Generic Intro to AB chem. (Sect. J)

1. Generic Intro to AB chem. (Sect. J) • Arrhenius theory: an acid forms H+ 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–. • Brønsted–Lowry theory 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.

2. Examples of Conjugate AB pairs

3. Couple of example problems… • Write a balanced chemical equation for the dissociation of each of the following (in H2O) • H2SO4, H2PO4-, • What is the conjugate base of each of the following? • HCO3-, CO32-, OH-, H2PO4-

4. Describing strength of Acids/Bases • Some acids are stronger than others…you know this • Where do we draw the line? • I suppose that depends on what line we draw… • Strong Acids/Bases dissociate COMPLETELY…100% • Use an arrow to designate the reaction • Kc values well above 103 • Weak Acids/Bases dissociate very little (~1% or less) • Use equilibrium arrows • Kc values well below 10-3

5. Ionization of HCl—strong/weak acids 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

6. Ionization of Ammonia This example also illustrates another Acid/Base definition…Lewis AB’s Electron pair acceptor—Lewis Acid; E pair donor—base.

7. Lewis Acid/Base interaction (2) A couple of things worth noting here. Arrow indicates movement of electrons! Helpful hint in Organic C of CO2 electron ‘starved’ due to pull of oxygens Illustrates the reaction of an “oxide” with water. In this case, formation of an acidic oxide (most elements on right form acidic oxides Metal oxides react to form bases (Na2O?) Amphoteric (goes both ways) Amphoteric oxides

8. Water Is Amphiprotic—goes both ways 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

9. 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, andnonelectrolytes—not just to pure water.

10. The pH Scale • Concentration of H3O+ can vary over a wide range in aqueous solution, from about 10 M to about 10–14 M. • A more convenient expression for H3O+ is pH. pH = –log [H3O+] and so [H3O+] = 10–pH • The “negative logarithm” function of pH is so useful that it has been applied to other species and constants. pOH = –log [OH–] and so [OH–] = 10–pOH pKw = –log Kw • At 25 °C, pKw = 14.00 • pKw = pH + pOH = 14.00

11. 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)

12. 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.

13. CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO–(aq) NH3(aq) + H2O(l) NH4+(aq) + OH–(aq) Ka and Kb—YOU’VE SEEN THIS!! The equilibrium constant for a Brønsted acid is represented by Ka, and that for a base is represented by Kb. Nothing more than a simple extension of what you already know!! [H3O+][CH3COO–] Ka = ––––––––––––––––– [CH3COOH] Notice that H2O is not included in either equilibrium expression. [NH4+][OH–] Kb = ––––––––––––– [NH3]

14. What’s happening in solution? • A solution of a weak acid or base… • Mostly acid molecules (or ions, like the phosphate example) • Very small [ ] of H+, and even SMALLER amounts of OH • Molecules or ions in ceaseless dynamic equilibrium

15. More on the acids Note the higher values of pKa. Higher values correspond to weaker acids. pKa’s more commonly used for acid strength

16. A graphic to relate Ka/weak/strong acids

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. • The strong acids are leveled to the same strength—to that of H3O+—when they are placed in water.

18. Weak acids, and how structure determines relative strength… • Yesterday morning, we touched on weak acids/bases • Some stronger than others…why? • Know they’re stronger b/c pKa values differ • We’ll cover the following situations for weak acids • Oxo acids • Binary Hydrides (not given much love in text) • Carboxylic Acids (pay attention—crops up again next year) • Problem—No ONE thing useful for predicting acidity. MUST use all of your knowledge (structure, trends etc)

19. Periodic Trends Binary Hydrides (HnX) • 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. • ANY 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.

20. Periodic Trends in Acid Strength

21. Strength of Oxoacids Acid strength increases with the electronegativity of the central atom, and with the number of terminal oxygen atoms.

22. Other ways to see this…

23. 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.

24. Select the stronger acid in each pair: Problems like this will obviously be on an exam… (a) nitrous acid, HNO2, and nitric acid, HNO3 (b) Cl3CCOOH and BrCH2COOH

25. Weak Bases…Again…not much love • We know many common strong bases (Metal—OH) • Weak bases? Not many…mostly R3N compounds • Why? They act as good Lewis Bases (lone pair on N) • Two main types—R group just a typical organic group • When R group involves a benzene group.

26. Strengths of Amines as Bases • Aromatic amines are much weaker bases than aliphatic amines. • This is due in part to the fact that the p electrons in the benzene ring of an aromatic molecule are delocalized and can involve the nitrogen atom’s lone-pair electrons in the resonance hybrid. • As a result, the lone-pair electrons are much less likely to accept a proton. • Electron-withdrawing groups on the ring further diminish the basicity of aromatic amines relative to aniline.

27. Select the weaker base in each pair: (Good Primer for your Organic Class this fall)

28. Equilibrium in Solutions of Weak Acids and Weak Bases These calculations are similar to the equilibrium calculations performed in Chapter 9. • An equation is written for the reversible reaction. • Data are organized, often in an ICE format. • Changes that occur in establishing equilibrium are assessed. • Simplifying assumptions are examined (the “5% rule”). • Equilibrium concentrations, equilibrium constant, etc. are calculated.

29. What is the pH of 0.00200 M ClCH2COOH(aq)?

30. Several Textbook sample problems What is the pH of 0.500 M NH3(aq)? 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 = ? Dimethylamine Dimethylammonium ion A Conceptual Example Without doing detailed calculations, indicate which solution has the greater [H3O+], 0.030 M HCl or 0.050 M CH3COOH.