Citrate-Insoluble Ph.osphoric Acid in Di- and Tricalcium Phosphates Some Factors Affecting Its Determination K. D. JACOB,L. F. RADER,JR., H. L. MARSHALL, AND K. C. BEESON, Fertilizer and Fixed Nitrogen Investigations, Bureau of Chemistry and Soils, Washington, D. C.
W
HEN ordinary superT H E W A T E R - I N S O L U B L E di- and triPHOSPHATE MATERIALS USED phosphate is treated calcium phosphates are becoming of increasing DICALCIUM with ammonia, a porimportance in the fertilizer industry because of S a m p l e 387 Was B a k e r and tion of the o r i g i n a l watertheir occurrence in ammoniated superphosphates. Adamson’s c. P. material. It soluble phosphoric acid (p,05) is converted into waterInvestigations with the pure salts show that the a p p e a r e d to be in the amori n s o l u b l e forms. Addition of solubilities of these compounds in ammonium phous condition and the analysis indicates that it was the anhymore than about per cent Of citrate solution are affected to a considerable d r o u s salt. S a m p l e 390 was ammonia results in the formaextent by the weight Of sample taken for analysis, K a h l b a u m ’ s c. p, crystalline tion of phosphates which are the presence of other salts, the p H of the citrate material c o n t a i n i n g approxiaho insoluble in neutral ammomately 2 moles of water of crysnium citrate solution as detersolution, and the temperature at which the salt t a l l i z a t i o n - The PhOWhoric mined by the official m e t h o d is heated before the citrate digestion. The soluacid-lime ratios i n d i c a t e that ( 1 ) ) t h e q u a n t i t y ?f ~i.tra!ebilities are aJected to a less extent by the.fineness b o t h m a t e r i a l s contained phosphoric lnof the particles, the gravity of the citrate solution, slightly lime than is theecreasing with the q u a n t i t y of retically required for the pure and the time of the citrate digestion. ammonia added, Recent invessalt. tigations (13, 15) indicate that the citrate-soluble and the citrate-insoluble phosphates TRICALCIUM PHOSPHATE.Samples 287, 1023, and 1093 formed under these conditions consist principally of dicalcium were sold as c. P. materials by Eimer and Amend, Baker and and tricalcium phosphates, respectively. In view of the fact Adamson, and Merck, respectively. The phosphoric acidthat heavily ammoniated superphosphate1 contains a rela- lime ratios indicate that these materials contained more lime tively large portion of its phosphoric acid in the form of di- than is theoretically required for the pure salt. Sample 1094 and tricalcium phosphates, information on the factors affect- was prepared by neutralizing a water suspension of pure caling the solubilities2 of these two compounds in ammonium cium hydroxide with a dilute solution of pure phosphoric citrate solutions is of particular importance, acid, evaporating to dryness on the steam bath, and finally Fresenius, Neubauer, and Luck ( 7 ) first proposed, in 1871, heating a t 900’ to 950” C., as described by Jacob and Reythe use of ammonium citrate solutions for the laboratory de- nolds (14). Sample 1095 was prepared by slowly adding a termination of the so-called “available” phosphoric acid in solution of pure trisodium phosphate to a solution containing water-insoluble phosphates. Since that time, numerous ex- a small excess of calcium nitrate. The precipitate was periments have been carried out on the solubilities of di- and washed with a saturated solution of tricalcium phosphate tricalcium phosphates in citrate solutions. The reported re- until the filtrate gave no test for nitrates, the salt finally sults, particularly those of the earlier investigators (6, 6, 8, being dried a t a temperature of 50” C. 9, 11, 16), are, however, very conflicting. As pointed out by The composition of the samples in given in Table I. Robinson (19) in his excellent review of the work on the subTABLEI. COXPOSITION OF PHOSPHATES ject published prior to 1919, this is no doubt owing largely to IGNITION PaOsthe variable and uncertain pH values of the citrate solutions SAMALKALIEE NITRATE Loss AT CaO used by the different investigators. The discrepancies are PLE Pa06 CaO Rz0 C1 N 1100’ C. RATIO % % % % % % also probably due in part to the fact that in many cases the DICALCIUM P H O E P B A T E ~ experiments were made with mixtures of phosphates rather 387 50.45 40.25 .. ,.... . 26.355 . . . . 1.253 than with the individual compounds. Furthermore, no 390 42.29 34.02 .... .. 1.243 TRICALCI‘CM PHOSPEtATEc single comprehensive investigation of the factors affecting 0.825 0.005 6.74 2.70 0.00 40.86 49.50 287 the citrate solubilities of these phosphates under defmitely 0.805 0.002 6.70 0.40 1023 40.44 50.25 1.65 0.816 0.009 5.44 0.88 0.00 41.82 61.22 1093 known and reproducible conditions seems to have been made. 0.835 0.000 0.27 0.00 0.00 1094 45.38 54.36 The present paper gives the results of a study of the citrate 0.853 0.103 10.97 0.56 0.00 1095 40.48 47.46 a Theoretical PaOs-CaO ratio = 1.267. solubility of tricalcium phosphate as affected by the weight b Loss a t 900’ C. of sample, the time of digestion with citrate solution, the size Theoretical PaOrCaO ratio = 0.845. of the phosphate particles, the pH and specific gravity of the METHOD OF DETERMININQ CITRATE-INSOLUBLE citrate solution, the presence of other compounds, and by PHOSPHORIC ACID heating the phosphate prior to the citrate digestion. The effect of several of these factors on the solubility of dicalcium When tricalcium phosphate is treated with citrate solution phosphate was also determined. according to the official method, clear filtrates are difficult 1 The term ammoniated superphosphate refere to ordinary auperphosphate to obtain because a portion of the phosphate usually passes that has been treated with anhydrous or aqueous ammonia. through the filter paper in the colloidal condition, and results 2 The term “solubility” a8 used in this paper refers to solubility as dein low values for citrate-insoluble phosphoric acid. Clear termined under oertain specified conditions which are not neoessarily equifiltrates may be obtained, however, by the use of short Paslibrium conditions. C
25
Vol. 4, No. 1
ANALYTICAL EDITION
26
proportional to the weight of sample when 0.5- to 2.0-gram samples were used. On the other hand, the weights of phosphoric acid dissolved from 2.0-gram samples of tricalcium phosphate were only 1.2 t o 1.7 times those dissolved from 0.5gram samples, indicating that as the weight of sample was increased the amount of phosphoric acid dissolved by the citrate solution rapidly approached the maximum amount that could be dissolved under the conditions of the determinations. The investigations of Zulkowski and Cedivoda ($3) indicate that the dissolution of calcium phosphates by citrate solution is due to reactions which result in the formation of ammonium phosphates and calcium citrates, or calcium ammonium citrates. As shown by the figures in Table 111,
teur-Chamberland filter tubes. Duplicate results thus obtained are usually in good agreement and, in the case of phosphates which filter clear through paper, check closely those obtained by the use of filter paper. These tubes were used in obtaining all the results for citrate-insoluble phosphoric acid given in the present paper. Except for the method of filtration and as noted otherwise, the official procedure was carefully followed on 100-mesh material. The neutral citrate solution was carefully prepared according to the official method ( I ) , using phenol red as an indicator. This solution had a true pH value of 6.95 at 20" C. as determined potentiometrically by means of the hydrogen electrode. I n all cases 100 cc. of the citrate solution were used.
OF SAMPLE ON CITRATE SOLUBILITY OF DI- AND TRICALCIUM PHOSPHATES TABLE11. EFFECTOF WEIGHT
(Samples digested for 30 minutes a t 65' C.)
-CITRATE-~NSOLUBLE PaOa AB PERCENTAQD OB:-,Sampel---Total PzosSAMPLE W t . of sample: Wt. of sample: 0.5 g. 1.0 g. 1.5 g. 2.0 g. 0.6 g. 1.0 g. 1.6 g. 2.0 g.
%
%
%
%
%
%
%
%
PzOs DISSOLVDD BY 100 cc. OF 0.5 g.
Mo.
CITRATE SOLN. Wt. of sample: 1.0 g. 1.5 g. 2 . 0 g.
FOR 0.6-QRAM SAMPLES
0.5 g.
Wt. of sample: 1.0 g. 1.5 g. 2.0 g.
Mo.
MQ.
Mg.
PzOs DISSOLVED BY 100 cc. O F CITRATD SOLN.ON BASISOF UNITY
DICALCIUM PHOSPHATE
387 390
0.00 0.00
0.00 0.00
2.48 0.36
8.12 2.33
0.0 0.0
0.0
0.0
4.9 0.9
16.1 6.5
252.3 211.5
504.5 422.9
719.6 629.0
846.6 799.2
1.00 1.00
2.00 2.00
2.85 2.97
3.36 3.78
172.6 199.8 168.3 209.5 210.0
186.9 204.3 186.9 216.3 257.7
193.8 236.6 195.4 214.6 282.0
1.00 1.00 1.00 1.00 1.00
1.31 1.22 1.21 1.19 1.26
1.42 1.26 1.35 1.23 1.55
1.47 1.44 1.41 1.22 1.69
TRICALCIUM PHOSPHATE
287 1023 1093 1094 1095
14.54 7.67 14.06 10.24 7.14
23.60 20.46 24.99 24.43 19.48
28.40 26.82 29.36 30.96 23.30
31.17 28.61 32.05 34.65 26.38
35.6 19.0 33.6 22.6 17.6
57.8 50.6 69.8 53.8 48.1
69.5 66.3 70.2 68.2 57.6
76.3 70.7 76.6 76.4 65.2
EFFECT OF WEIGHTOF SAMPLE The figures given in Table I1 show that the dicalcium phosphates were completely soluble in citrate solution when the weight of sample did not exceed 1.0 gram, but with larger samples a portion of the phosphoric acid was insoluble under the conditions of the experiments, the quantity increasing with the weight of sample. Although higher percentages of citrate-insoluble phosphoric acid were obtained from the 1.5- and 2.0-gram samples of the anhydrous dicalcium phosphate, sample 387, than from the hydrated material, sample 390, under corresponding conditions, i t will be noted that the anhydrous material originally contained a much higher percentage of total phosphoric acid, and that larger actual weights of phosphoric acid were dissolved from it. Haskins (IO) has shown that the percentage of citrate-insoluble phosphoric acid in commercial "precipitated" phosphate, obtained as a by-product of the manufacture of glue and gelatin from bones, is decreased t o a considerable extent by reducing the weight of the sample from 2.0 grams to 1.0 gram. Although this type of precipitated phosphate consists essentially of dicalcium phosphate, Haskin's failure to obtain complete solution of the phosphoric acid in 1.0-gram samples indicates that his materials may have contained some tricalcium phosphate. I n the case of tricalcium phosphate, there was a progressive and significant decrease in the percentage of citrate-insoluble phosphoric acid when the weight of sample was decreased by 0.5-gram steps from 2.0 grams to 0.5 gram, the change being greater when the weight of sample was decreased from 1.0 gram to 0.5 gram than from 2.0 grams to 1.0 gram. The results show that the citrate solubility of tricalcium phosphate varies somewhat with different samples, and that approximately 65 to 75 per cent of the total phosphoric acid is insoluble, under the conditions prescribed by the official method, when 2.0-gram samples are used. When the weight of sample is reduced, however, to 0.5 gram, only about 18 to 36 per cent of the total phosphoric acid is insoluble. In the case of dicalcium phosphate, the weight of phosphoric acid dissolved by 100 cc. of citrate solution was roughly
131.6 163.9 138.8 175.7 166.7
the dissolution of unignited di- and tricalcium phosphates is also accompanied by reactions that, where complete solution of the phosphates is not obtained, result in the formation of citrate-insoluble residues, which contain a lower ratio of phosphoric acid to lime than is present in the original phosphates. The residues contained less than 0.1 per cent of ammonia nitrogen. TABLE111. PaOrCaO RATIOS IN CITRATE-INSOLUBLE RESIDUES FROM DI- AND TRICALCIUM PHOSPHATES c
,
PzOs-CaO RATIOSO Citrate-Insoluble ResiduesWt. of sample: 0.6 g. 1.0 g. 1.5 g.
7
7-
SAMPLI
,Original material
2.0 g.
DICALCIUM PHOSPHATE
387 390
1.253 1.243
... ...
... .*.
1.127
...
1.189 0.910
TRICALCIUM PHOSPHATI
0.786 0.825 0.760 287 1023 0.805 0.775 0.757 0.781 0.767 1093 0.816 0.832 0.835 0.813 1094b 0.787 0.853 0.752 1095 a Theoretical ratio for dicalcium phosphate = phosphate = 0.845; for calcium hydroxyphosphate b Heated a t 900' to 950' C.
0.796 0.775 0.763 0.839 0.786 1.267; for 0.760.
0.792 0.754 0.779 0.830 0.783 tricaloium
Many years ago Warrington (22) observed that tricalcium phosphate hydrolyzes in water to give a product containing an excess of lime. Lorah, Tartar, and Wood (17) have recently made a more thorough investigation of this reaction and conclude that a product corresponding in composition to hydroxyapatite, 3Ca3(PO&.Ca (OH)g, is finally obtained by the prolonged treatment of tricalcium phosphate with large quantities of boiling water. Buch (4) observed that prolonged treatment of dicalcium phosphate with water also results in the formation of more basic phosphates. Hydrolysis of the calcium phosphates is a slow process in pure water, but is quite rapid in neutral ammonium citrate solution, The phosphoric acid-lime ratios in the residues from unignited tricalcium phosphate are close to the ratio required for the hydroxyphosphate. Hydrolysis does not seem to occur to a significant extent, however, in the case of some samples of tricalcium phosphate that have been ignited a t 900" to 950" C. (sample 1094). The results given in
January 15, 1932
INDUSTRIAL AND ENGINEERING CHEMISTRY
Table I V show that the citrate solubility of citrate-insoluble residues obtained from tricalcium phosphate is quite different from that of the original phosphate. The residue used in these experiments was obtained by treating 100 grams of tricalcium phosphate, sample 287, with 20 liters of citrate solution, as described in a previous paper (13). T.4BLE
IV.
SOLUBILITY OF TION OF
RESIDUZFROM
TOTAL PzOs-CaO P z O ~ RATIO Original material Citrate-insoluble residue
CITRATE
TRICALCIUM PHOSPHATE 287
EXTRAC-
PeOb DISSOLVED BY 100 cc. OF CITRATE SOLN. Wt. of sample: 0.5 g. 1.0 g. 1.5 g. 2 . 0 g.
% 40:&
Ma.
Ma.
Ma.
Ma.
0.825
13i.6
172:6
186-.9
193:s
38.68
0.774
59.2
69.9
73.2
86.6
27
of the citrate solution is very important in determining the solubiliby of tricalcium phosphate. TABLEV. EFFECT OF PH OF CITRATE SOLUTION ON SOLUBILITY OF TRICALCIUM PHOSPHATE 287 (Samples digested for 30 minutes at 66' C.) CITRATFI-INBOLUBLE Pa05 AS PERCENTAQH~ OF: PH OF Sp. GR. OF ----Sample--Total PzOb-7 CITRATE CITRATE Wt. of sample: Wt. of sample: SOLN. 0.5 g. 2.0 g. 0.5 g. 2.0 g. SOLN." % % % % 6.63 1.092 12.61 30.14 3019 73;s 6.76 1.091 13.81 30.67 33.8 75.1 6.91 1.092 14.45 31.13 35.4 76.2 7.09 1.092 15.88 31.73 38.9 77.7 1.091 7.27 16.62 32.04 40.7 78.4 (1 Values determined at 20' C. by means of hydrogen electrode.
EFFECT OF SPECIFIC GRAVITY OF CITRATESOLUTION
EFFECT OF PH OF CITRATESOLUTION Many investigators (19) have shown that the solubility of phosphates in citrate solution depends to a considerable extent on the reaction of the solution, the solubility of the calcium phosphates increasing with the acidity of the solution. This accounts for the discordant results reported by many of the early workers who had no accurate method for determining the true reaction of their solutions. The reaction of citrate solutions can now be accurately determined and controlled, however, by the use of potentiometric and colorimetric methods (29,20). Robinson (19) found that in the case of tricalcium phosphate, the percentage of total phosphoric acid present as citrate-insoluble phosphoric acid was 74.1, 79.7, 80.3, and 82.8, as determined by the official method using 2gram samples, when the pH values of the citrate solutions were 6.6, 7.0, 7.4, and 7.8, respectively. Robinson's solutions were all adjusted to a specific gravity of 1.09 at 20" C., and, consequently, 100 cc. of the solutions probably did not contain identical quantities of the citrate radical. I n order to obtain information on the effect of pH on the solubility of tricalcium phosphate when the concentration of citrate radical was maintained constant, several solutions containing exactly 188.13 grams of pure crystallized citric acid, CsH80y.H20,per liter3 a t 20" C. were prepared as follows: Five portions of 188.13 grams each of pure citric acid were weighed into separate liter flasks and dissolved in about 400 cc. of water. The solutions were roughly neutralized t o the desired pH values by the addition of strong ammonium hydroxide, care being taken not to overstep the particular pH values desired. The reactions of the solutions were then carefully adjusted colorimetrically to the desired pH values by the addition of small quantities of dilute ammonium hydroxide. Bromothymol blue was used as an indicator in preparin the solutions with pH values of 6.6 and 6.8, while phenol ref was used in preparing those with higher values. Ammonium citrate solution neutralized to a pH of 7.0 as determined colorimetrically using bromothymol blue has an actual pH value of about 6.8 as determined potentiometrically using the hydrogen electrode. Consequently, when this indicator was used the solutions were adjusted to apparent pH values 0.2 higher than those actually desired. Values obtained colorimetrically by the use of phenol red are only about 0.05 pH lower than those obtained with the hydrogen electrode. All the test portions were returned to the flasks and the solutions were diluted to 1 liter. The reactions of the solutions were finally determined potentiometrically using the hydrogen electrode. The figures given in Table V show that, with a constant concentration of citrate radical in the pH range 6.6 to 7.3, a difference of 0.2 pH in the reaction of the solution had an appreciable effect on the solubility of tricalcium phosphate. These results indicate that careful adjustment of the reaction
*
As an average of tests on eight carefully prepared solutions, Robinson (29)found that 1 liter of a oitrate solution having a pH of 7.0 and a specific gravity of 1.09 at 20° C. contains 172 grams of anhydrous, or 188.13 grams of crystallized, citric acid monohydrate.
The figures given in Table VI show that the solubility of 1gram samples of tricalcium phosphate, sample 287, was not affected to a significant extent by small changes in the specific gravity of the citrate solution within the gravity range of 1.090 to 1.094. Reducing the gravity of the solution to 1.OS8 and 1.086, respectively, resulted, however, in progressive and appreciable increases in the percentages of citrate-insoluble phosphoric acid. The citrate solutions used in these experiments were prepared as follows: Five liters of solution, neutral to phenol red and having a gravity greater than 1.094, were prepared. This solution was divided into five portions, and the gravities were adjusted t o the desired values at 20" C., as determined by the hydrometer.
At gravities of 1.086 to 1.094, addition of about 20 cc. of water per liter changes the gravity of neutral ammonium citrate solution by about 0.002, which corresponds to a change of about 2.0 per cent in the concentration of the solutions. Consequently, in the case of the 1-gram samples of tricalcium phosphate 287, a progressive reduction of 0.002 in the gravity of the citrate solution might be expected to produce a progressive increase of about 0.50 per cent in the amount of citrate-insoluble phosphoric acid. The results given in Tables I1 and VI show, however, that within the limits of the experiments the solubility of this compound is not a linear function of either the concentration or the total volume of the citrate solution. TABLE VI. EFFECT OF SPECIFIC GRAVITY OF CITRATE SOLUTION ON SOLUBILITY OF TRICALCIUM PHOSPHATE 287 (1-gram samples digested for 30 minutes at 65' C.) CITRATE-INSOLUBLE PzOs A S PERCENTAQE OF: CITRATE SOLN. Sample Total Pa06
SP.GR. OF l.OS6 1,088 1.090 1,092 1.094
%
%
25.29 24.91 24.44 24.41 24.32
61.9 61.0 59.8 59.7 59.5
Although the directions for the preparation of the official citrate solution specify that it shall have a specific gravity of 1.09 a t 20" C., no statement is given as to the temperature at which the 100-cc. aliquots used for the determination of insoluble phosphoric acid shall be drawn. The temperature of the citrate solution at the time the aliquot is drawn may range from 15" to 35" C. in different laboratories and in the same laboratory a t different seasons of the year. Tests on a neutral solution that had a gravity of 1.090+ a t 20' C., as determined by the hydrometer, showed that the gravity was 1.092 a t 15", 1.089 at 25", 1.088 at 30°, and 1.087 at 35" C. A change in gravity is of course accompanied by 8 change in the total weight of citrate radical in 100 cc. of the solution. Within the range of temperatures usually encountered in the laboratory this change in concentration will probably have no effect on the determination of citrate-insoluble phosphoric acid in such materials as straight superphosphates, double
ANALYTICAL EDITION
28
or triple superphosphates, and the usual types and grades of mixed fertilizers, but it may be of significance in the.case of pure tricalcium phosphate, highly ammoniated superphosphate, and bone materials. EFFECTOF PARTICLE SIZE A sample of 100 grams of tricalcium phosphate, sample 287, was separated into particle-size ranges of 20 to 50 mesh (0.864 to 0.254 mm.), 50 to 100 mesh (0.254 to 0.147 mm.), finer than 100 mesh, and finer than 200 mesh (0.074 mm.). The different particle sizes had the same content of total phosphoric acid. The figures given in Table VI1 show that under the conditions of the experiments the solubility of tricalcium phosphate increased as the size of particle was decreased to 100 mesh, Although the figures for citrate-insoluble phosphoric acid are scfmewhat higher in the 200-mesh material than in the 100-mesh material, the differences are probabIy of no significance. The effect of particle size became less pronounced as the weight of sample was increased from 0.5 gram to 2.0 grams.
Vol. 4, No. 1
TABLEIX. EFFECT OF CALCIUM SULFATE AND CALCIUM CARCITRATESOLUBILITY OF DICALCIUM PHOSPHATE 390
BONATE ON
(Samples digested for 30 minutea at 66' C.) CITRATE-INSOLUBLE PaOa AS PERCENTAGE OF: ----Sample-.-Total PiOsWT. OF WT. OF Added material: Added material: CaHP04.ADDED CaSO4.CaS0r.2HaO MATERIAL 2H10Q CaCOab 2Ha0 CaCOs Grams Grams % % % % 1.0 0.0 0100 0:oo 010 0:o 1.21 1.0 0.5 Trace 2.9 2.43 1.0 1.0 0.93 i:2 6.7 2.0 3.82 1.0 3.99 9.4 9.0 2.33 0.0 2.33 2.0 5.5 5.5 0.5 4.68 11.1 14.1 6.97 2.0 1.0 6.24 7.16 2.0 14.8 16.9 2.0 8.61 2.0 8.26 19.5 20.1 0 Synthetic gypsum. b Synthetic calcium carbonate.
TABLEX. EFFECT OF CALCIUM SULFATE AND CALCIUXCARBONATE ON CITRATE SOLUBILITY OF TRICALCIUM PHOSPHATE 1095
(Samples digested for 30 minutes at 65' C.) CITRATE-INSOLUBLE PaOa AS PERCENTAGE OF: ' WT. OF -Sample--Total PzO!WT. OF ADDED Added material: Added material: Caa(P04)a MATERIAL CaS04.2HzOa CaCOab CaSOa. 2H20 CaCOa Grams Grams % % % % 0.0 7.14 17.6 0.6 7.14 17.6 25.56 40.1 0.5 16.25 63.1 EFFECT OF PARTICLE SIZEON CITRATESOLUBILITY 0 . 6 30.74 57.1 1.0 0.6 23.12 76.9 OF TRICALCIUM PHOSPHATE 287 34.12 81.8 2.0 33.10 84.3 0.5 0.0 19.48 48.1 19.48 48.1 1.0 (Samples digested for 30 minutes at 65' C.) 0.6 30.35 61.6 24.88 75.0 1.0 CITRATE-INEOLUBLE PaOs AS 1 . 0 33.31 7 2 . 5 29.36 82.3 1 . 0 PERCENTAGE OF: 1.0 2.0 35.35 36.28 87.3 87.2 PARTICLE Srzm Sample Total Pa06 0.0 2.0 65.2 26.38 26.38 65.2 Mesh % % 0.5 2.0 76.7 33.81 30.66 83.5 1.0 2.0 82.6 38.64 33.40 88.0 20-60 21.89 63.6 2.0 2.0 91.2 37.64 36.90 92.7 60-100 18.44 46.1 < 100 14.64 35.6 a Synthetic gypsum. < 200 14.83 36.3 b Synthetic oalcium carbonate. 20-60 27.84 68.1 50-100 26.55 66.0 < 100 23.60 57.8 It is interesting to note that in the case of unignited tri< 200 24.07 68.9 20-50 32.39 79.3 calcium phosphate, sample 287 for instance, the actual weights 60-100 31.96 78.2 < 100 31.17 76.3 of phosphoric acid dissolved when 0.5- to 2.0-gram samples
TABLE VII.
WT. O F
CasPOn Grams 0.5 0.5 0.5 0.5 1.0 1.0 1.0 1.0 2.0 2.0 2.0
were digested for 1.5 hours depended upon the weight of sample, and showed no significant increases over the weights dissolved when the samples were digested for 1 hour. From The results given in Table VI11 show that when 0.5-gram this it may be concluded that with each particular weight of samples were used, the amount of citrate-insoluble phosphoric sample, the citrate solution dissolved during the course of 1 acid in unignited tricalcium phosflhate was decreased to a hour approximately the maximum weight of phosphoric acid significant extent by increasing the time of digestion from that it was capable of dissolving under the conditions of the 0.5 hour to 1.0 hour, but with 1.0- and 2.0-gram samples the experiments. The question then arises as to why greater effect was much less. I n general, increasing the time of di- weights of phosphoric acid were dissolved when Zgram gestion from 1.0 hour to 1.5 hours had no significant effect on samples were digested for 1hour than when 0.5- or 1-gram samthe amount of citrate-insolubl6 phosphoric acid. When the ples were digested under the same conditions. At present tricalcium phosphate was ignited a t 1000" C. there was, how- the reason for this is not entirely clear, but the data given ever, a progressive and significant decrease in the quantity of in Tables 11,111,and VI11 indicate that the action of citrate citrdte-insoluble phosphoric acid as the time of digestion was solutions on tricalcium phosphate involves not only the disincreased from 0.5 hour to 1.5 hours. solution of some tricalcium phosphate, but also the hydrolysis of a portion of the tricalcium phosphate to calcium hydroxyTABLE VIII. EFFECT OF TIME OF DIGESTION ON CITRATE SOLU- phosphate, and the dissolution of a certain amount of the BILITY OF TRICALCIUM PHOSPHATE last compound. Calcium hydroxyphosphate is much more (Samples digested at 65' C.) insoluble in citrate solution than is tricalcium phosphate, and CITRATB1-INSOLURLEP& AS Pa06 DISSOLVED it may be that a relatively greater proportion of the original WT. OF TIMEOF PERCENTAGE OF: BY 100 CC. O F OadPOah Total PaOa CITRATESOLN. tricalcium phosphate is converted into hydroxyphosphate _.,_ DIGESTION Sample Grams Hours % % Mg. when the weight of sample is decreased from 2.0 grams to CrtafPOds .. .- 1095 0.5 gram. 166.7 17.6 0.5 0.5 7.14 '
_I"\_
0.5 0.5 1.0 1.0 1.0
EFFECT OF TIMEOF DIGESTION
~
1.0 1.5 0.5 1.0 1.6
0.5 0.6 1.0 0.5 1.5 0.6 0.5 1.0 1.0 1.0 1.6 1.0 0.5 2.0 1.0 2.0 1.5 2.0 0.6 1.00 1.0 1.0' 1.6 1.0a 0 Ignited for 1 hour
11.6 4.69 9.9 4.00 48.1 19.48 44.0 17-81 44.9 18.19 Caa(PO4)a 287 35.6 14.64 31.0 12.66 30.8 12.58 57.8 23.60 56.9 23.26 66.7 23.16 76.3 31.17 74.6 30.46 74.1 30.26 64.0 26.17 61.2 24.99 68.6 23.94 at 1000D C. prior to oitrate digestion.
179.0 182.4 210.0 226.7 222.9
EFFECT OF OTHERCOMPOUNDS The results given in Tables IX and X show that gypsum
131.6 141.0 141.4 172.6 176.0 177.1 193.8 208.0 212.0 146.9 158.7 169.2
and calcium carbonate had a very pronounced effect in decreasing the citrate solubilities of the di- and tricalcium phosphates. Calcium carbonate had a greater effect than an equal weight of gypsum, and both had a much greater effect on the solubility of tricalcium phosphate than on dicalcium phosphate, Both gypsum and calcium carbonate had a more pronounced effect when 0.5-gram samples of tricalcium phosphate were used than when 1- and 2-gram sampIes were used. The effect of these compounds on the solubilities of di-
January 15, 1932
INDUSTRIAL AND ENGINEERING CHEMISTRY
and tricalcium phosphates is of considerable practical importance, since ammoniated superphosphate always contains large quantities of calcium sulfate, and calcium carbonate in the form of ground limestone is frequently added to mixed fertilizers. The figures given in Table XI show further that the solubility of tricalcium phosphate was also depressed to a significant extent by the presence of calcium fluoride, monocalcium phosphate, calcium chloride, magnesium carbonate, dolomitic limestone, and precipitated iron and aluminum oxides, and to a less extent by the presence of magnesium sulfate and diammonium phosphate. On the other hand, the solubility of tricalcium phosphate was increased by the presence of ammonium sulfate, monoammonium phosphate, sodium nitrate, potassium chloride, and potassium sulfate. Diammonium phosphate is an alkaline salt, and its effect in decreasing the solubility of tricalcium phosphate is no doubt largely due to an increase in the alkalinity of the citrate solution caused by its presence. Monoammonium phosphate, on the other hand, is an acid salt, and therefore it would be expected that its presence would increase the solubility of tricalcium phosphate. A 0.1 M solution of diammonium phosphate has a pH of 8.0, while a 0.1 M solution of monoammonium phosphate has a pH of 4.4 (21). Neutral alkali salts also seem to have a specific effect in increasing the solubility of tricalcium phosphate. Several investigators (18) have noted that the dissolution of calcium phosphates by water is favored by the presence of alkali metal salts. From the practical standpoint, the effect of water-soluble compounds is probably of no considerable significance, since in the official method of analysis they would be largely removed by washing prior to the citrate digestion.
29
precipitate is probably calcium citrate and i t dissolves almost completely during the final washing with hot water, TABLEXII. SOLUBILITY OF VARIOUSCOMPOUNDS IN NEUTRAL AMMONIUM CITRATE SOLUTION (Samples digested with 100 00. of citrate solution a t 65' C. for 30 minutes all materials ground to 100 mesh) WT. DISSOLVED BY 100 cc. MATERIAL OF CITRATESOLN. Crams
Approx. 3.0 Approx. 3.0 1.19 1.19 0.05 0.24 Approx. 3 .O 0.04
High-calcium limestone. Fluorspar Bureau of Standards Sample 79. C Prepared 'by treating Fe(NO;)s.QHnO with NHaOH and drying preoipitate on steam bath. a b
The depressing effect of calcium salts on the solubilities of di- and tricalcium phosphates seems to be due largely to the formation of calcium citrate, thus decreasing the quantity of citrate radical available for reaction with the calcium phosphates. The depressing effect of magnesium compounds and precipitated iron and aluminum oxides is probably due to similar reactions. The writers wish to call particular attention to the effect of iron on the solubility of calcium phosphates. A very pronounced greenish yellow color was noted in a citrate solution that had been prepared in an enamel-lined iron kettle, and when this solution was used on tricalcium phosphate much higher figures for citrate-insoluble phosphoric acid were obtained than when a colorless solution prepared in a glass conTABLE XI. EFFECTOF VARIOUSCOMPOUNDS ON CITRATE SOLU- tainer from the same lot of citric acid was used. The pH values and the specific gravities of the solutions were identical. BILITY OF TRICALCIUM PHOSPHATE 281 An analysis showed that 100 cc. of the colored solution con(1 gram of CasjPO& plus 1 gram of added material; samples digested for 30 minutes a t 65' C.) tained 0.16 gram of iron calculated as ferric oxide, which, CITRATECITRATEIN~OLUBLE IN~OLUBLEIowing to a crack in the enamel, had been dissolved from the PZOSA B PERPPOSAS PERkettle, whereas the colorless solution was free of iron. The ADDED CENTAQE OF: ADDED CENTAQE OF: results given in Table XI indicate that the low solubility obMATXIRIAL Total MATERIAL Total Sample PzOa Sample PzOa tained by use of the colored solution was due to its iron con% % % % tent. While small quantities of iron in the citrate solution None 23.60 57.8 MgCOaa CaS04.2Hz0,a syntdetic 33.89 82.9 will probably have no effect on the determination of citrates nthetic 30.60 74.9 Dolomitew 27.82 68.1 insoluble phosphoric acid in straight superphosphates and the CaJO4.ZHz0," bfgso4.7HzO 26.58 65.1 natural 33.74 82.6 (NHI)zSOI 20.76 50.8 usual types of mixed fertilizers, only iron-free solutions should NH4HzPOr 18.83 46.1 34.19 83.7 (NHa)zHPO4 26.87 65.8 be used on tricalcium phosphate, bone products, heavily NaNOa 20.71 50.7 ammoniated superphosphates, and other difficultly soluble natu'ral 34.83 85.2 KC1 20.64 50.5 CaFt a,c K2so4 20.54 50.3 materials. nalural 29.44 72.1 FezOaad 24.26 59.4 CaHa(POa)z.HzO 29.96 73.3 FezOaam 28.65 70.1 The foregoing results show that the citrate solubility of CaClz .2Hz0 31.13 76.2 AlzOsad 23.65 57.9 tricalcium phosphate in a mixed fertilizer may be considerably AlzOad 27.49 67.3 a Ground to 100 mesh. lower than that of the same weight of pure salt, not because of b High-caloium limestone. c Fluorspar, Bureau of Standards, Standard Sample 79. a change in the chemical nature of the phosphate but because d Baker and Adamson's c. P. material. of the effect of other compounds in reducing its solubility e Prepared by treating Fe(NOa)a.gHaO with NH4OH and drying preoipitate on steam bath. under a particular set of conditions. f Prepared b treating AlCla. 12Ha0 with NHaOH and drying precipitate on stea? bat%. u Containing 33.71 per cent CaO and 17.82 per oent MgO.
The figures given in Table XI1 show that in the absence of other compounds, 100 cc. of citrate solution dissolved, under the conditions of the experiments, approximately 3 grams of gypsum and of magnesium carbonate, 1.19 grams of calcium carbonate, 0.24 gram of dolomite, 0.05 gram of calcium fluoride, or 0.04 gram of precipitated ferric oxide, the solubility of the last compound depending upon its past history. A11 these compounds are very much more soluble in citrate solution than in water, owing a t least partly, no doubt, to chemical reaction with the citrate. The occurrence of chemical reactions is indicated by the fact that when 3 grams of gypsum are added to 100 cc. of citrate solution a t 65" C., the salt dissolves very rapidly and almost completely, but with further digestion a flocculent precipitate forms. This
EFFECT OF IGNITION According to Erlenmeyer and Antz (6), air-dried tricalcium phosphate is slightly more soluble in citrate solution than material which has been dried a t 50" C., and the latter is much more soluble than material which has been ignited. Barill6 (2) states that freshly prepared gelatinous tricalcium phosphate is more soluble than material that has been airor oven-dried, while the calcined material is insoluble. He also states that hydrated dicalcium phosphate is more soluble than the anhydrous salt, and that this material is rendered insoluble by calcination because of the formation of calcium pyrophosphate. Grupe and Tollens (8) note that citrate solution dissolves significant quantities of tricalcium phosphate if the salt is not dried a t too high a temperature. In order to obtain further information on this point, a study
Vol. 4, No. 1
ANALYTICAL EDITION
30
was made of the effect of ignition at various temperatures on the solubilities of di- and tricalcium phosphates. I n carrying out these experiments, the desired quantities of material were ignited in platinum dishes at the desired temperature for 1 hour, except those a t 105" C. which were heated to constant weight. The individual ignition residues were used for the citrate digestions. Consequently, the results obtained in a series of experiments with a given weight of original material are comparable because the same weight of total phosphoric acid was present in each case.
TEMPERATURE
OF IGNITION
- OC
about 800" C. and increased at higher temperatures. Although the absolute quantities of insoluble phosphoric acid were not the same in the two materials, the curves given in Figure 1 have similar slopes, indicating that the same factors were concerned in the changes in solubility of the two materials. The changes in the solubility of tricalcium phosphate on heating seem to be due principally to changes in the chemical constitution of the material as a result of the loss of water. Somewhat different results would probably be obtained on samples ignited to constant weight at the different temperatures. This point is being investigated further. It will be noted that in the case of t r i c a l c i u m p h o s p h a t e , sample 287 (Table XV), small but definite increases in loss of weight occurred as the temperature of ignition was increased from 800" to 1100" C. Bassett (3) has observed that t r i c a l c i u m p h o s p h a t e holds water very tenaciously, but he fails to state the temperatures at which his experiments were made. Unignited tricalcium phosphate passes to a certain extent into the colloidal condition when it is treated with citrate solution, and considerable time is required for filtering and washing the insoluble residue which is finally obtained in the form of a sticky mass. The insoluble material shows no tendency to creep up the sides of the vessel during either the citrate digestion or the washing with hot water. A similar behavior was noted in the case of samples ignited a t temperatures up to 400°C. On t,he other hand, the colloidal and sticky nature of the residues and the time required for filtering and washing progressively decreased, whereas the tendency of the insoluble material to creep up the sides of the vessel increased as the temperature of ignition was in-
Pure hydrated dicalcium phosphate, CaHP04.2H20, theoretically contains 20.93 per cent water of hydration and 5.23 per cent water of constitution, or a total of 26.16 per OF IGNITION ON CITRATE SOLUBILITY OF cent. The material used in the experiments, sample 390, TABLEXIV. EFFECT TRICALCITJH PHOSPHATE 1095 was a very pure crystalline salt containing 0.19 per cent free (Samples digested for 30 minutes at 65' C.) water in addition to the water of hydration and constitution. CITRATE-INSOLUBLE P2Os AS Table XI11 shows that removal of the water of hydration had PERCDNTAGB OF: WT. OF TEMP. O F IGNITION Sample Total P2Os Cas(P04)2 a small but definite effect in decreasing the solubility of di% % calcium phosphate. When the temperature was raised to Grams 0.5 . . c. . 7.14 17.6 0 . 5 1055 11.34 2 400" C., a portion of the water of constitution was also driven 0.5 200 15.39 3 88 .. 00 off and the solubility was very greatly decreased because of 0.6 400 18.94 46.8 0.5 1000 3.68 9.1 the formation of calcium pyrophosphate. Previous experi1.0 ... 19.48 48.1 1050 21.43 ments (1%')have shown that less than 10 per cent of the total 11 .. 00 200 25.37 65 22 .. 79 phosphoric acid in calcium pyrophosphate prepared by ig1.0 400 26.46 65.4 1.0 600 25.42 62.8 niting dicalcium phosphate a t 800' C. is soluble in neutral 1.0 700 15.69 38.8 1.0 800 12.15 30.0 ammonium citrate solution. 1.0 900 12.88 31.8 1.0 1000 16.20 40.0 TABLEXIII. EFFECT OF IGNITION ON CITRATE SOLUBILITY OF 2.0 ... 26.38 65.2 2.0 1055 29.18 72.1 DICALCIUM PHOSPHATE 390 (Samples digested for 30 minutes at 66' C.)
WT. OF TEMP.OF CaHP04.2HaO IQNITION Grams O c. 1.0
...
CITRATE-INSOLUBLE PZOSAS P E R C E N T OF: AG~ Sample
%
Total PZOK
%
a
31.48 32.36 26.50 Other samples heated for 1
77.8 79.9 65.5 hour.
LOSS IN WT.
%
0.00 0.0 0 06 0.1 i:i5b 200 0.23 0.5 17.91b 400 24.25 57.3 22,446 2.33 6.5 ... i65. 13.1 ... 5.53 2.0 200 9.70 22.9 ... 2.0 400 35.50 83.9 a Heated to oonstant weight. Total time of heating, 50 hours. Other samples heated for 1 hour. b Average of results obtained on 1- and 2-gram samples.
1.0 1.0 1.0 2.0 2.0
2.0 200 2.0 400 2.0 1000 Heated to oonstant weight,
105a
...
The figures given in Tables XIV and XV show that when tricalcium phosphate was ignited for 1 hour, the percentage of citrate-insoluble phosphoric acid steadily increased up to The an knition temperature Of about 6ooo *' phOsphoric acid then decreased rather sharply to a minimum at
ON CITRATE SOLUBILITY OF TABLExv. EFFECTOF IGNITION TRICALCIUM PHOSPHATE 287
(1-gram samples digested for 30 minutes at 65' C.) CITRATE-INSOLUBLE PZOSAS PERCBNTAGE OF: Total Pa01 LOSS Sample
TEMP.OF IGNITION
c.
%
23 60 ibba 23.97 2 6.81 200 29.18 400 30.01 600 700 26.39 23.33 800 900 24.35 1000 26.17 1100 28.56 a Heated to constant weight, samples heated for 1 hour.
'
%
IN WT.
%
57.8 ... 58.7 1.94 65.6 2.73 71.4 4.21 73.4 5.77 64.6 5.99 6.23 57.1 59.6 6.38 64.0 6.58 69.9 6.63 Total time of heating, 40 hours. Other
January 15, 1932
INDUSTRIAL AND ENGINEERING CHEMISTRY
creased from 400" to 1100" C. The extracts from samples ignited a t 1000" and 1100" C. filtered very rapidly and the insoluble residues had practically no colloidal and sticky properties. A comprehensive investigation of the constitution and properties of tricalcium phosphate and the complex calcium phosphates is now being made in the Fertilizer Technology Division of the Bureau of Chemistry and Soils. ACKNOWLEDGMENT The authors wish to express thanks to E. F. Snyder of the Soils Fertility Division for making the potentiometric pH measurements on the citrate solutions used in t,he determination of citrate-insoluble phosphoric acid. LITERATURE CITED (1) (2) (3) (4) (6)
Assocn. Official Agr. Chem., Methods, p. 4, 1925. Barill6, J . pharm. chim., [6], 27, 437 (1908). Bassett, J . Chem. Soc., 111, 620 (1917). Buch, 2. anorg. Chem., 52, 326 (1907). Dircks a n d Werenskiold, Landw. Vers.-Sta., 34, 426 (1887).
31
Erlenmeyer a n d Anta, Ber., 14, 1253 (1881). Fresenius, Neubauer, a n d Luck, 2. anal. Chem., 10, 149 (1871). Grupe a n d Tollens, Ber., 13, 1267 (1880). Grupe and Tollens, Ibid., 14, 754 (1881). Haskins, J . Assocn. Oficial Agr. Chem., 4, 64 (1920); 5, 97 (1921); 5, 460 (1922). Herzfeld a n d Feuerlein, 2.anal. Chem., 20, 191 (1881). Jacob, Beeson, Rader, a n d Ross, S. Assocn. O$ciaZ Agr. Chem., 14, 263 (1931). Jacob, Hill, Ross, a n d Rader, IND. ENG.CHIN., 22, 1392 (1930). Jacob a n d Reynolds, Ibid., 20, 1204 (1928). Keenen, Ibid., 22, 1378 (1930). Konig, 2. anal. Chem., 20, 49 (1881). Lorah, T a r t a r , a n d Wood. J. Am. Chem. Soc., 51, 1097 (1929). Mellor, "A Comprehensive Treatise on Inorganic a n d Theoretical Chemistry," Vol. 111, p. 876, Longmans, 1923. Robinson, Michigan Agr. Expt. Sta., Bull. 46 (1919). Robinson, J. Assocn. Oficial Agr. Chem., 5, 92 (1921); 5 , 443 (1922). Ross, Merr, and Jacob, IND.
[email protected]., 21, 286 (1929). Warrington, J . Chem. Soc. (London), 19, 296 (1866); 26, 983 (1873). Zulkowski and Cedivoda, Chem. Ind., 26, 1, 27 (1903).
RECEIVBDMay 20, 1931. Presented before the Divieion of Fertilizer Chemistry a t the 82nd Meeting of the American Chemical Society, Buffalo, N. Y.,August 31 to September 4, 1931.
Analysis of Beryllium Minerals FRANKLIN G. HILLS,Experimental Plant, Colorado School of Mines, Golden, Colo. LTHOUGH beryllium was discovered by Vauquelin in 1798, neither the metal, its minerals, nor its compounds have found any use in industry or the arts until quite recently. However, it has attracted considerable scientific interest since its discovery, Questions regarding its physical and chemical properties, its compounds, and methods of separation have received much attention and the answers have been found only after much study. I n making beryllium analyses the preliminary work offers no difficulties. It consists of a sodium carbonate fusion and the separation of silica in the usual way, then precipitation of the hydroxides with ammonia. This precipitate will carry the beryllium as hydroxide, together with iron, alumina, etc. As beryllium hydroxide is appreciably soluble in water, all washing should be done with a 2 per cent solution of ammonium nitrate. I n analyzing phosphate minerals, phosphorus must be separated as ammonium phosphomolybdate before proceeding with the other separations. The solution of the hydroxides in hydrochloric acid may be treated with hydrogen sulfide if second-group metals are present before proceeding with the other separations, but this is probably never necessary in ordinary mineral analysis. I n the author's work on the minerals mentioned above, three methods for the separation of aluminum and iron were used, beryllium in all cases being finally precipitated as the hydroxide with ammonia, and weighed as beryllium oxide after ignition. METHOD OF BERYLLIUM DEVELOPING CORPORATION H. S. Cooper, of the Beryllium Development Corporation of Cleveland, Ohio, reported (in a private communication to the author) a method which consists briefly in igniting the mixed hydroxides, fusing with sodium carbonate, leaching the melt with water, filtering, and then repeating the ignition, fusion, and leaching. Generally this second fusion and leaching will remove the last of the alumina, but with minerals high in
alumina a third fusion may be necessary. The last residue is ignited and fused with potassium pyrosulfate and dissolved in water. This should give a clear solution, except perhaps for a little silica which is removed by filtration. From this solution the hydroxides of iron and beryllium are precipitated with ammonia and collected on a filter, dissolved in dilute hydrochloric acid, and again precipitated with ammonia. The solution and re-precipitation are necessary to free the precipitate from alkali salts, which are readily adsorbed by beryllium hydroxide. The precipitate is ignited and weighed as beryllium oxide and ferric oxide. The precipitate is then dissolved in hydrochloric acid, the iron determined, and beryllium oxide found by difference. This method is obviously open to the criticism that obtaining the beryllium percentage by difference throws all the errors on the beryllium. Still it is a very satisfactory technical scheme. Another objection is that the method as outlined takes no account of manganese, which is a common constituent of beryllium minerals. If neglected, it is very largely carried through all the operations and will be reported as beryllium oxide. Manganese may be separated as follows: After making the bisulfate fusion, dissolving in water, and precipitating the hydroxides with ammonia, filter and dissolve the hydroxide precipitate with dilute nitric acid back into the beaker in which the precipitation was made. Evaporate to a small volume, but not to dryness, add 10 cc. of concentrated nitric acid, and precipitate the manganese with potassium chlorate. Filter out this precipitate and from the filtrate separate the iron and beryllium with ammonia as described above. This separation for manganese is never quite complete, so that traces of manganese will still be found with the beryllium. This small amount can be determined colorimetrically and deducted if necessary. With small amounts of manganese, oxidation sometimes goes to the permanganate condition.
