In the above tabulations no mention is made of the work of Fresenius, Neubauer, and Luck, on whose researches the citrate method is based, but an examination of their original paper shows that the temperature conditions are not carefully enough controlled to justify us in tabulating their results.[116] An attempt has been made to include in the above tables, work made under well-defined conditions, which will illustrate the various points under consideration. While each authority of value upon the subject is represented, no attempt has been made to include all the work done by any of them. One element that seems to have been generally overlooked in discussing the problem is that nearly all results have been obtained from one-half hour’s treatment of the material. This means simply the study of an incomplete reaction, and one which is interrupted while the solution is very rapidly going on. This, of course, is only clearly brought out by a comparison of long-time and short-time work in the various tables. In the opinion of Huston very much more work will have to be done before it can be assumed that we have any very clear knowledge of this subject, and very likely the final result will be that all kinds of goods cannot be examined by the same method. The fact that half a gram of dicalcium phosphate is instantly soluble in 100 cubic centimeters of citrate solution, at ordinary temperatures, while an equal amount of iron and aluminum phosphate is acted upon very slowly at ordinary temperatures will probably have to be taken into consideration, as well as the fact that dicalcium phosphate is less soluble in hot solutions of ammonium citrate than it is in cold solutions, while the reverse is true of the precipitated iron and aluminum phosphate.

At present, the only conclusion that can be safely drawn from the work, is that it would be unsafe to make any generalization upon the subject until more facts are at hand, except that the present methods are unscientific and, unsatisfactory. As the work progresses, new features present themselves, and in such a way as to show that they must be given careful consideration before drawing any final conclusions in the matter.

127. Arbitrary Determination of Reverted Phosphoric Acid.—The so-called reverted phosphoric acid, that is, the acid insoluble in water and soluble in a solution of ammonium citrate, is the most difficult constituent of commercial fertilizers from the point of view of the scientific analyst. A review of all the standard methods which have been given in the preceding pages for its determination must convince every careful observer that, as a rule, each process is based on arbitrary standards, and can give only concordant results when carried out under strictly unvarying conditions. For this reason there can be no just comparison between the results obtained by different methods, which vary from each other only in slight particulars. When, on the other hand, the processes are radically different, the deviations in data become more pronounced.

In such a condition of affairs the analyst is left to choose between methods. He must be guided in his choice not only by what seems to be the most scientific and accurate process, but also, to a certain extent, by the general practice of his professional brethren. For this country, therefore, it is strongly urged that the methods adopted by the Association of Official Agricultural Chemists, be followed in every detail.

By the phrase “reverted phosphoric acid,” was originally meant an acid once soluble in water, as CaH₄(PO₄)₂, and afterwards changed to a form insoluble in water, but soluble in ammonium citrate as Ca₂H₂(PO₄)₂. But in practice this has never been the true signification of the term. In the manufacture of acid and superphosphates there is formed, more or less of the dicalcium phosphate, either directly or after a time, and this salt which, in no sense can be called reverted, is entirely soluble in ammonium citrate. The iron and aluminum phosphates are also, to a certain degree, soluble in the same reagent. When an acid phosphate, containing various forms of calcium phosphate, is applied to a soil containing iron and alumina, the soluble parts of the compound tend to become fixed by union with those bases, or by precipitation as Ca₂H₂(PO₄)₂. But it is not alone reverted phosphate formed in this way, which the analyst is called on to determine in a fertilizer, although he may have occasion to treat it in soil analysis.

The expression “reverted phosphoric acid,” therefore, in practice not only includes a dicalcium phosphate, which once may have been the monocalcium salt, but also all of that salt originally existing in the superphosphate, and formed directly during its manufacture, as well as any iron and aluminum phosphates present which are soluble in ammonium citrate. The expression “citrate-soluble” is, therefore, to be preferred to “reverted” phosphoric acid.

In the reversion of the phosphoric acid in superphosphates the iron plays a far more important role than the aluminum sulfate. It was formerly supposed that the reversion took place as indicated in the following formula: 2CaH₄(PO₄)₂ + Fe₂O₃ = 2(CaHPO₄, FePO₄) + 3H₂O, while Wagner affirms that the reverted acid compounds consist of varying quantities of ferric oxid, aluminum oxid, phosphorus pentoxid, and calcium oxid, in various states of combination.[117] The more probable reaction is the following: 3CaH₄(PO₄)₂ + Fe₂(SO₄)₃ + 4H₂O = 2(FePO₄, 2H₃PO₄, 2H₂O) + 3CaSO₄. This reaction can be demonstrated by adding to a superphosphate solution one of a ferric salt. In addition to free phosphoric acid, iron phosphate is separated, which gradually passes into an insoluble form by the abstraction of water due to the crystallization of the gypsum. The alumina present in a superphosphate seems to have no direct influence on the process of reversion. Its phosphate salt is not acted on by the acid calcium phosphate. Even when a superphosphate solution is treated with alum no precipitation is produced, except on warming, and this disappears when the mass is again cold.

It is therefore not necessary in the process of manufacture to separate the alumina by digestion with a hot soda-lye before treating the mass with sulfuric acid.

In order to avoid the reversion of the phosphoric acid several plans have been proposed. One of the best is to use a little excess of sulfuric acid in the manufacture. This tends to hold the phosphoric acid in soluble form but is objectionable on account of drying, handling, and shipping the fertilizer. During the digestion, moreover, it is important that the temperature does not rise above 120°. Another method consists in adding to the dissolved rock a quantity of common salt chemically equivalent to its iron content. Ammonium sulfate also helps to hold the phosphoric acid water-soluble.

128. Influence of Movement.—The influence of time and temperature of digestion, and of variations in the composition of the ammonium citrate on the quantity of phosphoric acid dissolved by that reagent has been pointed out. Of great importance also in the process is the character of the movement to which the materials are subjected during the digestion. For this reason various mechanical devices have been constructed to secure uniformity of solution. Inasmuch as the temperature factor must also be faithfully observed, the best of these devices are so arranged as to admit of a uniform motion within a bath of water kept at the desired temperature which, by the Association method, is 65°.

129. Digestion Apparatus for Reverted Phosphates.—The digestion apparatus used by Huston consists of two wheels twenty-five centimeters in diameter, mounted on the same axis, having a clear space of four and one-half centimeters between them.[118] In the periphery of each wheel are cut twelve notches, which are to receive the posts bearing the rings through which the necks of the flasks pass. The posts are held in place by nuts which are screwed down on the faces of the wheel. Should it become necessary to take the apparatus apart, it is only necessary to loosen the nuts and the set screw holding one wheel to the shaft and all the parts can at once be removed. The posts extend ten centimeters beyond the face of the wheels, and the rings are four centimeters in internal diameter. Perforated plates, bearing a cross-bar, and held in place by strong spiral springs attached to the plate and the base of the posts, serve to hold the flasks in place. Each plate has a number stenciled through it for convenience in identifying the flasks when it is time to remove them. Attached to the outside of each post, close to the outer end, is a heavy wire which passes entirely around the apparatus, serving to keep the plates in place after they are removed from the flasks.

The apparatus is mounted on a substantial framework, thirty-six centimeters high and thirty centimeters wide at the base. The space in which the wheel revolves is fourteen centimeters wide. The base bars connecting the two sides are extended seven centimeters beyond one side, and serve for the attachment of lateral bracing. At the top of the framework, at one side, is attached a heavy bar forty-five centimeters long, which serves to carry the cog gearing which transmits the power. The upright shaft carries a cone pulley to provide for varying the speed. The usual speed is two revolutions a minute for the wheel carrying the flasks. The entire apparatus is made of brass. The details of construction are shown in Fig. 8. Round-bottomed flasks are used, and rubber stoppers are held in place by tying or by a special clamp shown at the lower right-hand of the figure.

When high temperatures are used, the plates and flasks are handled by the hooks shown at the left and right-hand upper corners of the figure.

When any other than room temperature is desired, the whole apparatus is immersed in water contained in the large galvanized tank forming the back-ground of the figures. The tank is seventy-five centimeters long, seventy-five centimeters high, and thirty centimeters wide. At one end, near the top, is an extension to provide space for heating the fluid in the flasks before introducing the solid in such cases as may be desired.

The apparatus is held in place by angle irons soldered to the bottom of the tank and a brace resting against the upright bar bearing the gear-wheels.

The water in the tank is heated by injecting steam, or by burners under the tank. As the tank holds about 300 pounds of water the work is not subject to sudden changes of temperature, and little trouble has been experienced in raising and maintaining the temperature of the water, especially when steam is used.

An electric motor, or a small water-motor with only a very moderate head of water, will furnish ample power.

130. Comparison of Results.—The following data show the results obtained by the digester as compared with those furnished by the official method, temperature and time of digestion being the same in each instance.

Ammonium Citrate Solution on Phosphates.

In comparing duplicates, the results from the use of the digester are found to be subject to less variation than those from the usual method.

131. Huston’s Mechanical Stirrer.—The stirring apparatus shown in Fig. 9 differs from those which have heretofore come into use, in requiring but a single belt to drive all the stirring rods, and in having all the parts protected from the laboratory fumes.[119] The details of the belt system are shown in the small diagram in the lower central part of the figure. The apparatus is mounted on a substantial wooden box, 200 centimeters long, thirty centimeters high, and eighteen centimeters wide. The driving pulleys, ten centimeters in diameter, are enclosed in the upper part of the case. The shafts on which these pulleys are mounted extend through the bottom of the enclosing box and carry a wooden disk, eleven centimeters in diameter, to prevent particles of foreign matter from falling into the beakers. The shafts extend two centimeters below these disks, and to the end of the shafts the bent stirring rods are attached by rubber tubing.

The board forming the support of the driving pulleys is extended two centimeters in front of the apparatus, and in this extension twelve notches are cut, in which are held the corks carrying the tubes which contain the solution to be used in precipitating the material in the beakers.

The ends of these tubes are drawn out to a fine point so as to deliver the liquid at the rate of about one drop per second.

The front of the apparatus is hinged and permits the whole to be closed when not in use, or during the precipitation.

The apparatus has proven extremely satisfactory in the precipitation of ammonium magnesium phosphate. The precipitate is very crystalline, and where the stirring is continued for some minutes, after the magnesia solution has all been added, no amorphous precipitate is observed on longer standing.

132. The Citrate Method Applied to Samples with Small Content of Phosphoric Acid.—It is well established that the citrate method does not give satisfactory results when applied to samples containing small percentages of phosphoric acid, especially when these are of an organic nature, as for instance, cottonseed cake-meal. In this laboratory attempts have been made to remedy this defect in the process so as to render the use of the method possible even in such cases.[120] Satisfactory results have been obtained by adding to the solution of the cake-meal a definite volume of a phosphate solution of known strength. Solutions of ordinary mineral phosphates are preferred for this purpose. The following example will show the application of the modified method:

In a sample of cake-meal, (cottonseed cake and castor pomace) the content of phosphoric acid obtained by the molybdate method, was 2.52 per cent.

Determined directly by the citrate method, the following data were obtained:

Allowing to stand thirty hours after adding magnesia mixture, 1.08 and 1.53 per cent in duplicates.

Allowing to stand seventy-two hours after adding magnesia mixture, 2.17 and 2.30 per cent in duplicates.

In each case fifty cubic centimeters of the solution were taken, representing half a gram of the sample.

In another series of determinations twenty-five cubic centimeters of the sample were mixed with an equal volume of a mineral phosphate solution, the value of which had been previously determined by both the molybdic and citrate methods. The fifty cubic centimeters thus obtained represented a quarter of a gram each of the cake-meal and mineral phosphates. The filtration followed eighteen hours after adding the magnesia mixture. The following data show the results of the determinations:

It is thus demonstrated that the citrate method can be applied with safety even to the determination of the phosphoric acid in organic compounds where the quantity present is less than three per cent. It is further shown that solutions of mineral phosphates varying in content of phosphoric acid from fifteen to thirty-two per cent may be safely used for increasing the content of that acid to the proper degree for complete precipitation. In cases where organic matters are present they should be destroyed by moist combustion with sulfuric acid as in the determination of nitrogen to be described in the next part.

133. Direct Precipitation of the Citrate-Soluble Phosphoric Acid.—The direct determination of citrate-soluble phosphoric acid by effecting the precipitation by means of magnesia mixture in the solution obtained from the ammonium citrate digestion, has been practiced for many years by numbers of European chemists, and the process has even obtained a place in the official methods of some European countries. Various objections have been urged, however, against the general employment of this method in fertilizer analysis on account of the inaccuracies in the results obtained in certain cases, and it has, therefore, been used to but a very limited extent in this country. Since it is impracticable to effect the precipitation with ammonium molybdate in the presence of citric acid the previous elimination or destruction of this substance has been recognized as essential to the execution of a process involving the separation of the phosphoric acid as phosphomolybdate.

It is evident from the data cited in the preceding paragraph, that great accuracy may be secured in this process by adding a sufficient quantity of a solution of a mineral phosphate and proceeding by the citrate method.

Ross has also proposed to estimate the acid soluble in ammonium citrate directly by first destroying the organic matter by moist combustion with sulfuric acid.[121] He recommends the following process:

After completion of the thirty minutes’ digestion of the sample with citrate solution, twenty-five cubic centimeters are filtered at once into a dry vessel. If the liquid be filtered directly into a dry burette, twenty-five cubic centimeters can be readily transferred to another vessel without dilution. After cooling, run twenty-five cubic centimeters of the solution into a digestion flask of 250-300 cubic centimeters capacity, add about fifteen cubic centimeters of concentrated sulfuric acid and place the flask on a piece of wire gauze over a moderately brisk flame; in about eight minutes the contents of the flask commence to darken and foaming begins, but this will occasion no trouble, if an extremely high, or a very low flame be avoided. In about twelve minutes the foaming ceases and the liquid in the flask appears quite black; about one grain of mercuric oxid is now added and the digestion is continued over a brisk flame. The operation can be completed in less than half an hour with ease, and in many cases, twenty-five minutes. After cooling, the contents of the flask are washed into a beaker, ammonia is added in slight excess, the solution is acidified with nitric, and after the addition of fifteen grams of ammonium nitrate, the process is conducted as usual.

In case as large an aliquot as fifty cubic centimeters of the original filtrate be used, ten cubic centimeters of sulfuric acid are added, and the digestion is conducted in a flask of 300-500 cubic centimeters capacity; after the liquid has blackened and foaming has progressed to a considerable extent, the flask is removed from the flame, fifteen cubic centimeters more of sulfuric acid are added, and the flask and contents are heated at a moderate temperature for two or three minutes; the mercuric oxid is then added and the operation completed as before described.

Following are some of the advantages offered by the method described:

(1) It dispenses with the necessity of the execution of the frequently tedious operation of bringing upon the filter and washing the residue from the ammonium citrate digestion, while the ignition of this residue together with the subsequent digestion with acid and filtration are also avoided.

(2) It affords a means for the direct estimation of that form of phosphoric acid which, together with the water-soluble, constitutes the available phosphoric acid, thus enabling the latter to be determined by making only two estimations.

(3) In connection with the advantages above mentioned it permits of a considerable saving of time, as well as of labor required in manipulation.

In addition to the tests with mercuric oxid, both potassium nitrate and potassium sulfate were used in the digestion to facilitate oxidation. With the former, several additions of the salt were necessary to secure a satisfactory digestion, and even then the time required was longer than with the mercury or mercuric oxid digestion. With potassium sulfate, the excessive foaming which took place interfered greatly with the execution of the digestion process.

134. Availability of Phosphatic Fertilizers.—There is perhaps no one question more frequently put to analysts by practical farmers than the one relating to the availability of fertilizing materials. The object of the manufacturer should be to secure each of the valuable ingredients of his goods in the most useful form. The ideal form in which phosphoric acid should come to the soil is one soluble in water. Even in localities where heavy rains may abound, there is not much danger of loss of soluble acid by percolation. As has before been indicated, the soluble acid tends to become fixed in all normal soils, and to remain in a state accessible to the rootlets of plants, and yet free from danger of leaching. For this reason, by most agronomists, the water-soluble acid is not regarded as more available than that portion insoluble in water, yet soluble in ammonium citrate.

In many of the States the statutes, or custom, prescribe that only the water and citrate-soluble acid shall be reckoned as available, the insoluble residue being allowed no place in the estimates of value. In many instances such a custom may lead to considerable error, as in the case of finely ground bones and some forms of soft and easily decomposable tricalcium phosphates. There are also, on the markets, phosphates composed largely of iron and aluminum salts, and these appear to have an available value often in excess of the quantities thereof soluble in ammonium citrate.

As a rule the apatites, when reduced to a fine powder and applied to the soil, are the least available of the natural phosphates. Next in order come the land rock and pebble phosphates which, in most soils, have only a limited availability. The soft fine-ground phosphates, especially in soils rich in humus, have an agricultural value, almost, if not quite equal to a similar amount of acid in the acid phosphates. Fine-ground bones also tend to give up their phosphoric acid with a considerable degree of readiness in most soils. Natural iron and aluminum phosphates, have also, as a rule, a high degree of availability. In each case the analyst must consider all the factors of the case before rendering a decision. Not only the relative solubility of the different components of the offered fertilizer in different menstrua must be taken into consideration, but also the character of the soil to which it is to be applied, the time of application, and the crop to be grown. By a diligent study of these conditions the analyst may, in the end, reach an accurate judgment of the merits of the sample.

135. Direct Weighing of the Molybdenum Precipitate.—It has already been stated that many attempts have, been made to determine the phosphoric acid by direct weighing as well as by titration, as in the Pemberton method. The point of prime importance in such a direct determination is to secure an ammonium phosphomolybdate mixture of constant composition. Unless this can be done no direct method, either volumetric or gravimetric, can give reliable results. Hanamann[122] proposes to secure this constant composition by varying somewhat the composition of the molybdate mixture and precipitating the phosphoric acid under definite conditions. The molybdate solution employed is prepared as follows:

The precipitation of the phosphoric acid is conducted in the cold with constant stirring. It is complete in half an hour. The ammonium phosphomolybdate is washed with a solution of ammonium nitrate and then with dilute nitric acid, dried, and ignited at less than a red heat. It should then have a bluish-black color throughout. Such a body contains 4.018 per cent of phosphoric anhydrid.

Twenty-five cubic centimeters of a sodium phosphate solution containing fifty milligrams of phosphoric acid, treated as above, gave a bluish-black precipitate weighing 1.249 grams, which, multiplied by 0.041018, equaled 50.018 milligrams of phosphorus pentoxid. The method should be tried on phosphates of various kinds and contents of phosphorus pentoxid before a definite judgment of its merits is formed.

CHEMISTRY OF THE MANUFACTURE
OF SUPERPHOSPHATES.

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:

2MgHP₄ + H₂SO₄ + 2H₂O = [MgH₄(PO₄)₂·2H₂O] + MgSO₄
and    MgHPO₄ + H₂SO₄ = H₃PO₄ + MgSO₄.

The magnesia, when in the form of oxid, is capable of producing a reversion of the monocalcium phosphate, as is shown below:

CaH₄(PO₄)₂ + MgO = CaMgH₂(PO₄)₂ + H₂O.

One part by weight of magnesia can render three and one-half parts of soluble monocalcium phosphate insoluble.

141. Determination of Quantity of Sulfuric Acid Necessary for Solution of a Mineral Phosphate.—The theoretical quantity of sulfuric acid required for the proper treatment of any phosphate may be calculated from its chemical analysis and by the formulas and reactions already given. For the experimental determination the method of Rümpler may be followed.[123]

Twenty grams of the fine phosphate are placed in a liter flask with a greater quantity of accurately measured sulfuric acid than is necessary for complete solution. The acid should have a specific gravity of 1.455 or 45° B. The mixture is allowed to stand for two hours at 50°. It is then cooled, the flask filled with water to the mark, well shaken, and the contents filtered. Fifty cubic centimeters of the filtrate are treated with tenth normal soda-lye until basic phosphate begins to separate. The excess of acid used is then calculated. Example: Twenty grams of phosphate containing 28.3 per cent of phosphoric acid, 10.0 per cent of calcium carbonate, 5.5 per cent of calcium fluorid, and 2.4 per cent of calcium chlorid were treated as above with sixteen cubic centimeters of sulfuric acid containing 10.24 grams of sulfur trioxid. In titrating fifty cubic centimeters of the filtrate obtained as described above, 10.4 cubic centimeters of tenth normal soda-lye were used, equivalent to 0.0416 gram of sulfur trioxid. Then 10.24 × 50 ÷ 1000 = 0.5120 = total sulfur trioxid in fifty cubic centimeters of the filtrate, and 0.5120 - 0.0416 = 0.4704 gram, the amount of sulfur trioxid consumed in the decomposition.

Therefore the sulfur trioxid required for decomposition is 47.04 per cent of the weight of the phosphate employed. One hundred parts of the phosphate would therefore require 47.04 parts of sulfur trioxid = to 73.6 parts of sulfuric acid of 1.710 specific gravity or 92.1 parts of 1.530 specific gravity.

A more convenient method than the one mentioned above consists in treating a small quantity of the phosphate, from one-half to one kilogram, in the laboratory, or fifty kilograms in a lead box, just as would be practiced on a large scale. A few tests with these small quantities, followed by drying and grinding will reveal to the skilled operator the approximate quantity and strength of sulfuric acid to be used in each case. The quantities of sulfuric acid as determined by calculation from analyses and by actual laboratory tests agree fairly well in most instances. There is, however, sometimes a marked disagreement. The general rule of practice is to use always an amount of sulfuric acid sufficient to produce and maintain water-soluble phosphoric acid in the fertilizer, but the sulfuric acid must not be used in such quantity as to interfere with the subsequent drying, grinding, and marketing of the acid phosphate.

For convenience the following table may be used for calculating the quantity of oil of vitriol needed for each unit of weight of material noted:

One Part by Weight of Each Substance Below Requires:

142. Phosphoric Acid Superphosphates.—If a mineral phosphate be decomposed by free phosphoric in place of sulfuric acid the resulting compound will contain about three times as much available phosphoric acid as is found in the ordinary acid phosphate. The reaction takes place according to the following formulas:

(1)  Ca₃(PO₄)₂ + 4H₃PO₄ + 3H₂O = 3[CAH₄(PO₄)₂·H₂O].

(2) Ca₃(PO₄)₃ + 2H₃PO₄ + 12H₂O = 3[Ca₂H₂(PO₄)₂·4H₂O].

In each case the water in the final product is probably united as crystal water with the calcium salts produced. The monocalcium salt formed in the first reaction is soluble in water and the dicalcium salt in the second reaction in ammonium citrate. Where fertilizers are to be transported to great distances there is a considerable saving of freight by the use of such a high-grade phosphate, which may, at times, contain over forty per cent of available acid. The phosphoric acid used is made directly from the mineral phosphate by treating it with an excess of sulfuric acid.

AUTHORITIES CITED IN
PART FIRST.