Chapter Ten

1. Fundamentals of General, Organic, and Biological Chemistry 5th Edition Chapter Ten Acids and Bases James E. Mayhugh Oklahoma City University 2007 Prentice Hall, Inc.

2. Outline • 10.1 Acids and Bases in Aqueous Solution • 10.2 Some Common Acids and Bases • 10.3 The Brønsted–Lowry Definition of Acids and Bases • 10.4 Water as Both an Acid and a Base • 10.5 Some Common Acid–Base Reactions • 10.6 Acid and Base Strength • 10.7 Acid Dissociation Constants • 10.8 Dissociation of Water • 10.9 Measuring Acidity in Aqueous Solution: pH • 10.10 Working with pH • 10.11 Laboratory Determination of Acidity • 10.12 Buffer Solutions • 10.13 Buffers in the Body • 10.14 Acid and Base Equivalents • 10.15 Titration • 10.16 Acidity and Basicity of Salt Solutions Chapter Ten

3. 10.1 Acids and Bases in Aqueous Solution • An acid is a substance that produces hydrogen ions, H+, when dissolved in water. (Arrhenius definition) • Hydronium ion: The H3O+ ion formed when an acid reacts with water. Chapter Ten

4. A base is a substance that produces hydroxide ions, OH-, when dissolved in water. (Arrhenius definition) • Bases can be metal hydroxides that release OH- ions when they dissolve in water or compounds that undergo reactions with water that produce OH- ions. Chapter Ten

5. 10.2 Some Common Acids and Bases • Sulfuric acid, H2SO4, is manufactured in greater quantity than any other industrial chemical. It is the acid use in the petroleum and pharmaceutical industry’s, and found in automobile batteries. • Hydrochloric acid, HCl, is “stomach acid” in the digestive systems of most mammals. • Phosphoric acid, H3PO4, is used to manufacture phosphate fertilizers. The tart taste of many soft drinks is due to the presence of phosphoric acid. • Nitric acid, HNO3, is a strong oxidizing agent that is used for many purposes. • Acetic acid, CH3CO2H, is the primary organic constituent of vinegar. Chapter Ten

6. Sodium hydroxide, NaOH, or lye, is used in the production of aluminum, glass, and soap. Drain cleaners often contain NaOH because it reacts with the fats and proteins found in grease and hair. • Calcium hydroxide, Ca(OH)2 ,or slaked lime, is made industrially by treating lime (CaO) with water. It is used in mortars and cements. An aqueous solution is often called limewater. • Magnesium hydroxide, Mg(OH)2,or milk of magnesia, is an additive in foods, toothpaste, and many over-the-counter medications. Many antacids contain magnesium hydroxide, can also be used in mortars and cements. • Ammonia, NH3, is used primarily as a fertilizer. A dilute solution of ammonia is frequently used around the house as a glass cleaner. Chapter Ten

7. Soap is manufactured by the reaction of vegetable oils and animal fats with the bases NaOH and KOH. Sodium hydroxide is also called caustic soda or lye and is the most commonly used of all bases.

8. 10.3 The Brønsted-Lowry Definition of Acids and Bases • ABrønsted–Lowry acid can donate H+ ions. • Monoprotic acids can donate 1 H+ ion, diprotic acids can donate 2 H+ ions, and triprotic acids can donate 3 H+ ions. Chapter Ten

9. Acetic acid is an organic acid, which donates a hydrogen ion in solution. Just one of the hydrogen's on acetic acid is acidic. The hydrogen attached to the electronegative oxygen atom is the acidic hydrogen on acetic acid.

10. A Brønsted–Lowry base acceptsH+ ions. • Putting the acid and base definitions together, an acid–base reaction is one in which a proton is transferred. The reaction need not occur in water. Chapter Ten

11. Which of the following would you expect to be a BrØnsted-Lowry acid or BrØnsted-Lowry base? • Fe3+ • H2CrO4 • NH3 • NO3–

12. Which of the following would you expect to be a BrØnsted-Lowry acid? • Fe3+ (referred to as a Lewis Acid) • H2CrO4 acid: 2H+ + CrO42- • NH3 base: NH4+ • NO3– base: HNO3

13. Which of the following would you expect to be a BrØnsted-Lowry base? • HNO2 • NH4+ • Ni2+ • PO43–

14. Which of the following would you expect to be a BrØnsted-Lowry base? • HNO2 • NH4+ • Ni2+ • PO43–

15. Conjugate acid–base pair: Two substances whose formulas differ by only a hydrogen ion, H+. • Conjugate base: The substance formed by loss of H+ from an acid. • Conjugate acid: The substance formed by addition of H+ to a base. Chapter Ten

17. Indentify the conjugate acid/base pairs HCO3-(aq) + H2O(l) ⇄ CO32-(aq) + H3O+(aq) HF(aq) + HPO42-(aq) ⇄ F-(aq) + H2PO4-(aq) Chapter Ten

18. Indentify the conjugate acid/base pairs acid base HCO3-(aq) + H2O(l) ⇄ CO32-(aq) + H3O+(aq) baseacid acid base HF(aq) + HPO42-(aq) ⇄ F-(aq) + H2PO4-(aq) base acid Chapter Ten

19. 10.4 Water as Both an Acid and a Base • Water is neither an acid nor a base in the Arrhenius sense because it does not produce appreciable concentrations of either H+ or OH-. In the Brønsted–Lowry sense water is bothan acid and a base. • In its reaction with ammonia, water donates H+ to ammonia to form the ammonium ion. Chapter Ten

20. When water reacts as a Brønsted–Lowry base, it acceptsH+from an acid like HCl. • Substances like water, which can react as either an acid or a base depending on the circumstances, are said to be amphoteric. Chapter Ten

21. 10.5 Some Common Acid-Base Reactions • Acids react with metal hydroxides to yield water and a salt in a neutralization reaction. The H+ ions and OH- ions are used up in the formation of water. HA(aq) + MOH(aq)H2O(l) +MA(aq) • Carbonate and bicarbonate ions react with acid by accepting H+ ions to yield carbonic acid, which is unstable, and rapidly decomposes to yield carbon dioxide gas and water. H2CO3(aq)H2O(l) + CO2(g) Chapter Ten

22. Acids react with ammonia to yield ammonium salts, most of which are water-soluble. • Living organisms contain a group of compounds called amines, which contain ammonia-like nitrogen atoms bonded to carbon. Amines react with acids just as ammonia does. Methylamine, an organic compound found in rotting fish, reacts with HCl: Chapter Ten

23. 10.6 Acid and Base Strength • Strong acid: An acid that gives up H+ easily and is essentially 100% dissociated in water. • Dissociation: The splitting apart of an acid in water to give H+ and an anion. • Weak acid: An acid that gives up H+ with difficulty and is less than 100% dissociated in water. • Weak base: A base that has only a slight affinity for H+ and holds it weakly. • Strong base: A base that has a high affinity for H+ and holds it tightly. Chapter Ten

24. The stronger the acid, the weaker its conjugate base; the weaker the acid, the stronger its conjugate base. • An acid–base proton transfer equilibrium always favors reaction of the stronger acid with the stronger base, and formation of the weaker acid and base. • The proton always leaves the stronger acid and always ends up in the weaker acid, whose stronger conjugate base holds the proton tightly. Chapter Ten

25. Chapter Ten

26. Shown below is the zwitterion of the amino acid alanine. The zwitterion can serve as • both an acid and a base. • only an acid. • only a base. • neither an acid not a base.

27. Shown below is the zwitterion of the amino acid alanine. The zwitterion can serve as • both an acid and a base. • only an acid. • only a base. • neither an acid not a base.

28. 10.7 Acid Dissociated Constants • The reaction of a weak acid with water, like any chemical equilibrium, can be described by an equilibrium equation. Ka = [H3O+][A-] [HA] • Strong acids have Ka >> 1 because dissociation is favored and weak acids have Ka << 1 because dissociation is not favored. • Donation of each successive H+ from a polyprotic acid is more difficult than the one before it, so Ka values become successively lower. Chapter Ten

29. Most organic acids, which contain the group -COOH, have Ka values near 10-5. Chapter Ten

30. The weaker the acid, the stronger its conjugate base. The reverse reaction is favored as indicated by the arrow size.

31. HC2H3O2 (acetic acid) Ka = 1.8 × 10-5

32. The equilibrium constant for the reaction shown below is called the base dissociation constant, Kb, and is a measure of the base strength of the acetate ion. For acetate ion Kb = 5.6  10–10. In this reaction the strongest acid/strongest base are • acetic acid/acetate ion. • acetic acid/hydroxide ion. • water/acetate ion. • water/hydroxide ion.

33. The equilibrium constant for the reaction shown below is called the base dissociation constant, Kb, and is a measure of the base strength of the acetate ion. For acetate ion Kb = 5.6  10–10. In this reaction the strongest acid/strongest base are • acetic acid/acetate ion. • acetic acid/hydroxide ion. • water/acetate ion. • water/hydroxide ion.

34. 10.8 Dissociation of Water • Like all weak acids, water is slightly dissociated into H+ and OH- ions. At 25oC, the concentration of each ion is 1.00 x 10-7 M in pure water. • The ion product constant for water, kw, is: kw = [H3O+][OH-] but [H2O] = 1, so much of it, so [H2O] kw = [H3O+][OH-]= 1.00 x 10-14 at 25oC. • Acidic solution: [H3O+] > 10-7 M and [OH-] < 10-7 M • Neutral solution: [H3O+] = 10-7 M and [OH-] = 10-7 M • Basic solution: [H3O+] < 10-7 M and [OH-] > 10-7 M Chapter Ten

36. What is the [H3O+] if a NaOH(aq) is 2.1 × 10-5 M at 25 °C? (Kw = 1.0 × 10-14) Kw = [H3O+] [OH-] 1.0 × 10-14 = [H3O+] 2.1 × 10-5 [H3O+] = 1.0 × 10-14 2.1 × 10-5 [H3O+] = 4.8 × 10-10 Chapter Ten

37. 10.9 Measuring Acidity in Aqueous Solution: pH • The concentrations of H3O+ or OH- in solution can vary over a wide range. A logarithmic scale can be easier to use. • pH = -log [H3O+ ] pOH = -log [OH- ] • [H3O+ ] = 10-pH [OH- ] = 10-pOH • Acidic solution: pH < 7 pOH > 7 • Neutral solution: pH = 7 pOH = 7 • Basic solution: pH > 7 pOH < 7 • pH + pOH = 14.00 at 25C Chapter Ten

38. Chapter Ten

40. A aqueous solution has a pH of 12.26 at 25 °C, what is the pOH? Is this acidic or basic? pH + pOH = 14 12.26 + pOH = 14 pOH = 14 - 12.26 = 1.74 A baking soda solution has [H3O+] of 1.3 × 10-8, what are the pH and pOH? Is this acidic or basic? pH + pOH = 14 pH = -log [H3O+ ] pH = -log 1.3 × 10-8 7.9 + pOH = 14 pH = -log 1.3 × 10-8 = 7.9 pOH = 14 - 7.9 = 6.1 Chapter Ten

41. A 0.10 M solution of ammonium ion, NH4+, has a pH of 5.6. Ammonium ion is a • strong acid. • weak acid. • strong base. • weak base.

42. A 0.10 M solution of ammonium ion, NH4+, has a pH of 5.6. Ammonium ion is a NH4+ + H2O ⇄ NH3 + OH- • strong acid. • weak acid. • strong base. • weak base.

43. 10.10 Working with pH • An antilogarithm has the same number of digits that the original number has to the right of the decimal point. • A logarithm contains the same number of digits to the right of the decimal point that the original number has. Chapter Ten

45. A aqueous solution has a pH of 12.26 at 25 °C. What is the [OH-]. Is the solution acidic or basic? pH + pOH = 14 12.26 + pOH = 14 pOH = 14 – 12.26 = 1.74 pH = -log [H3O+ ] or the inverse [H3O+ ] = 10-pH pOH = -log [OH- ] or the inverse [OH- ] = 10-pOH [OH- ] = 10-pOH [OH- ] = 10-1.74 = 1.82 × 10-2 Chapter Ten

46. 10.11 Laboratory Determination of Acidity (a) The color of universal indicator in solutions of known pH from 1 to 12. (b) Testing pH with a paper strip. Comparing the color of the strip with the code on the package gives the approximate pH. Chapter Ten

47. A much more accurate way to determine pH uses an electronic pH meter like the one shown below. • Electrodes are dipped into the solution, and the pH is read from the meter. Chapter Ten

48. 10.12 Buffer Solutions • Buffer: A combination of substances that act together to prevent a drastic change in pH; usually a weak acid and its conjugate base. • Rearranging the Ka equation shows that the value of [H3O+] depends on the ratio [HA]/[A-]. [H3O+] = Ka [HA]/[A-] • Most H3O+ added is removed by reaction with A- ,so [HA]increases and [A-] decreases. As long as these changes are small, the ratio [HA]/[A-] changes only slightly, and there is little change in the pH. Chapter Ten

49. When 0.010 mol of acid and 0.010 mol of base are added to 1.0 L of pure water and to 1.0 L of a 0.10 M acetic acid–0.10 M acetate ion buffer, the pH of the water varies between 12 and 2, while the pH of the buffer varies only between 4.85 and 4.68. Chapter Ten