Inorganic Chemistry / Transition metals and coordination chemistry

Reactions of transition-metal ions in solution

Ligand substitution, precipitation with hydroxide and ammonia, colours, and redox titrations with manganate(VII) and dichromate(VI).

A-Level

In a nutshell

When a transition-metal salt dissolves in water, the metal ion does not float about on its own: it is surrounded by water molecules bonded to it, forming a complex ion such as [\mathrm{Cu(H_2O)_6}]^{2+}, which is pale blue. Adding other substances changes those surroundings in ways you can see. Sodium hydroxide pulls protons off the water ligands and a coloured solid hydroxide drops out; concentrated hydrochloric acid swaps the water for chloride ions and the solution turns yellow-green; ammonia can do both. Because each metal gives its own pattern of colours and precipitates, these reactions are a quick way to identify ions. Transition-metal ions also change oxidation state easily, which is why purple manganate(VII) and orange dichromate(VI) are used to measure how much iron is in an iron tablet.

Ligand substitution

Water ligands can be replaced by other ligands. Ammonia is similar in size to water, so the coordination number stays at six: [\mathrm{Co(H_2O)_6}]^{2+} + 6\mathrm{NH_3} \rightarrow [\mathrm{Co(NH_3)_6}]^{2+} + 6\mathrm{H_2O}. Chloride is larger, so only four fit: [\mathrm{Cu(H_2O)_6}]^{2+} + 4\mathrm{Cl^-} \rightleftharpoons [\mathrm{CuCl_4}]^{2-} + 6\mathrm{H_2O}, turning blue to yellow-green and octahedral to tetrahedral. Replacing monodentate ligands by a multidentate one such as EDTA^{4-} is favoured by entropy, because the number of free particles rises (the chelate effect).

Precipitation with hydroxide and ammonia

Hydrated ions are acidic: [\mathrm{M(H_2O)_6}]^{3+} is more acidic than [\mathrm{M(H_2O)_6}]^{2+} because the smaller, more highly charged ion polarises O–H bonds more strongly. Adding \mathrm{OH^-} (or a little \mathrm{NH_3}, acting as a base) gives neutral hydroxides:

\mathrm{Cu(OH)_2(H_2O)_4} blue; \mathrm{Fe(OH)_2(H_2O)_4} green, darkening as air oxidises it; \mathrm{Fe(OH)_3(H_2O)_3} brown.
\mathrm{Al(OH)_3} is white and amphoteric: it redissolves in excess NaOH.
Excess ammonia dissolves copper hydroxide to give deep blue [\mathrm{Cu(NH_3)_4(H_2O)_2}]^{2+}.

Adding carbonate to \mathrm{M^{3+}} gives the hydroxide and \mathrm{CO_2} gas, not a carbonate, because the solution is acidic enough.

Redox titrations

In acid, \mathrm{MnO_4^-} + 8\mathrm{H^+} + 5e^- \rightarrow \mathrm{Mn^{2+}} + 4\mathrm{H_2O}, so 1 mol manganate(VII) oxidises 5 mol \mathrm{Fe^{2+}}. It is self-indicating: the endpoint is the first permanent pink. Dichromate(VI) gains 6 electrons per ion (\mathrm{Cr_2O_7^{2-}} \rightarrow 2\mathrm{Cr^{3+}}), so the ratio with \mathrm{Fe^{2+}} is 1 : 6; it needs an added indicator. Use dilute sulfuric acid, not hydrochloric, because manganate(VII) would oxidise chloride.

Key points

Ligand substitution with similar-sized ligands (H_2O, NH_3) keeps coordination number 6; larger Cl^- gives four-coordinate tetrahedral complexes.
The chelate effect is driven by an increase in entropy when multidentate ligands release many small ones.
Hydroxide or a little ammonia deprotonates aqua ions to give coloured neutral hydroxide precipitates; M^{3+} aqua ions are more acidic than M^{2+}.
Excess ammonia redissolves copper(II) hydroxide; excess NaOH redissolves amphoteric aluminium hydroxide.
MnO_4^- : Fe^{2+} is 1 : 5 and Cr_2O_7^{2-} : Fe^{2+} is 1 : 6 in acidic titrations.

Check yourself

5 questions

1What is seen when a few drops of sodium hydroxide are added to aqueous copper(II) sulfate?
Hydroxide removes protons from the aqua ligands, giving neutral blue Cu(OH)_2(H_2O)_4, which is insoluble.
2Why does adding concentrated HCl to [Cu(H_2O)_6]^{2+} change the coordination number from 6 to 4?
The larger chloride ligands cannot pack six around Cu^{2+}, so tetrahedral [CuCl_4]^{2-} forms.
3Which statement explains why [Fe(H_2O)_6]^{3+} is more acidic than [Fe(H_2O)_6]^{2+}?
The smaller, more highly charged Fe^{3+} draws electron density from coordinated water, weakening O–H bonds so protons are released more readily.
425.0 cm^3 of Fe^{2+} solution reacts with 20.0 cm^3 of 0.0200 mol dm^{-3} KMnO_4. What is the Fe^{2+} concentration?
Moles MnO_4^- = 4.00 × 10^{-4}; Fe^{2+} is five times this, 2.00 × 10^{-3} mol, in 0.0250 dm^3, giving 0.0800 mol dm^{-3}.
5Why is hydrochloric acid unsuitable for acidifying a manganate(VII) titration?
MnO_4^- is a strong enough oxidant to oxidise Cl^-, so the titre would be too large and the result wrong.

Recent preprints on this topic

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Matched by keyword from ChemRxiv via Crossref. Not peer-reviewed; read critically.