=136. Reactions with Phosphates.=—In this country the expressions
“acid” and “super” phosphates are used interchangeably. A more correct
use of the terms would designate by “acid” the phosphate formed
directly from tricalcium phosphate by the action of sulfuric acid,
while by “super” would be indicated a similar product formed by the
action of free phosphoric acid on the same materials. In Germany the
latter compound is called double phosphate.
The reaction which takes place in the first instance is represented by
the following formula:
3Ca₃(PO₄)₂ + 6H₂SO₄ + 12H₂O = 4H₃PO₄ + Ca₃(PO₄)₂ + 6(CaSO₄·2H₂O);
and 4H₃PO₄ + Ca₃(PO₄)₂ + 3H₂O = 3[CaH₄(PO₄)₂·H₂O].
A simpler form of the reaction is expressed as follows:
Ca₃(PO₄)₂ + 2H₂SO₄ + 5H₂O = CaH₄(PO₄)₂·H₂O + 2[CaSO₄·2H₂O].
If 310 parts, by weight, of fine-ground tricalcium phosphate be mixed
with 196 parts of sulfuric acid and ninety parts of water, and the
resulting jelly be quickly diluted with a large quantity of water, and
filtered, there will be found in the filtrate about three-quarters of
the total phosphoric as free acid. If, however, the jelly, at first,
formed as above, be left to become dry and hard, the filtrate, when the
mass is beaten up with water and filtered, will contain monocalcium
phosphate, CaH₄(PO₄)₂.
If the quantity of sulfuric acid used be not sufficient for complete
decomposition, the dicalcium salt is formed directly according to the
following reaction:
Ca₃(PO₄)₂ + H₂SO₄ + 6H₂O = Ca₂H₂(PO₄)₂·4H₂O + CaSO₄·2H₂O.
This arises, doubtless, by the formation, at first, of the regular
monocalcium salt and the further reaction of this with the tricalcium
compound, as follows:
CaH₄(PO₄)₂ + H₂O + Ca₃(PO₄)₂ + 7H₂O = 2[Ca₂H₂(PO₄)₂·4H₂O].
This reaction represents, theoretically, the so-called reversion of
the phosphoric acid. When there is an excess of sulfuric acid there is
a complete decomposition of the calcium salts with the production of
free phosphoric acid and gypsum. The reaction is represented by the
following formula:
Ca₃(PO₄)₂ + 3H₂SO₄ + 6H₂O = 2H₃PO₄ + 3[CaSO₄·2H₂O].
The crystallized gypsum absorbs the six molecules of water in its
molecular structure.
=137. Reactions with Fluorids.=—Since calcium fluorid is present in
nearly all mineral phosphates, the reactions of this compound must be
taken into consideration in a chemical study of the manufacture of acid
phosphates. When treated with sulfuric acid the first reaction which
takes place consists in the formation of hydrofluoric acid: CaF₂ +
H₂SO₄ = 2HF + CaSO₄. Since, however, there is generally some silica in
reach of the nascent acid, all, or a portion of it, combines at once
with this silica, forming silicon tetrafluorid: 4HF + SiO₂ = 2H₂O +
SiF₄. This compound, however, is decomposed at once in the presence of
water, forming hydrofluosilicic acid: 3SiF₄ + 2H₂O = SiO₂ + 2H₂SiF₆.
The presence of calcium fluorid in natural phosphates is extremely
objectionable from a technical point of view, both on account of the
increased consumption of oil of vitriol which it causes, but also by
reason of the injurious nature of gaseous fluorin compounds produced.
Each 100 pounds of calcium fluorid entails the consumption of 125.6
pounds of sulfuric acid.
=138. Reaction with Carbonates.=—Most mineral phosphates contain
calcium carbonate in varying quantities. This compound is decomposed on
treatment with sulfuric acid according to the reaction: CaCO₃ + H₂SO₄ =
CaSO₄ + H₂O + CO₂. When present in moderate amounts, calcium carbonate
is not an objectionable impurity in natural phosphates intended for
acid phosphate manufacture. The reaction with sulfuric acid which takes
place produces a proper rise in temperature throughout the mass, while
the escaping carbon dioxid permeates and lightens the whole mass,
assisting thus in completing the chemical reaction by leaving the
residual mass porous, and capable of being easily dried and pulverized.
Where large quantities of carbonate in proportion to the phosphate
are present the sulfuric acid used should be dilute enough to furnish
the necessary water of crystallization to the gypsum formed. For each
100 parts, by weight, of calcium carbonate, eighty parts of sulfuric
anhydrid are necessary, or 125 parts of acid of 1.710 specific gravity
= 60° Beaumé.
In some guanos a part of the calcium is found as pyrophosphate, and
this is acted upon by the sulfuric acid in the following way: Ca₂P₂O₇ +
H₂SO₄ = CaH₂P₂O₇ + CaSO₄.
=139. Solution of the Iron and Alumina Compounds.=—Iron may occur in
natural phosphates in many forms. It probably is most frequently met
with as ferric or ferrous phosphate, seldom as ferric oxid, and often
as pyrite, FeS₂. The iron also may sometimes exist as a silicate. The
alumina is found chiefly in combination with phosphoric acid, and as
silicate.
Where a little less sulfuric acid is employed, as is generally the
case, than is necessary for complete solution, the iron phosphate is
attacked as represented below:
3FePO₄ + 3H₂SO₄ = FePO₄·2H₂PO₄ + Fe₂(SO₄)₃.
When an excess of sulfuric acid is employed, the formula is reduced to
the simple one:
2FePO₄ + 3H₂SO₄ = 2H₃PO₄ + Fe₂(SO₄)₃.
A part of the iron sulfate formed reacts with the acid calcium
phosphate present to produce a permanent jelly-like compound, difficult
to dry and handle. As much as two per cent of iron phosphate, however,
may be present without serious interference with the commercial
handling of the product. By using more sulfuric acid as much as four
or five per cent of the iron phosphate can be held in solution. Larger
quantities are very troublesome from a commercial point of view. The
reaction of the ferric sulfate with monocalcium phosphate, is as
follows:
3CaH(PO₄)₂ + Fe₂(SO₄)₃ + 4H₂O = 2(FePO₄·2H₃PO₄·2H₂O) + 3CaSO₄.
Pyrite and the silicates containing iron are not attacked by sulfuric
acid, and these compounds are therefore left, in the final product, in
a harmless state. If the pyritic iron is to be brought into solution
aqua regia should be employed.
With sufficient acid the aluminum phosphate is decomposed with the
formation of aluminum sulfate and free phosphoric acid:
AlPO₄ + 3H₂SO₄ = Al₂(SO₄)₃ + 2H₃PO₄.
=140. Reaction with Magnesium Compounds.=—The mineral phosphates, as a
rule, contain but little magnesia. When present it is probably as an
acid salt, MgHPO₄. Its decomposition takes place in slight deficiency
or excess of sulfuric acid respectively as follows: