Today…

1. Today… • Turn in: • Nothing • Our Plan: • Notes – Acids, Bases, & pH • Worksheet #1 • Homework (Write in Planner): • Worksheet #1 due next class

2. Chapter Fifteen Acids, Bases, andAcid–Base Equilibria

3. Crash Course Introduction • http://www.youtube.com/watch?v=LS67vS10O5Y

4. pH pOH pH + pOH = 14 - log[H3O+1] - log[OH-1] 10-pOH 10-pH [OH-1] [H3O+1] Review - The Flow Chart • What is the pH if the [OH-1]is 3.8 x 10-4 M?

5. Sig Figs with pH • Significant figures with pH are unique! • pH has as many places AFTER THE DECIMAL as significant figures in the concentration. • Ex: If you have 3.4 x 10-3 M HCl, the pH is 2.47. • Try it out: The calculator tells you the pH of 1.789 x 10-5 M HNO3 is 4.747389659. What would you write as the pH?

6. Review • See if you remember how to do the math. Complete each problem. • [H3O+1] = 3.28 x 10-4, find pH • pH = 2.5, find [OH-1] • [OH-1] = 1.23 x 10-8, find [H3O+1] • pOH = 2.4, find pH • [H3O+1] = 6.980 x 10-3, find pOH

7. Definitions of Acid & Base

8. The Arrhenius Theory • Arrhenius theory: an acid forms H+(H3O+1) in water; and a base forms OH– in water. • But not all acid–base reactions involve water, and many bases (NH3, carbonates) do not contain any OH–…

9. The Brønsted–Lowry Theory • Brønsted–Lowrytheory defines acids and bases in terms of proton (H+) transfer. • A Brønsted–Lowry acid is a proton donor. • A Brønsted–Lowry base is a proton acceptor. • The conjugate base of an acid is the acid minus the proton it has donated. • The conjugate acid of a base is the base plus the accepted proton.

10. Ionization of HCl H2O is a base in this reaction because it accepts the H+ Conjugate acid of H2O HCl acts as an acid by donating H+ to H2O Conjugate base of HCl

11. Ionization of Ammonia

12. Water Is Amphiprotic H2O acts as an acid when it donates H+, forming the conjugate base ___ H2O acts as a base when it accepts H+, forming the conjugate acid ___ Amphiprotic: Can act as either an acid or as a base

13. Example 15.1 Identify the Brønsted–Lowry acids and bases and their conjugates in: (a) H2S + NH3 NH4+ + HS– (b) OH– + H2PO4– H2O + HPO42–

14. Let’s See if You’ve Got it… • For each acid, write its conjugate base: • HC2H3O2 • H3PO4 • HCO3-1

15. Strength of Acids & Bases

16. Strength • It is important to know whether or not an acid or base is considered strong or weak because it determines how we do calculations. • Today we will only do calculations for strong acids because strong acids COMPLETELY DISSOCIATE. • That means that their concentration is the SAME AS the [H3O+1].

17. Strong Acids • The “strong” acids—HCl, HBr, HI, HNO3, H2SO4, HClO4—are considered “strong” because they ionize completely in water. • The “strong” acids all appear above H3O+ in Table 15.1 on p. 620. • The strong acids are leveled to the same strength—to that of H3O+—when they are placed in water. • MEMORIZE THE 6 MOST COMMON STRONG ACIDS!

18. Periodic Trends in Acid Strength • The greater the tendency for HX (general acid) to transfer a proton to H2O, the more the forward reaction is favored and the stronger the acid. • A factor that makes it easier for the H+ to leave will increase the strength of the acid. • Acid strength is inversely proportional to H—X bond-dissociation energy. Weaker H—X bond => stronger acid. • Acid strength is directly proportional to anion radius. Larger X radius =>stronger acid.

19. Periodic Trends in Acid Strength

20. Strength of Carboxylic Acids • Carboxylic acids all have the –COOH group in common. • Differences in acid strength come from differences in the R group attached to the carboxyl group. • In general, the more that electronegative atoms appear in the R group, the stronger is the acid.

21. Helpful Note • Page 1 of the reference sheet that I gave you does a nice job of summarizing how acid strength increases.

22. Strength of Conjugate Acid–Base Pairs • A stronger acid can donate H+ more readily than a weaker acid. • The stronger an acid, the weaker is its conjugate base. • The stronger a base, the weaker is its conjugate acid. • An acid–base reaction is favored in the direction from the stronger member to the weaker member of each conjugate acid–base pair.

23. … the weaker the conjugate base. The stronger the acid … And the stronger the base … … the weaker the conjugate acid.

24. Example 15.2 Select the stronger acid in each pair: (a) nitrous acid, HNO2, and nitric acid, HNO3 (b) Cl3CCOOH and BrCH2COOH

25. A Little pH Theory

26. Self-Ionization of Water • Even pure water conducts some electricity. This is due to the fact that water self-ionizes: • The equilibrium constant for this process is called the ion product of water (Kw). • At 25 °C, Kw = 1.0 x 10–14 = [H3O+][OH–] • This equilibrium constant is very important because it applies to all aqueous solutions—acids, bases, salts, and nonelectrolytes—not just to pure water. Why did we leave out water?

27. The pH Scale Since pH is a logarithmic scale, cola drinks (pH about 2.5) are about ____ times as acidic as tomatoes (pH about 4.5)

28. Example 15.4 By the method suggested in Figure 15.5, a student determines the pH of milk of magnesia, a suspension of solid magnesium hydroxide in its saturated aqueous solution, and obtains a value of 10.52. What is the molarity of Mg(OH)2 in its saturated aqueous solution? The suspended, undissolved Mg(OH)2(s) does not affect the measurement.

29. 15.5 Example Is the solution 1.0 x 10–8 M HCl acidic, basic, or neutral?

30. That’s enough for one day… • Worksheet #1 is due next class!

31. Today… • Turn in: • Get out WS#1 to check • Our Plan: • Questions on WS#1 • Notes – Weak Acids • Practice from Reference Sheet • Investigation 14 Pre-Lab • Homework (Write in Planner): • Worksheet #2 due next class • Investigation 14 Pre-Lab

32. Quick Review • Given a 5.5 x 10-6 M H2SO4 solution: • Is H2SO4an acid or a base? • Is H2SO4strong or weak? • What is the pH of the solution? • What is the [OH-1]? • What is the conjugate base of H2SO4?

33. Quick Review - Answers • Acid • Strong • 5.26 • 1.8 x 10-9 M • HSO4-1

34. Equilibrium in Solutions of Weak Acids and Bases

35. pH of WEAK ACIDS • Weak acids do not dissociate completely • Some of the acid molecules are converted to hydronium ion, but some are not • To solve for hydronium or hydroxide, we use an equilibrium calculation (RICE Table)

36. pH of Weak Acids • To find pH, you will often have to write the equation for the dissociation of an acid or base, so let’s practice. • HNO2 is added to water • NH3 is added to water • Acetic acid is added to water

37. CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO–(aq) NH3(aq) + H2O(l) NH4+(aq) + OH–(aq) Ka and Kb The equilibrium constant for a Brønsted acid is represented by Ka, and that for a base is represented by Kb. [H3O+][CH3COO–] Ka = ––––––––––––––––– [CH3COOH] Notice that H2O is not included in either equilibrium expression. [NH4+][OH–] Kb = ––––––––––––– [NH3]

38. Acid/Base Strength and Direction of Equilibrium • In Table 15.1, HBr lies above CH3COOH in the acid column. • Since HBr is a stronger acid than CH3COOH, the equilibrium for the reaction: Weaker baseStronger base Weaker acid Stronger acid lies to the left. • We reach the same conclusion by comparing the strengths of the bases (right column of Table 15.1). • CH3COO– lies below Br– ; CH3COO– is the stronger base:

39. Equilibrium in Solutions of Weak Acids and Weak Bases These calculations are similar to the equilibrium calculations performed in Chapter 14. • An equation is written for the reversible reaction. • Data are organized, often in RICE format. • Changes that occur in establishing equilibrium are assessed. • Simplifying assumptions are examined (the “5% rule”). • Equilibrium concentrations, equilibrium constant, etc. are calculated.

40. The 5% Rule • EXPLAIN IT.

42. pH of WEAK ACIDS What is the pH of a solution of 0.1 M acetic acid if the Ka for acetic acid is 1.8 x 10-5? CH3COOH + H2O↔ CH3COO-1+ H3O+1

43. Example 15.6 Ordinary vinegar is approximately 1 M CH3COOH and as shown in Figure 15.6, it has a pH of about 2.4. Calculate the expected pH of 1.00 M CH3COOH(aq), and show that the calculated and measured pH values are in good agreement.

44. Example 15.7 (application of 5% rule) • What is the pH of 0.00200 M CH2ClCOOH (aq)?

45. Example 15.8 What is the pH of 0.500 M NH3(aq)?

46. Example 15.9 The pH of a 0.164 M aqueous solution of dimethylamine is 11.98. What are the values of Kb and pKb? The ionization equation is (CH3)2NH + H2O ↔ (CH3)2NH2+ + OH–Kb = ? DimethylamineDimethylammonium ion

47. Polyprotic Acids • A monoprotic acid has one ionizable H atom per molecule. • A polyprotic acid has more than one ionizable H atom per molecule. • Sulfuric acid, H2SO4 Diprotic • Carbonic acid, H2CO3 Diprotic • Phosphoric acid, H3PO4 Triprotic • The protons of a polyprotic acid dissociate in steps, each step having a value of Ka. • Values of Ka decrease successively for a given polyprotic acid. Ka1 > Ka2 > Ka3 , etc. • Simplifying assumptions may be made in determining the concentration of various species from polyprotic acids.

48. Example 15.11 Calculate the following concentrations in an aqueous solution that is 5.0 M H3PO4: (a) [H3O+] (b) [H2PO4–] (c) [HPO42–] (d) [PO43–]

49. Ions as Acids and Bases • HCl is a strong acid, therefore Cl– is so weakly basic in water that a solution of chloride ions (such as NaCl) is virtually neutral. • Acetic acid, CH3COOH, is a weak acid, so acetate ion, CH3COO–, is significantly basic in water. • A solution of sodium acetate (which dissociates completely into sodium and acetate ions in water) is therefore slightly basic: CH3COO– + H2O ↔CH3COOH + OH–

50. Carbonate Ion as a Base A carbonate ion accepts a proton from water, leaving behind an OH– and making the solution basic.