ACIDS & BASES

1. ACIDS & BASES

2. Arrhenius Theory 1. in aqueous solution 2. Acid: produces H+ 3. Base: produces OH-

3. Acid HA H3O+ + A- + O O - HA + + A H H H H H

4. HCl(g) + H2O H3O+(aq) + Cl-(aq) CH3COOH(l) + H2O = H3O+(aq) + CH3COO-(aq)

5. careless, but often seen HCl  H++ Cl- CH3COOHH++ CH3COO-

6. Base NaOH(s) Na+(aq) + OH-(aq)

7. Arrhenius acid/base reaction acid + base  H2O + a salt HA + MOH  HOH + MA

8. Monoprotic acid: HCl HCl(aq) + NaOH(aq)  H2O(l) + NaCl(aq) H+ + Cl- + Na+ + OH- H2O + Na+ + Cl- H+ + OH- H2O HCl 

9. Diprotic acid: H2SO4 H2SO4 (aq) + 2NaOH (aq) 2H2O(l) + Na2SO4 (aq) H+ + OH- H2O H2SO4 

10. Triprotic acid: H3PO4Polyprotic H3PO4(aq) + 3NaOH(aq) 3H2O(l) + Na3PO4(aq) H3PO4 + 3 OH- 3 H2O + PO43- H3PO4 

11. Bronsted-Lowry Theory 1. aqueous & nonaqueous solutions 2. Acid: species donating a proton HCl  H+ + Cl- H2SO4 H+ + HSO4- CH3COOH  H++ CH3COO-

12. Bronsted-Lowry Theory 3. Base: species accepting a proton OH- + H+ HOH H2O + H+ H3O+ NH3 + H+ NH4+

13. Conjugate acid-base pairs acid1+base1acid2+ base2 conjugate pairs HF+ HOH

14. Conjugate acid-base pairs acid1+base1acid2+ base2 conjugate pairs HF + HOH H3O+ + F-

15. ALL Arrhenius reactions are Bronsted-Lowry reactions HCl+ NaOH  H2O+ NaCl

16. NOT all Bronsted reactions are Arrhenius reactions CH3COOH + NH3 NH4+ + CH3COO-

17. Amphiprotic = AmphotericCan act as either an acid or a base HCl+ HOH H3O++ Cl- NH3 + HOH  NH4+ + OH- NH3 + OH- NH2- + HOH HOH+ HOHH3O++ OH-

18. ACID STRENGTH Relative ability of a compound to donate a proton Base strength is considered a result, not a cause

19. REVIEW Strong acid 100% dissociation Weak acid <100% dissociation Notice this is NOT related to concentration

20. Electronegativity is the most significant factor influencing the strength of acids & bases

21. HF > HCl > HBr > HI as acids in non-aqueous solvents, or as pure gases

22. Look at difference in electronegativities 2.1 H - F 4.0 2.1 H - Cl 3.0 2.1 H - Br 2.8 2.1 H - I 2.5

23. Most “ionic” is the most acidic Nonpolar Polar Ionic ED 0 ED 1.7 ED 4.0

24. However, as acids in aqueous solution HF < HCl = HBr = HI

25. 2.1 H - O 3.5 competition! 2.1 H - F 4.0 2.1 H - Cl 3.0 2.1 H - Br 2.8 2.1 H - I 2.5

26. Is methane acidic as a gas or in aqueous solution? 2.1 H - C 2.5

27. The strength of oxy-acids are also dependent on electronegativity.

28. Oxy-acids and bases have the same fundamental structure

29. NaOH: Na - O - H 0.9 3.5 2.1 HClO: Cl - O - H 3.0 3.5 2.1

30. In water, the more “ionic” bond dissociates, forming the acid or base

31. NaOH: Na - O - H 0.9 3.5 2.1 HClO: Cl - O - H 3.0 3.5 2.1

32. Are alcohols acids or bases? C - O - H 2.5 3.5 2.1

33. Acids in homologous series are of different strength

34. Acid Strength H2SO4 > H2SO3 HNO3 > HNO2 HClO4 > HClO3 > HClO2 > HClO

35. Structurally H2SO4 = O2S(OH)2 H2SO3 = OS(OH)2

36. Need to examine formal charge of central atom.

37. Acid Strength CH3COOH> CH3CH2OH CF3COOH > CH3COOH

38. Need to examine inductive effect of neighboring atoms.

39. pH pK Ka , Kb , Kw

40. 2H2OH3O++ OH-

41. Keq [H2O]2 = [H3O+][OH-] Kw = [H3O+ ][OH-] where Kw (25oC ) = 1 x 10-14

42. in a neutral solution [H3O+] = [OH-] 1 x 10-14 = [H3O+]2 = [OH-]2 [H3O+] = [OH-] = 1 x 10-7

43. pX = -log X pK = -log K pH = -log [H3O+] pOH = -log [OH-]

44. leveling effect of H2O limits [H3O+] & [OH-] to that controlled by H2O

45. upper limit [H3O+ ] = 1 lower limit [H3O+ ] = 1 x 10-14

46. pH scale acid neutralbase 7 14 0 highest [H3O+] on left lowest [H3O+] on right

47. [H3O+] and [OH-] must be considered together

48. Kw = [H3O+][OH-] -log Kw = -log {[H3O+][OH-]} -log Kw = {-log [H3O+]} + {-log[OH-]}

49. pKw = pH + pOH but Kw = 1 x 10-14 14 = pH + pOH

50. Relationship between conjugate acids & bases HA + H2OH3O++ A- A- + H2O HA + OH-