1. ACID-BASE TITRATIONS PART 3

2. WHAT DOES THE TITRATION GRAPH TELL? If we have a solid that dissolves: A 2 B (s)  2 A (aq) + B (aq) Then K sp is calculated by K sp = [A] 2 [B] ; at equilibrium (saturation)

3. STRONG ACID WITH STRONG BASE

4. WEAK BASE WITH STRONG ACID

12. GENERIC K SP If we have a solid that dissolves: A 2 B (s)  2 A (aq) + B (aq) Then K sp is calculated by K sp = [A] 2 [B] ; at equilibrium (saturation)

13. GENERAL K SP Solubility Product of a compound equals the product of the concentration of the ions involved in the equilibrium, each raised to the power of its coefficient in the equilibrium equation.

14. SOLUBILITY VS. K SP Solubility of a substance is the quantity that dissolves to form a saturated solution. Usually in g/L or mol/L K sp is the equilibrium constant between the ionic solid and the saturated solution. It has no units and is a measure of how much solid dissolves to form a saturated solution.

15. SOLUBILITY VS. K SP Solubility of a substance is the quantity that dissolves to form a saturated solution. Usually in g/L or mol/L K sp is the equilibrium constant between the ionic solid and the saturated solution. It has no units and is a measure of how much solid dissolves to form a saturated solution.

17. PRECIPITATION, ION SEPERATION AND K SP What is Q? (Not in terms of James Bond) Reaction Quotient (Q) : number obtained by substituting reactant and product concentrations (or partial pressures) at any point during a reaction into an equilibrium-constant expression.

18. PRECIPITATION, ION SEPERATION AND K SP What is Q? (Not in terms of James Bond) a A + b B   d D + e E Q = [D] d [E] e [A] a [B] b

19. PRECIPITATION, ION SEPERATION AND K SP Q and K sp What do they say? Q > K sp Precipitation occurs until Q = K sp Q = K sp at equilibrium (saturated) because Q = K sp Q < K sp solid dissolves until Q = K sp (think if the solid has not dissolved, then Q is 0 and needs to increase until Q = K sp

21. STRONG ACIDS Seven most common strong acids include six monoprotic acids and one diprotic acid HCl, HBr, HI, HNO 3, HClO 3, HClO 4, H 2 SO 4

24. STRONG ACIDS Exist entirely in solution as ions notice the yield sign HNO 3 (aq) + H 2 O (l)  H 3 O + (aq) + NO 3 - (aq) Or we can state this as: HNO 3 (aq)  H + (aq) + NO 3 - (aq)

25. HYDROLYSIS Ions ability to react with water to form H + ions and OH - ions.

26. HYDROLYSIS WITH ANIONS Ions ability to react with water to form H + ions and OH - ions. Look at the anion and see if it is a conjugate base of a strong acid. If yes, the tendency to abstract ions from water is negligible.

27. HYDROLYSIS WITH ANIONS If the anion is not a conjugate base of a strong acid, then it is a weak base and would make a weak acid by taking a H + ion and producing more OH - ions, thereby raising the pH (more basic).

28. HYDROLYSIS WITH ANIONS Complicated if anion has ionizable protons, such as H 2 PO 4 -, because it is amphiprotic, act as acid or base. Look at K a and K b of the ion. If K a >K b then it causes solution to be acidic. If K a <K b then causes solution to be basic.

29. HYDROLYSIS WITH CATIONS Ions ability to react with water to form H + ions and OH - ions. Polyatomic cation that containing one or more protons (H + ) will donate H + to water to make H 3 O +

30. HYDROLYSIS WITH CATIONS Cations of metals (not alkali and alkaline earth) attract water molecules and become hydrated (remember hydrates) and then cause a water molecule from the hydrate to lose a H + ion to another water molecule creating a hydronium ion.

31. ACID STRENGTH Depends on intermolecular forces. Polar bonds involving hydrogen allow H to act as proton donor (HCl) or Proton acceptor (NaH) Strength of bond. HF is so strong to not allow it to be a strong acid.

32. ACID STRENGTH Stability of the conjugate base. More stable conjugate bases are found in stronger acids. Acids with elements in the same group or family show stronger acids as the elements get bigger and bonds get weaker between the element and hydrogen.

33. ACID STRENGTH Acidity increases as the electronegativity increases in a period.

34. OXYACIDS Acids in which one or more O-H bonds are connected to a central atom (H 2 SO 4 ) Stronger acids are with similar ions with the same central atom but more oxygens. ClO - < ClO 2 - < ClO 3 - < ClO 4 -

35. CARBOXYLIC ACIDS Acids which contain a carboxyl group, COOH. Resonance of COO - allows for stability of conjugate base.

36. ARRHENIUS ACID- BASE Acids have Hydrogen ions in solution Bases have Hydroxide ions in solution

37. BRØNSTED-LOWRY ACID- BASE Acids are proton donors Bases are proton acceptors

38. LEWIS ACID- BASE Leweis Acids is an electron pair acceptor Lewis Bases are an electron pair donor. Everything already defined as a base (OH -, H 2 O, amine (N), or anion) are still bases.

39. LEWIS ACID- BASE Lewis Bases can also be an electron pair donor to things other than H + Ex. NH 3 + BF 3  NH 3 BF 3 Lewis Base Lewis Acid Electron Pair Acceptor is always referred to as a Lewis Acid if it is not usually called an acid

43. STRONG ACIDS Usually a strong acid is the single source of H + (aq) HNO 3 (aq)  H + (aq) + NO 3 - (aq) * If the [acid] is less than 10 -6 then we need to consider H + from H 2 O

44. STRONG BASES Seven most common soluble strong bases are the ionic hydroxides of the alkali metals and the ionic hydroxides of the alkaline earth metals NaOH, LiOH, KOH, etc. Ca(OH) 2, Sr(OH) 2, Mg(OH) 2, etc.

45. STRONG BASES Ionic hydroxides of the heavier alkaline earth metals Ca(OH) 2, Sr(OH) 2, Ba(OH) 2, Ra(OH) 2 have limited solubility and are used when solubility is not critical

46. STRONG BASES Strong basic solutions are also formed by substances that react with water to create OH - Most common substances are oxides Na 2 O, CaO, etc. Oxide ion reacts like: O 2- (aq) + H 2 O (l)  2 OH - (aq) Notice yield sign

47. WEAK ACIDS Acids that partially ionize in aqueous solutions Follow this generic reaction: HA (aq) + H 2 O (l)  H 3 O + (aq) + A - (aq) Equilibrium constant: K c = [H 3 O + ][A - ] = [H + ][A - ] [HA] [HA]

48. ACID-DISSOCIATION CONSTANT Equilibrium Constant to show ionization of an acid (weak acid). Larger K a indicates stronger acid Follow this generic reaction: HA (aq) + H 2 O (l)  H 3 O + (aq) + A - (aq) Equilibrium constant: K a = [H 3 O + ][A - ] = [H + ][A - ] [HA] [HA]

49. WEAK BASES Bases that react with water and aquire protons, creating conjugate acid and OH - Follow this generic reaction: B (aq) + H 2 O (l)  BH + (aq) + OH - (aq) Equilibrium constant: K c = [BH + ][OH - ] = [BH + ][OH - ] [B] [B]

50. BASE-DISSOCIATION CONSTANT Equilibrium Constant to show base reacting with H 2 O. Larger K a indicates stronger acid Follow this generic reaction: B (aq) + H 2 O (l)  H 3 O + (aq) + A - (aq) Equilibrium constant: K b = [BH + ][OH - ] = [BH + ][OH - ] [B] [B]

51. WEAK BASES Need one or more lone pair of electrons for a H + to bond Two categories: Neutral substances with a lone pair of electrons (usually nitrogen, amines) Anions of weak acids are weak bases (ClO - from NaClO)

52. RELATIONSHIP BETWEEN K A AND K B K a x K b = K w NH 4 + (aq) + H 2 O (l)  NH 3 (aq) + H 3 O + (aq) or: NH 4 + (aq)  NH 3 (aq) + H + (aq) and: NH 3 (aq) + H 2 O (l)  NH 4 + (aq) + OH - (aq) K a = [NH 3 ][H + ] K b = [NH 4 + ][OH - ] [NH 4 + ] [NH 3 ]

53. RELATIONSHIP BETWEEN K A AND K B When we add the acid and base reactions: NH 4 + (aq)  NH 3 (aq) + H + (aq) NH 3 (aq) + H 2 O (l)  NH 4 + (aq) + OH - (aq) H 2 O (l)  H + (aq) + OH - (aq) Then we multiply equilibrium constants: K a x K b = [NH 3 ][H + ] x [NH 4 + ][OH - ] [NH 4 + ] [NH 3 ]

54. RELATIONSHIP BETWEEN K A AND K B K a x K b = [H + ][OH - ] = K w So K a x K b = K w

55. CALCULATING K A FROM PH 0.10 M formic acid (HCOOH) has a pH of 2.38 at 25 °C. What is the K a ? 1 st Write reaction: HCOOH (aq)  H + (aq) + HCOO + (aq) To get the K a = [H + ][HCOO - ] [HCOOH]

56. CALCULATING K A FROM PH Calculate [H + ] from the pH of 2.38 pH = -log [H + ] 2.38 = -log [H + ] -2.38 = log [H + ] 10 -2.38 = [H + ] 4.2 x 10 -3 M = [H + ]

57. EXPRESS CHANGE IN [ ] BY LOOKING AT BCE COEFFICIENTS Concentrations in M = mol/L [HCOOH][H + ][HCOO - ] Initial [ ]0.1000 Change in [ ]-4.2 x 10 -3 4.2 x 10 -3 Equilibrium [ ] 0.10 - 4.2 x 10 -3 4.2 x 10 -3 BCE HCOOH (aq)  H + (aq)+ HCOO - (aq)

58. CALCULATING K A FROM PH Since equilibrium concentration for HCOOH is 0.10 – 4.2 x 10 -3 we get 0.10 – 0.0042 = 0.10 (with rounding) K a = [H + ][HCOO - ] [HCOOH] K a = (4.2 x 10 -3 ) (4.2 x 10 -3 ) = 1.8 x 10 -4 0.10

59. CALCULATING PH FROM K A What is pH of 0.30 M acetic acid (HCOOH) at 25 °C? 1 st Write reaction: CH 3 COOH (aq)  H + (aq) + CH 3 COO + (aq) To get the K a = [H + ][CH 3 COO - ] [CH 3 COOH] K a = 1.8 x 10 -5 from a table 16.2 p 667

60. EXPRESS CHANGE IN [ ] BY LOOKING AT BCE COEFFICIENTS Concentrations in M = mol/L [CH 3 COOH ] [H + ][CH 3 COO - ] Initial [ ]0.3000 Change in [ ]-x +x Equilibrium [ ]0.30 -x x x BCE CH 3 COOH (aq)  H + (aq)+ CH 3 COO - (aq)

61. CALCULATING PH FROM K A K a = [H + ][CH 3 COO - ] [CH 3 COOH] K a = (x) (x) = 1.8 x 10 -4 0.30 - x We can assume that x is negligible compared to 0.30 since K a is small and equilibrium should lie to the reactant side

62. CALCULATING PH FROM K A K a = (x) (x) = 1.8 x 10 -4 0.30 x 2 = (0.30) (1.8 x 10 -4 ) x 2 = 5.4 x 10 -6 x = √ 5.4 x 10 -6 x = 2.3 x 10 -3 [H + ] = 2.3 x 10 -3 M pH = -log(2.3 x 10 -3 ) = 2.64

63. PERCENT IONIZATION OF AN ACID Percent ionization = concentration ionized x 100 % original concentration Percent ionization = [H + ] equilibrium x 100 % [HA] initial

64. PERCENT IONIZATION OF AN ACID What is the percent ionization of a 0.035 M solution of HNO 2 that cantains 3.7 x 10 -3 M H + (aq)? Percent ionization = [H + ] equilibrium x 100 % [HA] initial Percent ionization = 3.7 x 10 -3 M x 100 % 0.035 M

66. ION-PRODUCT CONSTANT OF WATER K w = [H + ][OH - ] = 1.0 x 10 -14 (25 °C) If we know the K w and the concentration of either hydrogen ions or hydroxide ions we can calculate the missing variable

67. STRONG ACIDS Seven most common strong acids include six monoprotic acids and one diprotic acid HCl, HBr, HI, HNO 3, HClO 3, HClO 4, H 2 SO 4

69. STRONG ACIDS Seven most common strong acids include six monoprotic acids and one diprotic acid HCl, HBr, HI, HNO 3, HClO 3, HClO 4, H 2 SO 4

70. LEWIS ACID Contain H + in the formula When dissolved it increases the concentration of hydrogen ions

71. LEWIS BASE Contain OH - in the formula When dissolved it increases the concentration of hydroxide ions

72. BRONSTED-LOWRY ACID Acid is a H + donor Proton donor

73. BRONSTED-LOWRY BASE Base is a H + acceptor Proton acceptor

74. CONJUGATE ACID-BASE PAIR Acids form a conjugate base when they lose the proton Bases form a conjugate acid when they gain the proton

75. RELATIVE STRENGTHS OF ACIDS AND BASES Strong acids form conjugate bases that have negligible basicity Weak acids form conjugate bases that are weak bases Negligible acid in strength form conjugate bases that are strong bases

76. AUTOIONIZATION OF WATER Water can act as an acid or a base under the Bronsted-Lowry definition About 2 out of 10 9 water molecules ionize at a time. 2 H 2 O (l)  OH - (aq) + H 3 O + (aq)

77. ION-PRODUCT CONSTANT OF WATER Equilibrium constant for the autoionization of water is: K c = [H 3 O + ][OH - ] = 1.0 x 10 -14 (25 °C) This refers specifically to water so: K w = [H 3 O + ][OH - ] = 1.0 x 10 -14 (25 °C)

78. AUTOIONIZATION OF WATER Can be written like this in terms of one water molecule H 2 O (l)  OH - (aq) + H + (aq)

79. ION-PRODUCT CONSTANT OF WATER Equilibrium constant for the autoionization of water is: K w = [H 3 O + ][OH - ] = 1.0 x 10 -14 (25 °C) Or K w = [H + ][OH - ] = 1.0 x 10 -14 (25 °C) [H + ] and [H 3 O + ] can be used interchangeably

80. NEUTRAL [H + ] = [OH - ]

81. ION-PRODUCT CONSTANT OF WATER K w = [H + ][OH - ] = 1.0 x 10 -14 (25 °C) If we know the K w and the concentration of either hydrogen ions or hydroxide ions we can calculate the missing variable

82. PH SCALE Think of pH as the power of H + : pH = - log[H + ] = -log [H 3 O + ] Example: pH = -log(1.0 x 10 -7 ) = -(-7.00) = 7.00 In a logarithm, the numbers to the right are the only significant figures

83. POH SCALE Think of pOH as the power of OH - : pOH = - log[OH - ] = -log [OH - ] Example: pOH = -log(1.0 x 10 -7 ) = -(-7.00) = 7.00 In a logarithm, the numbers to the right are the only significant figures

84. POH SCALE -log[H + ] + (-log[OH - ] ) = -log K w pH + pOH = 14.00 ( 25 °C)