[Note 1:  If a selection of pure material for analysis is to be made, crystals which are cloudy are to be avoided on account of loss of water of crystallization; and also those which are red, indicating the presence of ferric iron.  If, on the other hand, the value of an average sample of material is desired, it is preferable to grind the whole together, mix thoroughly, and take a sample from the mixture for analysis.]

[Note 2:  When aqueous solutions of ferrous compounds are heated in the air, oxidation of the Fe^{++} ions to Fe^{+++} ions readily occurs in the absence of free acid.  The H^{+} and Oh^{-} ions from water are involved in the oxidation process and the result is, in effect, the formation of some ferric hydroxide which tends to separate.  Moreover, at the boiling temperature, the ferric sulphate produced by the oxidation hydrolyzes in part with the formation of a basic ferric sulphate, which also tends to separate from solution.  The addition of the hydrochloric acid prevents the formation of ferric hydroxide, and so far reduces the ionization of the water that the hydrolysis of the ferric sulphate is also prevented, and no precipitation occurs on heating.]

[Note 3:  The nitric acid, after attaining a moderate strength, oxidizes the Fe^{++} ions to Fe^{+++} ions with the formation of an intermediate nitroso-compound similar in character to that formed in the “ring-test” for nitrates.  The nitric oxide is driven out by heat, and the solution then shows by its color the presence of ferric compounds.  A drop of the oxidized solution should be tested on a watch-glass with potassium ferricyanide, to insure a complete oxidation.  This oxidation of the iron is necessary, since Fe^{++} ions are not completely precipitated by ammonia.

The ionic changes which are involved in this oxidation are perhaps most simply expressed by the equation

3Fe^{++} + no_{3}^{-}+ 4H^{+} —­> 3Fe^{+++} + 2H_{2}O + no,

the H^{+} ions coming from the acid in the solution, in this case either the nitric or the hydrochloric acid.  The full equation on which this is based may be written thus: 

6FeSO_{4} + 2HNO_{3} + 6HCl —­> 2Fe_{2}(so_{4})_{3} + 2FeCl_{3} + 2no + 4H_{2}O,

assuming that only enough nitric acid is added to complete the oxidation.]

[Note 4:  The ferric hydroxide precipitate tends to carry down some sulphuric acid in the form of basic ferric sulphate.  This tendency is lessened if the solution of the iron is added to an excess of Oh^{-} ions from the ammonium hydroxide, since under these conditions immediate and complete precipitation of the ferric hydroxide ensues.  A gradual neutralization with ammonia would result in the local formation of a neutral solution within the liquid, and subsequent deposition of a basic sulphate as a consequence of a local deficiency of Oh^{-} ions from the NH_{4}Oh and a partial hydrolysis of the ferric salt.  Even with this precaution the entire absence of sulphates from the first iron precipitate is not assured.  It is, therefore, redissolved and again thrown down by ammonia.  The organic matter of the filter paper may occasion a partial reduction of the iron during solution, with consequent possibility of incomplete subsequent precipitation with ammonia.  The nitric acid is added to reoxidize this iron.