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 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. • 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. 10.3 The Brønsted-Lowry Definition of Acids and Bases • A Brønsted–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

8. 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

9. 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

10. 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

11. 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

12. 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

13. 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

14. 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

15. 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

16. Chapter Ten

17. 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

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

19. 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-]= 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

20. 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

21. Chapter Ten

22. 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

23. 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

24. 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

25. 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

26. 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

27. 10. 13 Buffers in the Body • The pH of body fluids is maintained by three major buffer systems. The carbonic acid–bicarbonate system, the dihydrogen phosphate–hydrogen phosphate system, and a third system that depends on the ability of proteins to act as either proton acceptors or proton donors at different pH values. • The carbonic acid–bicarbonate system is the principal buffer in blood serum and other extracellular fluids. The hydrogen phosphate system is the major buffer within cells. Chapter Ten

28. A change in the breathing rate provides a quick adjustment in the bicarbonate buffer system. • When the CO2 concentration in the blood starts to rise, the breathing rate increases to remove CO2, thereby decreasing the acid concentration. • When the CO2 concentration in the blood starts to fall, the breathing rate decreases and acid concentration increases. Chapter Ten

29. Each day, acid produced in the body is excreted in the urine. The kidney returns HCO3- to the extracellular fluids, where it becomes part of the bicarbonate reserve. Chapter Ten

30. 10.14 Acid and Base Equivalents • We think in terms of ion equivalents when we are primarily interested in the ion itself rather than the compound that produced the ion. For similar reasons, we use acid or base equivalents. • One equivalent (Eq) of an acid is equal to the molar mass of the acid divided by the number of H+ ions produced per formula unit. Similarly, one equivalent of a base is the weight in grams that can produce one mole of OH- ions. Chapter Ten

31. One Eq of a monoprotic acid is the molar mass of the acid. One equivalent of a diprotic acid is the molar mass of the acid divided by 2 since each mole can produce two moles of H+. • One equivalent of any acid neutralizes one equivalent of any base; this is convenient when only the acidity or basicity of a solution is of interest rather than the identity of the acid or base. Chapter Ten

32. Because acid–base equivalents are so useful, clinical chemists sometimes express acid and base concentrations in normality rather than molarity. • The normality (N) of an acid or base solution is defined as the number of equivalents of acid or base per liter of solution. • For any acid or base, normality is always equal to molarity times the number of H+ or OH- ions produced per formula unit. • N of acid = (M of acid)  (# of H+ ions produced) • N of base = (M of base)  (# of OH- ions produced) Chapter Ten

33. 10.15 Titration • The pH of a solution gives the H+ concentration but not necessarily the total acid concentration. • The H+ concentration gives only the amount of acid that has dissociated into ions, whereas total acid concentration gives the sum of dissociated plus undissociated acid. • In a 0.10 M solution of acetic acid the total acid concentration is 0.10 M, yet the H+ concentration is only 0.0013 M because acetic acid is a weak acid that is only about 1% dissociated. • The total acid or base concentration of a solution can be found by carrying out a titration procedure. Chapter Ten

34. (a) A measured volume of the acid solution is placed in the flask along with an indicator. • (b) The base of known concentration is then added from a buret until the color change of the indicator shows that neutralization is complete. Chapter Ten

35. Reading the buret gives the volume of base. The molarity and volume of base yields the moles of base. • The coefficients in the balanced equation allow us to find the moles of acid neutralized. Dividing the moles of acid by the volume of the acid gives the acid molarity. Chapter Ten

36. 10.16 Acidity and Basicity of Salt Solutions • Salt solutions can be neutral, acidic, or basic, depending on the ions present, because some ions react with water to produce H+ and some ions react with water to produce OH-. • To predict the acidity of a salt solution, it is convenient to classify salts according to the acid and base from which they are formed in a neutralization reaction. • The general rule for predicting the acidity or basicity of a salt solution is that the stronger partner from which the salt is formed dominates. Chapter Ten

37. A salt formed from a strong acid and a weak base yields an acidic solution because the strong acid dominates. • A salt formed from a weak acid and a strong base yields a basic solution because the base dominates. • A salt formed from a strong acid and a strong base yields a neutral solution because neither dominates. Chapter Ten

38. Chapter Summary • According to the Brønsted–Lowry definition, an acid is a substance that donates a hydrogen ion and a base is a substance that accepts a hydrogen ion. • Thus, the generalized reaction of an acid with a base involves the reversible transfer of a proton. • A strong acidgives up a proton easily and is 100% dissociated in aqueous solution. • A weak acidgives up a proton with difficulty and is only slightly dissociated in water. • A strong baseaccepts and holds a proton readily, whereas a weak basehas a low affinity for a proton. Chapter Ten

39. Chapter Summary Cont. • The two substances that are related by the gain or loss of a proton are called a conjugate acid–base pair. • A proton-transfer reaction always takes place in the direction that favors formation of the weaker acid. • Water is amphoteric; it can act as an acid or a base. • The acidity or basicity of an aqueous solution is given by its pH, defined as the negative logarithm of the hydronium ion concentration. • A pH below 7 means an acidic solution; a pH equal to 7 means a neutral solution; and a pH above 7 means a basic solution. Chapter Ten

40. Chapter Summary Cont. • The pH of a solution can be controlled through the use of a bufferthat acts to remove either added H+ ions or added OH- ions. Most buffer solutions consist of roughly equal amounts of a weak acid and its conjugate base. • Acid (or base) concentrations are determined in the laboratory by titrationof a solution of unknown concentration with a base (or acid) solution of known strength until an indicator signals that neutralization is complete. Chapter Ten

41. End of Chapter 10 Chapter Ten