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Chapter 16

Chapter 16. Acids and Bases. Arrhenius Acid-Base Theory. Arrhenius, Svante August (1859-1927), Swedish chemist 1903 Nobel Prize in chemistry acid - hydrogen ion (proton) donor base - hydroxide ion donor. Bronsted-Lowrey Acid-Base Theory. acid - proton donor base - proton acceptor

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Chapter 16

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  1. Chapter 16 Acids and Bases

  2. Arrhenius Acid-Base Theory Arrhenius, Svante August • (1859-1927), Swedish chemist • 1903 Nobel Prize in chemistry acid - hydrogen ion (proton) donor base - hydroxide ion donor

  3. Bronsted-LowreyAcid-Base Theory acid - proton donor base - proton acceptor strong vs. weak acids and bases • strong - completely ionized • weak - partially ionized

  4. Brönsted-LowryAcid-Base Systems Acid-Base Conjugate Pairs conjugate base => what results when an acid looses its proton conjugate acid => what results when a base gains its proton

  5. Water’s Role as Acid or Base Water acting as a Base HA + H2O  H3O+ + A- acid base Water acting as an Acid B + H2O  BH+ + OH- base acid

  6. Conjugate Acid-Base Pairs

  7. Relative Strengths of Acid and Bases

  8. Strong vs. Weak Acids and Bases strong acid • completely ionized weak acid • partially ionized

  9. Mixtures of Strong and Weak Acids • the presence of the strong acid retards the dissociation of the weak acid

  10. Carboxylic Acids

  11. Amines

  12. Autoionization of Water H2O + H2O  H3O+ + OH- [H3O+][OH-] K = [H2O]2 Kw = K [H2O]2 = [H3O+][OH-] = 1.0  10-14

  13. Autoionization of Water Kw = [H3O+][OH-] = 1.0  10-14 for a neutral solution [H3O+] = [OH-] [H3O+][H3O+] = 1.0 x 10-14 [H3O+]2 = 1.0 x 10-14 [H3O+] = 1.0 x 10-7 -log([H3O+]) = -log(1.0  10-7) pH = -log([H3O+]) = 7.00

  14. Autoionization of Water a solution is considered acidic when [H3O+] > 1.0  10-7 M [OH-] < 1.0  10-7 M a solution is considered bacic when [H3O+] < 1.0  10-7 M [OH-] > 1.0  10-7 M

  15. pH Scale pH = - log10[H3O+] pH Scale @ -2 => @ +16

  16. pOH Scale pOH = - log10[OH-] pOH Scale @ -2 => @ +16

  17. Solutions of Strong Acids [H3O+] = Macid pH = - log10[H3O+] = - log10(Macid)

  18. pH of Aqueous Solutions

  19. Concentration Scales a solution is considered acidic when pH < 7 and pOH > 7 a solution is considered bacic when pH > 7 and pOH < 7 pKw = pH + pOH = 14.00

  20. EXAMPLE: What is the pH of a solution that has a [H3O+] = 0.435 M? pH = - log10[H3O+] = - log10(0.435) = 0.362

  21. EXAMPLE: What is the pH of a solution that has a [OH-] = 25M? pOH = - log10[OH-] = - log10(25) = - 1.40 = 14.00 - (- 1.40) pH = 14.00 - pOH = 15.40

  22. Measuring pH Arnold Beckman • inventor of the pH meter • father of electronic instrumentation

  23. pH Meter Photo by George Lisensky

  24. Acid Dissociation Constant HC2H3O2 + H2O  H3O+ + C2H3O2- [H3O+][C2H3O2-] K = [H2O][HC2H3O2] [H3O+][C2H3O2-] Ka = K  [H2O] = [HC2H3O2]

  25. Base Dissociation Constant NH3 + H2O  NH4+ + OH- [NH4+][OH-] K = [H2O][NH3] [NH4+][OH-] Kb = K  [H2O] = [NH3]

  26. Hydrated Metal Ions as Acids

  27. Ionization Constants for Acids

  28. Polyprotic Acids • acids where more than one hydrogen per molecule is released

  29. Polyprotic Acids

  30. EXAMPLE: Calculate the [H3O+] in an aqueous 0.140 M acetic acid solution. HC2H3O2 + H2O  H3O+ + C2H3O2- CHA = 0.140 M [H3O+]e[C2H3O2-]e Ka = = 1.75  10-5 M [HC2H3O2]e let [H3O+]e» [C2H3O2-]e [HC2H3O2]e= 0.140M - [H3O+]e

  31. EXAMPLE: Calculate the [H3O+] in an aqueous 0.140 M acetic acid solution. CHA = 0.140 M Ka = 1.75  10-5 M let [H3O+]e» [C2H3O2-]e [HC2H3O2]e= 0.140M - [H3O+]e using the quadratic equation - Ka + (Ka2 + 4KaCHA)1/2 [H3O+]e = 2 [H3O+]e = 1.56  10-3 M

  32. EXAMPLE: A saturated aqueous solution of caproic acid contains 11 g/L and has a pH = 2.94. What is its Ka? HC6H11O2 + H2O  H3O+ + C6H11O2- [H3O+]e[C6H11O2-]e Ka = = ? [HC6H11O2]e [H3O+]e = 10-pH = 10-2.94 = 1.1  10-3 M [C6H11O2-]e = [H3O+]e = 1.1  10-3 M [HC6H11O2]i = (11 g/L)(1 mol/116 g) = 9.5  10-2 M

  33. EXAMPLE: A saturated aqueous solution of caproic acid contains 11 g/L and has a pH = 2.94. What is its Ka? [C6H11O2-]e = [H3O+]e = 1.1  10-3 M [HC6H11O2]e = 9.4 x 10-2 M [H3O+]e[C6H11O2-]e Ka = [HC6H11O2]e (1.1  10-3 M)2 Ka = = 1.3  10-5 M 9.4  10-2 M

  34. Influence of Molecular Structure on Acid Strength Binary Hydrides • hydrogen & one other element • Bond Strengths • weaker the bond, the stronger the acid • Stability of Anion • higher the electronegativity, stronger the acid

  35. Influence of Molecular Structure on Acid Strength Oxyacids • hydrogen, oxygen, & one other element H-O-E • higher the electronegativity on E, stronger the acid as this weakens the bond between the O and H

  36. The Conjugate Partners of Strong Acids and Bases • the conjugate base of a strong acid has no net effect on the pH of a solution

  37. Weak Acids and Bases:Qualitative Aspects The Competition for Protons acid1 + base2 base1 + acid2

  38. Weak Acids and Bases:Qualitative Aspects The Leveling Effect of the Solvent in water both HClO4 and HCl are strong acids

  39. Weak Acids and Bases:Qualitative Aspects The Leveling Effect of the Solvent in glacial acetic acid HClO4 + CH3COOH(l) => CH3COOH2+ + ClO4- complete dissociation, thus strong acid HCl + CH3COOH(l) CH3COOH2+ + Cl- incomplete dissociation, equilibrium established, weak acid

  40. Solutions of Salts of Weak Acids A- + H2O  HA + OH- Kw [HA][OH-] Kh = K'b = = Ka [A-]

  41. Acid-Base Properties of Typical Ions

  42. Acid-Base Chemistryof Some Antacids

  43. Acid-Base in the Kitchen vinegar - acetic acid lemon juice (citrus juice) - citric acid baking soda - NaHCO3 milk - lactic acid baking powder - H2PO4- & HCO3-

  44. Household Cleaners

  45. Lewis Acids and Bases • acid => electron pair acceptor • base => electron pair donor the base donates a pair of electrons to the acid forming a coordinate covalent bond

  46. Lewis Acids and Bases Reactions H+ + NH3  NH4+ acid base Cu+2 + 4 NH3 [Cu(NH3)4+2] acid base

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