugate base: H 3 O + + A -  H 2 O + HA Note that this neutralization reaction is the reverse of the dissociation of a weak acid. Therefore, K eq for neutralization is = 1/K a This means K eq is large and the neutralization reaction goes essentially to completion

13. 16-3 Buffers Consider a buffer made from a weak acid and its conjugate base: If hydroxide ions are added, they react with the weak acid HA + OH -  H 2 O + A - This neutralization reaction also goes essentially to completion

14. 16-3 Buffers To prepare a buffer you need to add both the weak acid and a soluble salt of that acid in sufficient amounts. The aqueous reactions are as follows: CH 3 COOH(aq) + H 2 O(l) → CH 3 COO - (aq) + H 3 O + (aq) CH 3 COONa(aq) → CH 3 COO - (aq) + Na + (aq) Net: CH 3 COOH(aq) + H 2 O(l) → CH 3 COO - (aq) + H 3 O + (aq) ***BUT, note that in the net reaction, the [H 3 O + ] is not equal to the [CH 3 COO - ]

15. 16-3 Buffers The relationship between the pH of the solution and the concentration of the acid and its conjugate base is given by the Henderson-Hasselbach Equation

16. Henderson-Hasselbach Equation

17. 16-3 Buffers Note that when the concentration of the acid and the conjugate base are equal, pH = pK a The pH of buffer solutions is generally limited to the vicinity of the pK a for the acid Buffers are most effective in the range of pK a ± 1

18. 16-3 Buffers To make an effective buffer, the concentrations of the weak acid and its conjugate base must be at least 100 times the K a or K b Most pH buffers are made from solutions with concentrations that range from 0.1 to 1.0 M, so the Henderson-Hasselbach equation is a good approximation

19. 16-3 Buffers Calculate the pH of a solution that is 0.4 M sodium acetate and 0.2 M acetic acid. K a for acetic acid is 1.8 x 10 -5. A: pH = 5.04 Calculate the pH of a solution that is 0.5 M HF and 0.2 M sodium fluoride. K a for HF is 3.5 x 10 -4. A: pH = 3.05

20. 16-3 Buffers Calculate the number of grams of ammonium chloride that must be added to 500.0 mL of 0.32 M NH 3 to prepare a pH 8.50 buffer. K b for NH 3 is 1.8 x 10 -5. A: 48 g How many moles of sodium benzoate must be added to one liter of a 0.22 M solution of benzoic acid (pK a = 4.19) to prepare a liter of pH 4.50 buffer? A: 0.45 mol

21. 16-3 Buffers Concentrated buffers are generally more effective than dilute buffers The effectiveness of a buffer is measured by buffer capacity Buffer capacity = the amount of strong acid or base needed to change the pH of one liter of buffer by one unit

22. 16-3 Buffers Calculate the change in pH observed when 1.50 mL of 0.0670 M H 3 O + is added to: a) 100.0 mL of a pH 4.74 HCl solution b) 100.0 mL of a pH 4.74 buffer prepared from 0.120 M acetic acid and 0.120 M sodium acetate (the pK a for acetic acid is 4.74) A: a) pH decreases by 1.74 pH units b) pH decreases by 0.01 pH units

23. 16-3 Buffers Calculate the change in pH observed when 5.0 mL of 0.050 M NaOH are added to 100 mL of a pH 4.74 buffer prepared from 0.100 M acetic acid and 0.100 M sodium acetate (the pK a for acetic acid is 4.74) A: pH increases by 0.02 pH units

24. Acid/Salt Buffering Pairs Weak Acid Formula of the acid Example of a salt of the weak acid Hydrofluoric HF KF – Potassium fluoride Formic HCOOH KHCOO – Potassium formate Benzoic C 6 H 5 COOH NaC 6 H 5 COO – Sodium benzoate Acetic CH 3 COOH NaH 3 COO – Sodium acetate Carbonic H 2 CO 3 NaHCO 3 - Sodium bicarbonate Propanoic HC 3 H 5 O 2 NaC 3 H 5 O 2 - Sodium propanoate Hydrocyanic HCN KCN - potassium cyanide The salt will contain the anion of the acid, and the cation of a strong base (NaOH, KOH)

25. Base/Salt Buffering Pairs The salt will contain the cation of the base, and the anion of a strong acid (HCl, HNO 3 ) Base Formula of the base Example of a salt of the weak acid Ammonia NH 3 NH 4 Cl - ammonium chloride Methylamine CH 3 NH 2 CH 3 NH 3 Cl – methylammonium chloride Ethylamine C 2 H 5 NH 2 C 2 H 5 NH 3 NO 3 - ethylammonium nitrate Aniline C 6 H 5 NH 2 C 6 H 5 NH 3 Cl – aniline hydrochloride Pyridine C 5 H 5 N C 5 H 5 NHCl – pyridine hydrochloride

26. 16-4 Titrations of Weak Acids and Bases: Qualitative Aspects Objectives Separate the titration curve for a weak acid into regions in which a single equilibrium dominates Estimate the pH of mixtures of a weak acid and strong base or weak base and strong acid To sketch qualitative titration curves for acids of different strengths and concentrations

27. 16-4 Titrations of Weak Acids and Bases: Qualitative Aspects Typically divide the titration curve into four regions – Before any base added. – Part way to equivalence point. – Equivalence point. – Beyond equivalence point.

28. Fig. 16-7, p. 705

30. Fig. 16-8, p. 707 Titration of a Weak Acid with a Strong Base

31. Fig. 16-9, p. 708 Titration of a Weak base with a Strong Acid

32. 16-4 Titrations of Weak Acids and Bases: Qualitative Aspects The amount of base needed to neutralize an acid does not depend on whether you have a weak (acetic) or a strong (HCl) acid. It only depends on the amount of acid present The titration curves for all acids are indistinguishable beyond the equivalence point since excess hydroxide ion determines the pH of the solution.

33. 16-4 Titrations of Weak Acids and Bases: Qualitative Aspects The pH halfway to the equivalence point in the titration of a weak acid is equal to the pK a. The pH at the equivalence point is not 7.0 unless both the acid and the base are strong.

34. 16-5 Titrations of Weak Acids and Bases: Quantitative Aspects Objectives To determine the pH during the course of the titration of a weak acid with strong base To determine the pH during the course of the titration of a weak base with strong acid

35. p. 709

36. 16-5 Titrations of Weak Acids and Bases: Quantitative Aspects Calculate the pH in the titration of 25.0 mL of 0.500 M formic acid (Ka = 1.8 x10 -4 ) with 0.500 M NaOH after adding 0.0, 12.0, 25.0, and 40.0 mL of NaOH have been added and sketch the titration curve. pH = 2.02, 3.74, 8.57, 13.08

37. 16-6 Indicators Objectives To describe the acid-base chemistry of indicators To determine the pH during the course of the titration of a weak base with strong acid or a weak acid with a strong base

38. 16-6 Indicators Equivalence Point = the point in a titration when an acid and base are present in equal amounts End Point = point in the titration at which a color change occurs Indicator – substance that changes color to signal the end of a titration – an indicator should be chosen such that the end point and equivalence point are the same!

39. 16-6 Indicators Indicators should have the following attributes: – Change color abruptly as a function of pH – Color change should be observed readily by eye – Should not perturb the system – i.e. should not consume a significant amount of titrant or reactant

40. 16-6 Indicators Indicators are weak acid-base conjugate pairs in which the acid is a different color from the conjugate base For: HIn + H 2 O  H 3 O + + In - Then K in = [H 3 O + ][In - ] [HIn] – If pH < pK in, most of the indicator is in the acid form – If pH > pK in most of the indicator is in the base form

41. p. 715

42. 16-6 Indicators An indicator should be chosen to change color at or just beyond the equivalent point

43. pH Indicators and their ranges

44. Some Acid-Base Indicators Indicator pH Range in which Color Change Occurs Color Change as pH Increases Crystal violet Thymol blue Orange IV Methyl orange Bromcresol green Methyl red Chlorophenol red Bromthymol blue Phenol red Neutral red Thymol blue Phenolphthalein Thymolphthalein Alizarin yellow Indigo carmine 0.0 - 1.6 1.2 - 2.8 1.4 - 2.8 3.2 - 4.4 3.8 - 5.4 4.8 - 6.2 5.2 - 6.8 6.0 - 7.6 6.6 - 8.0 6.8 - 8.0 8.0 - 9.6 8.2 - 10.0 9.4 - 10.6 10.1 - 12.0 11.4 - 13.0 yellow to blue red to yellow red to yellow red to yellow yellow to blue red to yellow yellow to red yellow to blue yellow to red red to amber yellow to blue colourless to pink colourless to blue yellow to blue blue to yellow

45. Selection of Indicators

46. 16-7 Polyprotic Acids Objectives To define and write chemical equations for the dissociation of polyprotic acids To write the expressions for the equilibrium constants of polyprotic acids To calculate the pH of solutions that contain polyprotic acids To estimate pH of solutions of amphoteric species

47. 16-7 Polyprotic Acids Polyprotic acids provide more than one proton when they ionize – Ex: H 2 CO 3 +H 2 O  H 3 O + + HCO 3 - HCO 3 - + H 2 O  H 3 O + + CO 3 2- Ionization constants for the 1 st, 2 nd or 3 rd proton dissociations are denoted by K a1, K a2, or K a3 respectively.

48. Table 16-6, p. 718

49. 16-7 Polyprotic Acids Note that successive ionization constants get smaller often by a factor of more than 1000. To estimate the pH of a polyprotic acid, treat as if you had a mixture of weak acids as we did in Ch 15 – Use the largest value of K a to determine the pH as long as there is a difference of a factor of 100 between K a1 and K a2.

50. 16-7 Polyprotic Acids Calculate the pH of 0.33 M carbonic acid. A: pH = 3.42

51. 16-7 Polyprotic Acids Amphoteric Species have the properties of both acids and bases – Ex: HCO 3 - + H 2 O  H 3 O + + CO 3 2- HCO 3 - + H 3 O  H 2 CO 3 + H 2 O To estimate the pH of a solution of an amphoteric substance, compare K a and K b. If K a > K b, the solution is acidic. If K b > K a, the solution is basic.

53. 16-7 Polyprotic Acids Calculate the equilibrium constants for a solution that is 0.10 M sodium hydrogen ascorbate (NaHC 6 H 6 O 6 / ascorbic acid: K a1 = 8.0 x 10 -5, K a2 = 1.6 x 10 -12 ), and determine whether the solution is acidic or basic. A: K a = 1.6 x 10 -12 K b = 1.2 x 10 -10 slightly basic

54. p. 734

56. 16-8 Factors That Influence Solubility Objectives To determine how pH influences the solubility of precipitates To determine the effect of complex formation on the solubility of a precipitate

57. 16-8 Factors That Influence Solubility The solubility of a salt is affected by the acid- base properties of the anions and cations from which it is composed. As long as the anion or cation reacts with hydrogen or hydroxide ion, the solubility is influenced by pH Simultaneous Equilibria - dealing with both solubility equilibria and acid-base equilibria

58. 16-8 Factors That Influence Solubility Salt of Anions of Weak Acids – The solubility of a salt of a weak acid increases with the addition of acid BaF 2 (s)  Ba 2+ (aq) + 2 F - (aq) 2 [F - (aq) + H 3 O +  HF(aq) + H 2 O(aq)] BaF 2 (s) + 2H 3 O +  Ba 2+ (aq) +2HF(aq) + 2H 2 O(aq)

59. 16-8 Factors That Influence Solubility Salts of Transition-Metal Cations – Transition metal cations can react with substances that donate electrons – Ligands are the species that donate electrons. They are Lewis bases that have a lone electron pair that can form a covalent bond with an empty orbital belonging to the metallic cation – The resulting species is called a complex

60. NH 3, CN -, and H 2 O are Common Ligands

61. p. 723

62. Coordination Number  Coordination number refers to the number of ligands attached to the cation  2, 4, and 6 are the most common coordination numbers Coordination number Example(s) 2Ag(NH 3 ) 2 + 4CoCl 4 2- Cu(NH 3 ) 4 2+ 6Co(H 2 O) 6 2+ Ni(NH 3 ) 6 2+

63. 16-8 Factors That Influence Solubility Salts of Transition-Metal Cations – The solubility of a precipitate increases in the presence of a complexing agent Ag + (aq) + Cl - (aq)  AgCl(s) AgCl(s) + 2NH 3 (aq)  + 2Ag(NH 3 ) 2 + (aq) + Cl - (aq)

64. p. 722

65. Complex Ions and Solubility AgCl(s)  Ag + + Cl - K sp = 1.6 x 10 -10 Ag + + NH 3  Ag(NH 3 ) + K 1 = 2.1 x 10 3 Ag(NH 3 ) + NH 3  Ag(NH 3 ) 2 + K 2 = 8.2 x 10 3 AgCl + 2NH 3  Ag(NH 3 ) 2 + + Cl - K = K sp  K 1  K 2