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Chapter 14. Acids and Bases

Chapter 14. Acids and Bases. Early attributes of acids and bases (1600's) Acids Bases Taste sour Taste bitter Turn litmus red Turn litmus blue React with metals Feel slippery React with carbonates React with fats. Arrhenius Acids and Bases.

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Chapter 14. Acids and Bases

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  1. Chapter 14.Acids and Bases Early attributes of acids and bases (1600's) AcidsBases Taste sour Taste bitter Turn litmus red Turn litmus blue React with metals Feel slippery React with carbonates React with fats

  2. Arrhenius Acids and Bases Definitions of Svante Arrhenius, 1884 First working theory about acids and bases Acids contain hydrogen and produce H1+ ions in water. Bases contain hydroxide ions (OH1-) and are soluble in water.

  3. Acids and Bases

  4. Arrhenius Acids and Bases Acids are molecular compounds; a covalent bond attaches the hydrogen ion to the ad-jacent atom. Ionization, the separation of the molecule into ions, occurs when the molecule dis-solves in water.

  5. Arrhenius Acids and Bases Bases are ionic compounds; the hydroxide ion exists in the crystal structure of the solid compound. Dissociation occurs when the ionic solid dissolves in water, releasing the ions to move about.

  6. Arrhenius Acids and Bases Common Acids: HCl(aq), H2SO4, H3PO4, HNO3 HC2H3O2 = CH3COOH = acetic acid Common Bases: NaOH, KOH

  7. Bronsted-Lowry Acids and Bases The Arrhenius definition has some problems: It's restricted to water. It doesn't explain why solutions of some molecular compounds (NH3) and salts (Na2CO3) are basic. It doesn't explain why some salt solutions are acidic (aqueous Al3+, Fe3+ solutions).

  8. Bronsted-Lowry Acids and Bases Definitions of Brønsted and Lowry, 1923 Most widely used theory of acids and bases Acids are proton donors. Bases are proton acceptors. Reactions: HCl(aq) + H2O(l)  H3O1+(aq) + Cl1-(aq) H3O1+(aq) + NH3(aq) NH41+(aq) + H2O(l)

  9. Bronsted-Lowry Acids and Bases

  10. Bronsted-Lowry Acids and Bases Formation of water by the transfer of protons from H3O1+ ions to OH1 ions.

  11. Bronsted-Lowry Acids and Bases Works in solvents other than water Solves the base problem: NH3(aq) + H3O1+(aq)  NH41+(aq) + H2O(l) CO32-(aq) + H3O1+(aq)  HCO31-(aq) + H2O(l) Doesn't solve the acid problem; What is it with Al3+(aq) and Fe3+(aq)?

  12. Lewis Acids and Bases Definitions of Gilbert Lewis, 1923 Most general theory of acids and bases Acids are electron pair acceptors. Bases are electron pair donors.

  13. Bronsted-Lowry Acids and Bases Conjugate acid-base pairs: A reaction between and acid and a base produces a conjugate acid and a conjugate base HCl(aq) + H2O(l)  H3O1+(aq) + Cl1(aq) Acid Base Conjugate Conjugate Acid Base H3O1+(aq) + NH3(aq)  H2O(l) + NH41+(aq) Acid Base Conj. Conj. Base Acid

  14. Bronsted-Lowry Acids and Bases Choose the acid, base, conjugate acid, and conjugate base: HCOOH(aq) + NH3(aq)  HCOO1(aq) + NH41+ (aq) H2PO41(aq) + H2O(l) HPO42(aq) + H3O1+ (aq) H2O(l) + HPO42(aq)  + H3O1+ (aq) + PO43(aq)

  15. Bronsted-Lowry Acids and Bases Amphoteric substances can act as both acids and bases: HCOOH(aq) + H2O(l)  HCOO1(aq) + H3O1+ (aq) NH3 (aq) + H2O(l)  NH41+ (aq) + OH1(aq) HPO42(aq) + OH1(aq) PO43(aq) + H2O(l) HPO42(aq) + H3O1+ (aq)  H2PO42(aq) + H2O(l)

  16. Mono-, Di- and Triprotic Acids Monoprotic acids can transfer one proton CH3COOH + H2O  CH3COO1 + H3O1+ Diprotic acids can transfer two protons H2CO3 + H2O  HCO31 + H3O1+ HCO31 + H2O  CO32 + H3O1+ The first proton transfer is complete before the second one starts.

  17. Mono-, Di- and Triprotic Acids Triprotic acids can transfer three protons H3PO4 + H2O  H2PO41 + H3O1+ H2PO41 + H2O  HPO42 + H3O1+ HPO42 + H2O  PO43 + H3O1+ The first proton transfer is complete before the second one starts. The second proton transfer is complete before the third one starts.

  18. Strengths of Acids and Bases Acids differ in the extent of ionization when they are put in solution Strong acids ionize completely. There are only a few strong acids. Weak acids do not ionize completely. Most acids are weak acids. The equilibrium constant, Ka, is a measure of the strength of an acid.

  19. The Strong Acids FormulaName HCl(aq) Hydrochloric acid HBr(aq) Hydrobromic acid HI(aq) Hydriodic acid HNO3 Nitric acid HClO4 Perchloric acid HClO3 Chloric acid H2SO4 Sulfuric acid* *first proton

  20. Some Weak Acids FormulaNameKa HSO41 Hydrogen sulfate 1.2 x 102 C9H8O4 Acetylsalicylic acid 3.0 x 104 HCOOH Formic acid 1.8 x 104 HC3H5O3 Lactic acid 1.4 x 104 CH3COOH Acetic acid 1.8 x 105 H2CO3 Carbonic acid 4.3 x 107 H2S(aq) Hydrosulfuric acid 1.0 x 107 HCN(aq) Hydrocyanic acid 4.9 x 1010 C6H5OH Phenol 1.3 x 1010

  21. A comparison of the number of H3O1+ ions present in strong acid and weak acid solu-tions of equal concentration.

  22. The Strong Bases Soluble compounds that contain OH1- Group 1A Hydroxides Group 2A Hydroxides LiOH NaOH KOH Ca(OH)2 RbOH Sr(OH)2 CsOH Ba(OH)2

  23. The Weak Bases Ammonia: NH3 + H2O  NH41+ + OH1 Kb = 1.8 x 105 NH41+ + H2O  NH3 + H3O1+ Ka = 5.6 x 1010

  24. The Weak Bases Anions from weak acids: CH3COO1 + H2O  CH3COOH + OH1 CO32- + H2O  HCO31- + OH1

  25. Salts A salt is a compound containing a metal or polyatomic cation, and a nonmetal or polyatomic anion (except OH1). NaCl, NH4Cl, BaSO4, CaCO3, Al2(SO4)3 Neutralization reactions between an acid and a base produce a salt and water. HCl(aq) + NaOH(aq)  H2O(l) + NaCl(aq) 2 Al(OH)3(s) + 3 H2SO4(aq)  6 H2O + Al2(SO4)3 (aq)

  26. Hydrolysis of Salts Hydrolysis is a reaction of a substance with water. Salts may hydrolyze to form H3O1+ or OH1 along with other products.

  27. Hydrolysis of Salts The salt of a weak acid and a strong base gives a weakly basic aqueous solution. NaOH + HC2H3O2  NaC2H3O2 + H2O NaC2H3O2 + H2O  HC2H3O2 + OH1 + Na1+ Reestablishes equilibrium between acetate anion and acetic acid.

  28. Hydrolysis of Salts The salt of a weak base and a strong acid gives a weakly acidic aqueous solution. NH3 + HCl  NH4Cl NH4Cl + H2O  NH3 + H3O1+ + Cl1

  29. Hydrolysis of Salts The salt of a weak base and a weak acid can give a weakly acidic, neutral, or weakly ba-sic aqueous solution, depending on acid strengths. NH4C2H3O2 + H2O  HC2H3O2 + NH3 The salt of a strong acid and a strong base give a neutral solution. NaCl + H2O  Na1+ + Cl1 + H2O

  30. Hydrolysis of Salts Some metal ions, if they're small and have a high charge, give acidic solutions. Al3+ + 2 H2O  AlOH2+ + H3O1+ Keq = 1.4 x 105 Fe3+ + 2 H2O  FeOH2+ + H3O1+ Keq = 6.3 x 103 Cr3+ + 2 H2O  CrOH2+ + H3O1+ Keq = 1.6 x 104

  31. Net Ionic Equations 2 Al(OH)3(s) + 3 H2SO4(aq)  6 H2O(l) + Al2(SO4)3(aq) 2 HCl(aq) + CaCO3(s)  2 CaCl2(aq) + CO2(g) + H2O(l) Ionic Equations show dissolved ionic sub-stances as ions rather than as compounds. Net Ionic Equations show only the participat-ing species. "Spectator" ions are not shown.

  32. Net Ionic Equations Ionic Equations 2 Al(OH)3(s) + 6 H1+(aq) + 3 SO42(aq)  6 H2O(l) + 2 Al3+(aq) + 3 SO42(aq) 2 H1+(aq) + 2 Cl2(aq) + CaCO3(s)  Ca2+(aq) + 2 Cl2(aq) + CO2(g) + H2O(l) Net Ionic Equations (NIE's) 2 Al(OH)3(s) + 6 H1+ 6 H2O(l) + 2 Al3+ 2 H1+ + CaCO3(s)  Ca2+ + CO2(g)+ H2O(l)

  33. Self-Ionization of Water The self-ionization of water is an acid-base reaction in which one water molecule trans-fers a proton to another. 2 H2O  H3O1+ + OH1Kw = 1.0 x 1014

  34. Self-Ionization of Water 2 H2O  H3O1+ + OH1Kw = 1.0 x 1014 Kw = ion product constant for water Kw = 1.0 x 1014 = [H3O1+] [OH1] In pure water, [H3O1+] = [OH1] = 1.0 x 107M

  35. The relationship between [H3O1+] and [OH1] in aqueous solution is an inverse proportion; when [H3O1+] is increased, [OH1] decreases, and vice versa.

  36. Self-Ionization of Water An acidic solution has [H3O1+] > 1.0 x 107 M [OH1] < 1.0 x 107 M A basic solution has [OH1] > 1.0 x 107 M [H3O1+] < 1.0 x 107 M A neutral solution has [H3O1+] = [OH1] = 1.0 x 107 M

  37. Self-Ionization of Water Examples: In a 0.015 M solution of HCl, what is the concentration of OH1? Is the solution acidic or basic? [OH1] is 4.0 x 105. What is [H3O1+]? Is the solution acidic or basic?

  38. The pH Scale [H3O1+] can vary over a wide range, and is often low. Often, you need scientific notation to express it. This isn't always convenient. A simpler way to write [H3O1+] is pH pH = log [H3O1+] [H3O1+] = 10pH

  39. Common (base 10) Logarigthms A logarithm is the power to which a base, such as 10, must be raised to produce a given number. Number Logarithm 0.010 = 1.0 x 10-2 -2.00 1.0 = 1.0 x 100 0.00 10 = 1.0 x 101 1.00 CoefficientExponent Characteristic Mantissa

  40. Common (base 10) Logarigthms What happens if the coefficient of the number isn’t 1.0? Number Logarithm 0.050 = 5.0 x 10-2 -1.30 5.0 = 5.0 x 100 0.70 50 = 5.0 x 1011.70 CoefficientExponent Characteristic Mantissa

  41. Common (base 10) Logarigthms How logarithms simplify mathematics: 2.594 x 103 x 6.022 x 1023 = 1.562 x 1027 log(2.594e3) + log(6.022 e23) = log(1.562e27) 3.4140 + 23.7797 = 27.1937 antilog(27.1937) = 1027.1937 = 1.562 x 1027

  42. Common (base 10) Logarigthms How logarithms simplify mathematics: Slide rules use logarithmic scales for multiplication and division.

  43. The pH Scale An acidic solution has [H3O1+] > 1.0 x 107 M pH < 7.0 A basic solution has [OH1] > 1.0 x 107 M pH > 7.0 A neutral solution has [H3O1+] = [OH1] = 1.0 x 107 M pH = 7.0

  44. A pH meter is used to measure pH values. The pH of vinegar is 2.32 (left). The pH of milk of magnesia in water is 9.39 (right).

  45. The pH Scale Give the pH for [H3O1+] = 0.010 M = 4.2 x 103 M = 1.0 x 107 M = 6.8 x 1010 M = 1.0 x 1012 M Are the solutions acidic or basic?

  46. The pH Scale Give [H3O1+] for pH = 3.00 = 4.50 = 6.85 = 7.00 = 10.75 Are the solutions acidic or basic?

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