Copper

This page looks at some aspects of copper chemistry required for UK A-level exams. The chemistry of copper(II) is mainly summarised from elsewhere on the site, with links available to more detailed explanations. The page also covers some simple copper(I) chemistry.

The simplest ion that copper forms in solution is the typical blue hexaaquacopper(II) ion – [Cu(H₂O)₆]²⁺.

Hydroxide ions (from, say, sodium hydroxide solution) remove hydrogen ions from the water ligands attached to the copper ion.

Once a hydrogen ion has been removed from two of the water molecules, you are left with a complex with no charge – a neutral complex. This is insoluble in water and a precipitate is formed.

In the test-tube, the colour change is:

The ammonia acts as both a base and a ligand. With a small amount of ammonia, hydrogen ions are pulled off the hexaaqua ion exactly as in the hydroxide ion case to give the same neutral complex.

That precipitate dissolves if you add an excess of ammonia.

The ammonia replaces water as a ligand to give tetraamminediaquacopper(II) ions. Notice that only 4 of the 6 water molecules are replaced.

The colour changes are:

You simply get a precipitate of what you can think of as copper(II) carbonate.

If you add concentrated hydrochloric acid to a solution containing hexaaquacopper(II) ions, the six water molecules are replaced by four chloride ions.

The reaction taking place is reversible.

Because the reaction is reversible, you get a mixture of colours due to both of the complex ions.

You may find the colour of the tetrachlorocuprate(II) ion variously described as olive-green or yellow. If you add water to the green solution, it returns to the blue colour.

Copper(II) ions oxidise iodide ions to iodine, and in the process are themselves reduced to copper(I) iodide.

The initial mucky brown mixture separates into an off-white precipitate of copper(I) iodide under an iodine solution.

If you pipette a known volume of a solution containing copper(II) ions into a flask, and then add an excess of potassium iodide solution, you get the reaction we have just described.

You can find the amount of iodine liberated by titration with sodium thiosulfate solution.

As the sodium thiosulfate solution is run in from a burette, the colour of the iodine fades. When it is almost all gone, you add some starch solution. This reacts reversibly with iodine to give a deep blue starch-iodine complex which is much easier to see.

You add the last few drops of the sodium thiosulfate solution slowly until the blue colour disappears.

If you trace the reacting proportions through the two equations, you will find that for every 2 moles of copper(II) ions you had to start with, you need 2 moles of sodium thiosulfate solution.

If you know the concentration of the sodium thiosulfate solution, it is easy to calculate the concentration of the copper(II) ions.

Copper(I) chemistry is limited by a reaction which occurs involving simple copper(I) ions in solution. This is a good example of disproportionation – a reaction in which something oxidises and reduces itself.

Copper(I) ions in solution disproportionate to give copper(II) ions and a precipitate of copper.

Any attempt to produce a simple copper(I) compound in solution results in this happening.

For example, if you react copper(I) oxide with hot dilute sulfuric acid, you might expect to get a solution of copper(I) sulfate and water produced. In fact you get a brown precipitate of copper and a blue solution of copper(II) sulfate because of the disproportionation reaction.

We've already seen that copper(I) iodide is produced as an off-white precipitate if you add potassium iodide solution to a solution containing copper(II) ions. The copper(I) iodide is virtually insoluble in water, and so the disproportionation reaction doesn't happen.

Similarly copper(I) chloride can be produced as a white precipitate (reaction described below). Provided this is separated from the solution and dried as quickly as possible, it remains white. In contact with water, though, it slowly turns blue as copper(II) ions are formed.

The disproportionation reaction only occurs with simple copper(I) ions in solution.

Forming copper(I) complexes (other than the one with water as a ligand) also stabalises the copper(I) oxidation state.

For example, both [Cu(NH₃)₂]⁺ and [CuCl₂]^- are copper(I) complexes which don't disproportionate.

The chlorine-containing complex is formed if copper(I) oxide is dissolved in concentrated hydrochloric acid.

You can think of this happening in two stages. First, you get copper(I) chloride formed:

But in the presence of excess chloride ions from the HCl, this reacts to give a stable, soluble copper(I) complex.

You can get the white precipitate of copper(I) chloride (mentioned above) by adding water to this solution. This reverses the last reaction by stripping off the extra chloride ion.