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3.5.5 Reactions of Inorganic Compounds in Aqueous Solution - Acidity or hydrolysis reactions

Specification

Students should:
  • understand the equilibria:

[M(H2O)6]2+ + H2O M(H2O)5(OH)]+ + H3O+

[M(H2O)6]3+ + H2O [M(H2O)5(OH)]2+ + H3O+

  • to show generation of acidic solutions with M3+, and very weakly acidic solutions with M2+
  • understand that the acidity of [M(H2O)6]3+ is greater than that of [M(H2O)6]2+ in terms of the (charge/size ratio) of the metal ion
  • be able to describe and explain the simple test-tube reactions of M2+ (aq) ions, limited to M = Fe, Co and Cu, and of M3+ (aq) ions, limited to M = Al, Cr and Fe, with the bases OH-, NH3 and CO32-
  • know that MCO3 is formed but that M2(CO3)3 is not formed
  • know that some metal hydroxides show amphoteric character by dissolving in both acids and bases (e.g. hydroxides of Al3+ and Cr3+)
  • know the equilibrium reaction 2CrO42- + 2H+ Cr2O72- + H2O

The acidity of hexaaqua ions

As explained in the previous section, the acidity of the hexaaqua salts is due to polarisation of the hydrogen-oxygen bonds, driven by the high charge density of the transition metal ion.

Polarisation of aqua ligands

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Test tube reactions of some transition metal aqueous ions

Iron(II)

This exists in solution as the hexaaquairon(II) complex ion, [Fe(H2O)6]2+.

It is a pale green solution, which slowly oxidises on standing in air to give a yellow/brown solution of hexaaquairon(III) ions.

[IMG]

Reaction with hydroxide ions

The pale green solution forms a gelatinous dark green precipitate with both sodium hydroxide and ammonia solution.

[Fe(H2O)6]2+ + 2OH- [Fe(H2O)4(OH)2] + 2H2O

The precipitate forms a brown colouration near the surface where the oxygen of the air oxidises the iron(II) hydroxide to iron(III) hydroxide.

After the precipitation of iron(II) hydroxide, the charge density of the iron(II) ion is not sufficent to allow the hydroxide ions to abstract more protons from the water molecules. There is no change on adding excess hydroxide ions.

Reaction with ammonia

Ammonia solution provides both ammonia molecules, that can act as ligands, and hydroxide ions, which can abstract protons from polarised water molecules. The ammonia does not coordinate to the iron(II) ion in this case and all that is seen is the dark green gelatinous precipitate of iron(II) hydroxide. There is no change on adding excess ammonia solution.

[Fe(H2O)6]2+ + 2OH- [Fe(H2O)4(OH)2] + 2H2O

Reaction with carbonate ions

Iron carbonate is precipitated as a light green powdery precipitate when a soluble carbonate is added to an iron(II) solution. The weak acidity of the iron(II) solution is insufficient to decompose the carbonate ion.

[Fe(H2O)6]2+ + CO32- FeCO3 + 6H2O

Cobalt(II)

This exists in solution as the hexaaquacobalt(II) complex ion, [Co(H2O)6]2+.

It is a red solution.

Reaction with hydroxide ions

Cobalt(II) solutions produce a blue cobalt(II) hydroxide precipitate when sodium hydroxide solution is added.

[Co(H2O)6]2+ + 2OH- [Co(H2O)4(OH)2] + 2H2O

There is no change on adding excess sodium hydroxide solution. On standing, the cobalt(II) hydroxide precipitate turns grey at the surface as the cobalt(II) ions become oxidised to cobalt(III) ions.

Reaction with ammonia

The ammonia solution provides both ammonia molecules and hydroxide ions. The hydroxide ions produce a blue precipitate with cobalt(II) solutions.

[Co(H2O)6]2+ + 2OH- [Co(H2O)4(OH)2] + 2H2O

On addition of excess ammonia ...

Reaction with carbonate ions

When a soluble carbonate is added to a solution of cobalt(II) ions a precipitate of cobalt(II) carbonate is formed.

[Co(H2O)6]2+ + CO32- CoCO3 + 6H2O

Copper(II)

This exists in solution as the hexaaquacopper(II) complex ion, [Cu(H2O)6]2+.

It is a blue solution.

Reaction with hydroxide ions

Copper(II) salt solutions deposit a light blue precipitate when hydroxide ions are added:

[Cu(H2O)6]2+ + 2OH- [Cu(H2O)4(OH)2] + 2H2O

There is no change on addition of excess hydroxide ions.

Reaction with ammonia

Copper(II) salt solutions deposit a light blue precipitate when ammonia solution is added. The ammonia solution provides ammonia molecules and hydroxide ions, both of which can abstract hydrogen ions from the water ligands:

[Cu(H2O)6]2+ + 2OH- [Cu(H2O)4(OH)2] + 2H2O

On addition of excess ammonia a ligand exchange reaction takes place and four of the water ligands are removed from the complex. The light blue precipitate dissolves to form a deep blue solution of dihydroxytetramminecopper(II) ions:

[Cu(H2O)4(OH)2] + 4NH3 [Cu(NH3)4(H2O)2]2+ + 2OH-

The two remaining water ligands are much further from the copper ion than the ammonia ligands. For this reason the formula is sometimes written as [Cu(NH3)4]2+ and called the copper(II) tetrammine ion.

Reaction with carbonate ions

The carbonate ions produce a precipitate of light green copper(II) carbonate

[Cu(H2O)6]2+ + CO32- CuCO3 + 6H2O

Aluminium

This exists in solution as the hexaaquaaluminium complex ion, [Al(H2O)6]3+.

It is a colourless solution which is highly acidic by hydrolysis, due to the high charge density of the Al3+ ion.

[Al(H2O)6]3+ [Al(H2O)5OH]2+ + H+

There is also a degree of further hydrolysis:

[Al(H2O)5OH]2+ [Al(H2O)4(OH)2]+ + H+

Reaction with hydroxide ions

Addition of hydroxide ions to the equilibria above pull them to the right hand side and also abstract a third proton forming a neutral species:

[Al(H2O)4(OH)2]+ + OH- [Al(H2O)3(OH)3] + H2O

The aluminium hydroxide appears as a white fine precipitate.

Addition of excess hydroxide ions abstracts a futher proton:

[Al(H2O)3(OH)3] + OH- [Al(H2O)2(OH)4]- + H2O

This is now a complex ion, the aluminate ion, with a negative charge and can dissolve once more. Hence, we see the precipitate redissolving on addition of excess NaOH(aq).

This reaction is indicative of the amphoteric nature of aluminium.

Reaction with ammonia

Ammonia solution provides hydroxide ions, which precipitate aluminium hydroxide. However, on addition of excess ammonia solution the precipitate does not redissolve.

[Al(H2O)4(OH)2]+ + OH- [Al(H2O)3(OH)3] + H2O

Reaction with carbonate ions

The hexaaquaaluminium dissociation produces a solution that is acidic enough to decompose carbonate ions. There is an acid base reaction.

[Al(H2O)6]3+ [Al(H2O)5OH]2+ + H+

CO32- + 2H+ CO2 + H2O

As the carbon dioxide is produced bubbles of gas are seen (effervescence) and eventually, as the carbonate ions strip the hydrogen ions from the complex a precipitate of aluminium hydroxide is formed.

The overall reaction may be written as:

2[Al(H2O)6]3+ + 3CO32- 2[Al(H2O)3(OH)3] + 3CO2 + 3H2O

Chromium(III)

This exists in solution as the hexaaquachromium(III) complex ion, [Cr(H2O)6]3+.

It is a green solution that, like aluminium solutions, is acidic by hydrolysis due to the high charge density of the chromium(III) ion.

Reaction with hydroxide ions

Hydroxide ions react with the available hydrogen ions formed in the dissociation of the hexaaquachromium(III) complex. This produces a green precipitate of chromium(III) hydroxide.

[Cr(H2O)6]2+ + 3OH- [Cr(H2O)3(OH)3] + 3H2O

On addition of excess hydroxide ions the precipitate redissolves as more hydrogen ions are abstracted from the water ligands. This gives complexes of variable formula, with up to six hydroxy complexes:

[Cr(H2O)3(OH)3] + 3OH- [Cr(OH)6]3- + 3H2O

Reaction with ammonia

Ammonia solution provides both ammonia molecules and hydroxide ions capable of abstracting hydrogen ions from the hexaaqua complex. This results in the formation of the hydroxide as a precipitate:

[Cr(H2O)6]2+ + 3OH- [Cr(H2O)3(OH)3] + 3H2O

However, on addition of excess ammonia solution, a ligand exchange reaction can take place resulting in partial redissolving of the hydroxide precipitate:

[Cr(H2O)3(OH)3] + 6NH3 [Cr(NH3)3]3+ + 3H2O + 3OH-

Reaction with carbonate ions

The strongly acidic chromium(III) complex can decompose carbonate ions:

[Crl(H2O)6]3+ [Cr(H2O)5OH]2+ + H+

CO32- + 2H+ CO2 + H2O

The eventually all of the chromium(III) gets precipitated as the hydroxide. The overall reaction:

2[Cr(H2O)6]3+ + 3CO32- 2[Cr(H2O)3(OH)3] + 3CO2 + 3H2O

Iron(III)

This exists in solution as the hexaaquairon(III) complex ion, [Fe(H2O)6]3+.

It is a yellow solution.

Reaction with hydroxide ions

The hydroxide ions give a dark brown gelatinous precipitate of iron(III) hydroxide:

[Fe(H2O)6]3+ + 3OH- [Fe(H2O)3(OH)3] + 3H2O

There is no further change on addition of excess hydroxide ions.

Reaction with ammonia

Ammonia provides hydroxide ions and ammonia molecules, both of which can abstract hydrogen ions from the water ligands, forming a precipitate of iron(III) hydroxide:

[Fe(H2O)6]3+ + 3OH- [Fe(H2O)3(OH)3] + 3H2O

On addition of excess ammonia there is no change.

Reaction with carbonate ions

The strongly acidic iron(III) complex can decompose carbonate ions:

[Fe(H2O)6]3+ [Fe(H2O)5OH]2+ + H+

CO32- + 2H+ CO2 + H2O

The eventually all of the iron(III) gets precipitated as the hydroxide. The overall reaction:

2[Fe(H2O)6]3+ + 3CO32- 2[Fe(H2O)3(OH)3] + 3CO2 + 3H2O

Addition of excess carbonate ions causes no further change.

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Summary of ionic reactions

With metal 2+ salt solution

The initial products of adding hydroxide ions, ammonia solution to metal complexes in the 2+ oxidation state are always the same. The metal hydroxide is formed as a precipitate. The above reactions can be written in a much simpler form:

M2+(aq) + 2OH-(aq) M(OH)2(s)

On addition of excess hydroxide there is no further change, but in certain cases, eg copper(II), addition of excess ammonia causes ligand exchange and dissolves the precipitate.

Carbonate ions produce metal(II) carbonate precipitates:

M2+(aq) + CO32-(aq) MCO3(s)

With metal 3+ salt solution

Once again a hydroxide precipitate is formed by all of the ions on addition of hydroxide ions or ammonia solution, but carbonate ions are decomposed by the acidic nature of the hydrolysed metal(III) complexes and so form a hydroxide precipitate.

Addition of excess hydroxide ions redissolves the amphoteric aluminium hydroxide and chromium(III) hydroxide, but has no effect on iron(III) hydroxide.

Addition of excess ammonia solution to the hydroxides partially redissolves chromium(III) hydroxide, but has no effect on aluminium hydroxide or iron(III) hydroxide.

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Chromium(VI) - 4s0 3d0

In the +6 state chromium is a strong oxidising agent, particularly in acidic solution. There are two common series of salts, the chromate(VI) salts and the dichromate(VI) salts. Chromate(VI) is stable in basic solution while dichromate(VI) is stable in acidic solution:

Cr2O72- + OH- 2CrO42- + H+

This is not a redox reaction. The equilibrium can be moved from one side to the other by addition of acid or base. Add base and the OH- ions reacts with the hydrogen ions in the equilibrium above, pulling the equilibrium to the right hand side. Add acid and the extra hydrogen ions push the equilibrium to the left hand side.

The dissolved dichromate(VI) ion is orange and the chromate(VI) ion is yellow.

potassium dichromate(VI)   potassium chromate(VI)

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