Prof HM Marques C301 717-6737 hmarques@chem.wits.ac.za

1. Chemistry II Inorganic Chemistry Part 2 Prof HM Marques C301 717-6737 hmarques@chem.wits.ac.za Chapters 4 and 5

2. Chapter 4 – Acids and Bases

3. Revise from Chem I Bronsted-Lowry definition of an acid & a base Lewis definition of an acid & a base

4. Ligands Complex ion

5. Lewis base Lewis acid

6. Base Acid Water is amphiprotic (amphoteric) – it can act as either an acid or a base HCl(g) + H2O(l)  H3O+ (aq) + Cl- (aq) NH3(g) + H2O(l)  OH- (aq) + NH4+(aq)

8. (H3O)+(H2O)3 or H9O4+

9. HA(aq) + H2O (l)  H3O+(aq) + A-(aq) Acid dissociation constant Measure of the strength of the acid

10. Acid strength decreases Perchloric acid (HClO4) Ka = 1010pKa = -10 Sulphuric acid (H2SO4) 102 -2 Phosphoric (H3PO4) 7.5 x 10-3 1.92 Hydrocyanic acid (HCN) 4.9 x 10-10 9.31 The higher the pKa, the weaker the acid

12. Henderson-Hasselbalch equation • When [Base] = [Acid], pH = pKa • At any pH, it can be shown (see Tut) that

13. B(aq) + H2O (l)  OH-(aq) + HB+(aq) Basicity constant Measure of the strength of the base The higher the pKb, the weaker the base

14. H2O (l) + H2O (l)  OH-(aq) + H3O+(aq) Autoprotolysis constant of water = 1.00 x 10-14 at 25.0 oC

15. Oxalic acid is a polyprotic acid: (H2Ox) H2Ox + H2O  HOx- + H3O+Ka1 = 5.9 x 10-2 HOx- + H2O  Ox2- + H3O+Ka2 = 6.4 X 10-5 We will calculate the species distribution as a function of pH

16. Fractional abundance, so 0  x 1

17. Similarly,

19. HOx- H2Ox Ox2- Fractional abundance

20. H2Ox 0.93 HOx- Ox2- 0.05 0.02 Species distribution at pH 3 93% HOx– ; 5% Ox2- ; 2% H2Ox

21. The acidity of a proton …then the X-H bond can split heterolytically with X retaining the electron pair… For a proton to be acidic it must be attached to an electronegative element (O, F, Cl, Br, I; to a lesser extent N, S) R––X: + H+ R––X––H If X is electronegative… …and delivering H+ to a Lewis base

22. Organic acids have acidic and non-acidic protons Non-acidic protons because C is not electronegative enough Acidic proton because H bonded to electronegative O

23. AQUA ACIDS Acidic proton on a water molecule coordinated to a metal ion If metal is able to polarise the M-O bond towards it… …that will cause the H-O bond to be polarised towards O… …releasing H+ to be accepted by a Lewis base.

24. AQUA ACIDS Acidic proton on a water molecule coordinated to a metal ion

25.   [Fe(OH2)6]3+ [Fe(OH2)4(OH)2]+ + H+ [Fe(OH2)5(OH)]2+ + H+ Aqua acids are in principle polyprotic acids etc.

26. HYDROXOACIDS Acidic proton on a hydroxyl group bonded or coordinated to a central atom

27. OXOACIDS Acidic proton on a hydroxyl group bonded or coordinated to a central atom on which there is an oxo (=O) group

28. Aqua acids, hydroxoacids and oxoacids may be successive stages of the deprotonation of an aqua acid

29. Aqua acids Central atom in lower oxidation states s block, d block metals in lower (+1, +2, +3) oxidation states, metals on the left of the p block

30.  H+ Strengths of aqua acids This can be rationalised using an electrostatic (ionic) model H radius of the metal ion of charge n+ diameter of coordinated water molecule

31. Work done = [Potential at r = ] – [Potential at r = (r+ + d)]

32. -RT ln K = -nFE(r +d) + -2.303 nFE(r +d) + -log K = -2.303 nFE(r +d) RT + pKa = RT For this process, G = -RT ln K = -nFE RT ln K = nF(E)

33. -2.303 nFE(r +d) + pKa = RT

34. -2.303 nFE(r +d) + pKa = RT

35. pKa should become smaller, and acidity should increase with • an increase in the charge on the ion • a decrease of the size of the ion

36. The term (z/r+) is also known as the ionic potential  Alternatively we could say by adding the charge on the proton and the diameter of water. The term (z2/(r+ + d)) is called the electrostatic parameter

37. pKa gets smaller and acidity increases as electrostatic parameter increases... How good is this electrostatic model (for gas phase) in solution? See Fig. 4.3

38. Model quite poor for many of the d block ions; their acidity is often much higher than predicted by the model Model quite good for s block ions, some d block ions, and the lanthanides How good is this electrostatic model (for gas phase) in solution?

39. Major reason for failure of the model: bonding between the metal and its ligands often not purely ionic, and there is some covalency in metal-ligand bonds. Model is worst for metals like Sn2+ and Hg2+ that form very covalent complexes.

40. The more covalent the M-O bond… …the more the O-H bond is polarised towards O… …the more readily H+ is lost, and the more acidic the compound

41. Oxoacids • Formed by • electronegative elements top right of periodic table (e.g., N, P, S, Cl) • elements in high oxidation state (e.g., Te, I, As, Se)

42. H3PO3, phosphorus acid, is O1(PH)(OH)2 See Table 4.2 General formula: OpE(OH)q

43. Pauling’s rules Empirical rules that allow one to estimate the pKa of oxoacids General formula: OpE(OH)q • pKa 8 – 5p • For p > 1, each successive pKa increases by about 5 units

44. Estimate the pKas of H3AsO4 Example Actual values: 2.3, 6.9, 11.5 *Estimates are usually good to within 1-2 units

45. The strengths of oxoacids can be varied by substitution: CF3 and F are more electron withdrawing than OH; these acids are stronger acids than H2SO4

46. ...polarises the OH bond, making the proton more acidic Electron withdrawl... The strengths of oxoacids can be varied by substitution:

47. The strengths of oxoacids can be varied by substitution: CF3 and F are more electron withdrawing than OH; these acids are stronger acids than H2SO4 NH2 and CH3 are electron donating – hence these acids are weaker acid than H2SO4

48. Oxides Oxides of non-metals are acidic. When dissolved in water, they bind water and release a proton SO3(g) + H2O(l) → H2SO4(aq) → H+(aq) + HSO4–(aq) SO3 is the anhydride of H2SO4

49. When dissolved in water, they accept a proton from water, producing an alkaline solution MgO(s) + H2O(l) → Mg(OH)2(s) Mg2+(aq) + 2OH–(aq) Acidic and basic oxides neutralised each other CaO + SO2→ CaSO3 Acidic oxides are neutralised by bases SO2 + NaOH → Na+HSO3– Oxides of metals are basic.

50. Fig. 4.4 Oxides of the elements in the boundary region between metals and non-metals often show amphoteric behaviour.