=119. Time Required for Precipitation of Phosphoric Acid.=—The length
of time required for the complete precipitation of the phosphoric acid
by molybdate mixture is perhaps much less than generally supposed.
At 65° the precipitation, as shown by de Roode, is complete in five
minutes.[100] In a given case the weight of pyrophosphate obtained
after five minutes was 0.0676 gram, and exactly the same weight was
found after twenty-four hours. In view of these facts analysts would
often be able to save time by omitting the delay usually demanded by
the setting aside of the yellow precipitate for a few hours in order
to secure a complete separation of the phosphoric acid. In the method
of the official chemists it is directed that the digestion at 65° be
continued for one hour, and this time may possibly be shortened with
advantage. In all cases, however, where there is any doubt in regard to
the complete separation, some of the molybdate solution should be added
to the filtrate and, with renewed digestion, it should be noted whether
any additional precipitate be formed.
=120. Examination of the Pyrophosphate.=—In fertilizer control it is
not usually thought necessary to examine the magnesium pyrophosphate
for impurities. Among those most likely to be found is silica. It
is proper, in all cases where accuracy is required, to dissolve the
precipitate in nitric acid, boil for some time to convert the pyro-
into orthophosphate, and reprecipitate with molybdate and magnesia
mixture. This treatment will separate the silica which remains
practically insoluble after the first ignition. It has been observed
by some analysts that the results obtained by the official method are
a trifle too high and also that on re-solution the second precipitate
of pyrophosphate weighs less than the first.[101] The difference in
most cases is very little but it may become a quantity of considerable
magnitude in samples where soluble silica is found in notable
quantities. The danger of contamination with iron, alumina, and arsenic
has already been mentioned but it is not of sufficient importance to
warrant further attention.
=121. Iodin in Phosphates.=—The presence of iodin has been detected in
many natural phosphates and is of interest in the discussion of the
problem of their origin.[102] A qualitative test for the detection
of iodin may be applied in the following manner: Some finely ground
phosphate is mixed with strong sulfuric acid and the gases arising from
the reaction are aspired into some carbon disulfid or chloroform. The
violet coloration arising indicates the presence of iodin. The gases
carrying the iodin may also be brought into contact with starch-paste
producing the well-known blue color.
The quantity of iodin present in a phosphate is rarely more than one or
two-tenths of one per cent. It can be determined as a silver salt, in
the absence of chlorin or by any of the standard methods found in works
on qualitative analysis.
Iodin is quite a constant constituent of Florida phosphates.
For a quantitative determination, the sample is treated with an excess
of strong sulfuric acid in a closed flask and during the decomposition
a stream of air is aspired through the flask and caused to bubble
through absorption bulbs containing sodium hydroxid in solution.
The temperature of the decomposition may be raised to about 200°.
After the solution of the sample the sodium iodid formed is oxidized
by heating with potassium permanganate, acidulated and mixed with a
solution of potassium iodid to hold the free iodin in solution. The
free iodin is determined in the usual way by titration with standard
sodium thiosulfate solution. The reactions preparatory to the titration
are represented by the following formulas:
2KI + H₂SO₄ = K₂SO₄ + 2HI.
2HI + H₂SO₄ = 2H₂O + SO₂ + 2I.
6I + 6NaOH = NaIO₃ + 5NaI + 3H₂O.
NaI + 2KMnO₄ + H₂O = NaIO₃ + 2KOH + 2MnO₂.
HIO₃ + 5HI = 6I + 3H₂O.
The titration is represented by the following reaction:
2Na₂S₂O₂ + 2I = 2NaI + 2NaI + Na₂S₄O₄.
The decinormal solution of sodium thiosulfate may be used. Grind the
crystals of the salt to a fine powder, dry between blotting papers,
and use 24.8 grams of the dried salt per liter. The quantity of iodin
found in phosphates is so minute that it is hardly worth while to make
a quantitative determination of it.
=122. Occurrence of Chromium in Phosphates.=—In some phosphates a small
quantity of chromium has been found. In a sample of phosphate from the
Island of Los Roques in the Caribbean Sea, Gilbert found three-fourths
per cent of chromium oxid (Cr₂O₃). The phosphates containing chromium
have a greenish color and are characterized by great insolubility in
solutions containing organic acids. The chromium is to be determined by
the usual methods described in mineral analysis.
=123. Estimation of Vanadium.=—In the complete analysis of basic slags
it becomes necessary to determine the presence of vanadium and its
quantity. The method used in this laboratory for the purpose is the
volumetric process of Lindemann.[103] It is conducted as follows:
Dissolve four grams of the finely powdered slag in sixty cubic
centimeters of dilute sulfuric acid (1: 4), boil for a few minutes,
cool, make the volume up to 100 cubic centimeters, filter, and take
an aliquot part for the determination. Add decinormal potassium
permanganate solution in slight excess to secure the oxidation of the
vanadium to vanadium pentoxid. Add, drop by drop, a weak solution of
ferrous sulfate until the pink color just disappears. Prepare a ferrous
sulfate solution by dissolving 2.183 grams of piano wire in sulfuric
acid and making the volume to one liter. Titrate the vanadic mixture
with this solution until a drop of the clear liquor removed and brought
in contact with potassium ferricyanid shows a distinctive blue-green
color.
One cubic centimeter of the ferrous sulfate solution is equivalent to
0.002 gram of vanadium, 0.002888 gram of vanadium dioxid, and 0.003648
gram of vanadium pentoxid. The ferrous sulfate solution may also be
made and standardized by any of the approved methods in common use.
The method described by Blair, designed especially for the estimation
of vanadium in iron and steel, is conducted in the following
manner:[104] Five grams of the drillings are dissolved in fifty cubic
centimeters of nitric acid of 1.24 specific gravity. The solution is
evaporated to dryness in a porcelain dish and heated thereafter until
the nitrates are nearly decomposed. After cooling, the dried mass is
transferred to a mortar and finely ground with thirty grams of dry
sodium carbonate and three grams of sodium nitrate. The finely ground
materials are placed in a platinum dish and fused for an hour at a
high temperature. Spread the fused mass over the sides of the dish
while cooling, and afterwards dissolve in hot water, filter, and wash
until the volume is a little over half a liter. Add nitric acid to
decompose carbonates, but not completely, and boil to get rid of carbon
dioxid, being careful to keep the mass always slightly alkaline. Add
nitric acid, drop by drop, until slightly in excess, and then sodium
carbonate to marked alkalinity, boil, and filter. Add a slight excess
of nitric acid to the filtrate, and the development of a yellow color
will indicate the presence of vanadic acid. Add to the solution a
small quantity of mercurous nitrate and then an excess of mercuric
oxid, suspended in water to render the solution neutral and insure the
complete precipitation of mercurous vanadate. The mercurous salt also
precipitates phosphoric, chromic, tungstic, and molybdic acids which
may be present. Boil, filter, and wash the precipitate with hot water,
dry, and ignite. Fuse the residue with sodium carbonate and a little
nitrate. Dissolve the fused mass, after cooling, in a little water and
filter. Add to the filtrate, ammonium chlorid in excess, from three to
five grams for each 100 cubic centimeters of the solution, and allow
to stand, with occasional stirring, for some time. Ammonium vanadate,
insoluble in a saturated solution of ammonium chlorid, separates as
a white powder. It is necessary to keep the solution alkaline, and a
drop of ammonia should be added from time to time for this purpose. The
appearance of a yellowish tint at any time indicates that the solution
has become acid, and this acidity must be corrected, or else the
results will be too low. Separate the ammonium vanadate by filtration;
wash first with a saturated solution of ammonium chlorid containing a
little free ammonia, and then with alcohol. Dry, ignite, and moisten
with a few drops of nitric acid; again ignite to obtain the compound as
vanadium pentoxid. This compound contains 56.22 per cent of vanadium.
The method of Rosenheim and Holversheet may also be used.[105] It is
based on the preliminary precipitation of the vanadic acid as a barium
or lead salt. The substance supposed to contain vanadium is first
brought into solution in such a manner as to secure it as vanadic acid,
which is then precipitated with barium chlorid or lead acetate. The
precipitate is boiled with hydrochloric acid and potassium bromid, and
the liberated bromin determined by the quantity of iodin set free from
potassium iodid. In the absence of bodies, such as molybdic acid, which
are reduced by sulfurous acid or hydrogen sulfid, the vanadic acid
may also be determined by reducing it with one of these reagents and,
after removing the excess by boiling, titrating the vanadium tetroxid
with potassium permanganate. When vanadic and phosphoric acids occur
together the former may be first reduced to tetroxid with sulfurous
acid, and after expelling excess of this reagent, the phosphoric acid
may be separated with molybdate solution and removed by filtration.
When the amount of vanadic acid is large the phosphoric acid should
be separated rapidly at 55°-60°, using a considerable excess of the
molybdate; or the vanadic acid may first be determined in the solution
volumetrically by the bromin process above described, and afterwards
the phosphoric acid obtained by evaporating to dryness with a little
sulfuric acid, taking the residue up with water, reducing the vanadic
with sulfurous acid and precipitating the phosphoric acid with
molybdate solution as described above.
=124. Fluorin in Bones.=—Fluorin is not only a constituent of mineral
phosphates but also of bones. According to the researches of Carnot
there is often as much as one-half per cent of calcium fluorid in bones
and teeth.[106] Gabriel has suggested a means of determining a minimum
limit of fluorin in bones and teeth by the development of etchings in
comparison with known quantities of pure calcium fluorid. The minimum
quantity of calcium fluorid necessary to produce a distinct etching, in
known conditions, having been determined, the test is applied to known
weights of ignited bone or teeth. He concludes from his results, that
the ash of bones and teeth often contains less than one-tenth per cent
of fluorin. Since, however, there is a loss of fluorin from calcium
fluorid, on ignition, the whole of the fluorin may not have been
available in the tests described.
=125. Note on the Separation of Iron and Aluminum Phosphates from
the Calcium Compound.=—There are many points of difference noted in
the descriptions given by authors of the deportment of the iron, and
aluminum phosphates in presence of a large excess of the calcium salt.
Especially is this true of the statements made by Hess and Glaser[107]
in paragraphs =34= and =35=. The subject is of such importance, from an
analytical point of view, as to merit a careful study.
In this laboratory a thorough investigation of the mutual deportment
of these three phosphates has been made by Brown with the following
results:[108] When a mixture containing a known weight of the salts
was treated exactly as Hess directs, in no case was there a complete
separation of the iron aluminum phosphate from the calcium salt. In
order to discover the cause of the failure, pure solutions of calcium
and iron aluminum phosphates were treated under identical conditions
by the necessary reagents. Fifty cubic centimeters of a solution of
calcium phosphate, containing about one gram of the salt, were treated
with 100 cubic centimeters of water and fifty cubic centimeters of
the commercial ammonium acetate containing 150 grams of the salt in a
liter. An immediate precipitate was produced at ordinary temperature,
and on heating to 60° it became abundant. The addition of ammonium
chlorid, phosphate, and nitrate in successive portions, does not
prevent the precipitation. Making the solution more dilute lessens the
difficulty when twenty cubic centimeters of a ten per cent solution
of ammonium phosphate are first added, followed by the usual quantity
of ammonium acetate; a clear crystalline precipitate is sometimes
observed. Experience also shows that the trouble is not due to an
excess of the ammonium acetate.
In treating a solution of iron aluminum phosphate, in similar
circumstances, with the ammonium acetate, it is found that a complete
precipitation takes place.
Since diluting the solution of the calcium salt diminishes its tendency
to form a precipitate with the ammonium acetate the true method of
separation seems to lie in that direction. The calcium salt is held
completely in solution when the separation is made in the following way.
The solution containing the mixed phosphates is diluted so as to
contain not more than one gram thereof in half a liter. To this is
added one drop of dimethylanilin orange, and afterwards ammonium
hydroxid, until a very slight precipitate is formed. The mixture is
heated to 70° and from twenty to twenty-five cubic centimeters of a
twenty-five per cent solution of acid ammonium acetate are added,
enough to change the rose color of the indicator to orange. The iron
aluminum phosphate is separated by filtration and washed with a hot
five per cent solution of ammonium nitrate.
The washed precipitate shows no impurity due to calcium, as proved
by dissolving it, reprecipitating and filtering, adding ammonium
hydroxid to the filtrate, and heating for a long time. Sometimes a
slight troubling of the clear liquid may be observed which may be due
to a slight solubility of the iron aluminum phosphate in washing, an
accident that may occur if the temperature be allowed to fall below
70°, but no weighable amount of material is obtained. If due to
calcium phosphate, a greater dilution in the first precipitation will
remove even this mere trace of that salt. In the above conditions the
contamination of the iron aluminum precipitate with calcium phosphate
may be entirely avoided. We had also undertaken here the problem of
separating the phosphoric acid by the citrate method, followed by a
destruction of the citric acid in the filtrate by combustion with
sulfuric acid according to the kjeldahl process, and final separation
of the iron and alumina in the residues when our attention was called
to substantially the same process as described by Jean.[109] The method
merits a further critical examination.
=126. Phosphoric Acid Soluble in Ammonium Citrate.=—There is no other
point connected with the determination of phosphoric acid which has
excited so much discussion and about which there is such difference of
opinion as the solubility of phosphates in ammonium citrate. It was
clearly established by Huston, in 1882, that the ammonium citrate, as
used in fertilizer analysis, would attack normal tricalcium phosphate
as it exists in bones.[110]
In a raw bone, finely ground, containing 20.28 per cent of phosphoric
acid, the following quantities were found to be soluble in a neutral
ammonium citrate solution of 1.09 specific gravity.
Time of digestion, thirty minutes.
Temperature 30° 40° 50° 60°
Per cent P₂O₅ dissolved 2.76 4.01 3.39 5.88
From this it appears that the quantity of acid dissolved increases with
the temperature of digestion with the exception of the number obtained
at 50°. When the time of digestion was increased there was also found
a progressive increase in the amount of acid passing into solution. At
40° for forty-five minutes the per cent dissolved was 4.97, and at 40°
for one hour it was 5.92. These early determinations had the effect
of calling attention to the thoroughly empirical process which was in
use, in many modified forms, by agricultural chemists, the world over
for determining so-called reverted phosphoric acid in fertilizers.
Since the publication of the paper above named many investigations have
been undertaken by Huston and others relating to this matter.[111]
The general results of these studies, tabulated by Huston, are given
below.[112]