1. Frederick A. Bettelheim William H. Brown Mary K. Campbell Shawn O. Farrell www.cengage.com/chemistry/bettelheim William H. Brown Beloit College Chapter 8 Acids and Bases

2. 8-2 Arrhenius Acids and Bases In 1884, Svante Arrhenius proposed these definitions: Acid: Acid: A substance that produces H + ions in aqueous solution. Base: Base: A substance that produces OH - ions in aqueous solution. Today we know that H + reacts immediately with a water molecule to give a hydronium ion.

3. 8-3 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.

4. 8-4 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: However, some substances are bases but do not contain hydroxide ions. Instead, some bases produce OH - by reacting with water molecules, here shown for ammonia:

5. 8-5 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. The Arrhenius theory is limited  a more comprehensive theory is needed

6. 8-6 Brønsted-Lowry Acids and 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.

7. 8-7 Brønsted-Lowry Acids and Bases Brønsted-Lowry definitions do not require water as a reactant.

8. 8-8 Brønsted-Lowry Acids and Bases We can use curved arrows to show the transfer of a proton from acetic acid to ammonia: acid base conjugate acid conjugate base

9. 8-9 Brønsted-Lowry Acids and Bases 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  Monoprotic = 1 proton (1 H atom) H 2 CO 3  diprotic = 2 protons (2 H atoms) H 3 PO 4  triprotic =3 protons (3H atoms)

10. 8-10 Brønsted-Lowry Acids and 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 can function as either an acid or a base.

11. 8-11 Brønsted-Lowry Acids and 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. One of the most important amphiprotic substances 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, CH 3 COOH, for example, gives up only one proton  the hydrogen atom attached to the O atom.

12. 8-12 Brønsted-Lowry Acids and 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).

13. 8-13 Acids and their Conjugate Bases

14. 8-14 Acid and Base Strength Strong acid: Strong acid: An acid that is dissociated 100 % in water. One that reacts completely or almost completely with water to form H 3 O + ions.

15. 8-15 Acid and Base Strength Strong base: Strong base: One that reacts completely or almost completely with water to form OH - ions. (dissociated 100 % in water

16. 8-16 Here are the six most common strong acids and the four most common strong bases:

17. 8-17 Acid and Base Strength Weak acid: Weak acid: A substance that dissociates only partially in water to produce H 3 O + ions before equilibrium is reached Acetic acid, for example, is a weak acid. In water, only 4 out every 1000 molecules are converted to acetate ions: Hydrofluoric acid is also only partially ionized

18. 8-18 Acid and Base Strength Weak base: Weak base: A substance that only partially reacts with water to produce OH - ions. Ammonia, for example, is a weak base:

19. 8-19 Acid Strength vs. Concentration Acid strength refers to the degree to which an acid dissociates in water to form hydronium ions (H 3 O + )  Strong acids dissociate 100%, weak acids only slightly  Just because an acid is weak doesn’t mean it won’t do damage.  Vinegar a weak acid, but if it is concentrated, it is very corrosive.  Even low concentrations of strong acids can be corrosive.

20. 8-20 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?

21. 8-21 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 and equilibrium lies on the side of the weaker acid and weaker base.

22. 8-22 Acid-Base Equilibria Identify the two acids and bases, and their relative strengths. The position of this equilibrium lies to the right.

23. 8-23 Acid-Base Equilibria Example: Example: Predict the position of equilibrium for this acid- base reaction:

24. 8-24 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.

25. 8-25 Acid Ionization Constants When a weak acid, HA, dissolves in water The equilibrium constant expression, K eq, for its 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

26. 8-26 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 (next screen), 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.

27. 8-27 K a and pK a for Some Weak Acids

28. 8-28 Properties of Acids and Bases Neutralization: Acids and bases react with each other in a process called neutralization; these reactions are discussed in Section 8.9. 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.

29. 8-29 Properties of Acids and 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 follows. Omitting spectator ions gives this net ionic equation.

30. 8-30 Properties of Acids and Bases Reaction with metal oxides Strong acids react with metal oxides to give water plus a salt.

31. 8-31 Properties of Acids and 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.

32. 8-32 Properties of Acids and 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.

33. 8-33 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.

34. 8-34 Self-Ionization of Water ion product of water, K w Combining 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:

35. 8-35 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:

36. 8-36 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.

37. 8-37 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.

38. 8-38 pH and pOH pH of some common materials

39. 8-39 Acid-Base Titrations Titration: Titration: An analytical procedure in which a solute in a solution of known concentration reacts according to 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. 8-40 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. equivalence point 3. There must be a clear-cut change in a measurable property at the equivalence point (when the reagents have combined exactly). 4. We must have accurate measurements of the amount of each reactant.

41. 8-41 Acid-Base Titrations Table 8.2 Some acid-base indicators

42. 8-42 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 equivalence point of the titration. Requirement 4: Requirement 4: We use volumetric glassware.

43. 8-43 Acid-Base Titrations Experimental measurements Doing the calculations

44. 8-44 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 is an acid or base “shock absorber.” A pH buffer is commonly referred to simply as 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.

45. 8-45 pH Buffers How does an acetate buffer resist 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:

46. 8-46 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.

47. 8-47 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 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.

48. 8-48 pH Buffers Buffer capacity Buffer capacity: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.

49. 8-49 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.

50. 8-50 Henderson-Hasselbalch Eq. Henderson-Hasselbalch equation: Henderson-Hasselbalch equation: A mathematical relationship between: pH pK a of the weak acid, HA The concentrations of HA and its conjugate base A -. It is derived in the following way: taking the logarithm of this equation gives

51. 8-51 Henderson-Hasselbalch Eq. 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:

52. 8-52 Henderson-Hasselbalch Eq. Example: Solution 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? Solution The equilibrium we are dealing with and its pK a are Substituting these values in the H-H equation gives

53. 8-53 Biochemical Buffers The original laboratory buffers for use in biochemical studies 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 severe limitations: ---they often change their pH too much if the solution is diluted or the temperature is changed. ---they often permeate cells in solution thereby changing the chemistry in the interior of the cells. To overcome these short comings, N.E. Good developed a series of buffers that consist of zwitterions, molecules that to not readily permeate cell membranes.

54. 8-54 Biochemical Buffers Typical Good buffers are:

55. 8-55 Chapter 8 Acids and Bases End Chapter 8