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Acids and bases in Inorganic Chemistry

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

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Acids and bases in Inorganic Chemistry

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

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

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

  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 Hydronium ion Write an expression for Keq (the equilibrium constant)

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

  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 (Ka) is defined.

  11. pKa The strength of an acid can be measured by the value of Ka, or alternatively pKa where “p” means -log10

  12. Polyprotic acids Acid Formula Ka pKa phosphoric acid H3PO4 K1 = 7.1 x 10-3 2.12 H2PO4- K2 = 6.2 x 10-8 7.21 HPO42- K3 = 4.6 x 10-13 12.34

  13. Polyprotic acids Acid Formula Ka pKa phosphoric acid H3PO4 K1 = 7.1 x 10-3 2.12 H2PO4- K2 = 6.2 x 10-8 7.21 HPO42- K3 = 4.6 x 10-13 12.34

  14. Predicting the acidity of oxo-acids See Shriver & Atkins p.120 - 122 • Pauling’s rules: • For a formula EOp(OH)q, • Ka = 105p-8 which gives: pKa = 8-5p • If q > 1, successive Ka values for additional ionisations are 10-5 smaller (or pKa 5 units larger) Example: H2SO4 or SO2(OH)2 p=2 and q=2 K1 = 105x2-8 = 102 and pKa = 8 - (5 x 2) = -2 K2 = 102 x 10-5 = 10-3 (Experimental value 10-1.9)

  15. Predicting the acidity of oxo-acids • Pauling’s rules: • For a formula EOp(OH)q, • Ka = 105p-8 which gives: pKa = 8-5p • If q > 1, successive Ka values for additional ionisations are 10-5 smaller (or pKa 5 units larger) p=1 and q=2 Work out K1 and K2 for H2SO3 (or SO(OH)2) K1 = 10(5 x 1)-8 = 10-3 (Experimental value = 1.2 x 10-2) K2 = 10-3 x 10-5 = 10-8 (Experimental value = 6.6 x 10-3)

  16. 0 of 136 Predicting the acidity of oxo-acids p=1 and q=2 Work out K1 and K2 for H2SO3 (or SO(OH)2) K1 = 10(5 x 1)-8 = 10-3 K2 = 10-3 x 10-5 = 10-8 • I got it right • I got it wrong, but now understand it • I’m not sure about this

  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 Conjugate Acid Acid Base Conjugate Base Conjugate Acid Acid Base Conjugate Base

  19. 0 of 136 Which is the conjugate base? Base Conjugate Acid NH3 + H2O NH4+ + OH- • NH4+ • OH- • NH3 • H2O Conjugate Base Acid

  20. 0 of 136 Which is the correct term for describing the role of water in this equilibrium? Base Conjugate Acid NH4+ + H2O H3O+ + NH3 • Conjugate acid • Acid • Base • Conjugate base Acid Conjugate Base

  21. 0 of 136 Which is the conjugate acid? Base Conjugate Acid OEt- + CH3COOH EtOH + CH3COO- • OEt- • CH3COOH • EtOH • CH3COO- Acid Conjugate Base

  22. Acidity of oxides in water Classifications of oxides: Acidic: e.g. SO3 + H2O H+ + HSO4- Basic: e.g. CaO + H2O Ca2+ + 2 OH- Amphoteric: in acid: Al2O3 + 6 H3O+ + 3 H2O 2 [Al(OH2)6]3+ and in base: Al2O3 + 2 OH- + 3 H2O 2 [Al(OH)4]- See Shriver & Atkins p.122 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

  23. Acidity of oxides in water Amphoteric: in acid: Al2O3 + 6 H3O+ + 3 H2O 2 [Al(OH2)6]3+ and in base: Al2O3 + 2 OH- + 3 H2O 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: 2H2O H3O+ + OH- An acid is a species which increases the concentration of H3O+, while a base is a species which increases the concentration of OH-. Also, species which react with the solvent to give changes in H3O+ 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- + H2O  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. 0 of 136 Liquid Ammonia The self-ionisation reaction is: 2NH3 NH4+ + NH2- • I got it right • I got it wrong, but now understand it • I don’t get it

  27. Liquid Ammonia An acid will increase the concentration of the cationic species The self-ionisation reaction is: 2NH3 NH4+ + NH2- A base will increase the concentration of the anionic species Ammonium salts (e.g. NH4Cl) are acids because they increase the concentration of NH4+. Alkali metal amides (e.g. NaNH2) are bases because they increase the concentration of NH2-. Species which can protonate H2O to form H3O+ (e.g. HCl) will protonate NH3 to give NH4+.

  28. The self-ionisation reaction is: 3HF H2F+ + HF2- 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. H2SO4 + HF  H2F+ + HSO4- But, some species which are only weakly acidic in water can act as bases in HF: e.g. CH3COOH + 2HF  CH3COOH2+ + HF2-

  29. The self-ionisation reaction is: 3HF  H2F+ + HF2- Liquid HF An acid will increase the concentration of the cationic species The self-ionisation reaction is: 3HF H2F+ + HF2- A base will increase the concentration of the anionic species Fluoride ion donors (e.g. NaF) become bases: F- + HF  HF2- Fluoride ion acceptors (e.g. SbF5) become acids: e.g. SbF5 + 2HF  SbF6- + H2F+

  30. [ ] • -

  31. See Shriver & Atkins p.125 - 134 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

  32. Lewis acids and bases Bronsted acids and bases are also Lewis acids and bases. Consider: NH3 + H2O  NH4+ + OH- NH3 is a Bronsted base as it accepts a proton. At the same time: NH3 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: [ ] • - SbF5 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

  35. Other examples of donor-acceptor complexes

  36. Other examples of donor-acceptor complexes

  37. [AB] a K = [A] x [B] “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?

  38. Hard acids and bases These tend to be small and not easily polarised. e.g Hard acids: H+, Li+, Na+, Mg2+, Ca3+, Al3+, BF3 Hard bases: F-, OH-, O2-, H2O, NH3, NO3- Soft acids and bases These tend to be large and more easily polarised. e.g Soft acids: Cu+, Ag+, Hg+, Hg2+, Cd2+, BH3 Soft bases: H-, I-, CN-, SCN-, CO, PR3

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

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