Acids and Bases
Bronsted-Lowry Acids and Bases:
- An acid is a proton donor
- A base is a proton acceptor
Lewis Acid and Bases:
- An acid is an electron pair acceptor
- A base is an electron pair donor
- Ligands which form bonds to transition metal ions are acting as Lewis bases.
- The metal ions are acting as Lewis acids.
- All Bronsted-Lowry acids are Lewis bases aswell.
Iron (II) and Iron (III) Aqua Complexes
Solutions of Fe 2+ are not noticably acidic but solutions of Fe3+ are, because...
- Fe3+ ions are smaller and more highly charged so attract electrons more easily.
- In [Fe(H20)6]3+ the Fe strongly attracts electrons from the oxygen in the water ligands, weakening the O-H bond in water.
- This means it will readily release a H+ ion making the solution acidic.
- [Fe(H2O)6]3+ -> [Fe(H2O)5(OH)]2+ + H+
Generally aqua-ions of M3+ are more acidic than M2+.
Reactions of M2+ ions and M3+ ions
If you add a base (such as OH-) it will remove protons from the aqueous complex.
- [M(H2O)6]3+(aq) + OH- --> [M(H2O)5(OH)]2+(aq) + H2O(aq)
- [M(H2O)5(OH)]2+(aq) + OH- --> [M(H2O)4(OH)2]+(aq) H2O(aq)
- [M(H2O)4(OH)2]+(aq) + OH- --> M(H2O)3(OH)3 (s) + H2O(aq)
- [M(H2O)6]2+(aq) + OH- --> [M(H2O)5(OH)]+ +H2O(aq)
- [M(H2O)5(OH)]+(aq) + OH- --> M(H2O)4(OH)2 (s) + H2O(aq)
M(H2O)4(OH)2 is M(OH)2 and is uncharged, insoluble and a precipitate.
Ligand Substitution Reactions
H2O ligands in metal aqua ions can be replaced by other ligands, this is either because...
- the other ligands form stronger co-ordinate bonds (better Lewis bases)
- they are present in higher concentration and an equilibrium is displaced
- Water and ammonia are similar sized and uncharged so ligand exchange occurs without a change in co-ordination number or charge.
- Ammonia is a better ligand as the lone pair electrons is more easily donated due to the nitrogen not being as electronegative as the oxygen.
- Carbonates of 2+ transition metal ions exist while those of 3+ do not.
- [Fe(H2O)6]2+(aq) + CO32- --> FeCO3(aq) + 6H2O(aq)
Test for Iron Ions:
- In dilute solutions the 2+ and 3+ precipitates are hard to tell apart
- Add alkali as the precipitate colours are more obvious
Amphoteric - shows both acidic and basic properties
Examples include aluminium hydroxide and chromium hydroxide:
- Al(H2O)3(OH)3 + HCl -> [Al(H2O)6]3+ + 3Cl-
- Al(H2O)3(OH)3 + OH- -> [Al(OH)4]- + 3H2O
- Cr(H2O)3(OH)3 + 3H3O+ -> [Cr(H2O)6]3+ + 3H2O
- Cr(H2O)3(OH)3 + 3OH- -> [Cr(OH)6]3- + 3H2O