The redox reaction between copper metal and concentrated nitric acid is complex, and depends on the concentration of the nitric acid and the temperature of the reaction mixture. Initially, with the commercial concentrated acid (about 70% HNO₃), the reaction is

Cu + 4H⁺ + 2NO₃^- à Cu²⁺ + 2NO₂ + 2H₂O.

The mixture becomes green owing to complexes of Cu²⁺ and NO₂ being formed, and brown fumes are evolved. When the acid becomes more dilute as the reaction proceeds, reduction to NO is likely instead. None of this affects the eventual outcome, which is a solution of copper(II) aquo ions and nitrate ions together with the aquo ions of the other metals found in the alloy. If this is brass, the diluted reaction product will contain [Cu(H₂O)₆]²⁺ and [Zn(H₂O)₆]²⁺ ions. The latter do not affect the titration.

However, excess of nitrate ions in a strongly acidic solution do; since they are oxidising agents under these conditions, and iodide ions are good reducing agents, any excess of nitric acid will oxidise the iodide ions to iodine and will give an erroneous endpoint:

2NO₃^- + 4H⁺ + 2I^- à I₂ + 2NO₂ + 2H₂O.

The acidic solution is therefore neutralised with sodium carbonate solution. Since this will cause some precipitation of the metal hydroxides, these are redissolved with a dilute solution of ethanoic acid. This gives a solution acidic enough to keep the metal ions in solution, but not so much so that oxidation of iodide ions occurs.

An excess of iodide ions react with copper(II) ions thus:

2Cu²⁺(aq) + 4I^-(aq) à 2CuI (s) + I₂ (aq)

the iodine strictly being present as I₃^-(aq). You might like to look up the standard electrode potentials for the Cu²⁺|Cu⁺ and ½I₂|I^- half-reactions, and then suggest why the reaction occurs.

The reaction of iodine with thiosulphate ions is

I₂ + 2S₂O₃^2- à S₄O₆^2- + 2I^-.

1 Weigh accurately about 2.5 g of brass into a 250 cm³ beaker, and add about 10 cm³ of concentrated nitric acid (very corrosive!). Allow the beaker to stand in the fume cupboard, and add more nitric acid as necessary to enable complete solution of the metal. You should aim to use as little nitric acid as possible; no more than 20 cm³ will be needed.

2 Transfer all of the solution so obtained to a 250 cm³ graduated flask, washing the beaker into the flask with a little pure water.

3 Add 2 mol dm^-3 sodium carbonate solution drop by drop to the solution in the graduated flask until a faint precipitate just persists. Then add 2 mol dm^-3 ethanoic acid solution drop by drop until the precipitate just redissolves. Then add pure water to the mark and mix well.

4 Pipette 25.0 cm³ of this solution into a 250 cm³ conical flask, add about 10 cm³ of a 10% solution of potassium iodide, and swirl. Copper(I) iodide will precipitate. Titrate the liberated iodine with standard 0.100 mol dm^-3 sodium thiosulphate solution until the iodine colour is quite pale (the mixture is of course not clear) and add a little starch indicator. Continue titration until the blue colour is discharged and the mixture is a creamy-white colour. Repeat to obtain three consistent titres.