1. Ch. 15: Acids and Bases Dr. Namphol Sinkaset Chem 201: General Chemistry II

2. I. Chapter Outline I.Introduction II.Acid/Base Definitions III.Acid Strength, K a, and p Functions IV.Weak Acid Equilibria V.Basic Solutions VI.Ions and Salts – Acid or Base? VII.Polyprotic Acids VIII.Acid Strength & Molecular Structure

3. I. Acids and Bases Acid-base reactions are a huge category within chemistry. Acids have a sour taste, dissolve many materials, and neutralize bases. Bases have a bitter taste, a slippery feel, and the ability to neutralize acids. We have found many uses for acids and bases.

4. I. Common Acids

5. I. Some Acid Structures

6. I. Common Bases

7. II. Acid/Base Definitions There are actually many different definitions for what acids and bases are. Which definition we use depends on what we are trying to describe. We cover 3 definitions:  Arrhenius  Brønsted-Lowry  Lewis

8. II. Arrhenius Definitions An Arrhenius acid is a substance that produces H + (H 3 O + ) ions in aqueous solution. An Arrhenius base is a substance that produces OH - ions in aqueous solution. Under this definition, acids and bases neutralize each other via H + (aq) + OH - (aq)  H 2 O (l).

9. II. Arrhenius Acid/Base

10. II. Brønsted-Lowry Definitions These definitions are more widely applicable. A Brønsted-Lowry acid is a proton donor. A Brønsted-Lowry base is a proton acceptor. Recall that H + is a naked proton. Notice these definitions do not depend on aqueous solutions.

11. II. Brønsted-Lowry Acids/Bases Under this definition, acids and bases always occur in pairs.  HCl (aq) + H 2 O (l)  H 3 O (aq) + Cl - (aq)  NH 3(aq) + H 2 O (l)  NH 4 + (aq) + OH - (aq) Substances that can act as either an acid or base are called amphoteric.

12. II. Conjugate Acid/Base Pairs In an acid/base reaction, the acid becomes a base, and the base becomes an acid. Conjugate acid-base pairs are two substances related to each other by a transfer of a proton.

13. II. Practice Problem Identify the conjugate acid-base pairs in the reactions below.  HNO 3(aq) + H 2 O (l)  H 3 O + (aq) + NO 3 - (aq)  C 5 H 5 N (aq) + H 2 O (l)  C 5 H 5 NH + (aq) + OH - (aq)

14. II. Lewis Definitions The Lewis definitions of acids and bases are very broad and widely applicable. The focus is on electron pairs. A Lewis acid is an electron pair acceptor. A Lewis base is an electron pair donor.

15. II. Lewis Acids The Lewis definition of an acid allows substances w/out H to be classified as acids. A Lewis acid has an empty orbital (or can rearrange e-’s to create an empty orbital) that can accept an e- pair. e.g. BF 3 + :NH 3  F 3 B:NH 3

16. II. Example Lewis Acids

17. II. Cationic Lewis Acids Small, highly charged metal ions have lost e-’s and therefore can act as Lewis acids.

18. III. Acid Strength Recall the difference between a strong electrolyte and a weak electrolyte. Strong and weak acids are defined accordingly.  A strong acid completely ionizes in solution.  A weak acid partially ionizes in solution. Thus, acid strength depends on the equilibrium:  HA (aq) + H 2 O (l)  H 3 O + (aq) + A - (aq)

19. III. Strong Acids Strong acids ionize completely; thus, a 1.25 M HCl solution will have [H 3 O + ] = 1.25 M and [Cl - ] = 1.25 M.

20. III. Weak Acids Weak acids establish an equilibrium between their ionized and nonionized forms.

21. III. Weak Acids The strength of a weak acid depends on the attraction between H + and A -, i.e. how badly HA wants to stay together. In general, the stronger the acid, the weaker the conjugate base. HA (aq) + H 2 O (l)  H 3 O + (aq) + A - (aq)

22. III. Weak Acid Strength Is there a way to characterize the strength of a weak acid? Since weak acids set up an equilibrium, we can calculate equilibrium constants to see how much they dissociate. Of course, higher values of K would mean more dissociation, more H 3 O +, and thus a stronger acid.

23. III. Acid Dissociation Constants Since these are acid equilibria, we give the constants a new name: acid dissociation constants, K a. For the generic weak acid equilibrium HA (aq) + H 2 O (l)  H 3 O + (aq) + A - (aq) :

24. III. Sample Problem Write the acid dissociation equation for (CH 3 ) 2 NH 2 + and its K a expression.

25. III. Sample Problem Which acid is stronger: HF with K a of 3.5 x 10 -4 or HCN with K a of 4.9 x 10 -10 ? Which conjugate base is stronger?

26. III. Water Reacts w/ Itself! We said before that water is amphoteric; it can also react with itself in an acid/base reaction. This is called an autoionization reaction.

27. III. Water Autoionization As you see, this is an equilibrium, so we can write an equilibrium expression which we call the ion product constant for water, K w. K w is sometimes called the water dissociation constant. At 25 °C, K w = 1.00 x 10 -14.

28. III. Pure Water In pure water, the autoionization is the only source of H + and OH -. When [H + ] = [OH - ], the solution is neutral.

29. III. Acidic/Basic Solutions In an acidic solution, additional H 3 O + ions exist, increasing [H 3 O + ]. In a basic solution, additional OH - ions exist, increasing [OH - ]. However, in all aqueous solutions, the product of hydronium and hydroxide concentrations always equals K w.

30. III. Sample Problem Calculate the [H 3 O + ] concentration of a solution that has [OH - ] = 1.5 x 10 -2 M at 25 °C. Is the solution acidic or basic?

31. III. The pH Scale pH is simply another way to specify the acidity of a solution.  pH = -log [H 3 O + ]  pH 7 is basic. The log function has the following significant figure rule:

32. III. The p Function The p function can be used on anything. A p function is the mathematical operation of taking the negative log.  pOH = -log [OH - ]  pK a = -log K a  pNa + = -log [Na + ]

33. III. pH/pOH Relationship

34. III. Sample Problems a)Calculate the H 3 O + concentration for a solution that has a pH of 8.37. b)Calculate the H 3 O + concentration for a solution with a pOH of 3.42. c)Calculate the pK a for phenol which has a K a = 1.3 x 10 -13.

35. IV. Weak Acid Equilibria Equilibrium problems involving weak acids are similar to calculations we’ve done before. Problems look for either K a, pH, or equilibrium concentrations. In these problems, we ignore the contribution from the autoionization of water. However, they are more complicated because information can be given in a variety of different ways!

36. IV. Sample Problem In a 0.100 M solution of lactic acid, the pH is 2.44 at 25 °C. Calculate the K a and pK a of lactic acid at this temperature.

37. IV. Sample Problem Nicotinic acid is a weak acid w/ pK a = 4.85. Calculate [H 3 O + ] and the pH of a 0.050 M solution of nicotinic acid.