As you can see, first, group 2 solution is heated up to its boiling point. This is necessary to eliminate sulfide (S2-) from Na2S which has been used in group 2 analysis. The solubility of hydrogen sulfide (gas) decreases and it leaves the solution. This moves the reaction below to the right:
Why are we doing this? Because we are going to buffer to pH 9. Remind that sulfide ion concentration related to the pH:
Decreasing [H+] the concentration of S2- grows. Therefore, if the pH is acidic the sulfide ions would precipitate all of the Group IV cations, ruining the whole analysis. That's why we are looking for quantitative elimination of the sulfide ion. How can we check that H2S is leaving the pot? Just with a piece of paper soaked of lead acetate. When H2S reaches the paper, this is blackened because of lead sulphide formation.
Preliminary analysis ( ferrous salts presence)
Before group 3 analysis there's something we should check. Let's take 2-3 drops of the solution from group 2 (see the scheme above). We're going to control if is there any Fe 2+ inside.
We can carry out a symple test:
Add 1-2 drops of potassium ferricyanide. Fe 2+ (if present) precipitates immediately ferrous ferricyanide, blue.
If Fe3+ was present instead, the solution will turn brownish → Fe[Fe(CN)6], ferric ferricyanide. But this is unlikely, because usually sulfide oxidizes iron (II) to iron (III).
(E ° Fe3+ / Fe2+ = 0.77 V | E° S°/S2- = - 0.4 V)
- Treatment with concentrated HNO3
If we detect Fe2+ then we should treat the solution from group 2 with 4-5 drops of concentrated HNO3, in order to oxidize Fe2+ to Fe3+. If we do not oxidise the iron (II), this would form later complexes with ammonia like [Fe(NH3)6]2+ and would not be precipitated quantitatively as hydroxide.
3Fe2+ + NO3- + 4H+ 3Fe3+ + NO + 2H2O