The PowerPoint Presentation is a set of slides associated with a discussion on Acid-Base Equilibrium.

1. Acid Base Equilibrium

2. Definition of an Acid An Arrhenius acid is a substance that dissociates in aqueous medium to form hydronium ions, H3O+ . Example:

3. Definition of an Acid A Brønsted-Lowery acid is a substance that is a proton donor. Example:

4. Definition of an Acid A Lewis acid is a substance that is an electron pair acceptor. Example:

5. Definition of a Base An Arrhenius base is a substance that ionizes in aqueous medium to produce hydroxide ions. Example:

6. Definition of a Base A Brønsted-Lowery base is a substance that is a proton acceptor. Example:

7. Definition of a Base A Lewis base is an electron pair donor. Example:

8. Conjugate Acid-Base

9. Relative Strengths of Acids and Bases Proton transfer reactions proceed from the stronger acid-base pair to the weaker acid-base pair. For the relative strengths of acids and bases refer to figure 16.4 on page 672 of Brown, LeMay, Bursten and Murphy. Write a balanced equation for the reaction that occurs between ammonium chloride and sodium carbonate. Decide if the equilibrium lies predominantly toward the reactants or the products.

10. A Weak Acid

11. A Weak Base

12. The Values of K A large vaalue of K means that the reaction favors the formation of product (s). A small value of K means that the reaction favors the reactants. If Kais greater than 1, then the acid is strong and dissociates 100%. If Kb is greater than 1, then the base is strong and dissociates 100%.

13. The Values of K If Ka is 10-16 – 1, then the acid is weak. If Kb is 10-16 – 1, then the base is weak. If Ka is less than 10-16 , then the acid is very weak. If Kb is less than 10-16 , then the base is very weak. Refer to Appendix D, Aqueous Equilibrium Constants, page 1115-1116 for the values of K for acids and bases.

14. Cationic Acid (1) Auto ionization of water: (2) Ammonia acting as a Brønsted-Lowry base:

15. Cationic Acid equation (3) Add equation (1) and the reverse of equation (2) to obtain equation (3) equation (1) equation (2)

16. Cationic Acid

17. Cationic Acid

18. Cationic Acid

19. Cationic Base

20. Anionic Base (1) Auto ionization of water: (2) HCN acting as a Brønsted-Lowry acid:

21. Anionic Base equation (3) Add equation (1) and the reverse of equation (2) to obtain equation (3)

22. Anionic Base

23. Anionic Base

24. Anionic Base

25. Application Calculate Kh for CH3CH(OH)COO- given that the Kafor lactic acid, CH3CH(OH)COOH, is 1.4 x 10-4.

30. pH and pOH

31. pH Solutions with pH less than 7.00 at 25oC are basic. Solutions with pH greater than 7.00 at 25oC are acidic. Solutions with pH equal to 7.00 at 25oC are neutral.

32. Application Calculate the pH of a solution made by dissolving 0.7000 g of NaOH in sufficient water to produce a volume of 500.0 mL.

33. NaOH, a strong base, is 100% ionized in aqueous medium t o form OH-(aq).

36. Application If [H3O+] in vinegar is 1.6 x 10-3, calculate its pH.

38. Application The pH of seawater is 8.30. Calculate the [H3O+] and [-OH] of seawater.

40. Determining pH Indicator: A substance that changes colors in some known pH range.

41. Conjugate phenolphalein in basic medium pink Phenolphalein in acidic medium colorless

42. Calculating Ka from pH and the Initial Concentration of an Acid

43. Calculating Ka from pH and the Initial Concentration of an Acid

44. A 0.10 M solution of aqueous lactic acid, CH3CH(OH)COOH, has a pH equal to 2.43 at 25oC. Calculate the Ka for lactic acid at 25oC.

46. Calculating Equilibrium Concentrations and pH from Ka and the Initial Concentration of an Acid

47. Calculating Equilibrium Concentrations and pH from Ka and the Initial Concentration of an Acid

48. Calculate the equilibrium concentrations of the hydronium ion and the benzoate ion of a 0.020 M solution of benzoic acid. Calculate the pH a 0.020 M solution of benzoic. ao Ka = 6.3 x 10-5

49. = 1.1 x 10-3 M

51. Calculate the hydronium ion concentration and pH of a 0.010 M formic acid solution. (Ka = 1.8 x 10-4)

54. If the Acid is less than 15% ionized, then approximations can be made In the previous problem,

55. Therefore, we could have made the following Approximation: 0.010 is greater than x; consequently 0.010 – x = 0.010

57. What if we change the initial concentration of the formic acid? Calculate the hydronium ion concentration and pH of a 0.010 M formic acid solution. (Ka = 1.8 x 10-4)

58. Let’s make an approximation: Assume 0.0010 is greater than x Therefore, the assumption cannot be made, and The quadratic equation must be used.

62. Calculating the pH of an Aqueous Solution of a Weak Base

65. Calculate the pH of a 0.010 M aqueous solution of pyridine, C5H5N. Pyridine is A weak base with Kb = 1.7 x 10-9.

66. Calculating the pH of an Aqueous Solution of the Salt of a Weak Acid

67. Calculating the pH of an Aqueous Solution of the Salt of a Weak Acid

69. Calculate the pH of a 0.015 M solution of sodium hypochlorite, NaClO. The Ka for hypochlorous acid is 3.5 x 10-8)

71. Calculating the pH of an Aqueous Solution of the Salt of a Weak Base

74. Calculate the pH of a 0.50 M solution of ammonium chloride, NH4Cl. The Ka for ammonia is 1.8 x 10-5)

77. Common Ion Effect Addition of a common ion to the equilibrium. Example: adding NaA to an acid solution of HA ao x x y y y

80. Calculating the pH and the Equilibrium Concentrations of Species in Solution form from the Reactions of Anionic Bases that Produce Multi-Species in Solution

81. Species that form Multiple Species in Solution

82. Species that form Multiple Species in Solution

83. Species that form Multiple Species in Solution

84. Species that form Multiple Species in Solution

85. Species that form Multiple Species in Solution

86. Species that form Multiple Species in Solution

87. Species that form Multiple Species in Solution

88. Calculate the equilibrium concentrations of all species in solution for 0.10 M solution of Na2CO3. H2CO3+ H2O HCO3- + H3O+ Ka1 = 4.3 x 10-7 HCO3- + H2O CO32- + H3O+ Ka2 = 5.6 x 10-11

94. Calculate the pH of the solution