1. Acids and Bases Chapter 19 DHS Chemistry

2. Definitions Bronsted – Lowry Arrhenius

3. Bronsted-Lowry Acid Proton (hydrogen ion) donor. Proton = H + HCl + H 2 0  Cl - + H 3 0 + donoracceptor Acid

4. Bronsted-Lowry Base Proton (hydrogen ion) acceptor. Proton = H + HCl + +NH 3  NH 4 + + Cl - donoracceptor Base

5. Arrhenius Acid  Produces H + /H 3 O + ions  Ex. HCl + H 2 0  H 3 0 + + Cl - acid

6. Arrhenius Base  Produces H + /H 3 O + ion  Ex. NH 3 + H 2 O  NH 4 + + OH - Base

7. Properties

8. Acid Properties Typically acids start with ________ in their formulas. Like... HCl and H 2 SO 4 hydrogen

9. Acid Properties ____ taste turn blue litmus paper red (pink) Sour

10. Acid Properties react with bases to produce salt and water (neutralization)

11. electrolytes (may be strong or weak) © Prentice Hall Properties - Acids

12. Acid Properties contain hydrogen and most react with active metals to produce hydrogen gas

13. Base Properties Most bases (not all) have OH in their chemical formulas. Like... NaOH and KOH

14. Base Properties Bitter taste Antacids help neutralizes stomach acid. They are bases.

15. Base Properties Feel slippery

16. turn red litmus paper blue react with an acid to produce salt and water (neutralization) electrolytes (may be strong or weak) Properties - Bases

17. Litmus Paper Aciase Relue B D

18. Common Ones

19. A. Common Acids AcidFormulaMajor Uses sulfuric acid H 2 SO 4 Car batteries, production of metals phosphoric acid H 3 PO 4 Found in soft drinks Nitric acid HNO 3 Production of explosives hydrochloric acid HCl Cleaning of metals Most commonly produced in the world

20. Common Bases BaseFormulaMajor Uses ammonia NH 3 Refrigerant, household cleaners sodium hydroxide NaOH Drain cleaner potassium hydroxide KOH To increase pH of acidic soils

21. Strength/Electrolytic Behavior of Acids and Bases

22. Strong Acids Ionize almost completely (almost 100%) Strong electrolytes (when dissolved) Strong electrolyte = bright light bulb

23. Examples of Strong Acids HCl, HBr, HI, HNO 3, H 2 SO 4 These acids ionize almost completely in water. *memorize those that are highlighted

24. Weak Acids Ionize partially Weak electrolytes (when dissolved) Weak electrolyte = dim light bulb

25. Examples of Weak Acids HF, HCN, HC 2 H 3 O 2, H 2 CO 3 These acids ionize almost completely in water. *memorize those that are highlighted

26. Strong Bases Ionize almost completely (almost 100%) Strong electrolytes (when dissolved) Strong electrolyte = bright light bulb

27. Strong Bases NaOH, KOH, Ca(OH) 2, LiOH, CsOH, RbOH, Sr(OH) 2 These bases ionize almost completely in water. *memorize those that are highlighted

28. Weak Bases Ionize partially Weak electrolytes (when dissolved) Weak electrolyte = dim light bulb

29. Weak bases CH 3 N 2, NH 3, NaCN, Mg(OH) 2 These bases ionize almost completely in water. *memorize those that are highlighted

30. Some substances can act as both an acid and a base – these substances are considered amphoteric. EX: HCl + H 2 0  H 3 0 + + Cl -  water as a base acid base NH 3 + H 2 0  NH 4 + + OH -  water as an acid base acid H 2 0 + H 2 0  H 3 0 + + OH -  water as an acid pure water has H 2 O, H 3 O +, OH - & base

31. Polyprotic acids Acids that can contain multiple hydrogens to donate. Match the terms with the probable acid monoproticH 2 SO 4 diproticH 3 PO 4 polyproticHCl

32. Acids that contain 1 ionizable hydrogen are monoprotic acids. Ex. HCl, HBr Polyprotic acids

33. Acids that contain 2 ionizable hydrogens are diprotic acids. Ex. H 2 SO 4 Polyprotic acids

34. Acids that contain 3 ionizable hydrogens are triprotic/polyprotic acids. Ex. H 3 PO 4 Polyprotic acids

35. II. Acid and Base Reactions

36. A. Reactions Involving Acids 1. Acids with Active Metals –Acids will react with active metals (metals more reactive than hydrogen) to release hydrogen gas –General Form: H + HX  MX + H 2 EX: Mg(s) + 2HNO 3 (aq)  Mg(NO 3 ) 2 (aq) + H 2 (g) Active metalacid salt/ionic compound hydrogen gas

37. Recall from your lab Questions/Analysis 2. Write a balanced reaction for the reaction between the zinc and hydrochloric acid. Be sure to include states of all substances. Zn(s) + 2HCl(aq)  ZnCl 2 (aq) + H 2 (g)

38. 2. Nonmetallic Oxide and Water acids can be produced from the reaction of a nonmetallic oxide (ex. SO 3 ) and water General Form: NMO + water  HX EX: SO 3 (g) + H 2 O(l)  H 2 SO 4 (aq) acid anhydride water acid (a substance that produces an acid when combined with water is called an acid anhydride.)

39. 3. Acids with Carbonates acids will react with carbonates (ex. NaCO 3 ) to produce a salt and water and carbon dioxide gas EX: Na 2 CO 3 (s) + 2HCl(aq)  2NaCl(aq) + CO 2 (g)+ H 2 0(l) carbonateacid water Carbon dioxide Salt/ ionic compound

40. Recall from your lab Questions/Analysis 3. Write a balanced reaction for the reaction between calcium carbonate and hydrochloric acid. Include states of all substances. CaCO 3 (s) + 2HCl(aq)  CaCl 2 (aq) + CO 2 (g) + H 2 O(l)

41. B. Reactions Involving Bases

42. 1. Metallic Oxide and Water a metallic oxide (ex. Na 2 O) and water will combine to produce a base General Form: MO + H 2 O  Base metal oxide

43. 1. Metallic Oxide and Water General Form: MO + H 2 O  Base metal oxide EX: Na 2 O(s) + H 2 O(l)  2NaOH(aq) (a substance that produces a base when combined with water is called a basic anhydride.) Metal oxide water base (basic anhydride)

44. C. Neutralization Reactions The reaction of an acid with a base produces water and a class of compounds called salts. HA + BOH  B A + H 2 O ex. HCl(aq) + NaOH(aq)  NaCl (aq) + H 2 O(l) acid basesaltwater

45. Ex. The salt is highlighted in each case. HCl + NaOH → NaCl + H 2 O HC 2 H 3 O 2 + KOH  KC 2 H 3 O 2 + H 2 O H 2 SO 4 + 2 NH 4 OH  (NH 4 ) 2 SO 4 + H 2 O HC 2 H 3 O 2 + NH 4 OH  NH 4 C 2 H 3 O 2 + H 2 O Strong acid plus strong base Weak acid plus strong base Strong acid plus weak base Weak acid plus weak base

46. Recall from your lab Questions/Analysis 4. Write a balanced reaction for the reaction between hydrochloric acid and sodium hydroxide. Include states of all substances. HCl(aq) + NaOH(aq)  NaCl(aq) + H 2 O(l)

47. Salts are compounds consisting of an anion from an acid and a cation from a base. If you mix a solution of a strong acid with a strong base, a neutral solution results.

48. Reactions in which an acid and a base react in aqueous solution are called neutralization reactions. All neutralization reactions are double-replacement reactions

49. Practice 1.ID the type, complete, and balance these reactions involving acids and bases. a) hydrochloric acid + Al b) sulfuric acid + zinc 6HCl(aq) + 2Al(s)  2AlCl 3 (aq) + 3H 2 (g) H 2 SO 4 (aq) + Zn(s)  ZnSO 4 (aq) + H 2 (g) Acid + active metal

50. Practice 1.ID the type, complete, and balance these reactions involving acids and bases. c) nitric acid + potassium hydroxide d) calcium oxide + water Acid + base (neutralization) Metal oxide + water HNO 3 (aq) + KOH(aq)  KNO 3 (aq) + H 2 O(l) CaO(s) + H 2 O(l)  Ca(OH) 2 (aq)

51. III. Strength of Acids and Bases

52. A. Ionization acids will form ions (electrolytes) when dissolved in water in a process called ionization. EX: HCl(aq) + H 2 O(l)  H 3 O + (aq) + Cl - (aq)

53. Bases Bases also will form ions (electrolytes) when dissolved in a process called dissociation. EX:NaOH(s)  Na + (aq) + OH - (aq) H2OH2O

54. One way to measure the strength of an acid or a base is to measure how much of the original molecule remains after it has been added to water. If little remains, the acid or base is strong. If a lot of the original molecule remains, the acid or base is weak.

55. For example, when HCl is added to water, nearly all of the HCl molecules are converted to ions. HCl + H 2 O  H 3 O + +Cl - Strong Acid

56. When acetic acid is added to water, most of the acetic acid molecules remain as molecules and only a small portion of the molecules are converted to ions. Weak Acid

57. Do not confuse the terms strong and weak with the terms concentrated and dilute. Strength refers to what % of the original molecules convert to ions in water. Concentrated or dilute refer to how many total moles there are in water.

58. Strong acids (lots of H + ions): HCl, HBr, HI, HNO 3, H 2 SO 4 Highlight the acids that will ionize almost completely in water. Weak acids (some H + ions): HF, HCN, HC 2 H 3 O 2, H 2 CO 3 Strong bases (lots of OH - ions): NaOH, KOH, Ca(OH) 2, LiOH, CsOH, RbOH, Sr(OH) 2 Highlight the bases that are very soluble in water. Weak bases (some OH - ions): CH 3 N 2, NH 3, NaCN, Mg(OH) 2

59. B. pH Scale

60. Hydronium vs Hydroxide HydroniumH 3 O + –A hydrogen ion in water –H + + H 2 O  H 3 O + –H + and H 3 O + used interchangeably –For acids HydroxideOH - –For bases

61. 1. Background Any aqueous solution contains both hydronium ions and hydroxide ions. This stems from the fact that water will ionize to a very small amount:

62. Note: for pure water, the number of hydronium ions is equal to the number of hydroxide ions. H 2 O(l) + H 2 O(l)  H 3 O + (aq) + OH - (aq)

63. when an acid is added to water, the number of hydronium ions increases HCl(aq) + H 2 O(l)  H 3 O + (aq) + Cl - (aq)

64. when a base is added to water, the number of hydroxide ions increases H 2 O(l) NaOH(s)  Na + (aq) + OH - (aq)

65. it is the ratio of hydronium ions to hydroxide ions that determines whether a solution will be an acid, a base, or neutral

66. acid: contains more hydronium ions than hydroxide ions base: contains more hydroxide ions than hydronium ions neutral: the # of hydronium ions is equal to the # of hydroxide ions

68. 2. pH Scale a measure of the number of hydronium or hydroxide ions is the pH scale it is based on the concentration of hydrogen and hydroxide ions in solution pH is defined as the negative logarithm of the hydrogen ion concentration

69. [ ] = concentration (molarity) notice pH is based on a log (base 10) scale pH = -log([H + ]) pH= -log([H 3 O + ])

70. pH Scale the typical pH scale runs from 0 to 14.

72. a pH of 7 is considered neutral which means that the concentration of hydrogen ions and the concentration of hydroxide ions are equal

73. as you go down on the pH scale (< 7), solutions are considered acidic solutions with pH’s greater than 7 are considered basic

74. Summary of the pH scale pHCategoryConcentration of ions < 3strong acidmany H 3 O + 3-7Weak acidH 3 O + > OH - 7neutralH 3 O + = OH - 7-11Weak baseH 3 O + < OH - > 11strong basemany OH -

75. Practice 1. Determine whether the following are a strong acid, a weak acid, a strong base, a weak base, neutral solution a. pH = 2.5 b. lots of hydroxide ions, hardly any hydronium ions c. little more hydroxide ions than hydronium ions Strong acid strong Base Weak base

76. There is also something called pOH, which is a measure of the concentration of hydroxide ions. pOH = -log ([OH - ]

77. pH and pOH are related by the following: 14 = pH + pOH

78. Again, for strong, single hydroxide bases, [OH-] = molarity of the base. Ex: What is the pOH of KOH if the pH is 14? pH + pOH = 14 14 + pOH = 14 pOH = 0

80. Practice 1. Determine whether the following are a strong acid, a weak acid, a strong base, a weak base, neutral solution pOH = 2.5 12.0 M NaOH pH = 11.5 strong base pOH = 1.07 pH = 12.9 strong base

81. IV. Titrations

82. A. Titrations The concentration of an acid (or base) in solution can be determined by performing a neutralization reaction. acid + base  salt + water

83. An indicator is used to show when neutralization has occurred. An indicator is a substance that forms different colors in different pH solutions. Phenolphthalein is a common indicator used in acid- base titrations. It will change from colorless in acidic environments to pink in basic environments.

84. Acid Base Indicators (pH sensitive) Litmus paper

85. Acid Base Indicators (pH sensitive) pH paper

86. Acid Base Indicators (pH sensitive) Universal Indicator

87. Acid Base Indicators (pH sensitive) Phenolphthalein

88. Acid Base Indicators (pH sensitive) Red Cabbage Juice

89. The solution of known concentration is called the standard solution. The standard solution is added using a buret. The process of adding a known amount of solution of known concentration to determine the concentration of another solution is called titration. The point at which the indicator changes color is the end point of the titration.

91. Informal titration This can also be done less “formally” using any volume measures. The results won’t be as accurate, but it gets you close. Example: (drops, substitute for mL)

92. Steps in a formal titration 1.A measured volume of a solution of unknown concentration (acid or base) is added to an Erlenmeyer flask. 2.A solution of known molarity (acid or base) is added to a buret. Known M V measured Known V Unknown M

93. 3. Several drops of an indicator are added to the unknown solution 4. Measured volumes of a solution of known molarity (acid or base) are mixed into unknown solution until the indicator just barely changes color to pink. Known M V measured Known V Unknown M

94. The end point Phenolphthalein indicator Clear = Acid Pink = Base

95. How to read a buret

96. This can also be done less “formally” using any volume measures. The results won’t be as accurate, but it gets you close. Example: (drops, substitute for mL)

97. B. Solving Titration Problems Remember, in order for the solution to be neutral, V = volume M = molarity (M) A = acid B = base If the ratio of H + to OH - is 1:1, then M A V A = M B V B (similar to dilutions)

98. 1) It takes 26.23 mL of a 1.008 M NaOH solution to neutralize 35.28 mL of a monoprotic acid solution. What is its molarity? M A = ? MM B = 1.008 M V A = 35.28 mLV B = 26.23 mL M A V A = M B V B M A (35.28 mL) = (1.008 M) (26.23 mL) M B = 0.749 M

99. 1.008 M Base 35.28 mL of ?? M acid Keep adding base until there is a color change. Volume of base added = 26.23 mL

100. 2) If 15.50 mL of Ca(OH) 2 solution were neutralized with 23.40 mL of 0.533 M H 2 SO 4, what is the concentration of the Ca(OH) 2 ? M A = 0.533 MM B = ? M V A = 23.40 mLV B = 15.50 mL M A V A = M B V B (0.533 M)(23.40 mL) = M B (15.50 mL) 0.805 M = M B

101. Ex 3: 25.00 mL of 0.720 M nitric acid (HNO 3 ) is used to completely neutralize a 1.0 M NaOH solution. What volume of NaOH is present? M A = 0.720 MM B = 1.0M V A = 25mLV B = ?? M A V A = M B V B (0.720 M)(25 mL) = (1.0 M) V B 18.00 mL = V B

102. Practice Box Answers 1)0.385 M H 3 PO 4 2)5.76 x 10 -3 M NaOH 3)0.840 M NaOH 4)55.6 mL H 2 CO 3

103. 1.What is the molarity of phosphoric acid if 15.0 mL of the solution is completely neutralized by 38.5 mL of 0.150 M Al(OH) 3 ? H 3 PO 4 Al(OH) 3 M A = ??M B = 0.150M V A = 15mLV B = 38.5mL M A V A = M B V B M A (15mL) = (0.150M)(38.5mL) M A = 0.385M H 3 PO 4

104. 2. It takes 26.23 mL of a 0.01 M NaOH solution to neutralize a 45.56 mL of a HCl solution. What is the concentration of the acid? HCl NaOH M A = ?M B = 0.01M V A = 45.56 mLV B = 26.23 mL M A V A = M B V B M A (45.56 mL) = (0.01 M)(26.23 mL) M A = 5.76 x 10 -3 M HCl

105. 3. What is the molarity of potassium hydroxide if 20.0 mL of the solution is neutralized by 28.0 mL of 0.60 M HCl? HCl NaOH M A = 0.60MM B = ?? V A = 28.0mLV B = 20.0mL M A V A = M B V B (0.60M)(28.0mL) = M B (20.0mL) 0.840M NaOH = M B

106. 4. How many mL of 0.45 M HCl must be added to 25.0 mL of 1.00 M KOH to make a neutral solution? HCl KOH M A = 0.45 MM B = 1.00 M V A = ??V B = 25.0 mL M A V A = M B V B M A (0.45 mL) = (1.00 M)(25.0 mL) M A = 55.6 mL HCl