1. 9 9-1 © 2006 Thomson Learning, Inc. All rights reserved General, Organic, and Biochemistry, 8e Bettelheim, Brown, Campbell & Farrell

2. 9 9-2 © 2006 Thomson Learning, Inc. All rights reserved Chapter 9 Acids and Bases Acids and Bases

3. 9 9-3 © 2006 Thomson Learning, Inc. All rights reserved Arrhenius Acids and Bases In 1884, Svante Arrhenius proposed these definitions: Acid:Acid: a substance that produces H 3 O + ions in aqueous solution. Base:Base: a substance that produces OH - ions in aqueous solution. This definition of an acid is a slight modification of the original Arrhenius definition, which was that an acid produces H + in aqueous solution Today we know that H + reacts immediately with a water molecule to give a hydronium ion

4. 9 9-4 © 2006 Thomson Learning, Inc. All rights reserved Arrhenius Acids and Bases When HCl, for example, dissolves in water, its reacts with water to give hydronium ion and chloride ion We use curved arrows to show the change in position of electron pairs during this reaction; the first arrow shows the formation of an O-H bond, the second shows the breaking of an H-Cl bond:

5. 9 9-5 © 2006 Thomson Learning, Inc. All rights reserved Arrhenius Acids and Bases With bases, the situation is slightly different. Many bases are metal hydroxides such as KOH, NaOH, Mg(OH) 2, and Ca(OH) 2. These compounds are ionic solids and when they dissolve in water, their ions become hydrated and separate: Not all bases are hydroxides; these bases produce OH - by reacting with water molecules, here shown for ammonia:

6. 9 9-6 © 2006 Thomson Learning, Inc. All rights reserved Arrhenius Acids and Bases We use curved arrows to show the transfer of a proton from water to ammonia. The first arrow shows the formation of an N-H bond, the second shows breaking of an H-O bond.

7. 9 9-7 © 2006 Thomson Learning, Inc. All rights reserved Acid and Base Strength Strong acid:Strong acid: one that reacts completely or almost completely with water to form H 3 O + ions. Strong base:Strong base: one that reacts completely or almost completely with water to form OH - ions. Here are the six most common strong acids and the four most common strong bases:

8. 9 9-8 © 2006 Thomson Learning, Inc. All rights reserved Acid and Base Strength Weak acid: Weak acid: a substance that dissociates only partially in water to produce H 3 O + ions. acetic acid, for example, is a weak acid; in water, only 4 out every 1000 molecules are converted to acetate ions: Weak base: Weak base: a substance that dissociates only partially in water to produce OH - ions. ammonia, for example, is a weak base:

9. 9 9-9 © 2006 Thomson Learning, Inc. All rights reserved Brønsted-Lowry Acids & Bases Acid:Acid: a proton donor. Base:Base: a proton acceptor. Acid-base reaction:Acid-base reaction: a proton transfer reaction. Conjugate acid-base pair:Conjugate acid-base pair: any pair of molecules or ions that can be interconverted by transfer of a proton.

10. 9 9-10 © 2006 Thomson Learning, Inc. All rights reserved Brønsted-Lowry Acids & Bases Brønsted-Lowry definitions do not require water as a reactant:

11. 9 9-11 © 2006 Thomson Learning, Inc. All rights reserved Brønsted-Lowry Acids & Bases We can use curved arrows to show the transfer of a proton from acetic acid to ammonia:

12. 9 9-12 © 2006 Thomson Learning, Inc. All rights reserved

13. 9 9-13 © 2006 Thomson Learning, Inc. All rights reserved Brønsted-Lowry Acids & Bases Note the following about the conjugate acid-base pairs in the table 9.2: 1. An acid can be positively charged, neutral, or negatively charged; examples of each type are H 3 O +, H 2 CO 3, and H 2 PO 4 -. 2. A base can be negatively charged or neutral; examples are OH -, Cl -, and NH 3. 3. Acids are classified a monoprotic, diprotic, or triprotic depending on the number of protons each may give up; examples are HCl, H 2 CO 3, and H 3 PO 4.

14. 9 9-14 © 2006 Thomson Learning, Inc. All rights reserved Brønsted-Lowry Acids & Bases Carbonic acid, for example can give up one proton to become bicarbonate ion, and then the second proton to become carbonate ion: 4. Several molecules and ions appear in both the acid and conjugate base columns; that is, each can function as either an acid or a base.

15. 9 9-15 © 2006 Thomson Learning, Inc. All rights reserved Brønsted-Lowry Acids & Bases The HCO 3 - ion, for example, can give up a proton to become CO 3 2-, or it can accept a proton to become H 2 CO 3. amphiprotic.A substance that can act as either an acid or a base is said to be amphiprotic. The most important amphiprotic substance in Table 9.2 is H 2 O; it can accept a proton to become H 3 O +, or lose a proton to become OH -. 5. A substance cannot be a Brønsted-Lowry acid unless it contains a hydrogen atom, but not all hydrogen atoms in most compounds can be given up. Acetic acid, for example, gives up only one proton.

16. 9 9-16 © 2006 Thomson Learning, Inc. All rights reserved Brønsted-Lowry Acids & Bases 6. There is an inverse relationship between the strength of an acid and the strength of its conjugate base The stronger the acid, the weaker its conjugate base. HI, for example, is the strongest acid in Table 9.2, and its conjugate base, I -, is the weakest base in the table. CH 3 COOH (acetic acid) is a stronger acid that H 2 CO 3 (carbonic acid); conversely, CH 3 COO - (acetate ion) is a weaker base that HCO 3 - (bicarbonate ion).

17. 9 9-17 © 2006 Thomson Learning, Inc. All rights reserved Acid-Base Equilibria We know that HCl is a strong acid, which means that the position of this equilibrium lies very far to the right. In contrast, acetic acid is a weak acid, and the position of its equilibrium lies very far to the left. But what if the base is not water? How can we determine which are the major species present?

18. 9 9-18 © 2006 Thomson Learning, Inc. All rights reserved Acid-Base Equilibria To predict the position of an acid-base equilibrium such as this, we do the following: Identify the two acids in the equilibrium; one on the left and one on the right. Using the information in Table 9.2, determine which is the stronger acid and which is the weaker acid. Also determine which is the stronger base and which is the weaker base; remember that the stronger acid gives the weaker conjugate base, and the weaker acid gives the stronger conjugate base. The stronger acid reacts with the stronger base to give the weaker acid and weaker base; equilibrium lies on the side of the weaker acid and weaker base.

19. 9 9-19 © 2006 Thomson Learning, Inc. All rights reserved Acid-Base Equilibria Identify the two acids and bases, and their relative strengths: the position of this equilibrium lies to the right:

20. 9 9-20 © 2006 Thomson Learning, Inc. All rights reserved Acid-Base Equilibria Example: Example: predict the position of equilibrium for this acid-base reaction:

21. 9 9-21 © 2006 Thomson Learning, Inc. All rights reserved Acid-Base Equilibria Example: Example: predict the position of equilibrium in this acid-base reaction: Solution: Solution: the position of this equilibrium lies to the right.

22. 9 9-22 © 2006 Thomson Learning, Inc. All rights reserved Acid Ionization Constants When a weak acid, HA, dissolves in water the equilibrium constant expression, K eq, for this ionization is: Because water is the solvent and its concentration changes very little when we add HA to it, we treat [H 2 O] as a constant equal to 1000 g/L or 55.5 mol/L. We combine the two constants to give a new constant, which we call an acid ionization constant, K a :

23. 9 9-23 © 2006 Thomson Learning, Inc. All rights reserved Acid Ionization Constants K a for acetic acid, for example is 1.8 x 10 -5. Because the acid ionization constants for weak acids are numbers with negative exponents, we commonly express acid strengths as pK a where: The value of pK a for acetic acid is 4.75. Values of K a and pK a for some weak acids are given in Table 9.3. As you study the entries in this table, note the inverse relationship between values of K a and pK a. The weaker the acid, the smaller its K a, but the larger its pK a.

24. 9 9-24 © 2006 Thomson Learning, Inc. All rights reserved

25. 9 9-25 © 2006 Thomson Learning, Inc. All rights reserved Properties of Acids & Bases Neutralization: Acids and bases react with each other in a process called neutralization; these reactions are discussed in Section 9.6A. Reaction with metals: Strong acids react with certain metals (called active metals) to produce a salt and hydrogen gas, H 2 reaction of a strong acid with a metal is a redox reaction; the metal is oxidized to a metal ion and H + is reduced to H 2.

26. 9 9-26 © 2006 Thomson Learning, Inc. All rights reserved Properties of Acids & Bases Reaction with metal hydroxides: Reaction of an acid with a metal hydroxide gives a salt plus water; The reaction is more accurately written as: Omitting spectator ions gives this net ionic equation:

27. 9 9-27 © 2006 Thomson Learning, Inc. All rights reserved Properties of Acids & Bases Reaction with metal oxides Strong acids react with metal oxides to give water plus a salt:

28. 9 9-28 © 2006 Thomson Learning, Inc. All rights reserved Properties of Acids & Bases Reaction with carbonates and bicarbonates strong acids react with carbonates to give carbonic acid, which rapidly decomposes to CO 2 and H 2 O: strong acids react similarly with bicarbonates:

29. 9 9-29 © 2006 Thomson Learning, Inc. All rights reserved Properties of Acids & Bases Reaction with ammonia and amines Any acid stronger than NH 4 + is strong enough to react with NH 3 to give a salt. In the following reaction, the salt formed is ammonium chloride, which is shown as it would be ionized in aqueous solution: In Ch 16 we study amines, compounds in which one or more hydrogens of NH 3 are replaced by carbon groups.

30. 9 9-30 © 2006 Thomson Learning, Inc. All rights reserved Self-Ionization of Water Pure water contains a very small number of H 3 O + ions and OH - ions formed by proton transfer from one water molecule to another The equilibrium constant expression for this reaction is: We can treat [H 2 O] as a constant = 55.5 mol/L

31. 9 9-31 © 2006 Thomson Learning, Inc. All rights reserved Self-Ionization of Water ion product of water, K wCombining these constants gives a new constant called the ion product of water, K w. In pure water, the value of K w is 1.0 x 10 -14. In pure water, H 3 O + and OH - are formed in equal amounts (remember the balanced equation for the self- ionization of water). This means that in pure water:

32. 9 9-32 © 2006 Thomson Learning, Inc. All rights reserved Self-Ionization of Water The equation for the ionization of water applies not only to pure water but also to any aqueous solution. The product of [H 3 O + ] and [OH - ] in any aqueous solution is equal to 1.0 x 10 -14. For example, if we add 0.010 mol of HCl to 1.00 liter of pure water, it reacts completely with water to give 0.010 mole of H 3 O +. In this solution, [H 3 O + ] is 0.010 or 1.0 x 10 -2. This means that the concentration of hydroxide ion is:

33. 9 9-33 © 2006 Thomson Learning, Inc. All rights reserved pH and pOH Because hydronium ion concentrations for most solutions are numbers with negative exponents, we commonly express these concentrations as pH, where: pH = -log [H 3 O + ] We can now state the definitions of acidic and basic solutions in terms of pH: Acidic solution:Acidic solution: one whose pH is less than 7.0. Basic solution:Basic solution: one whose pH is greater than 7.0. Neutral solution:Neutral solution: one whose pH is equal to 7.0.

34. 9 9-34 © 2006 Thomson Learning, Inc. All rights reserved pH and pOH Just as pH is a convenient way to designate the concentration of H 3 O +, pOH is a convenient way to designate the concentration of OH -. pOH = -log[OH - ] the ion product of water, K w, is 1.0 x 10 -14 taking the logarithm of this equation gives: pH + pOH = 14 Thus, if we know the pH of an aqueous solution, we can easily calculate its pOH.

35. 9 9-35 © 2006 Thomson Learning, Inc. All rights reserved pH and pOH pH of some common materials

36. 9 9-36 © 2006 Thomson Learning, Inc. All rights reserved Acid-Base Titrations Titration: Titration: an analytical procedure in which a solute in a solution of known concentration reacts with a known stoichiometry with a substance whose concentration is to be determined: In this chapter, we are concerned with titrations in which we use an acid (or base) of known concentration to determine the concentration of a base (or acid) in another solution.

37. 9 9-37 © 2006 Thomson Learning, Inc. All rights reserved Acid-Base Titrations An acid-base titration must meet these requirements: 1. We must know the equation for the reaction so that we can determine the stoichiometric ratio of reactants to use in our calculations. 2. The reaction must be rapid and complete. end point 3. There must be a clear-cut change in a measurable property at the end point (when the reagents have combined exactly). 4. We must have accurate measurements of the amount of each reactant.

38. 9 9-38 © 2006 Thomson Learning, Inc. All rights reserved Acid-Base Titrations As an example, let us use 0.108 M H 2 SO 4 to determine the concentration of a NaOH solution Requirement 1:Requirement 1: we know the balanced equation: Requirement 2:Requirement 2: the reaction between H 3 O + and OH - is rapid and complete. Requirement 3:Requirement 3: we can use either an acid-base indicator or a pH meter to observe the sudden change in pH that occurs at the end point of the titration. Requirement 4:Requirement 4: we use volumetric glassware.

39. 9 9-39 © 2006 Thomson Learning, Inc. All rights reserved Acid-Base Titrations experimental measurements doing the calculations

40. 9 9-40 © 2006 Thomson Learning, Inc. All rights reserved pH Buffers pH buffer: pH buffer: a solution that resists change in pH when limited amounts of acid or base are added to it. A pH buffer as an acid or base “shock absorber.” A pH buffer is common called simply a buffer. The most common buffers consist of approximately equal molar amounts of a weak acid and a salt of the weak acid; that is, approximately equal molar amounts of a weak acid and a salt of its conjugate base. For example, if we dissolve 1.0 mole of acetic acid and 1.0 mole of its conjugate base (in the form of sodium acetate) in water, we have an acetate buffer.

41. 9 9-41 © 2006 Thomson Learning, Inc. All rights reserved pH Buffers How does an acetate buffer resists changes in pH? If we add a strong acid, such as HCl, added H 3 O + ions react with acetate ions and are removed from solution: If we add a strong base, such as NaOH, added OH - ions react with acetic acid and are removed from solution:

42. 9 9-42 © 2006 Thomson Learning, Inc. All rights reserved pH Buffers The effect of a buffer can be quite dramatic Consider a phosphate buffer prepared by dissolving 0.10 mole of NaH 2 PO 4 (a weak acid) and 0.10 mole of Na 2 HPO 4 (the salt of its conjugate base) in enough water to make 1 liter of solution.

43. 9 9-43 © 2006 Thomson Learning, Inc. All rights reserved pH Buffers Buffer pH Buffer pH If we mix equal molar amounts of a weak acid and a salt of its conjugate base, the pH of the solution will be equal to the pK a of the weak acid. If we want a buffer of pH 9.14, for example, we can mix equal molar amounts of boric acid (H 3 BO 3 ), pK a 9.14, and sodium dihydrogen borate (NaH 2 BO 3 ), the salt of its conjugate base

44. 9 9-44 © 2006 Thomson Learning, Inc. All rights reserved pH Buffers Buffer capacity Buffer capacity is the amount of hydronium or hydroxide ions that a buffer can absorb without a significant change in pH. Buffer capacity depends both its pH and its concentration

45. 9 9-45 © 2006 Thomson Learning, Inc. All rights reserved Blood Buffers The average pH of human blood is 7.4. any change greater than 0.10 pH unit in either direction can cause illness. To maintain this pH, the body uses three buffer systems: carbonate buffer:carbonate buffer: H 2 CO 3 and its conjugate base, HCO 3 - phosphate buffer:phosphate buffer: H 2 PO 4 - and its conjugate base, HPO 4 2- proteins:proteins: discussed in Chapter 21.

46. 9 9-46 © 2006 Thomson Learning, Inc. All rights reserved Henderson-Hasselbalch Eg. Henderson-Hasselbalch equation: Henderson-Hasselbalch equation: a mathematical relationship between: pH, pK a of the weak acid, HA concentrations of HA and its conjugate base A - It is derived in the following way: taking the logarithm of this equation gives

47. 9 9-47 © 2006 Thomson Learning, Inc. All rights reserved Henderson-Hasselbalch Eg. multiplying through by -1 gives: -log K a is by definition pK a, and -log [H 3 O + ] is by definition pH; making these substitutions gives: rearranging terms gives:

48. 9 9-48 © 2006 Thomson Learning, Inc. All rights reserved Henderson-Hasselbalch Eg. Example: Example: what is the pH of a phosphate buffer solution containing 1.0 mole of NaH 2 PO 4 and 0.50 mole of Na 2 HPO 4 dissolved in enough water to make 1.0 liter of solution

49. 9 9-49 © 2006 Thomson Learning, Inc. All rights reserved Henderson-Hasselbalch Eg. Example: Example: what is the pH of a phosphate buffer solution containing 1.0 mole of NaH 2 PO 4 and 0.50 mole of Na 2 HPO 4 in enough water to make one liter of solution Solution Solution the equilibrium we are dealing with and its pK a are substituting these values in the H-H equation gives

50. 9 9-50 © 2006 Thomson Learning, Inc. All rights reserved Biochemical Buffers The original buffers were made from simple acids and bases, such as acetic acid, phosphoric acid, and citric acid and their conjugate bases. Many of these, however, have limitations: ---they often changed their pH too much if the solution was diluted or the temperature changed. ---they often permeated cells in solution thereby changing the chemistry in the interior of the cell. To overcome these shortcomings, N.E. Good developed a series of buffers that consist of zwitterions, molecules that do not readily permeate cell membranes.

51. 9 9-51 © 2006 Thomson Learning, Inc. All rights reserved Biochemical Buffers Typical Good buffers are:

52. 9 9-52 © 2006 Thomson Learning, Inc. All rights reserved End Chapter 9 Chapter 9 Acids and Bases