1. Chapter 10 Acids, Bases, and Salts

2. Chapter 10 Table of Contents Copyright © Cengage Learning. All rights reserved 2 10.1Arrhenius Acid-Base Theory 10.2Brønsted-Lowry Acid-Base Theory 10.3Mono-, Di-, and Triprotic Acids 10.4Strengths of Acids and Bases 10.5Ionization Constants for Acids and Bases 10.6Salts 10.7Acid-Base Neutralization Reactions 10.8Self-Ionization of Water 10.9The pH Concept 10.10The pKa Method for Expressing Acid Strength 10.11The pH of Aqueous Salt Solutions 10.12Buffers 10.13The Henderson-Hasselbalch Equation 10.14Electrolytes 10.15Equivalents and Milliequivalents of Electrolytes 10.16Acid-Base Titrations

3. Arrhenius Acid-Base Theory Section 10.1 Copyright © Cengage Learning. All rights reserved 3 Arrhenius acid: hydrogen-containing compound that produces H + ions in solution.  Example: HNO 3 → H + + NO 3 – Arrhenius base: hydroxide-containing compound that produces OH – ions in solution.  Example: NaOH → Na + + OH –

4. Arrhenius Acid-Base Theory Section 10.1 Copyright © Cengage Learning. All rights reserved 4 Ionization The process in which individual positive and negative ions are produced from a molecular compound that is dissolved in solution. –Arrhenius acids

5. Arrhenius Acid-Base Theory Section 10.1 Copyright © Cengage Learning. All rights reserved 5 Dissociation The process in which individual positive and negative ions are released from an ionic compound that is dissolved in solution. –Arrhenius Bases

6. Arrhenius Acid-Base Theory Section 10.1 Copyright © Cengage Learning. All rights reserved 6 Difference Between Ionization and Dissociation

7. Section 10.2 Brønsted-Lowry Acid-Base Theory Copyright © Cengage Learning. All rights reserved 7 Brønsted-Lowry acid: substance that can donate a proton (H + ion) to some other substance; proton donor. Brønsted-Lowry base: substance that can accept a proton (H + ion) from some other substance; proton acceptor. HCl + H 2 O  Cl  + H 3 O + acid base

8. Section 10.2 Brønsted-Lowry Acid-Base Theory Copyright © Cengage Learning. All rights reserved 8 Brønsted-Lowry Reaction To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERECLICK HERE

9. Section 10.2 Brønsted-Lowry Acid-Base Theory Copyright © Cengage Learning. All rights reserved 9 Acid in Water HA(aq) + H 2 O(l) H 3 O + (aq) + A - (aq) acid base conjugate conjugate acid base

10. Section 10.2 Brønsted-Lowry Acid-Base Theory Copyright © Cengage Learning. All rights reserved 10 Acid Ionization Equilibrium To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERECLICK HERE

11. Section 10.2 Brønsted-Lowry Acid-Base Theory Copyright © Cengage Learning. All rights reserved 11 Amphiprotic Substance A substance that can either lose or accept a proton and thus can function as either a Brønsted-Lowry acid or a Brønsted-Lowry base.  Example: H 2 O, H 3 O + H 2 O, OH –

12. Section 10.3 Mono-, Di-, and Triprotic Acids Copyright © Cengage Learning. All rights reserved 12 Monoprotic Acid An acid that supplies one proton (H + ion) per molecule during an acid-base reaction. HA + H 2 O A  + H 3 O +

13. Section 10.3 Mono-, Di-, and Triprotic Acids Copyright © Cengage Learning. All rights reserved 13 Diprotic Acid An acid that supplies two protons (H + ions) per molecule during an acid-base reaction. H 2 A + H 2 O HA  + H 3 O + HA  + H 2 O A 2  + H 3 O +

14. Section 10.3 Mono-, Di-, and Triprotic Acids Copyright © Cengage Learning. All rights reserved 14 Triprotic Acid An acid that supplies three protons (H + ions) per molecule during an acid-base reaction. H 3 A + H 2 O H 2 A  + H 3 O + H 2 A  + H 2 O HA 2  + H 3 O + HA 2  + H 2 O A 3  + H 3 O +

15. Section 10.3 Mono-, Di-, and Triprotic Acids Copyright © Cengage Learning. All rights reserved 15 Polyprotic Acid An acid that supplies two or more protons (H + ions) during an acid-base reaction. Includes both diprotic and triprotic acids.

16. Section 10.4 Strengths of Acids and Bases Copyright © Cengage Learning. All rights reserved 16 Strong Acid Transfers ~100% of its protons to water in an aqueous solution. Ionization equilibrium lies far to the right. Yields a weak conjugate base.

17. Section 10.4 Strengths of Acids and Bases Copyright © Cengage Learning. All rights reserved 17 Commonly Encountered Strong Acids

18. Section 10.4 Strengths of Acids and Bases Copyright © Cengage Learning. All rights reserved 18 Weak Acid Transfers only a small % of its protons to water in an aqueous solution. Ionization equilibrium lies far to the left. Weaker the acid, stronger its conjugate base.

19. Section 10.4 Strengths of Acids and Bases Copyright © Cengage Learning. All rights reserved 19 Differences Between Strong and Weak Acids in Terms of Species Present

20. Section 10.4 Strengths of Acids and Bases Copyright © Cengage Learning. All rights reserved 20 Bases Strong bases: hydroxides of Groups IA and IIA.

21. Section 10.5 Ionization Constants for Acids and Bases Copyright © Cengage Learning. All rights reserved 21 Acid Ionization Constant The equilibrium constant for the reaction of a weak acid with water. HA(aq) + H 2 O(l) H 3 O + (aq) + A - (aq)

22. Section 10.5 Ionization Constants for Acids and Bases Copyright © Cengage Learning. All rights reserved 22 Acid Strength, % Ionization, and K a Magnitude Acid strength increases as % ionization increases. Acid strength increases as the magnitude of K a increases. % Ionization increases as the magnitude of K a increases.

23. Section 10.5 Ionization Constants for Acids and Bases Copyright © Cengage Learning. All rights reserved 23 Base Ionization Constant The equilibrium constant for the reaction of a weak base with water. B(aq) + H 2 O(l) BH + (aq) + OH – (aq)

24. Section 10.6 Salts Copyright © Cengage Learning. All rights reserved 24 Ionic compounds containing a metal or polyatomic ion as the positive ion and a nonmetal or polyatomic ion (except hydroxide) as the negative ion. All common soluble salts are completely dissociated into ions in solution.

25. Section 10.7 Acid-Base Neutralization Reactions Copyright © Cengage Learning. All rights reserved 25 Neutralization Reaction The chemical reaction between an acid and a hydroxide base in which a salt and water are the products. HCl + NaOH → NaCl + H 2 O H 2 SO 4 + 2 KOH → K 2 SO 4 + 2 H 2 O

26. Section 10.7 Acid-Base Neutralization Reactions Copyright © Cengage Learning. All rights reserved 26 Formation of Water

27. Section 10.8 Self-Ionization of Water Copyright © Cengage Learning. All rights reserved 27 Self-Ionization Water molecules in pure water interact with one another to form ions. H 2 O + H 2 O H 3 O + + OH – Net effect is the formation of equal amounts of hydronium and hydroxide ions.

28. Section 10.8 Self-Ionization of Water Copyright © Cengage Learning. All rights reserved 28 Self-Ionization of Water

29. Section 10.8 Self-Ionization of Water Copyright © Cengage Learning. All rights reserved 29 Ion Product Constant for Water At 24°C: K w = [H 3 O + ][OH – ] = 1.00 × 10 –14 No matter what the solution contains, the product of [H 3 O + ] and [OH – ] must always equal 1.00 × 10 –14.

30. Section 10.8 Self-Ionization of Water Copyright © Cengage Learning. All rights reserved 30 Relationship Between [H 3 O + ] and [OH – ]

31. Section 10.8 Self-Ionization of Water Copyright © Cengage Learning. All rights reserved 31 Three Possible Situations [H 3 O + ] = [OH – ]; neutral solution [H 3 O + ] > [OH – ]; acidic solution [H 3 O + ] < [OH – ]; basic solution

32. Section 10.8 Self-Ionization of Water Copyright © Cengage Learning. All rights reserved 32 Exercise Calculate [H 3 O + ] or [OH – ] as required for each of the following solutions at 24°C, and state whether the solution is neutral, acidic, or basic. a) 1.0 × 10 –4 M OH – b) 2.0 M H 3 O +

33. Section 10.8 Self-Ionization of Water Copyright © Cengage Learning. All rights reserved 33 Exercise Calculate [H 3 O + ] or [OH – ] as required for each of the following solutions at 24°C, and state whether the solution is neutral, acidic, or basic. a) 1.0 × 10 –4 M OH – 1.0 × 10 –10 M H 3 O + ; basic b) 2.0 M H 3 O + 5.0 × 10 –15 M OH – ; acidic

34. Section 10.9 The pH Concept Copyright © Cengage Learning. All rights reserved 34 pH = –log[H 3 O + ] A compact way to represent solution acidity. pH decreases as [H + ] increases. pH range between 0 to 14 in aqueous solutions at 24°C.

35. Section 10.9 The pH Concept Copyright © Cengage Learning. All rights reserved 35 Exercise Calculate the pH for each of the following solutions. a) 1.0 × 10 –4 M H 3 O + b)0.040 M OH –

36. Section 10.9 The pH Concept Copyright © Cengage Learning. All rights reserved 36 Exercise Calculate the pH for each of the following solutions. a) 1.0 × 10 –4 M H 3 O + pH = 4.00 b)0.040 M OH – pH = 12.60

37. Section 10.9 The pH Concept Copyright © Cengage Learning. All rights reserved 37 Exercise The pH of a solution is 5.85. What is the [H 3 O + ] for this solution?

38. Section 10.9 The pH Concept Copyright © Cengage Learning. All rights reserved 38 Exercise The pH of a solution is 5.85. What is the [H 3 O + ] for this solution? [H 3 O + ] = 1.4 × 10 –6 M

39. Section 10.9 The pH Concept Copyright © Cengage Learning. All rights reserved 39 pH Range pH = 7; neutral pH > 7; basic –Higher the pH, more basic. pH < 7; acidic –Lower the pH, more acidic.

40. Section 10.9 The pH Concept Copyright © Cengage Learning. All rights reserved 40 Relationships Among pH Values, [H 3 O + ], and [OH – ]

41. Section 10.10 The pK a Method for Expressing Acid Strength Copyright © Cengage Learning. All rights reserved 41 pK a = –log K a pK a is calculated from K a in exactly the same way that pH is calculated from [H 3 O + ].

42. Section 10.10 The pK a Method for Expressing Acid Strength Copyright © Cengage Learning. All rights reserved 42 Exercise Calculate the pK a for HF given that the K a for this acid is 6.8 × 10 –4.

43. Section 10.10 The pK a Method for Expressing Acid Strength Copyright © Cengage Learning. All rights reserved 43 Exercise Calculate the pK a for HF given that the K a for this acid is 6.8 × 10 –4. pK a = 3.17a

44. Section 10.11 The pH of Aqueous Salt Solutions Copyright © Cengage Learning. All rights reserved 44 Salts Ionic compounds. When dissolved in water, break up into its ions (which can behave as acids or bases). Hydrolysis – the reaction of a salt with water to produce hydronium ion or hydroxide ion or both.

45. Section 10.11 The pH of Aqueous Salt Solutions Copyright © Cengage Learning. All rights reserved 45 Types of Salt Hydrolysis 1.The salt of a strong acid and a strong base does not hydrolyze, so the solution is neutral.  KCl, NaNO 3

46. Section 10.11 The pH of Aqueous Salt Solutions Copyright © Cengage Learning. All rights reserved 46 Types of Salt Hydrolysis 2.The salt of a strong acid and a weak base hydrolyzes to produce an acidic solution.  NH 4 Cl NH 4 + + H 2 O → NH 3 + H 3 O +

47. Section 10.11 The pH of Aqueous Salt Solutions Copyright © Cengage Learning. All rights reserved 47 Types of Salt Hydrolysis 3.The salt of a weak acid and a strong base hydrolyzes to produce a basic solution.  NaF, KC 2 H 3 O 2 F – + H 2 O → HF + OH – C 2 H 3 O 2 – + H 2 O → HC 2 H 3 O 2 + OH –

48. Section 10.11 The pH of Aqueous Salt Solutions Copyright © Cengage Learning. All rights reserved 48 Types of Salt Hydrolysis 4.The salt of a weak acid and a weak base hydrolyzes to produce a slightly acidic, neutral, or slightly basic solution, depending on the relative weaknesses of the acid and base.

49. Section 10.11 The pH of Aqueous Salt Solutions Copyright © Cengage Learning. All rights reserved 49 Neutralization “Parentage” of Salts

50. Section 10.12 Buffers Copyright © Cengage Learning. All rights reserved 50 Key Points about Buffers Buffer – an aqueous solution containing substances that prevent major changes in solution pH when small amounts of acid or base are added to it. They are weak acids or bases containing a common ion. Typically, a buffer system is composed of a weak acid and its conjugate base.

51. Section 10.12 Buffers Copyright © Cengage Learning. All rights reserved 51 Buffers Contain Two Active Chemical Species 1.A substance to react with and remove added base. 2.A substance to react with and remove added acid.

52. Section 10.12 Buffers Copyright © Cengage Learning. All rights reserved 52 Adding an Acid to a Buffer To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERECLICK HERE

53. Section 10.12 Buffers Copyright © Cengage Learning. All rights reserved 53 Buffers To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERECLICK HERE

54. Section 10.12 Buffers Copyright © Cengage Learning. All rights reserved 54 Addition of Base [OH – ion] to the Buffer HA + H 2 O H 3 O + + A – The added OH – ion reacts with H 3 O + ion, producing water (neutralization). The neutralization reaction produces the stress of not enough H 3 O + ion because H 3 O + ion was consumed in the neutralization. The equilibrium shifts to the right to produce more H 3 O + ion, which maintains the pH close to its original level.

55. Section 10.12 Buffers Copyright © Cengage Learning. All rights reserved 55 Addition of Acid [H 3 O + ion] to the Buffer HA + H 2 O H 3 O + + A – The added H 3 O + ion increases the overall amount of H 3 O + ion present. The stress on the system is too much H 3 O + ion. The equilibrium shifts to the left consuming most of the excess H 3 O + ion and resulting in a pH close to the original level.

56. Section 10.13 The Henderson-Hasselbalch Equation Copyright © Cengage Learning. All rights reserved 56 Henderson-Hasselbalch Equation

57. Section 10.13 The Henderson-Hasselbalch Equation Copyright © Cengage Learning. All rights reserved 57 Exercise What is the pH of a buffer solution that is 0.45 M acetic acid (HC 2 H 3 O 2 ) and 0.85 M sodium acetate (NaC 2 H 3 O 2 )? The K a for acetic acid is 1.8 × 10 –5. pH = 5.02

58. Section 10.14 Electrolytes Copyright © Cengage Learning. All rights reserved 58 Acids, bases, and soluble salts all produce ions in solution, thus they all produce solutions that conduct electricity. Electrolyte – substance whose aqueous solution conducts electricity.

59. Section 10.14 Electrolytes Copyright © Cengage Learning. All rights reserved 59 Example: table sugar (sucrose), glucose Nonelectrolyte – does not conduct electricity

60. Section 10.14 Electrolytes Copyright © Cengage Learning. All rights reserved 60 Example: strong acids, bases, and soluble salts Strong Electrolyte – completely ionizes/dissociates

61. Section 10.14 Electrolytes Copyright © Cengage Learning. All rights reserved 61 Example: weak acids and bases Weak Electrolyte – incompletely ionizes/dissociates

62. Section 10.15 Equivalents and Milliequivalents of Electrolytes Copyright © Cengage Learning. All rights reserved 62 The molar amount of that ion needed to supply one mole of positive or negative charge. 1 mole K + = 1 equivalent 1 mole Mg 2+ = 2 equivalents 1 mole PO 4 3– = 3 equivalents Equivalent (Eq) of an Ion

63. Section 10.15 Equivalents and Milliequivalents of Electrolytes Copyright © Cengage Learning. All rights reserved 63 1 milliequivalent = 10 –3 equivalent Milliequivalent

64. Section 10.15 Equivalents and Milliequivalents of Electrolytes Copyright © Cengage Learning. All rights reserved 64 Concentrations of Major Electrolytes in Blood Plasma

65. Section 10.15 Equivalents and Milliequivalents of Electrolytes Copyright © Cengage Learning. All rights reserved 65 Exercise The concentration of Ca 2+ ion present in a sample is 5.3 mEq/L. How many milligrams of Ca 2+ ion are present in 180.0 mL of the sample?

66. Section 10.15 Equivalents and Milliequivalents of Electrolytes Copyright © Cengage Learning. All rights reserved 66 Exercise The concentration of Ca 2+ ion present in a sample is 5.3 mEq/L. How many milligrams of Ca 2+ ion are present in 180.0 mL of the sample? 19 mg Ca 2+ ion

67. Section 10.16 Acid-Base Titrations Copyright © Cengage Learning. All rights reserved 67 A neutralization reaction in which a measured volume of an acid or a base of known concentration is completely reacted with a measured volume of a base or an acid of unknown concentration. For a strong acid and base reaction: H + (aq) + OH – (aq)  H 2 O(l)

68. Section 10.16 Acid-Base Titrations Copyright © Cengage Learning. All rights reserved 68 Titration Setup

69. Section 10.16 Acid-Base Titrations Copyright © Cengage Learning. All rights reserved 69 A compound that exhibits different colors depending on the pH of its solution. An indicator is selected that changes color at a pH that corresponds as nearly as possible to the pH of the solution when the titration is complete. Acid-Base Indicator

70. Section 10.16 Acid-Base Titrations Copyright © Cengage Learning. All rights reserved 70 Indicator – yellow in acidic solution; red in basic solution

71. Section 10.16 Acid-Base Titrations Copyright © Cengage Learning. All rights reserved 71 Concept Check For the titration of sulfuric acid (H 2 SO 4 ) with sodium hydroxide (NaOH), how many moles of sodium hydroxide would be required to react with 1.00 L of 0.500 M sulfuric acid to reach the endpoint?

72. Section 10.16 Acid-Base Titrations Copyright © Cengage Learning. All rights reserved 72 Concept Check For the titration of sulfuric acid (H 2 SO 4 ) with sodium hydroxide (NaOH), how many moles of sodium hydroxide would be required to react with 1.00 L of 0.500 M sulfuric acid to reach the endpoint? 1.00 mol NaOH