General, Organic, and Biochemistry, 7e

1. General, Organic, and Biochemistry, 7e Bettelheim, Brown, and March

2. Chapter 8 Acids and Bases

3. Arrhenius Acids and Bases • In 1884, Svante Arrhenius proposed these definitions • acid: a substance that produces H3O+ ions aqueous solution • 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. 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

5. 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 merely separate • other bases are not hydroxides; these bases produce OH- by reacting with water molecules

6. Arrhenius Acids and Bases • we use curved arrows to show the transfer of a proton from water to ammonia

7. Acid and Base Strength • Strong acid: one that reacts completely or almost completely with water to form H3O+ ions • 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. Acid and Base Strength • Weak acid: a substance that dissociates only partially in water to produce H3O+ ions • acetic acid, for example, is a weak acid; in water, only 4 out every 1000 molecules are converted to acetate ions • Weak base: a substance that dissociates only partially in water to produce OH- ions • ammonia, for example, is a weak base

9. Brønsted-Lowry Acids & Bases • Acid: a proton donor • Base: a proton acceptor • Acid-base reaction: a proton transfer reaction • Conjugate acid-base pair: any pair of molecules or ions that can be interconverted by transfer of a proton

10. Brønsted-Lowry Acids & Bases • Brønsted-Lowry definitions do not require water as a reactant

11. Brønsted-Lowry Acids & Bases • we can use curved arrows to show the transfer of a proton from acetic acid to ammonia

13. Brønsted-Lowry Acids & Bases • Note the following about the conjugate acid-base pairs in the table 1. an acid can be positively charged, neutral, or negatively charged; examples of each type are H3O+, H2CO3, and H2PO4- 2. a base can be negatively charged or neutral; examples are OH-, Cl-, and NH3 3. acids are classified a monoprotic, diprotic, or triprotic depending on the number of protons each may give up; examples are HCl, H2CO3, and H3PO4

14. 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. Brønsted-Lowry Acids & Bases • the HCO3- ion, for example, can give up a proton to become CO32-, or it can accept a proton to become H2CO3 • a substance that can act as either an acid or a base is said to be amphiprotic • the most important amphiprotic substance in Table 8.2 is H2O; it can accept a proton to become H3O+, 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. 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 8.2, and its conjugate base, I-, is the weakest base in the table • CH3COOH (acetic acid) is a stronger acid that H2CO3 (carbonic acid); conversely, CH3COO- (acetate ion) is a weaker base that HCO3- (bicarbonate ion)

17. 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. 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 8.2, 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. Acid-Base Equilibria • identify the two acids and bases, and their relative strengths • the position of this equilibrium lies to the right

20. Acid-Base Equilibria • Example: predict the position of equilibrium in this acid-base reaction

21. Acid-Base Equilibria • Example: predict the position of equilibrium in this acid-base reaction • Solution: the position of this equilibrium lies to the right

22. Acid Ionization Constants • when a weak acid, HA, dissolves in water • the equilibrium constant, Keq, for this ionization is • because water is the solvent and its concentration changes very little when we add HA to it, we treat [H2O] 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, Ka

23. Acid Ionization Constants • Ka 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 pKa where • the value of pKa for acetic acid is 4.75 • values of Ka and pKa for some weak acids are given in Table 8.3 • as you study the entries in this table, note the inverse relationship between values of Ka and pKa • the weaker the acid, the smaller its Ka, but the larger its pKa

25. Properties of Acids & Bases • Neutralization • acids and bases react with each other in a process called neutralization; these reactions are discussed in Section 8.10 • Reaction with metals • strong acids react with certain metals (called active metals) to produce a salt and hydrogen gas, H2 • 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 H2

26. 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. Properties of Acids & Bases • Reaction with metal oxides • strong acids react with metal oxides to give water plus a salt

28. Properties of Acids & Bases • Reaction with carbonates and bicarbonates • strong acids react with carbonates to give carbonic acid, which rapidly decomposes to CO2 and H2O • strong acids react similarly with bicarbonates

29. Properties of Acids & Bases • Reaction with ammonia and amines • any acid stronger than NH4+ is strong enough to react with NH3 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 NH3 are replaced by carbon groups

30. Self-Ionization of Water • pure water contains a very small number of H3O+ ions and OH- ions formed by proton transfer from one water molecule to another • the equilibrium expression for this reaction is • we can treat [H2O] as a constant = 55.5 mol/L

31. Self-Ionization of Water • combining these constants gives a new constant called the ion product of water, Kw • in pure water, the value of Kw is 1.0 x 10-14 • in pure water, H3O+ and OH- are formed in equal amounts (remember the balanced equation for the self-ionization of water) • this means that in pure water

32. 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 [H3O+] and [OH-] in any aqueous solution is equal to 1.0 x 10-14 • for example, if we add 0.010 mole of HCl to 1 liter of pure water, it reacts completely with water to give 0.010 mole of H3O+ • in this solution, [H3O+] is 0.010 or 1.0 x 10-2 • this means that the concentration of hydroxide ion is

33. 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 [H3O+] • we can now state the definitions of acidic and basic solutions in terms of pH • acidic solution: one whose pH is less than 7.0 • basic solution: one whose pH is greater than 7.0 • neutral solution: one whose pH is equal to 7.0

34. pH and pOH • just as pH is a convenient way to designate the concentration of H3O+, pOH is a convenient way to designate the concentration of OH- pOH = -log[OH-] • the ion product of water, Kw, 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. pH and pOH • pH of some common materials

36. pH of Salt Solutions • When some salts dissolve in pure water, there is no change in pH from that of pure water • Many salts, however, are acidic or basic and cause a change the pH when they dissolve • We are concerned in this section with basic salts and acidic salts

37. pH of Salt Solutions • Basic salt: the salt of a strong base and a weak acid; when dissolved in water, it raises the pH • as an example of a basic salt is sodium acetate • when this salt dissolves in water, it ionizes; Na+ ions do not react with water, but CH3COO- ions do • the position of equilibrium lies to the left • nevertheless, there are enough OH- ions present in 0.10 M sodium acetate to raise the pH to 8.88

38. pH of Salt Solutions • Acidic salt: the salt of a strong acid and a weak base; when dissolved in water, it lowers the pH • an example of an acidic salt is ammonium chloride • chloride ion does not react with water, but the ammonium ion does • although the position of this equilibrium lies to the left, there are enough H3O+ ions present to make the solution acidic

39. Acid-Base Titrations • 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

40. Acid-Base Titrations • An acid-base titration must meet these requirement 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 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

41. Acid-Base Titrations • As an example, let us use 0.108 M H2SO4 to determine the concentration of a NaOH solution • requirement 1: we know the balanced equation • requirement 2: the reaction between H3O+ and OH- is rapid and complete • 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: we use volumetric glassware

42. Acid-Base Titrations • experimental measurements • doing the calculations

43. pH Buffers • 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

44. pH Buffers • How an acetate buffer resists changes in pH • if we add a strong acid, such as HCl, added H3O+ 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

45. pH Buffers • The effect of a buffer can be quite dramatic • consider a phosphate buffer prepared by dissolving 0.10 mole of NaH2PO4 (a weak acid) and 0.10 mole of Na2HPO4 (the salt of its conjugate base) in enough water to make 1 liter of solution

46. pH Buffers • 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 pKa of the weak acid • if we want a buffer of pH 9.14, for example, we can mix equal molar amounts of boric acid (H3BO3), pKa 9.14, and sodium dihydrogen borate (NaH2BO3), the salt of its conjugate base

47. pH Buffers • Buffer capacity depends both its pH and its concentration

48. Blood Buffers • The average pH of human blood is 7.4 • any change larger than 0.10 pH unit in either direction can cause illness • To maintain this pH, the body uses three buffer systems • carbonate buffer: H2CO3 and its conjugate base, HCO3- • phosphate buffer: H2PO4- and its conjugate base, HPO42- • proteins: discussed in Chapter 21

49. Henderson-Hasselbalch Eg. • Henderson-Hasselbalch equation: a mathematical relationship between • pH, • pKa of the weak acid, HA • concentrations HA, and its conjugate base, A- • It is derived in the following way • taking the logarithm of this equation gives

50. Henderson-Hasselbalch Eg. • multiplying through by -1 gives • -log Ka is by definition pKa, and -log [H3O+] is by definition pH; making these substitutions gives • rearranging terms gives