1. Acids and Bases Chapter 15 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

2. Properties of: AcidsBases 1.Sour taste 2.Feels like water 3.Conducts electricity 4.Turns litmus paper red 5.Reacts with bases to produce salts and water 6.Reacts with some metals and releases hydrogen gas o Can you think of a reaction that this occurs? 6. React with carbonates to form carbon dioxide 1.Bitter taste 2.Feels slippery 3.Conducts electricity 4.Turns litmus paper blue 5.Reacts with acids to produce salts and water 6.Not Reactive with metals

3. Binary Acids Contains only two different elements –Hydrogen and an electronegative element (usually a halogen) Nomenclature: hydro - _________ - ic acid REVIEW ACID NAMING!!!

4. Binary Nomenclature

5. Oxyacid Contains hydrogen, oxygen, and a third element (hydrogen with a polyatomic ion) Nomenclature:

6. Polyprotic Acids Monoprotic: –Donates 1 proton –Example: Diprotic: –Donates 2 protons –Example: Triprotic: – Donates 3 protons – Example:

7. STRONG ACIDS AND BASES AND DEFINITIONS

8. Strength of Acids and Bases Strong Acids HCl HBr HI HNO 3 H 2 SO 4 HClO 4 HClO 3 Strong Bases LiOH NaOH KOH RbOH CsOH Ca(OH) 2 Sr(OH) 2 Ba(OH) 2 MEMORIZE THESE!

9. Strong Acids and Bases are strong electrolytes Weak Acids and Bases are weak electrolytes HNO 3 (aq) + H 2 O (l) H 3 O + (aq) + NO 3 - (aq) HF (aq) + H 2 O (l) H 3 O + (aq) + F - (aq) KOH (s) K + (aq) + OH - (aq) H2OH2O NO 2 - (aq) + H 2 O (l) OH - (aq) + HNO 2 (aq)

10. Arrhenius Definition Arrhenius Acid: –Substance that produces H + (H 3 O + ) in water Arrhenius Base: –Substance that produces OH - in water

11. Acid example But, NH 3 is a base and does not have (OH - )!

12. Bronsted-Lowry Definition Bronsted-Lowry Acid: –Proton donor Bronsted-Lowry Base: –Proton acceptor Forms conjugate acid-base pairs

13. acid is a proton donor base is a proton acceptor acidbase conjugate base conjugate acid

14. Conjugate Acid-Base Pairs and Strengths H 3 O + is the strongest acid that can exist in aqueous solution. The OH - ion is the strongest base that can exist in aqueous solution. The stronger the acid, the weaker the conjugate base. The stronger the base, the weaker the conjugate acid.

16. H+H+ H O H + OH - acidbase N H H H H+H+ + acidbase N H H H H + Lewis Acid and Base Lewis Acid: –Substance that can accept a pair of electrons Lewis Base: –Substance that can donate a pair of electrons

17. Lewis Acids and Bases N H H H acidbase F B F F + F F N H H H No protons donated or accepted! 15.12

18. Acid and Base Definitions

19. ACID STRENGTH

20. Acid Strength Acid Strength depends on how easily you can produce H + ions H X H + + X -

21. H X H + + X - The bigger X is, the weaker the HX bond is HF << HCl < HBr < HI Binary Acid Strength DOWN a Group Then the stronger the acid!

22. H X H + + X - The greater electronegativity of X, the more polar the HX bond, the easier to break the HX bond H 3 P < H 2 S << HCl Binary Acid Strength ACROSS a Period Then the stronger the acid!

23. Z O HZ O-O- + H + -- ++ Depends on: Electronegaitivity of “Z” Number of oxygens (oxidation state) on “Z” Oxyacid Strength

24. Oxyacids having different central atoms (Z) and that have the same number of oxygens. Acid strength increases with increasing electronegativity of Z H O Cl O O H O Br O O Cl is more electronegative than Br HClO 3 > HBrO 3 1. Electronegativity

25. Oxyacids having the same central atom (Z) but different number of oxygen atoms. Acid strength increases with the number oxygen atoms around Z. HClO 4 > HClO 3 > HClO 2 > HClO 2. Oxidation Numbers The more oxygens = stronger acid

26. Polyprotic Acids The strength of a polyprotic acid and its anions decreases with increasing negative charge. Increasing negative charge makes it more difficult to release H +. Increase # of H’s = stronger acid H 3 PO 4 › H 2 PO 4 - › HPO 4 2-

27. TYPES OF SALTS

28. Neutral Solutions: Salts containing an alkali metal or alkaline earth metal ion (except Be 2+ ) and the conjugate base of a strong acid (e.g. Cl -, Br -, and NO 3 - ). NaCl (s) Na + (aq) + Cl - (aq) H2OH2O Neutral Salts Comes from a strong base (NaOH) Comes from a strong acid (HCl)

29. Basic Salts Basic Solutions: Salts derived from a strong base and a weak acid. NaC 2 H 3 O 2 (s) Na + (aq) + C 2 H 3 O 2 - (aq) H2OH2O Comes from a strong base (NaOH) Comes from a weak acid (HC 2 H 3 O 2 )

30. Acid Solutions: Salts derived from a strong acid and a weak base. NH 4 Cl (s) NH 4 + (aq) + Cl - (aq) H2OH2O Acidic Salts Comes from a weak base (NH 3 ) Comes from a strong acid (HCl)

31. Basic Oxides Metal oxides that react with water to form a base Na 2 O + H 2 O 2 NaOH Acidic Oxides Nonmetal oxides that react with water to form acids: SO 3 + H 2 O  H 2 SO 4

32. Amphoteric Oxides React with acids to produce salt and water Al 2 O 3 (s) + 6 HCl (aq) 2 AlCl 3 (aq) + 3 H 2 O (l) React with bases to produce salt Al 2 O 3 (s) + 2 NaOH (aq) + 3 H 2 O (l) 2 Na Al(OH) 4 (aq) Intermediate element oxides (below stair step) that can act as both an acid and a base

33. Oxides Summary Basic OxidesMetallic elements Amphoteric OxidesIntermediate elements Acidic OxidesNonmetallic elements Note: There are some exceptions but you will not be responsible for knowing these

34. PH CALCULATIONS

35. O H H+ O H H O H HH O H - + [] + H 2 O (l) H + (aq) + OH - (aq) H 2 O + H 2 O H 3 O + + OH - acid conjugate base base conjugate acid autoionization of water Water is Amphoteric acts as both an acid and a base

36. H 2 O (l) H + (aq) + OH - (aq) K c = [H + ][OH - ] [H 2 O] [H 2 O] = constant K c = K w = [H + ][OH - ] ion-product constant (K w ): product of the molar concentrations of H + and OH - ions at a particular temperature. At 25 0 C K w = [H + ][OH - ] = 1.0 x 10 - 14 Ion Product Constant of Water

37. What is the concentration of OH - ions in a HCl solution whose hydrogen ion concentration is 1.3 M? K w = [H + ][OH - ] = 1.0 x 10 - 14 [H + ] = 1.3 M [OH - ] = KwKw [H + ] 1 x 10 -14 1.3 = = 7.7 x 10 -15 M Given: Solution:

38. Calculating pH pH = - log [H + ] [H + ] = [OH - ] [H + ] > [OH - ] [H + ] < [OH - ] Solution Is neutral acidic basic [H + ] = 1 x 10 -7 [H + ] > 1 x 10 -7 [H + ] < 1 x 10 -7 pH = 7 pH < 7 pH > 7 At 25 0 C pH[H + ] 15.3

39. Significant Figures There must be as many significant figures to the right of the decimal as there are in the number whose logarithm was found. –Example: [H 3 O + ] = 1 × 10 −7 one significant figure pH = 7.0

40. pOH = −log [OH – ] pH + pOH = 14.0 Calculating pOH

41. The pH of rainwater collected in a certain region of the northeastern United States on a particular day was 4.82. What is the H + ion concentration of the rainwater? pH = - log [H + ] [H + ] = 10 -pH = 10 -4.82 = 1.5 x 10 -5 M The OH - ion concentration of a blood sample is 2.5 x 10 -7 M. What is the pH of the blood? pH + pOH = 14.00 pOH = -log [OH - ]= -log (2.5 x 10 -7 )= 6.60 pH = 14.00 – pOH = 14.00 – 6.60 = 7.40

42. The Circle of pH pH pOH [ H 3 O + ] [ OH - ] -log [H 3 O + ] 10 (-pH) 10 (-pOH) -log [OH - ] [ H 3 O + ] [ OH - ] = 10 -14 pH+ pOH = 14

43. What is the pH of a 2 x 10 -3 M HNO 3 solution? HNO 3 is a strong acid – 100% dissociation. HNO 3 (aq) + H 2 O (l) H 3 O + (aq) + NO 3 - (aq) pH = -log [H + ] = -log [H 3 O + ] = -log(0.002) = 2.7 Start End 0.002 M 0.0 M

44. What is the pH of a 1.8 x 10 -2 M Ba(OH) 2 solution? Ba(OH) 2 is a strong base – 100% dissociation. Ba(OH) 2 (aq) Ba 2+ (aq) + 2OH - (aq) Start End 0.018 M 0.036 M0.0 M pH = 14.00 – pOH = 14.00 + log(0.036) = 12.56

45. HA (aq) + H 2 O (l) H 3 O + (aq) + A - (aq) HA (aq) H + (aq) + A - (aq) K a = [H + ][A - ] [HA] K a = acid ionization constant KaKa weak acid strength Weak Acids (HA) and Acid Ionization Constants

47. What is the pH of a 0.5 M HF solution (at 25 0 C)? HF (aq) H + (aq) + F - (aq) K a = [H + ][F - ] [HF] = 7.1 x 10 -4 HF (aq) H + (aq) + F - (aq) Initial (M) Change (M) Equilibrium (M) 0.500.00 -x-x+x+x 0.50 - x 0.00 +x+x xx

48. What is the pH of a 0.5 M HF solution (at 25 0 C)? Continued… K a = x2x2 0.50 - x = 7.1 x 10 -4 Ka  Ka  x2x2 0.50 = 7.1 x 10 -4 0.50 – x  0.50 K a << 1 x 2 = 3.55 x 10 -4 x = 0.019 M [H + ] = [F - ] = 0.019 M pH = -log [H + ] = 1.72 [HF] = 0.50 – x = 0.48 M

49. 0.50 – x  0.50 K a << 1 When the percent ionization is less than or equal to 5% x = 0.019 0.019 M 0.50 M x 100% = 3.8% Less than 5% Approximation ok. When can I use the approximation? Percent ionization = [H + ] [HA] 0 x 100% [HA] 0 = initial concentration To calculate percent ionization:

50. Does the approximation always work? What is the pH of a 0.050 M HF solution (at 25 0 C)? Ka  Ka  x2x2 0.050 = 7.1 x 10 -4 x = 0.0060 M 0.006 M 0.050 M x 100% = 12% More than 5% Approximation not ok. Must solve for x exactly using quadratic equation.

51. K a = x2x2 0.05 - x = 7.1 x 10 -4 x 2 + 0.00071x – 3.55 x 10 -5 = 0 ax 2 + bx + c =0 -b ± b 2 – 4ac  2a2a x = x = 0.0056x = - 0.0063 HF (aq) H + (aq) + F - (aq) Initial (M) Change (M) Equilibrium (M) 0.050.00 -x-x+x+x 0.05 - x 0.00 +x+x xx [H + ] = x = 0.0056 M pH = -log[H + ] = 2.25

52. Solving weak acid ionization problems: 1.Identify the major species that can affect the pH. In most cases, you can ignore the autoionization of water. Ignore [OH - ] because it is determined by [H + ]. 2.Use ICE to express the equilibrium concentrations in terms of single unknown x. 3.Write K a in terms of equilibrium concentrations. Solve for x by the approximation method. If approximation is not valid, solve for x exactly. 4.Calculate concentrations of all species and/or pH of the solution.

53. What is the pH of a 0.122 M monoprotic acid whose K a is 5.7 x 10 -4 ? HA (aq) H + (aq) + A - (aq) Initial (M) Change (M) Equilibrium (M) 0.1220.00 -x-x+x+x 0.122 - x 0.00 +x+x xx K a = x2x2 0.122 - x = 5.7 x 10 -4 Ka  Ka  x2x2 0.122 = 5.7 x 10 -4 0.122 – x  0.122 K a << 1 x 2 = 6.95 x 10 -5 x = 0.0083 M 0.0083 M 0.122 M x 100% = 6.8% More than 5% Approximation not ok.

54. K a = x2x2 0.122 - x = 5.7 x 10 -4 x 2 + 0.00057x – 6.95 x 10 -5 = 0 ax 2 + bx + c =0 -b ± b 2 – 4ac  2a2a x = x = 0.0081x = - 0.0081 HA (aq) H + (aq) + A - (aq) Initial (M) Change (M) Equilibrium (M) 0.1220.00 -x-x+x+x 0.122 - x 0.00 +x+x xx [H + ] = x = 0.0081 M pH = -log[H + ] = 2.09

55. NH 3 (aq) + H 2 O (l) NH 4 + (aq) + OH - (aq) K b = [NH 4 + ][OH - ] [NH 3 ] K b is the base ionization constant KbKb weak base strength Solve weak base problems like weak acids except solve for [OH-] instead of [H + ]. Weak Bases and Base Ionization Constants

57. HA (aq) H + (aq) + A - (aq) A - (aq) + H 2 O (l) OH - (aq) + HA (aq) KaKa KbKb H 2 O (l) H + (aq) + OH - (aq) KwKw K a K b = K w Ionization Constants of Conjugate Acid-Base Pairs

59. Acid-Base Equilibria and Solubility Equilibria Chapter 16 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

60. Must have: 1.A weak acid or a weak base AND 2. The salt of the weak acid or weak base buffer solution: ability to resist changes in pH upon the addition of small amounts of either acid or base. Buffer Solutions

61. Buffer Example Add strong acid H + (aq) + CH 3 COO - (aq) CH 3 COOH (aq) Add strong base OH - (aq) + CH 3 COOH (aq) CH 3 COO - (aq) + H 2 O (l) Consider an equal molar mixture of CH 3 COOH and CH 3 COONa

62. Which of the following are buffer systems? (a) KF/HF (b) KBr/HBr (c) Na 2 CO 3 /NaHCO 3 (a) HF is a weak acid and F - is its conjugate base buffer solution (b) HBr is a strong acid not a buffer solution (c) CO 3 2- is a weak base and HCO 3 - is it conjugate acid buffer solution

63. 16.3 What are some examples of buffers?

64. TITRATIONS

65. Titrations Titration: when a solution of accurately known concentration is added gradually added to another solution of unknown concentration until the chemical reaction between the two solutions is complete. Slowly add base to unknown acid UNTIL The indicator changes color (pink) Indicator – substance that changes color at (or near) the equivalence point

66. Titration Points equivalence point: point at which the two solutions used in a titration are present in chemically equivalent amounts end point: point in a titration at which an indicator changes color

67. Common Indicators and pH Ranges

68. Which indicator do I choose? pH less than 7 –Indicators that change color at pH lower than 7 are used to determine the equivalence point of strong-acid/weak- base titrations. –strong-acid/weak-base titration = acidic. pH at 7 –Indicators that undergo transition at about pH 7 are used to determine the equivalence point of strong-acid/strong base titrations. –strong acids/strong bases = salt solution with a pH of 7.

69. The titration curve of a strong acid with a strong base. 16.5

70. Which indicator(s) would you use for a titration of HNO 2 with KOH ? Weak acid titrated with strong base. At equivalence point, will have conjugate base of weak acid. At equivalence point, pH > 7 Use cresol red or phenolphthalein

71. NaOH (aq) + HCl (aq) H 2 O (l) + NaCl (aq) OH - (aq) + H + (aq) H 2 O (l) 0.10 M NaOH added to 25 mL of 0.10 M HCl Strong Acid-Strong Base Titrations

72. CH 3 COOH (aq) + NaOH (aq) CH 3 COONa (aq) + H 2 O (l) CH 3 COOH (aq) + OH - (aq) CH 3 COO - (aq) + H 2 O (l) At equivalence point (pH > 7): 16.4 Weak Acid-Strong Base Titrations

73. HCl (aq) + NH 3 (aq) NH 4 Cl (aq) At equivalence point (pH < 7): 16.4 H + (aq) + NH 3 (aq) NH 4 + (aq) Strong Acid-Weak Base Titrations

74. Strong Acid-Strong Base Calcs For monoprotic acids and bases: M A V A = M B V B

75. Normality The value for normality is always larger than that of molarity. Normality takes into consideration the number of hydrogens/hydroxides a substance will produce when dissociated. N A V A = N B V B

76. Normality Example 25.0 mL of 1.5 M sample of H 3 PO 4 is neutralized by 22.5 mL of Zn(OH) 2. What is the concentration of Zn(OH) 2 ?

77. Antacids Video with Methyl Orange

78. KSP

79. AgCl (s) Ag + (aq) + Cl - (aq) K sp = [Ag + ][Cl - ] K sp is the solubility product constant MgF 2 (s) Mg 2+ (aq) + 2F - (aq) K sp = [Mg 2+ ][F - ] 2 Ag 2 CO 3 (s) 2Ag + (aq) + CO 3 2 - (aq) K sp = [Ag + ] 2 [CO 3 2 - ] Ca 3 (PO 4 ) 2 (s) 3Ca 2+ (aq) + 2PO 4 3 - (aq) K sp = [Ca 2+ ] 3 [PO 3 3 - ] 2 Solubility Equilibrium

80. 16.6

81. Molar solubility (mol/L):number of moles of solute dissolved in 1 L of a saturated solution. Solubility (g/L) is the number of grams of solute dissolved in 1 L of a saturated solution. Solubility Definitions

82. What is the solubility of silver chloride in g/L ? K sp = 1.6 x 10 -10 AgCl (s) Ag + (aq) + Cl - (aq) K sp = [Ag + ][Cl - ] = 1.6 x 10 -10 Initial (M) Change (M) Equilibrium (M) 0.00 +s+s +s+s ss K sp = s 2 s = K sp = 1.3 x 10 -5  [Ag + ] = 1.3 x 10 -5 M [Cl - ] = 1.3 x 10 -5 M Solubility of AgCl = 1.3 x 10 -5 mol AgCl 1 L soln 143.35 g AgCl 1 mol AgCl x = 1.9 x 10 -3 g/L

83. 16.6

84. Q vs. K sp Dissolution of an ionic solid in aqueous solution: Q = K sp Saturated solution Q < K sp Unsaturated solution No precipitate Q > K sp Supersaturated solution Precipitate will form

85. If 2.00 mL of 0.200 M NaOH are added to 1.00 L of 0.100 M CaCl 2, will a precipitate form? The ions present in solution are Na +, OH -, Ca 2+, Cl -. Only possible precipitate is Ca(OH) 2 (solubility rules). Is Q > K sp for Ca(OH) 2 ? [Ca 2+ ] 0 = 0.100 M[OH - ] 0 = 4.0 x 10 -4 M K sp = [Ca 2+ ][OH - ] 2 = 8.0 x 10 -6 Q = [Ca 2+ ] 0 [OH - ] 0 2 = 0.10 x (4.0 x 10 -4 ) 2 = 1.6 x 10 -8 Q < K sp No precipitate will form From CaCl 2 : 0.100 M CaCl 2 and 1.00 L = 0.100 mol Ca 2+ / 1.002 L From NaOH: 0.200 M NaOH and 0.002 L = 0.0004 mol OH - / 1.002 L

86. common ion effect: shift in equilibrium caused by the addition of a compound having an ion in common with the dissolved substance. The presence of a common ion suppresses the ionization of a weak acid or a weak base. Consider mixture of CH 3 COONa (strong electrolyte) and CH 3 COOH (weak acid). CH 3 COONa (s) Na + (aq) + CH 3 COO - (aq) CH 3 COOH (aq) H + (aq) + CH 3 COO - (aq) common ion

87. The Common Ion Effect and Solubility The presence of a common ion decreases the solubility of the salt. What is the molar solubility of AgBr in pure water? AgBr (s) Ag + (aq) + Br - (aq) K sp = 7.7 x 10 -13 s 2 = K sp s = 8.8 x 10 -7

88. Solubility with Common Ion NaBr (s) Na + (aq) + Br - (aq) [Br - ] = 0.0010 M AgBr (s) Ag + (aq) + Br - (aq) [Ag + ] = s [Br - ] = 0.0010 + s  0.0010 K sp = 0.0010 x s s = 7.7 x 10 -10 s = 8.8 x 10 -7 s = 7.7 x 10 -10 > Pure waterSolution without common ion > with common ion What is the molar solubility of AgBr in 0.0010 M NaBr?

89. Solubility with and without common ion s = 8.8 x 10 -7 s = 7.7 x 10 -10 > without common ion > with common ion AgBr in water vs. AgBr in NaBr Solution

90. pH and Solubility The presence of a common ion decreases the solubility. Insoluble bases dissolve in acidic solutions Insoluble acids dissolve in basic solutions Mg(OH) 2 (s) Mg 2+ (aq) + 2OH - (aq) K sp = [Mg 2+ ][OH - ] 2 = 1.2 x 10 -11 K sp = (s)(2s) 2 = 4s 3 4s 3 = 1.2 x 10 -11 s = 1.4 x 10 -4 M [OH - ] = 2s = 2.8 x 10 -4 M pOH = 3.55 pH = 10.45 removeadd

91. If the normal pH is 10.45… If you lower the pH less than 10.45 Lower [OH - ] OH - (aq) + H + (aq) H 2 O (l) Increase solubility of Mg(OH) 2 If you raise the pH greater than 10.45 Raise [OH - ] Decrease solubility of Mg(OH) 2 Mg(OH) 2 (s) Mg 2+ (aq) + 2OH - (aq) removeadd