1. Chapter 16 “Acids, Bases, and Salts” Adapted & Modified from: Pre-AP Chemistry Charles Page High School Stephen L. Cotton

2. Acid-Base Theories OBJECTIVES: –Define the properties of acids and bases.

3. Acid-Base Theories OBJECTIVES: –Compare and contrast acids and bases as defined by the theories of: a) Arrhenius, b) Brønsted-Lowry, and c) Lewis.

4. Properties of Acids  taste sour (don’t try this at home).  conduct electricity.  Can be strong or weak electrolytes in aqueous solution  React with metals to form H 2 gas.  Change the color of indicators (for example: blue litmus turns to red).  React with bases (metallic hydroxides) to form water and a salt.

5. Properties of Acids  pH of less than 7 (more on this concept of pH in a later lesson)  react with carbonates and bicarbonates to produce a salt, water, and carbon dioxide gas  How do you know if a chemical is an acid? –It usually starts with Hydrogen. –HCl, H 2 SO 4, HNO 3, etc. (but not water!)

6. Naming Acids  Acids are ionic compounds (cation + anion)  H + is the cation  All acids start with H  When acids dissolve into ion pairs, they produce and H +  HCl → H + + Cl -  HNO 3 → H + + NO 3 -

7. Naming Acids When the anion ends in —ide, the name begins with the prefix hydro-, the stem anion has the suffix —ic and is followed by the word acid. HCl = hydrochloric acid. When the anion name ends in —ite, the acid name ending changes to —ous, followed by acid. HClO 2 = Chlorous acid. When the anion name ends in —ate, the acid name ending changes to —ic, followed by cid. HClO 3 = Chloric acid

8. Acids Affect Indicators, by changing their color Blue litmus paper turns red in contact with an acid (and red paper stays red).

9. Acids have a pH less than 7 (p.515)

10. Acids React with Active Metals Acids react with active metals to form salts and hydrogen gas: HCl (aq) + Mg (s) → MgCl 2(aq) + H 2(g) This is a single-replacement reaction

11. 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

12. Acids Neutralize Bases HCl + NaOH → NaCl + H 2 O - Neutralization reactions ALWAYS produce a salt (which is an ionic compound) and water. -Of course, it takes the right proportion of acid and base to produce a neutral salt

13. Sulfuric Acid = H 2 SO 4  Highest volume production of any chemical in the U.S. (approximately 60 billion pounds/year)  Used in the production of paper  Used in production of fertilizers  Used in petroleum refining; auto batteries

14. Nitric Acid = HNO 3  Used in the production of fertilizers  Used in the production of explosives  Nitric acid is a volatile acid – its reactive components evaporate easily  Stains proteins yellow (including skin!)

15. Hydrochloric Acid = HCl  Used in the “pickling” of steel  Used to purify magnesium from sea water  Part of gastric juice, it aids in the digestion of proteins  Sold commercially as Muriatic acid

16. Phosphoric Acid = H 3 PO 4  A flavoring agent in sodas (adds “tart”)  Used in the manufacture of detergents  Used in the manufacture of fertilizers  Not a common laboratory reagent

17. Acetic Acid = HC 2 H 3 O 2 (also called Ethanoic Acid, CH 3 COOH)  Used in the manufacture of plastics  Used in making pharmaceuticals  Acetic acid is the acid that is present in household vinegar

18. Properties of Bases (metallic hydroxides)  React with acids to form water and a salt.  Taste bitter.  Feel slippery (don’t try this either).  Can be strong or weak electrolytes in aqueous solution  Change the color of indicators (red litmus turns blue).

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

20. Bases Affect Indicators Red litmus paper turns blue in contact with a base (and blue paper stays blue). Phenolphthalein turns purple in a base.

21. Bases have a pH greater than 7 (p.515)

22. Bases Neutralize Acids Milk of Magnesia contains magnesium hydroxide, Mg(OH) 2, which neutralizes stomach acid, HCl. 2 HCl + Mg(OH) 2 MgCl 2 + 2 H 2 O Magnesium salts can cause diarrhea (thus they are used as a laxative) and may also cause kidney stones.

23. Acid-Base Theories

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

25. Polyprotic Acids?  Some compounds have more than one ionizable hydrogen to release  HNO 3 nitric acid - monoprotic  H 2 SO 4 sulfuric acid - diprotic - 2 H +  H 3 PO 4 phosphoric acid - triprotic - 3 H +  Having more than one ionizable hydrogen does not mean stronger!

26. 2. Brønsted-Lowry - 1923  Acid is hydrogen-ion donor (H + or proton)  Base is hydrogen-ion acceptor.  Acids and bases always come in pairs.  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)  Water is a base; makes hydronium ion.

27. Acids and bases come in pairs  A “conjugate base” is the remainder of the original acid, after it donates it’s hydrogen ion  A “conjugate acid” is the particle formed when the original base gains a hydrogen ion  Thus, a conjugate acid-base pair is related by the loss or gain of a single hydrogen ion.

28. Acids and bases come in pairs  General equation is: HA (aq) + H 2 O (l) ↔ H 3 O + (aq) + A - (aq) Acid + Base ↔ Conjugate acid + Conjugate base Acid + Base ↔ Conjugate acid + Conjugate base  NH 3 + H 2 O ↔ NH 4 1+ + OH 1- base acid c.a. c.b. base acid c.a. c.b.  HCl + H 2 O ↔ H 3 O 1+ + Cl 1- acid base c.a. c.b. acid base c.a. c.b.  Amphoteric – a substance that can act as both an acid and base- as water shows

29. 3. Lewis Acids and Bases  Gilbert Lewis focused on the donation or acceptance of a pair of electrons during a reaction  Lewis Acid - electron pair acceptor  Lewis Base - electron pair donor  Most general of all 3 definitions; acids don’t even need hydrogen

30. Hydrogen Ions and Acidity OBJECTIVES: –Describe how [H 1+ ] and [OH 1- ] are related in an aqueous solution.

31. Hydrogen Ions and Acidity OBJECTIVES: –Classify a solution as neutral, acidic, or basic given the hydrogen-ion or hydroxide-ion concentration.

32. Hydrogen Ions and Acidity OBJECTIVES: –Convert hydrogen-ion concentrations into pH values and hydroxide-ion concentrations into pH values.

33. Hydrogen Ions and Acidity OBJECTIVES: –Describe the purpose of an acid-base indicator.

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

35. Acid-Base Characteristics  Solution Type Relationship  Acid [H+ ] > [OH-]  Neutral [H+ ] = [OH-]  Base [H+ ] < [OH-]

36. The pH concept – from 0 to 14  pH = pouvoir hydrogene (Fr.) “hydrogen power”  definition: pH = -log[H + ]  in neutral pH = -log(1 x 10 -7 ) = 7  in acidic solution [H + ] > 10 -7  pH < -log(10 -7 ) –pH < 7 (from 0 to 7 is the acid range) –in base, pH > 7 (7 to 14 is base range)

37. The pH concept – from 0 to 14  pH is the negative logarithm of the hydrogen ion concentration.  pH = -log [H+]  Remember that sometimes H3O+ is written, so pH = -log [H3O +]

38. pH and pOH  pH = -log [H+]  pOH = -log [OH-]  pH + pOH = 14  pH< 7 Acid  pH> 7 Base  pH = 7 Neutral

39. pH and pOH pH = -log [H + ] [H+]pHAcid or Base 6.5 x 10-11 3.3 x 10-3

40. pH and pOH pH = -log [H + ] [H+]pHAcid or Base 6.5 x 10-1110.18 3.3 x 10-32.48

41. pH and pOH pH = -log [H + ] [H+]pHAcid or Base 6.5 x 10-1110.18Base 3.3 x 10-32.48Acid

42. pH and pOH pOH = -log [OH - ] [OH-]pOHpHAcid or Base 3.4 x 10-6 8.3 x 10-8

43. pH and pOH pOH = -log [OH - ] [OH-]pOHpHAcid or Base 3.4 x 10-6 5.47 8.3 x 10-8 7.08

44. pH and pOH pOH = -log [OH - ] [OH-]pOHpHAcid or Base 3.4 x 10-6 5.478.53 8.3 x 10-8 7.086.92

45. pH and pOH pOH = -log [OH - ] [OH-]pOHpHAcid or Base 3.4 x 10-6 5.478.53Base 8.3 x 10-8 7.086.92Acid

46. Complete the Table pHAcid or Base pOH[H+][OH-} 3.80 13.50 6.67

47. Complete the Table pHAcid or Base pOH[H+][OH-} 3.80A10.201.6E-46.3E-11 13.50B.503.2E-143.2E-1 6.67A7.342.2E-74.6E-8

48. Calculating pOH  pOH = -log [OH - ]  [H + ] x [OH - ] = 1 x 10 -14 M 2  pH + pOH = 14  Thus, a solution with a pOH less than 7 is basic; with a pOH greater than 7 is an acid  Not greatly used like pH is.

52. Measuring pH  Why measure pH? Everyday solutions we use - everything from swimming pools, soil conditions for plants, medical diagnosis, soaps and shampoos, etc.  Sometimes we can use indicators, other times we might need a pH meter

53. How to measure pH with wide-range paper 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.

54. Some of the many pH Indicators and their pH range ( Fig. 16.4, p.520)

55. Acid-Base Indicators  Although useful, there are limitations to indicators: –usually given for a certain temperature (25 o C), thus may change at different temperatures –what if the solution already has a color, like paint? – the ability of the human eye to distinguish colors is limited

56. Acid-Base Indicators  A pH meter may give more definitive results –some are large, others portable –works by measuring the voltage between two electrodes; typically accurate to within 0.01 pH unit of the true pH

57. Strengths of Acids and Bases OBJECTIVES: –Define strong acids and weak acids.

58. Strengths of Acids and Bases OBJECTIVES: –Describe how an acid’s strength is related to the value of its acid dissociation constant.

59. Strength  Acids and Bases are classified acording to the degree to which they ionize in water: –Strong are completely ionized in aqueous solution; this means they ionize 100 % –Weak ionize only slightly in aqueous solution  Strength is very different from Concentration

60. Strength  Strong – means it forms many ions when dissolved (100 % ionization)  Mg(OH) 2 is a strong base- it falls completely apart (nearly 100% when dissolved). –But, not much dissolves- so it is not concentrated

61. Strong and Weak Acids/Bases Generally divide acids and bases into STRONG or WEAK ones. STRONG ACID: HNO 3 (aq) + H 2 O (l) ---> H 3 O + (aq) + NO 3 - (aq) HNO 3 is about 100% dissociated in water.

62. Weak acids are much less than 100% ionized in water. One of the best known is acetic acid = CH 3 CO 2 H Strong and Weak Acids/Bases

63. Strong Base: 100% dissociated in water. NaOH (aq) ---> Na + (aq) + OH - (aq) NaOH (aq) ---> Na + (aq) + OH - (aq) Strong and Weak Acids/Bases Other common strong bases include KOH and Ca(OH) 2. CaO (lime) + H 2 O --> Ca(OH) 2 (slaked lime) Ca(OH) 2 (slaked lime) CaO

64. Weak base: less than 100% ionized in water One of the best known weak bases is ammonia NH 3 (aq) + H 2 O (l)  NH 4 + (aq) + OH - (aq) Strong and Weak Acids/Bases

65. pH testing There are several ways to test pH –Blue litmus paper (red = acid) –Red litmus paper (blue = basic) –pH paper (multi-colored) –pH meter (7 is neutral, 7 base) –Universal indicator (multi-colored) –Indicators like phenolphthalein –Natural indicators like red cabbage, radishes

66. pH indicators Indicators are dyes that can be added that will change color in the presence of an acid or base. Some indicators only work in a specific range of pH Once the drops are added, the sample is ruined Some dyes are natural, like radish skin or red cabbage

67. Strength vs. Concentration  The words concentrated and dilute tell how much of an acid or base is dissolved in solution - refers to the number of moles of acid or base in a given volume  The words strong and weak refer to the extent of ionization of an acid or base  Is a concentrated, weak acid possible?

68. Neutralization Reactions OBJECTIVES: –Define the products of an acid-base reaction.

69. Neutralization Reactions OBJECTIVES: –Explain how acid-base titration is used to calculate the concentration of an acid or a base.

70. Neutralization Reactions OBJECTIVES: –Explain the concept of equivalence in neutralization reactions.

71. Neutralization Reactions OBJECTIVES: –Describe the relationship between equivalence point and the end point of a titration.

72. Acid-Base Reactions  Acid + Base  Water + Salt  Properties related to every day: –antacids depend on neutralization –farmers adjust the soil pH –formation of cave stalactites –human body kidney stones from insoluble salts

73. Acid-Base Reactions  Neutralization Reaction - a reaction in which an acid and a base react in an aqueous solution to produce a salt and water: HCl (aq) + NaOH (aq)  NaCl (aq) + H 2 O (l) H 2 SO 4(aq) + 2KOH (aq)  K 2 SO 4(aq) + 2 H 2 O (l)

74. Titration  Titration is the process of adding a known amount of solution of known concentration to determine the concentration of another solution  Remember? - a balanced equation is a mole ratio  The equivalence point is when the moles of hydrogen ions is equal to the moles of hydroxide ions (= neutralized!)

76. Titration  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

77. 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

78. Neutralization  The solution of known concentration is called the standard solution –added by using a buret  Continue adding until the indicator changes color –called the “end point” of the titration –Sample Problem 19.7, page 616

79. Setup for titrating an acid with a base

80. TitrationTitration 1. Add solution from the buret. 2. Reagent (base) reacts with compound (acid) in solution in the flask. 3. Indicator shows when exact stoichiometric reaction has occurred. (Acid = Base) This is called NEUTRALIZATION. This is called NEUTRALIZATION.

81. PROBLEM: You have 50.0 mL of 3.0 M NaOH and you want 0.50 M NaOH. What do you do? Add water to the 3.0 M solution to lower its concentration to 0.50 M Dilute the solution!

82. PROBLEM: You have 50.0 mL of 3.0 M NaOH and you want 0.50 M NaOH. What do you do? But how much water do we add? 3.0 M NaOH 0.50 M NaOH H 2 O ConcentratedDilute

83. PROBLEM: You have 50.0 mL of 3.0 M NaOH and you want 0.50 M NaOH. What do you do? How much water is added? The important point is that ---> moles of NaOH in ORIGINAL solution = moles of NaOH in FINAL solution moles of NaOH in FINAL solution

84. PROBLEM: You have 50.0 mL of 3.0 M NaOH and you want 0.50 M NaOH. What do you do? Amount of NaOH in original solution = M V = M V = (3.0 mol/L)(0.050 L) = 0.15 mol NaOH Amount of NaOH in final solution must also = 0.15 mol NaOH Volume of final solution = (0.15 mol NaOH) / (0.50 M) = 0.30 L or 300 mL

85. PROBLEM: You have 50.0 mL of 3.0 M NaOH and you want 0.50 M NaOH. What do you do? Conclusion: add 250 mL of water to 50.0 mL of 3.0 M NaOH to make 300 mL of 0.50 M NaOH. 3.0 M NaOH 0.50 M NaOH H 2 O ConcentratedDilute

86. A shortcut A shortcut M 1 V 1 = M 2 V 2 Preparing Solutions by Dilution

87. You try this dilution problem You have a stock bottle of hydrochloric acid, which is 12.1 M. You need 400 mL of 0.10 M HCl. How much of the acid and how much water will you need?

88. 35.62 mL of NaOH is neutralized with 25.2 mL of 0.0998 M HCl by titration to an equivalence point. What is the concentration of the NaOH? LAB PROBLEM #1: Standardize a solution of NaOH — i.e., accurately determine its concentration.