1.14 What Happened to p K b ?

1. 1.14What Happened to pKb?

2. About pKa and pKb • A separate “basicity constant” Kb is not necessary. • Because of the conjugate relationships in the Brønsted-Lowry approach, we can examine acid-base reactions by relying exclusively on pKa values.

3. N •• H H N •• H Example • Which is the stronger base, ammonia (left) or pyridine (right)? • Recall that the stronger the acid, the weaker the conjugate base. • Therefore, the stronger base is the conjugate of the weaker acid. • Look up the pKa values of the conjugate acids of ammonia and pyridine in Table 1.7.

4. + N H Example H H + weaker acid N H pKa = 9.3 H pKa = 5.2 stronger acid Therefore, ammonia is a stronger base than pyridine

5. 1.15How Structure Affects Acid Strength

6. The Main Ways Structure Affects Acid Strength • The strength of the bond to the atom from which the proton is lost. • The electronegativity of the atom from which the proton is lost. • Changes in electron delocalization on ionization.

7. HF HCl HBr HI pKa 3.1 -3.9 -5.8 -10.4 weakest acid strongest acid strongest H—X bond weakest H—X bond Bond Strength • Bond strength is controlling factor when comparing acidity of hydrogen halides.

8. Bond Strength • Recall that bond strength decreases in a group in going down the periodic table. • Generalization: Bond strength is most important factor when considering acidity of protons bonded to atoms in same group of periodic table (as in HF, HCl, HBr, and HI). • Another example: H2S (pKa = 7.0) is a stronger acid than H2O (pKa = 15.7).

9. The Main Ways Structure Affects Acid Strength • The strength of the bond to the atom from which the proton is lost. • The electronegativity of the atom from which the proton is lost. • Changes in electron delocalization on ionization.

10. Electronegativity • Electronegativity is controlling factor when comparing acidity of protons bonded to atoms in the same row of the periodic table.

11. CH4 NH3 H2O HF pKa 60 36 15.7 3.1 weakest acid strongest acid least electronegative most electronegative Electronegativity

12. R R . . . . . . H H + – . . O + + H A H A O Electronegativity • The equilibrium becomes more favorable as A becomes better able to bear a negative charge. • Another way of looking at it is that H becomes more positive as the atom to which it is attached becomes more electronegative.

13. Bond strength versus Electronegativity • Bond strength is more important when comparing acids in which the proton that is lost is bonded to atoms in the same group of the periodic table. • Electronegativity is more important when comparing acids in which the proton that is lost is bonded to atoms in the same row of the periodic table.

14. pKa In many acids the acidic proton is bonded to oxygen. HO—H 15.7 CH3O—H 15.2 CH3CH2O—H 16 Alcohols (RO—H) resemble water (HO—H) in their acidity. (CH3)2CHO—H 17 (CH3)3CO—H 18 Acidity of Alcohols

15. Acidity of Alcohols • Electronegative substituents can increase the acidity of alcohols by drawing electrons away from the —OH group. CH3CH2OH CF3CH2OH pKa 16 11.3 weaker stronger

16. d+ F H C C F O H F H Inductive Effect • The greater acidity of CF3CH2OH compared to CH3CH2OH is an example of an inductive effect. • Inductive effects arise by polarization of the electron distribution in the bonds between atoms.

17. Electrostatic Potential Maps • The greater positive character of the proton of the OH group of CF3CH2OH compared to CH3CH2OH is apparent in the more blue color in its electrostatic potential map. CH3CH2OH CF3CH2OH

18. O O CH3C CF3C O O H H Another example of the inductive effect pKa 4.7 0.50 weaker stronger

19. The Main Ways Structure Affects Acid Strength • The strength of the bond to the atom from which the proton is lost. • The electronegativity of the atom from which the proton is lost. • Changes in electron delocalization on ionization.

20. R R . . . . . . H H + – . . O + + H A H A O Electron Delocalization • Ionization becomes more favorable if electron delocalization increases in going from right to left in the equation. • Resonance is a convenient way to show electron delocalization.

21. H O •• •• H •• O •• •• •• H O + H O N •• •• – •• + •• •• – O + + H O N O O •• •• •• •• •• •• – H •• Nitric Acid + pKa = -1.4

22. •• O •• – •• •• + O N O •• •• •• – •• Nitric Acid • Nitrate ion is stabilized by electron delocalization.

23. •• – O •• •• •• – + O N •• •• O •• •• •• •• – O O •• •• •• – •• •• •• + + O N O N O O •• •• •• •• •• •• – – •• •• Nitric Acid • Negative charge is shared equally by all three oxygens.

24. •• H O H O C •• – •• + O + H O C •• •• •• H CH3 Acetic Acid •• H O •• •• O + •• •• •• H CH3 pKa = 4.7

25. •• O •• – •• O C •• •• CH3 Acetic Acid • Acetate ion is stabilized by electron delocalization.

26. •• – O •• •• O C •• •• CH3 Acetic Acid • Negative charge is shared equally by both oxygens. •• O •• – •• O C •• •• CH3

27. 1.16Acid-Base Equilibria

28. Generalization • The equilibrium in an acid-base reaction is favorable if the stronger acid is on the left and the weaker acid is on the right. Stronger acid + Stronger base Weaker acid + Weaker base

29. H H – .. . . . . . . . . . . Br .. pKa = -5.8 stronger acid pKa = -1.7 weaker acid H H .. . . Br H O .. Example of a strong acid + O + H + The equilibrium lies to the side of the weaker acid. (To the right)

30. •• •• O O •• •• H H – + O OCCH3 O—H •• •• •• •• H H pKa = 4.7 weaker acid pKa = -1.7 stronger acid •• •• H—OCCH3 •• •• Example of a weak acid + + The equilibrium lies to the side of the weaker acid. (To the left)

31. Important Points • A strong acid is one that is stronger than H3O+.A weak acid is one that is weaker than H3O+. • A strong base is one that is stronger than HO–.A weak base is one that is weaker than HO–. • The strongest acid present in significant quantities when a strong acid is dissolved in water is H3O+. The strongest acid present in significant quantities when a weak acid is dissolved in water is the weak acid itself.

32. •• •• H—O Phenol pKa = 10 stronger acid WaterpKa = 15.7 weaker acid •• •• •• Predicting the Direction of Acid-Base Reactions – •• •• – H—OC6H5 OC6H5 + + H—O—H •• •• •• The equilibrium lies to the side of the weaker acid. (To the right) Phenol is converted to phenoxide ion by reaction with NaOH.

33. O O – •• •• – H—OC6H5 OC6H5 + + HOCO HOCO—H •• •• •• •• •• •• Phenol pKa = 10 weaker acid Carbonic acidpKa = 6.4 stronger acid •• •• Predicting the Direction of Acid-Base Reactions The equilibrium lies to the side of the weaker acid. (To the left) Phenol is not converted to phenoxide ion by reaction with NaHCO3.

34. 1.17Lewis Acids and Lewis Bases

35. Definitions • Arrhenius • An acid ionizes in water to give protons. A base ionizes in water to give hydroxide ions. • Brønsted-Lowry • An acid is a proton donor. A base is a proton acceptor. • Lewis • An acid is an electron pair acceptor. A base is an electron pair donor.

36. Lewis acid + Lewis base – + + + + + A A A A A—B A—B A—B A—B B B B B •• •• •• •• – – + + + – Lewis Acid-Lewis Base Reactions • The Lewis acid and the Lewis base can be either a neutral molecule or an ion.

37. CH2CH3 CH2CH3 + – O •• •• O F3B •• CH2CH3 CH2CH3 Example: Two Neutral Molecules F3B + Lewis acid Lewis base Product is a stable substance. It is a liquid witha boiling point of 126°C. Of the two reactants,BF3 is a gas and CH3CH2OCH2CH3 with a boiling point of 34°C.

38. + H3C—Br + •• •• – – •• •• •• •• •• •• H—O H—O H—OCH3 H—O—H •• •• •• •• •• •• •• •• •• •• •• •• •• •• Lewis base Lewis acid – – Br Br •• •• + H—Br + Example: Ion + Neutral molecule Reaction is classified as a substitution. But noticehow much it resembles a Brønsted acid-base reaction.

39. + H3C—Br + •• •• – – •• •• •• •• •• •• H—O H—O H—OCH3 H—O—H •• •• •• •• •• •• •• •• •• •• •• •• •• •• Lewis base Lewis acid – – Br Br •• •• + H—Br + Example: Ion + Neutral molecule Brønsted acid-base reactions are a subcategory ofLewis acid-Lewis base reactions.

40. End of Chapter 1