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- 1. Acid Base Equilibrium<br />
- 2. Definition of an Acid<br />An Arrhenius acid is a substance that dissociates in aqueous medium to form hydronium ions, H3O+ .<br />Example:<br />
- 3. Definition of an Acid<br />A Brønsted-Lowery acid is a substance that is a proton donor.<br />Example:<br />
- 4. Definition of an Acid<br />A Lewis acid is a substance that is an electron pair acceptor.<br />Example:<br />
- 5. Definition of a Base<br />An Arrhenius base is a substance that ionizes in aqueous medium to produce hydroxide ions.<br />Example:<br />
- 6. Definition of a Base<br />A Brønsted-Lowery base is a substance that is a proton acceptor.<br />Example:<br />
- 7. Definition of a Base<br />A Lewis base is an electron pair donor.<br />Example:<br />
- 8. Conjugate Acid-Base<br />
- 9. Relative Strengths of Acids and Bases<br />Proton transfer reactions proceed from the stronger acid-base pair to the weaker acid-base pair.<br />For the relative strengths of acids and bases refer to figure 16.4 on page 672 of Brown, LeMay, Bursten and Murphy.<br />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.<br />
- 10. A Weak Acid<br />
- 11. A Weak Base<br />
- 12. The Values of K<br />A large vaalue of K means that the reaction favors the formation of product (s).<br />A small value of K means that the reaction favors the reactants.<br />If Kais greater than 1, then the acid is strong and dissociates 100%.<br />If Kb is greater than 1, then the base is strong and dissociates 100%.<br />
- 13. The Values of K<br />If Ka is 10-16 – 1, then the acid is weak.<br />If Kb is 10-16 – 1, then the base is weak.<br />If Ka is less than 10-16 , then the acid is very weak.<br />If Kb is less than 10-16 , then the base is very weak.<br />Refer to Appendix D, Aqueous Equilibrium Constants, page 1115-1116 for the values of K for acids and bases.<br />
- 14. Cationic Acid<br />(1) Auto ionization of water: <br />(2) Ammonia acting as a Brønsted-Lowry base: <br />
- 15. Cationic Acid<br />equation (3)<br />Add equation (1) and the reverse of equation (2) <br />to obtain equation (3)<br />equation (1)<br />equation (2)<br />
- 16. Cationic Acid<br />
- 17. Cationic Acid<br />
- 18. Cationic Acid<br />
- 19. Cationic Base<br />
- 20. Anionic Base<br />(1) Auto ionization of water: <br />(2) HCN acting as a Brønsted-Lowry acid: <br />
- 21. Anionic Base<br />equation (3)<br />Add equation (1) and the reverse of equation (2) <br />to obtain equation (3)<br />
- 22. Anionic Base<br />
- 23. Anionic Base<br />
- 24. Anionic Base<br />
- 25. Application<br /> Calculate Kh for CH3CH(OH)COO- given that the Kafor lactic acid, CH3CH(OH)COOH, is 1.4 x 10-4.<br />
- 26.
- 27.
- 28.
- 29.
- 30. pH and pOH<br />
- 31. pH<br />Solutions with pH less than 7.00 at 25oC are basic.<br />Solutions with pH greater than 7.00 at 25oC are acidic.<br />Solutions with pH equal to 7.00 at 25oC are neutral.<br />
- 32. Application<br /> 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.<br />
- 33. NaOH, a strong base, is 100% ionized in aqueous medium t<br />o form OH-(aq). <br />
- 34.
- 35.
- 36. Application<br />If [H3O+] in vinegar is 1.6 x 10-3, calculate its pH.<br />
- 37.
- 38. Application<br /> The pH of seawater is 8.30. Calculate the [H3O+] and [-OH] of seawater.<br />
- 39.
- 40. Determining pH<br /> Indicator: A substance that changes colors in some known pH range.<br />
- 41. Conjugate phenolphalein in <br />basic medium<br />pink<br />Phenolphalein in acidic medium<br />colorless<br />
- 42. Calculating Ka from pH and the Initial Concentration of an Acid<br />
- 43. Calculating Ka from pH and the Initial Concentration of an Acid<br />
- 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.<br />
- 45.
- 46. Calculating Equilibrium Concentrations and pH from Ka and the Initial Concentration of an Acid<br />
- 47. Calculating Equilibrium Concentrations and pH from Ka and the Initial Concentration of an Acid<br />
- 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.<br />ao<br />Ka = 6.3 x 10-5<br />
- 49. = 1.1 x 10-3 M<br />
- 50.
- 51. Calculate the hydronium ion concentration and pH of a 0.010 M formic acid solution. (Ka = 1.8 x 10-4)<br />
- 52.
- 53.
- 54. If the Acid is less than 15% ionized, then approximations can be made<br />In the previous problem,<br />
- 55. Therefore, we could have made the following <br />Approximation:<br />0.010 is greater than x; consequently <br />0.010 – x = 0.010<br />
- 56.
- 57. What if we change the initial concentration <br />of the formic acid?<br />Calculate the hydronium ion concentration and pH of a 0.010 M formic acid solution. (Ka = 1.8 x 10-4)<br />
- 58. Let’s make an approximation: <br />Assume 0.0010 is greater than x<br />Therefore, the assumption cannot be made, and<br />The quadratic equation must be used.<br />
- 59.
- 60.
- 61.
- 62. Calculating the pH of an Aqueous Solution of a Weak Base<br />
- 63.
- 64.
- 65. Calculate the pH of a 0.010 M aqueous <br />solution of pyridine, C5H5N. Pyridine is <br />A weak base with Kb = 1.7 x 10-9.<br />
- 66. Calculating the pH of an Aqueous Solution of the Salt of a Weak Acid<br />
- 67. Calculating the pH of an Aqueous Solution of the Salt of a Weak Acid<br />
- 68.
- 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)<br />
- 70.
- 71. Calculating the pH of an Aqueous Solution of the Salt of a Weak Base<br />
- 72.
- 73.
- 74. Calculate the pH of a 0.50 M solution of ammonium chloride, NH4Cl. The Ka for ammonia is 1.8 x 10-5)<br />
- 75.
- 76.
- 77. Common Ion Effect<br />Addition of a common ion to the equilibrium.<br />Example: adding NaA to an acid solution of HA<br />ao<br />x<br />x<br />y<br />y<br />y<br />
- 78.
- 79.
- 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 <br />
- 81. Species that form Multiple Species in Solution<br />
- 82. Species that form Multiple Species in Solution<br />
- 83. Species that form Multiple Species in Solution<br />
- 84. Species that form Multiple Species in Solution<br />
- 85. Species that form Multiple Species in Solution<br />
- 86. Species that form Multiple Species in Solution<br />
- 87. Species that form Multiple Species in Solution<br />
- 88. Calculate the equilibrium concentrations of all species in solution for 0.10 M solution of Na2CO3.<br />H2CO3+ H2O HCO3- + H3O+ Ka1 = 4.3 x 10-7<br />HCO3- + H2O CO32- + H3O+ Ka2 = 5.6 x 10-11<br />
- 89.
- 90.
- 91.
- 92.
- 93.
- 94. Calculate the pH of the solution<br />

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