1. Acids and bases in Inorganic Chemistry By the way: You will be allowed to bring molecular modelling kits into exams. You can find a link to the retailer’s website on Blackboard.

2. What am I talking about? 1.An acid 2.A base 3.A buffer 4.An indicator “It can be described as a proton acceptor” 0 of 136

3. What am I talking about? 1.An acid 2.A base 3.A buffer 4.An indicator “This is a substance which changes colour at a particular pH value” 0 of 136

4. What am I talking about? 1.An acid 2.A base 3.A buffer 4.An indicator “It can be described as a proton donor” 0 of 136

5. Some definitions... Svante Arrhenius acids produce H + ions in aqueous solutions bases produce OH - ions in aqueous solutions water required, so only allows for aqueous solutions only protic acids are allowed; required to produce hydrogen ions only hydroxide bases are allowed

6. Some definitions... Johannes Nicolaus Brønsted - Thomas Martin Lowry acids are proton donors bases are proton acceptors aqueous solutions are permissible bases besides hydroxides are permissible only protic acids are allowed

7. Some definitions... Gilbert Newton Lewis acids are electron pair acceptors bases are electron pair donors least restrictive of acid-base definitions

8. Brønsted-Lowry acids and bases Write an expression for K eq (the equilibrium constant) Hydronium ion

9. Brønsted-Lowry acids and bases A. I got it right B. I got it wrong, but now understand it C. I’m not sure about this 0 of 136

10. Brønsted-Lowry acids and bases As the concentration of water remains relatively constant at 55.56 moldm -3 for dilute solutions of acids, and the effect of the equilibrium is negligible, a new equilibrium constant, the acidity constant (K a ) is defined.

11. pKa The strength of an acid can be measured by the value of K a, or alternatively pK a where “p” means -log 10

12. Polyprotic acids Acid Formula K a pK a phosphoric acid H 3 PO 4 K 1 = 7.1 x 10 -3 2.12 H 2 PO 4 - K 2 = 6.2 x 10 -8 7.21 HPO 4 2- K 3 = 4.6 x 10 -13 12.34

13. Polyprotic acids Acid Formula K a pK a phosphoric acid H 3 PO 4 K 1 = 7.1 x 10 -3 2.12 H 2 PO 4 - K 2 = 6.2 x 10 -8 7.21 HPO 4 2- K 3 = 4.6 x 10 -13 12.34

14. Predicting the acidity of oxo- acids Pauling’s rules: For a formula EO p (OH) q, K a = 10 5p-8 which gives: pK a = 8-5p If q > 1, successive K a values for additional ionisations are 10 -5 smaller (or pK a 5 units larger) Example: H 2 SO 4 or SO 2 (OH) 2 p=2 and q=2 K 1 = 10 5x2-8 = 10 2 and pK a = 8 - (5 x 2) = -2 K 2 = 10 2 x 10 -5 = 10 -3 (Experimental value 10 -1.9 ) See Shriver & Atkins p.120 - 122

15. Predicting the acidity of oxo- acids Pauling’s rules: For a formula EO p (OH) q, K a = 10 5p-8 which gives: pK a = 8-5p If q > 1, successive K a values for additional ionisations are 10 -5 smaller (or pK a 5 units larger) Work out K 1 and K 2 for H 2 SO 3 (or SO(OH) 2 ) K 1 = 10 (5 x 1)-8 = 10 -3 (Experimental value = 1.2 x 10 -2 ) K 2 = 10 -3 x 10 -5 = 10 -8 (Experimental value = 6.6 x 10 -3 ) p=1 and q=2

16. Work out K 1 and K 2 for H 2 SO 3 (or SO(OH) 2 ) K 1 = 10 (5 x 1)-8 = 10 -3 K 2 = 10 -3 x 10 -5 = 10 -8 p=1 and q=2 Predicting the acidity of oxo-acids A. I got it right B. I got it wrong, but now understand it C. I’m not sure about this 0 of 136

17. Conjugate Acids and Bases Consider the following system...when hydrofluoric acid is added to water, there is an increase in the hydrogen ion concentration... But the reaction does not go to completion. The four species are in equilibrium - so which is the acid and which is the base?

18. Conjugate Acids and Bases AcidBase Conjugate Base Conjugate Acid AcidBase Conjugate Base Conjugate Acid

19. Which is the conjugate base? A. NH 4 + B. OH - C. NH 3 D. H 2 O NH 3 + H 2 O NH 4 + + OH - Conjugate Base Conjugate Acid Acid Base 0 of 136

20. Which is the correct term for describing the role of water in this equilibrium? NH 4 + + H 2 O H 3 O + + NH 3 Conjugate Base Conjugate Acid Acid Base A. Conjugate acid B. Acid C. Base D. Conjugate base 0 of 136

21. Which is the conjugate acid? OEt - + CH 3 COOH EtOH + CH 3 COO - Conjugate Base Conjugate Acid Acid A. OEt - B. CH 3 COOH C. EtOH D. CH 3 COO - Base 0 of 136

22. Acidity of oxides in water Classifications of oxides: Acidic: e.g. SO 3 + H 2 O H + + HSO 4 - Basic: e.g. CaO + H 2 O Ca 2+ + 2 OH - Amphoteric: in acid: Al 2 O 3 + 6 H 3 O + + 3 H 2 O 2 [Al(OH 2 ) 6 ] 3+ and in base: Al 2 O 3 + 2 OH - + 3 H 2 O 2 [Al(OH) 4 ] - The elements in circles have amphoteric oxides even in their highest oxidation states. The elements in boxes have acidic oxides in their maximum oxidation states and amphoteric oxides in their lower oxidation states See Shriver & Atkins p.122

23. Acidity of oxides in water Amphoteric: in acid: Al 2 O 3 + 6 H 3 O + + 3 H 2 O 2 [Al(OH 2 ) 6 ] 3+ and in base: Al 2 O 3 + 2 OH - + 3 H 2 O 2 [Al(OH) 4 ] -

24. Acids and Bases in Other Solvent Systems 1) Solvent systems similar to water, in which self-ionisation involves proton transfer. In the case of water: 2H 2 O H 3 O + + OH - An acid is a species which increases the concentration of H 3 O +, while a base is a species which increases the concentration of OH -. Also, species which react with the solvent to give changes in H 3 O + or OH - concentration can similarly be classified as an acids or bases e.g the OEt - ion in NaOEt would be a base, because: OEt - + H 2 O  EtOH + OH -

25. Acids and Bases in Other Solvent Systems Liquid ammonia is another solvent which, like water, can undergo self-ionisation. Write an equation for the self-ionisation of ammonia.

26. A. I got it right B. I got it wrong, but now understand it C. I don’t get it Liquid Ammonia The self-ionisation reaction is: 2NH 3 NH 4 + + NH 2 - 0 of 136

27. Liquid Ammonia Ammonium salts (e.g. NH 4 Cl) are acids because they increase the concentration of NH 4 +. Alkali metal amides (e.g. NaNH 2 ) are bases because they increase the concentration of NH 2 -. Species which can protonate H 2 O to form H 3 O + (e.g. HCl) will protonate NH 3 to give NH 4 +. An acid will increase the concentration of the cationic species A base will increase the concentration of the anionic species The self-ionisation reaction is: 2NH 3 NH 4 + + NH 2 -

28. The self-ionisation reaction is: 3HF H 2 F + + HF 2 - Liquid HF An acid will increase the concentration of the cationic species A base will increase the concentration of the anionic species Some species which are acids in water will also be acids in liquid HF: e.g. H 2 SO 4 + HF  H 2 F + + HSO 4 - But, some species which are only weakly acidic in water can act as bases in HF: e.g. CH 3 COOH + 2HF  CH 3 COOH 2 + + HF 2 -

29. The self-ionisation reaction is: 3HF  H 2 F + + HF 2 - Liquid HF An acid will increase the concentration of the cationic species A base will increase the concentration of the anionic species Fluoride ion donors (e.g. NaF) become bases: F - + HF  HF 2 - Fluoride ion acceptors (e.g. SbF 5 ) become acids: e.g. SbF 5 + 2HF  SbF 6 - + H 2 F + The self-ionisation reaction is: 3HF H 2 F + + HF 2 -

30. [] -

31. Lewis acids and bases When we talk about Lewis acidity and basicity, the thing being donated or accepted is an electron pair (not a proton!). A Lewis acid is an electron pair acceptor. A Lewis base is an electron pair donor. This definition extends the range of interactions which can be described as “acid-base” reactions to include a wide range of donor-acceptor behaviour. Examples you know about: Dative covalent (co-ordinate) bonding Transition metal – ligand complexes See Shriver & Atkins p.125 - 134

32. Lewis acids and bases Bronsted acids and bases are also Lewis acids and bases. Consider: NH 3 + H 2 O  NH 4 + + OH - NH 3 is a Bronsted base as it accepts a proton. At the same time: NH 3 is a Lewis base as it donates a pair of electrons.

33. Lewis acids and bases Bronsted acids and bases are also Lewis acids and bases. Consider: - [] SbF 5 is a Lewis acid as it accepts a pair of electrons from a fluoride ion. The fluoride ion donates a pair of electrons so is a Lewis base.

34. Other examples of donor-acceptor complexes

37. “Hard” and “soft” acids and bases The strength of the donor-acceptor interaction in Lewis acid-base complexes can be measured quantitatively in terms of the stability constant K. A + B AB High K indicates a strong interaction Low K indicates a weak interaction Why does the interaction change for different complexes? K = [AB] a [A] x [B]

38. Hard acids and bases These tend to be small and not easily polarised. e.g Hard acids: H +, Li +, Na +, Mg 2+, Ca 3+, Al 3+, BF 3 Hard bases: F -, OH -, O 2-, H 2 O, NH 3, NO 3 - Soft acids and bases These tend to be large and more easily polarised. e.g Soft acids: Cu +, Ag +, Hg +, Hg 2+, Cd 2+, BH 3 Soft bases: H -, I -, CN -, SCN -, CO, PR 3

39. Hard acids and bases Remember this: Hard acids tend to bind to hard bases (high ionic character) Soft acids tend to bind to soft bases (high covalent character)

40. Hard acids and bases Hard – Hard  High K Soft – Soft  High K Hard – Soft  Low K A + B AB K = [AB] a [A] x [B]