1. Acids, Bases, and Salts By: Chris Squires

2. Effects of Acid Rain on Marble (marble is calcium carbonate CaCO 3 ) George Washington: BEFORE acid rain AFTER acid rain

3. Acids have a pH less than 7

4. OPERATIONAL ACIDS Reacts with metals to form H 2 gas. HCl (aq) + Mg (s) → MgCl 2(aq) + H 2(g) HCl (aq) + Mg (s) → MgCl 2(aq) + H 2(g)

5. Taste Sour

6. Conducts electricity.

7. Neutralizes BASES

9. Acids Turn Litmus Red Blue litmus paper turns red in contact with an acid

10. Acids React with Carbonates and Bicarbonates HCl + NaHCO 3 NaCl + H 2 O + CO 2 Hydrochloric acid + sodium bicarbonate salt + water + carbon dioxide An old-time home remedy for relieving an upset stomach

11. OPERATIONAL BASES

12. BASES ARE SLIPPERY

13. BASES are Corrosive

14. BASES

15. BASES OPERATE

16. Bases have a pH greater than 7

17. Bases Turn Litmus Blue Red litmus paper turns blue in contact with a base (and blue paper stays blue).

18. Properties of Bases (metallic hydroxides) n React with acids forms water and salt. n Taste bitter & feel slippery n Electrolytes in aqueous n Change the color of indicators (bases go Blue, with litmus).

19. Examples of Bases  Sodium hydroxide, NaOH (lye for drain cleaner; soap)  Potassium hydroxide, KOH (alkaline batteries)  Magnesium hydroxide, Mg(OH) 2 (Milk of Magnesia)

20. Theories Acid Base H+ H+H+

21. Acid-Base Theories n a) Arrhenius, n b) Brønsted-Lowry n c) Lewis (not covered).

22. Svante Arrhenius (1859-1927)

23. 1. Arrhenius Definition - 1887 n Acids produce (H + ) in solution (HCl → H + + Cl - ) n Bases produce OH - in water. (NaOH → Na + + OH - ) n Limited to aqueous solutions. n Only one kind of base (hydroxides) n NH 3 (ammonia) could not be an Arrhenius base: no OH - produced.

24. 3 flaws with Arrhenius 1. A proton (H + ) does not exist ALONE in water, it makes H 3 O + (aq) 2. NH 3 (g) is a base when dissolved in water; but does not contain the OH - (aq) 3. Amphiprotics like HCO 3 - (aq) appear to contain the H + ion BUT ARE bases

25. Johannes Brønsted Thomas Lowry (1879-1947) (1874-1936) Denmark England

26. Brønsted-Lowry - 1923 n A broader definition than Arrhenius n Acid is a H +,base is H + acceptor. n HCl is an acid. –When it dissolves in water, it gives it’s proton to water. HCl (g) + H 2 O (l) ↔ H 3 O + (aq) + Cl - (aq) HCl (g) + H 2 O (l) ↔ H 3 O + (aq) + Cl - (aq) n Water is a base; makes hydronium ion.

27. A Brønsted–Lowry acid… …must have removable proton. HCl, H 2 O, H 2 SO 4 A Brønsted–Lowry base must accept this H+ NH 3, H 2 O, CO 3 2- and usually has negative charge

28. Acids and bases come in pairs n A “conjugate base” is the remainder of the original acid, after it donates it’s hydrogen ion n A “conjugate acid” is the particle formed when the original base gains a hydrogen ion

29. Conjugate Acids and Bases: n From the Latin word conjugare, meaning “to join together.” n Reactions between acids and bases always yield their conjugate bases and acids.

30. If can also be BOTH Acid and Base…...it is amphiprotic. HCO 3 – HSO 4 – H2OH2OH2OH2O

31. Who Theory: Acid When Arrheniusincreases H + 1880’s Brønstedproton donor1923 Lewis electron pair acceptor same1923 Three THEORETICAL definitions

32. Hydrogen Ions from Water n Water ionizes, or falls apart into ions: H 2 O ↔ H + + OH - n Called the “self ionization” of water n Occurs to a very small extent: [ H + ] = [OH - ] = 1 x 10 -7 M n Since they are equal a neutral solution n K w = [H + ] x [OH - ] = 1 x 10 -14 n 14 = pH + pOH

34. pH and Significant Figures n [H + ] = 0.0010 M = 1.0 x 10 -3 M 0.0010 has 2 sig figs 0.0010 has 2 sig figs n the pH = 3.00, with the two numbers to the right of the decimal corresponding to the 2 sig figs

35. STRONG V WEAK

36. STRONG VS WEAK

37. Strength In any acid-base reaction, the equilibrium favors the reaction that moves the proton to the stronger base. HCl ( aq ) + H 2 O ( l )  H 3 O + ( aq ) + Cl – ( aq) Cl – equilibrium lies so far to the right K is not measured ( K >>1), so is not a base, it is neutral

38. Strength n Strong acids completely dissociate in water. –Their conjugate bases are actually so weak as not to be able to act as a base at all.

39. Strength n Acids & Bases are classified according to the degree they ionize in water: –Strong are completely ionized, 100% –Weak ionize only slightly, like 1% Strength & Concentration Night V Day

40. Strong Acid Dissociation (makes 100 % ions)

41. Weak Acid Dissociation (only partially ionizes)

43. STRONG V WEAK ACIDS

44. LOL

45. Measuring strength n Ionization is reversible for WEAK acids: HA + H 2 O ↔ H + + A - n This makes an equilibrium n K a = [H + ][A - ] [HA] n Strong acid = more products & large K a (Note that water is NOT shown, (Note that the arrow goes both directions.)

46. Calculating pH from K a Calculate the pH of a 0.30 M solution of acetic acid, C 2 H 3 O 2 H, at 25°C. K a for acetic acid at 25°C is 1.8  10 -5. K a for acetic acid at 25°C is 1.8  10 -5.

47. Calculating pH from K a Use the ICE table: [C 2 H 3 O 2 ], M[H 3 O + ], M[C 2 H 3 O 2 − ], M Initial0.3000 Change–x–x+x+x+x+x Equilibrium0.30 – xxx

48. Calculating pH from K a Use the ICE table: [C 2 H 3 O 2 ], M[H 3 O + ], M[C 2 H 3 O 2 − ], M Initial0.3000 Change–x–x+x+x+x+x Equilibrium0.30 – x ≈ 0.30 xx Simplify: how big is x relative to 0.30?

49. Calculating pH from K a Now, (1.8  10 -5 ) (0.30) = x 2 5.4  10 -6 = x 2 2.3  10 -3 = x pH = = 2.64

50. Calculating K a from the pH n The pH is 2.38 for a 0.10 M solution of formic acid, HCOOH,. Calculate K a K a, requires equilibrium concentrations. We can find [H 3 O + ], which is the same as [HCOO − ], from pH, so pH = X

51. Calculating K a from the pH pH = –log [H 3 O + ] – 2.38 = log [H 3 O + ] 10 -2.38 = 10 log [H 3 O + ] = [H 3 O + ] 4.2  10 -3 = [H 3 O + ] = [HCOO – ] = X

52. Calculating K a from pH All values In Mol/L [HCOOH][H 3 O + ][HCOO − ] Initially0.1000 Change –4.2  10 -3 +4.2  10 -3 Equilibrium 0.10 – 4.2  10 -3 = 0.0958 = 0.10 4.2  10 -3

53. Calculating K a from pH [4.2  10 -3 ] [0.10] K a = = 1.8  10 -4

54. Calculating Percent Ionization [A - ] eq = [H 3 O + ] eq = 4.2  10 -3 M [A - ] eq = [H 3 O + ] eq = 4.2  10 -3 M [A - ] eq + [HCOOH] eq = [HCOOH] initial = 0.10 M [A - ] eq + [HCOOH] eq = [HCOOH] initial = 0.10 M %rx = [4.2  10 -3 / 0.10] (x 100) %rx = [4.2  10 -3 / 0.10] (x 100) = 4.2%

55. What about bases? n Strong bases dissociate completely. n MOH + H 2 O ↔ M + + OH - (M = a metal) n Base dissociation constant = K b n K b = [M + ][OH - ] [MOH] n Stronger base = more dissociated ions are produced, thus a n larger K b and a small Ka.

56. Weak Bases The Kb for this reaction is The Kb for this reaction is

57. pH of Basic Solutions What is the pH of a 0.15M solution of NH 3 ? [NH 4 + ] [OH − ] [NH 3 ] K b = = 1.8  10 -5

58. pH of Basic Solutions [NH 3 ][NH 4 + ][OH − ] Initial0.1500 Equil 0.15 - x  0.15 xx

59. pH of Basic Solutions Ignore X (1.8  10 -5 ) (0.15) = x 2 2.7  10 -6 = x 2 1.6  10 -3 = x 2 ( x ) 2 (0.15-X) 1.8  10 -5 = pOH = –log (1.6  10 -3 ) pOH = 2.80 pH = 11.20

60. Strength vs. Concentration n The words concentrated and dilute tell how much of an acid or base is dissolved in solution - refers to the # of moles of acid or base n Is a concentrated, weak acid possible?

61. Acid and Base Strength Acetate is a stronger base than H 2 O, LOWER on table, The stronger base “wins” the proton. HC 2 H 3 O 2 ( aq ) + H 2 OH 3 O + ( aq ) + C 2 H 3 O 2 – ( aq)

62. TITRATIONS

64. Titration n Titration is the process of adding a known amount of solution of known concentration to determine the concentration of another solution n The equivalence point is when the moles of hydrogen ions is equal to the moles of hydroxide ions (= neutralized!)

65. Titration n The concentration of acid (or base) in solution can be determined by performing a neutralization reaction –An indicator is used to show when neutralization has occurred –Often we use phenolphthalein- because it is colorless in neutral and acid; turns pink in base

66. Steps - Neutralization reaction #1. A measured volume of acid of unknown concentration is added to a flask #2. Several drops of indicator added #3. A base of known concentration is slowly added, until the indicator changes color; measure the volume

67. Neutralization n The solution of known concentration is called the standard solution –added by using a buret n Continue adding until the indicator changes color –called the “end point” of the titration

68. strong acid (HCl) v. strong base (NaOH)

69. Very little pH change during the initial 20cm 3 Very sharp change in pH over the addition of less than half a drop of NaOH pH 1 as 0.1M HCl (strong acid) strong acid (HCl) v. strong base (NaOH)

70. Steady pH change Sharp change in pH over the addition of less than half a drop of NaOH Curve levels off at pH 13 due to excess 0.1M NaOH (a strong alkali) pH 4 due to 0.1M CH 3 COOH (weak monoprotic acid) w. acid (CH 3 COOH) v s.base (NaOH)

71. strong acid (HCl) v. strong base (NaOH) Any of the indicators listed will be suitable - they all change in the ‘vertical’ portion PHENOLPHTHALEIN LITMUS METHYL ORANGE

72. weak (CH 3 COOH) v. weak (NH 3 ) Types Steady pH change pH 4 due to 0.1M CH 3 COOH (weak monoprotic acid) NO SHARP CHANGE IN pH Curve levels off at pH 10 due to excess 0.1M NH 3 (a weak alkali)

73. Polyprotic Acids? n Some compounds have more than one ionizable hydrogen to release n HNO 3 nitric acid - monoprotic n H 2 SO 4 sulfuric acid - diprotic - 2 H + n H 3 PO 4 - triprotic - 3 H + n IMPT for TITRATING BUT NOT FOR pH calculating…why?

74. Polyprotic Acids n Here are three successive ionizations of phosphoric acid: n Here are three successive ionizations of phosphoric acid: n The first dissociation constants for phosphoric acid is greater than the second, 100,000 times greater n This means nearly all the H+ ions in the solution comes from the first step of dissociation. H 3 PO 4(s) + H 2 O (l) H 3 O + (aq) + H 2 PO 4 − (aq) K a1 = 7.25×10 −3 H 2 PO 4 − (aq) + H 2 O (l) H 3 O + (aq) + HPO 4 2− (aq) K a2 = 6.31×10 −8 HPO 4 2− (aq) + H 2 O (l) H 3 O + (aq) + PO 4 3− (aq) K a3 = 3.98×10 −13

75. Other pH curves - Other pH curves - acid v. carbonate Sodium carbonate reacts with hydrochloric acid in two steps... Step 1 Na 2 CO 3 + HCl ——> NaHCO 3 + NaCl Step 2 NaHCO 3 + HCl ——> NaCl + H 2 O + CO 2 Overall Na 2 CO 3 + 2HCl ——> 2NaCl + H 2 O + CO 2 First rapid pH change around pH = 8.5 due to the formation of NaHCO 3. Can be detected using phenolphthalein Second rapid pH change around pH = 4 due to the formation of acidic CO 2. Can be detected using methyl orange. There are two sharp pH changes

76. Polyprotic acids (H 3 PO 4 ) There are three sharp pH changes Each successive addition of NaOH is the same as equal number of moles are involved. Phosphoric acid is triprotic; it reacts with sodium hydroxide in three steps... Step 1 H 3 PO 4 + NaOH ——> NaH 2 PO 4 + H 2 O Step 2 NaH 2 PO 4 + NaOH ——> Na 2 HPO 4 + H 2 O Step 3 Na 2 HPO 4 + NaOH ——> Na 3 PO 4 + H 2 O

77. Other pH curves - polyprotic acids (H 3 PO 4 ) pH of H 3 PO 4 = 1.5 Phosphoric acid is triprotic; it reacts with sodium hydroxide in three steps... Step 1 H 3 PO 4 + NaOH ——> NaH 2 PO 4 + H 2 O Step 2 NaH 2 PO 4 + NaOH ——> Na 2 HPO 4 + H 2 O Step 3 Na 2 HPO 4 + NaOH ——> Na 3 PO 4 + H 2 O

78. Other pH curves - polyprotic acids (H 3 PO 4 ) pH of NaH 2 PO 4 = 4.4 pH of H 3 PO 4 = 1.5 Phosphoric acid is triprotic; it reacts with sodium hydroxide in three steps... Step 1 H 3 PO 4 + NaOH ——> NaH 2 PO 4 + H 2 O Step 2 NaH 2 PO 4 + NaOH ——> Na 2 HPO 4 + H 2 O Step 3 Na 2 HPO 4 + NaOH ——> Na 3 PO 4 + H 2 O

79. Other pH curves - polyprotic acids (H 3 PO 4 ) pH of Na 2 HPO 4 = 9.6 pH of NaH 2 PO 4 = 4.4 pH of H 3 PO 4 = 1.5 Phosphoric acid is triprotic; it reacts with sodium hydroxide in three steps... Step 1 H 3 PO 4 + NaOH ——> NaH 2 PO 4 + H 2 O Step 2 NaH 2 PO 4 + NaOH ——> Na 2 HPO 4 + H 2 O Step 3 Na 2 HPO 4 + NaOH ——> Na 3 PO 4 + H 2 O

80. Other pH curves - polyprotic acids (H 3 PO 4 ) pH of Na 3 PO 4 = 12 pH of Na 2 HPO 4 = 9.6 pH of NaH 2 PO 4 = 4.4 pH of H 3 PO 4 = 1.5 Phosphoric acid is triprotic; it reacts with sodium hydroxide in three steps... Step 1 H 3 PO 4 + NaOH ——> NaH 2 PO 4 + H 2 O Step 2 NaH 2 PO 4 + NaOH ——> Na 2 HPO 4 + H 2 O Step 3 Na 2 HPO 4 + NaOH ——> Na 3 PO 4 + H 2 O

81. Salt Hydrolysis n A salt is an ionic compound that: –formed in a neutralization reaction –some neutral; others acidic or basic

82. Salt Hydrolysis n To see if the resulting salt is acidic or basic, check the by dissociating the salt NaCl, a neutral salt (NH 4 ) 2 SO 4, acidic salt CH 3 COOK, basic salt

83. Adding a “salt” to water n If the salt is from a strong acid and base then the pH is 7, for example KNO 3. n A salt formed between a strong acid and a conjugate of a weak base is an acid salt, example NH 4 Cl, pH =acidic. n A salt formed between a conjugate of a weak acid and a strong base is a basic salt, for example NaCH 3 COO, and the pH will be basic.

84. Salt solutions: hydrolysis NaCl ( aq ) NH 4 Cl ( aq ) NaClO ( aq )

85. Buffers n Buffers are solutions in which the pH remains relatively constant, even when small amounts of acid or base are added –made from a pair of chemicals: a weak acid and one of it’s salts; or a weak base and one of it’s salts

87. Buffers n A buffer system is better able to resist changes in pH than pure water n It is a pair of chemicals (conjugates) –one chemical neutralizes any acid added, while the other chemical would neutralize any base

88. Buffers n Example: Ethanoic (acetic) acid and sodium ethanoate (also called sodium acetate) n The buffer capacity is the amount of acid or base that can be added before a significant change in pH

89. Buffers n The two buffers that are crucial to maintain the pH of human blood are: 1. carbonic acid (H 2 CO 3 ) & hydrogen carbonate (HCO 3 - ) 2. dihydrogen phosphate (H 2 PO 4 1- ) & monohydrogen phoshate (HPO 4 2- )

90. Aspirin (which is a type of acid) sometimes causes stomach upset; thus by adding a “buffer”, it does not cause the acid irritation. Bufferin is one brand of a buffered aspirin that is sold in stores.

91. Indicators 1. Moisten the pH indicator paper strip with a few drops of solution, by using a stirring rod. 2.Compare the color to the chart on the vial – then read the pH value.

92. Some of the many pH Indicators and their pH range

93. How Do We Measure pH? –An indicator Compound that changes color in solution.