2 FeCl3 + 3 Na2CO3 + 6 H2O ⥂
2 Fe(OH)3 ↓ + 3 H2CO3 + 6 NaCl
3 H2CO3 ⇄ 3 H2O + 3 CO2 ↑.
Since the bivalent metal ions are precipitated by sodium carbonate as carbonates and the trivalent ones as hydroxides, the reagent, obviously, cannot be used to separate the two groups. But barium carbonate is so little soluble in water that it will not precipitate manganous, zinc, nickel, cobalious and ferrous carbonates[391] from solutions of their chlorides or nitrates. We have, for instance, ZnCl2 + BaCO3 ↓ ⥃ BaCl2 + ZnCO3. Barium carbonate has, however, the same effect on ferric chloride (exp.) and on the other chlorides of the trivalent group, as has sodium carbonate, i.e. it precipitates their hydroxides. By means of barium carbonate [p195] we can, therefore, precipitate the hydroxides of the aluminium group without precipitating the ions of the zinc group. The separation is carried out in a, practically, neutral medium (free carbonic acid in excess is evolved; barium carbonate alone, when treated with water, is slightly alkaline) and thus avoids the error of facilitating the precipitation of the bivalent metals in the shape of salts of the acidic forms of the trivalent metals, i.e. as aluminates, chromites, and so forth. Manganous salts are liable to oxidation to manganic salts, when exposed to the air, especially in alkaline, neutral or slightly acid solutions, and prolonged exposure of the barium carbonate mixture to the air may result in the precipitation of manganic hydroxide, Mn(OH)3, with the other trivalent hydroxides. Provision is made for its detection in the systematic analysis.
Analysis of the Aluminium Group.
We have, in this instance, the case of a very weak, difficultly soluble acid, aluminium hydroxide, forming a salt with a weak, soluble base, ammonium hydroxide. The conditions determining the solubility of aluminium hydroxide in ammonium hydroxide, as an aluminate NH4AlO2, may be shown as follows: for the acid ionization of aluminium hydroxide, Al(OH)3 ⇄ AlO2− + H+ + H2O (p. [172]); the solubility-product for a saturated solution is [AlO2−] × [H+] = KAc.S.P.. Further, from [H+] × [HO−] = KHOH, we find [H+] = KHOH / [HO−]. Then [AlO2−] = [HO−] × KAc.S.P. / KHOH, which shows that the solubility of aluminium hydroxide, as aluminate, is proportional to the concentration [HO−] of the hydroxide-ion in the solution. For NH4OH we have [NH4+] × [HO−] / ([NH3] + [NH4OH]) = 0.000,018 (p. [161]), and consequently, [HO−] = 0.000,018 × ([NH3] + [NH4OH]) / [NH4+]. Then [HO−] is the smaller, the smaller the excess of ammonium hydroxide used (which is approximately equal to ([NH3] + [NH4OH])) and the greater the concentration [NH4+] of the ammonium-ion, i.e. of the added ammonium salt. The solubility of Al(OH)3, as aluminate, in ammonium hydroxide and ammonium chloride is, therefore, directly proportional to the excess of ammonium hydroxide, and indirectly proportional to the concentration of the ammonium salt present.[392]
The Favorable Conditions for a Maximum Precipitation of an Amphoteric Hydroxide.
[Al3+] = KBas.S.P. × [HO−]−3.
I
For the sake of a certain simplicity in the result, we will, for the moment, consider aluminium hydroxide to ionize as an acid according to Al(HO)3 ⇄ AlO33− + 3 H+, which would resemble the basic ionization. Then we would have [AlO33−] × [H+]3 = K′Ac.S.P., and, using the relation [H+] = KHOH / [HO−], we have
[AlO33−] = [HO−]3 × K′Ac.S.P. × KHOH−3.
II