1. Ch. 14: Acids and Bases 14.1 Intro to Acids and Bases 14.2 Acid Strength 14.3 pH Scale

2. Acids and Bases: Characteristics Acids: Taste Sour, (Many of our foods are acid), react with metals to form Hydrogen gas. pH <7.0. Bases: Taste Bitter (many poisons in the world are bases, we are very sensitive to bitter tastes). Soaps, many cleaners and drain cleaners are basic. pH > 7.0. AKA: Alkaline

3. Common household substances that contain acids and bases. Vinegar is a dilute solution of acetic acid. Drain cleaners contain strong bases such as sodium hydroxide.

4. Definitions Arrhenius – Acids: produce H+ in aqueous solution – Bases: produces OH- in aqueous solution – very limited Bronsted-Lowry – Acids: H+ donor – Bases: H+ acceptor – more general

5. Video Clip Distinguishing between Acids and Bases

6. Models of Acids and Bases More Brønsted-Lowry:. Uses the Water molecule in the reaction. Conjugate Acid – Base pairs. One of each pair on each side of the reaction. Example: HCl + H 2 O  Cl  + H 3 O + acid basebase acid

7. Dissociation Reactions When dissolving in water, the ions of the acid or base will separate (dissociate). The reaction is written as: HCl H+ + Cl- OR (in the Bronsted-Lowry version) HCl + H 2 O H 3 O + + Cl- Hydronium Ion = H + and H 2 O

8. Conjugate Acid/Base Pairs HA(aq) + H 2 O(l)  H 3 O + (aq) + A  (aq) conj conj conj conj acid 1 base 2 acid 2 base 1 conjugate base: everything that remains of the acid molecule after a proton is lost. conjugate acid: formed when the proton is transferred to the base.

9. Figure 14.1: The reaction of HCl and H 2 O.

11. Figure 14.3: The reaction of NH 3 with HCl to form NH 4 + and Cl-.

12. Acid equilibrium expression =Acid Dissociation Constant, K a Water is both a solvent and a reactant. The concentration of water is much greater than the other species. So the concentration of water is considered a constant [55M]. The water concentration is part of the acid equilibrium constant, K a.

13. Ka Generic For a general reaction: Where A- is any Anion. HA + H 2 O H 3 O + + A- K a = [H 3 O + ] [A-] OR [HA] K a = [H + ] [A-] [HA]

14. Acid dissociation constant equilibrium expression where H+ is removed to form conjugate base so for: HA + H 2 O ⇄ H 3 O + + A -

15. Video Clip Classifying Acids and Bases by Strength.

16. Figure 14.4: Graphic representation of the behavior of acids of different strengths in aqueous solution. (a) A strong acid. (b) A weak acid.

17. Acid Strength 4 Its equilibrium position lies far to the right. (H + + NO 3 - )  Yields a weak conjugate base. (NO 3  ) 4 Ka is considered to be very large (or greater than 10 -1 ) Strong Acid:

18. Acid Strength (continued) 4 Its equilibrium lies far to the left. (CH 3 COOH)  Yields a much stronger (it is relatively strong) conjugate base than water. (CH 3 COO  ) 4 Ka values are less than 10 -1 Weak Acid:

19. Figure 14.6: (a) A strong acid HA is completely ionized in water. (b) A weak acid HB exists mostly as undissociated HB molecules in water. Note that the water molecules are not shown in this figure.

21. Figure 14.5: The relationship of acid strength and conjugate base strength for the reaction

22. Water as an Acid and a Base Water is amphoteric (it can behave either as an acid or a base). H 2 O + H 2 O  H 3 O + + OH  conj conj acid 1 base 2 acid 2 base 1 K w = 1  10  14 at 25°C Notice how small this number is. Very few water molecules will dissociate.

23. Water Pure water is neutral because is has equal numbers of H+ and OH-. In pure water, [H+] = [OH-] = 1 x 10 -7 Kw = [H+] [OH-] = 1.0 x 10 -14 Get used to noticing the powers on the ten.

24. Figure 14.7: Two water molecules react to form H 3 O+ and OH 2.

25. Strong Acids to Memorize All other acids are weak HCl, HBr, HI HNO 3 H 2 SO 4 (first H+) HClO 3, HClO 4, HIO 4

26. Classic and common Weak Acids and Weak Bases Weak Acids: HC 2 H 3 O 2 (CH 3 COOH, acetic/ethanoic) HF Hydrofluoric acid Weak Base: NH 3 (aq) = NH 4 OH Ammonia Ammonium Hydroxide Other amines: CH 3 NH 2 Methylamine CH 3 CH 2 NH 2 Ethylamine

27. Strong Bases to Memorize All other bases are weak NaOH KOH LiOH CsOH RbOH

28. Water equilibrium In acid or base solutions, the water equilibrium responds (Le Chatelier’s) to the addition of H+ or OH- to maintain Kw = 1 x 10 -14 Example: 1.0 x 10 -3 M HCl (strong acid = 100% dissociates) = 1.0 x 10 -3 M H+ Kw = [H+] [OH-] = [1.0 x 10 -3 ] [OH] Kw = 1.0 x 10 -14 = (1.0 x 10 -3 ) [OH] Solve for [OH]

29. Water Calculations You will use Kw = 1.0 x 10 -14 = [H+] [OH-] and other information, to solve for concentrations (Molarity) of all the “species” in a reaction vessel [H+], [OH-], [A-], [H 2 O] and whether there is more H+(acid) or OH- (base) Pg 661, Sample exercise 14.3

31. The pH Scale pH   log[H + ] pH in water ranges from 0 to 14. K w = 1.00  10  14 = [H + ] [OH  ] pK w = 14.00 = pH + pOH As pH rises, pOH falls (sum = 14.00).

32. Figure 14.8: The pH scale and pH values of some common substances.

33. Figure 14.9: pH meters are used to measure acidity.

34. Sig. Figs. With pH Since pH is a log function: the number in front of the decimal represents the power on the ten (powers of ten don’t count for sig figs) So… we turn to the numbers after the decimal (in a pH #) to show the sig. figs. Example: if [H+] = 3.45 x 10 -7 (there are 3 sig figs) pH = 6.462, 3 places after the decimal. The 6 is close to the original power of ten.