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INDUSTRIAL AND ENGINEERING CHEMISTRY
Vol. 23. No. 11
Action of Sulfur Dioxide on Phosphates of Calcium' A. E. Hughes2 and F. K. Cameron UNIVERSITY OF NORTH CAROLINA, CHAPELHILL, N. C ,
The vapor pressures have been determined for the systems sulfur dioxide-water and sulfur dioxide-water-tricalcium phosphate, between the temperatures 0" and 90" C. The results of previous investigators are confirmed, but the range of observations has been much extended. Tricalcium phosphate and phosphate rock, when treated with saturated aqueous solution of sulfur dioxide, at 25" to 100' C. at pressures of from 3 to 10 atmospheres can be completely taken into solution. The phosphoric acid (P20b)goes into solution more rapidly; if, however, on the establishment of equilibrium, the liquid phase be so small that a solid phase persists, then it has been found that calcium oxide (CaO) is present in excess of the
.. .. .. .. ECAUSE of possible importance for the production of commercial fertilizers, the action of aqueous sulfur dioxide on the phosphates of calcium has long attracted inventors The patent literature is voluminous (SS), but contains little of real importance. The investigations described in the academic journals are few and the results generally inconclusive, Gerland (18) in 1871, when the commercial production of phosphatic fertilizers was an infant industry, found dicalcium phosphate to be quite soluble in aqueous sulfurous acid, and to form, as he thought, a double salt. But Rotondi (SI) pointed out that he must have produced a solid mixture probably containing calcium sulfite and dicrtlcium phosphate. Designolles (13),Tirelli (SG), Sestini ( S / t ) , Defries ( I W ) , and Faber ( I J ) , as well as many others, studied the reaction with the production of fertilizers as their ultimate aim but obtained no practical results. Mebane, Dobhins, and Cameron (68) studied the system Ca0-P2Os-SO2-H?O a t 0" and 25" C. under atmospheric pressure, the results explaining the failures of previous investigators to obtain practical results a t ordinary temperatures and pressures. They found a small range of concentrations of the liquid phase which, on evaporation, would yield monoor dicalcium phosphates, or mixtures, but the dose control of these concentrations, as well as the precipitation of a large part of the phosphorus as a more basic salt forbade a reasonable expectation of commercial development. Mieg (29) called attention to the desirability of treating phosphates with sulfur dioxide, and Thilo (35) has shown that dicalcium phosphate and calcium sulfite are solid products of the reaction. In the simpler system CaO-S02-Hz0, Aufsalzteil ( I ) claims that the disulfite as well as normal sulfite is a stable solid. This system is being investigated by Campbell (IO). No evidence has been adduced for the presence of the disulfite of calcium when phosphoric acid is added to the system. The present investigation is an exploration of the effect of sulfur dioxide, with and without the presence of water, on the more basic calcium phosphates over a range of temperature and pressure which might be commercially practicable. Since the purpose of the investigation is a commercial one, the detailed study of the reactions encountered must
B
1Received June 2 1931. f Robert Ober Fellow, University of North Carolina. The investigation, in port described in this paper, was made possible by a grant from G. Ober Sons and Company, through the courtesy of the president, Gustavus
Ober, Jr.
phosphoric acid in the solid phase. All the phosphoric acid can be made available by such treatment. Direct measurements of the volume of sulfur dioxide absorbed by tricalcium phosphate have been made over a wide range of temperatures and at several pressures. This method of dry treatment has been applied to phosphate rock up to 1100" C. The reactions involved are very complex, and much sulfur dioxide is decomposed. The maximum phosphoric acid rendered available is about two-thirds of the total present. No procedure utilizing sulfur dioxide promises, under present conditions, better economy than the standard methods of treating phosphates with sulfuric acid. ,I
.
.. . .
be postponed for a more propitious opportunity. It is believed, nevertheless, that they have been followed sufficiently to meet the needs of this inquiry. Two general lines of investigation were followed. I n the first, water was present in relatively large amounts; i. e., the phosphate was treated with an aqueous liquid solution, saturated with respect to sulfur dioxide. I n order to insure saturation, there was present some liquid sulfur dioxide. I n the second line of inquiry water was present, from vanishing amounts to somewhat more than sufficient to saturate gaseous sulfur dioxide. The purpose, in all cases. was the conversion of the phosphoric acid (I'?O,) into an available form. Available phosphoric acid was determined by the procedure of the Association of Official Agricultural Chemists (19) as the sum of that extracted by a standard proportion of water and that extracted from the residue by a standard aqueous solution of ammonium citrate. That a metastable aqueous solution of monocalcium phosphate is readily obtainable is well known; and that dicalcium phosphate is soluble in ammonium citrate in the proportions specified, has been confirmed by t!ie authors. Without regard to questionable theoretical implications, the procedure has value as a police regulation in the sale of phosphatic fertilizers and is accepted by the industry. Materials Used
PHOSPHATE ROCK-A 200-pound (90.7-kg.) sample of Florida rock, ground to pass 100 mesh, was thoroughly mixed. An analysis gave the following results: V" CaO
PnOs Si02 Fez03 AI&; MgO
1
46'.>7 32 35 14.20
2.91 Trace
Fluorine and carbon dioxide, both present, were not determined. TRICALCIUN PHOSPHATE-Attempts to prepare this product by the method of Lorah, Tartar, and Wood ( d 7 ) , adding aqueous ammonia to a solution of monocalcium phosphate a t 100" C. until the mother liquor is slightly alkaline, invariably yielded a precipitate with a greater ratio of calcium to phosphorus than corresponds to the theoretical formula. Consequently, a large sample obtained from the Baker Chemical Company was employed. It was obviously a precipitate from water solution and showed no crystals under the microscope,
INDUSTRIAL A X D ENGINEERIIZ'G CHEMISTRY
Kovember, 1931
and the calcium-phosphorus ratio closely approximated that required by the formula. No appreciable amounts of impurities were recognized.
__ THM+K?S7AT
0"
CUYCR F?8 C a V D W R
C W N O WfflETfR
W F A n N T , lNP
Figure 1-Apparatus for Determining Vapor Pressure of Sulf u r Dioxide, Dry or Wet
SULFURDIoxIDE-This gas was obtained from the Baker Chemical Company in iron cylinders. It was dried by bubbling through concentrated sulfuric acid. When it was desirable to purify it, as in the vapor-pressure measurements, the dried gas was cooled to a temperature below -5" C., when i t liquefied, and any oxygen, nitrogen, chlorine, or carbon dioxide present was removed in the attending vapor phase. OxYGEPi-The gas was a commercial product obtained in the usual steel cylinders. No attempt was made to purify it beyond separating it from water and mechanical contaminations by passing it through concentrated sulfuric acid and plugs of cotton or glass wool. The analytical reagents used were the usual ones of the laboratory. Analytical Methods
PHOSPHORIC Acm-This was precipitated with ammonium molybdate from the warm aqueous solution containing iiitric acid; the precipitate was separated by filtration and carefully washed. The precipitate was then suspended in water and titrated with standard aqueous sodium hydroxide. Usually it has been found desirable to add an excess of the alkali and, after standing until solution was completed, to titrate back with standard acid. LIME OR CALCIUMOxIDE-This was determined in the combined filtrate and washings from the phosphoric acid precipitate The solution was brought to boiling to insure the removal of any phosphorus To the clear boiling solution, crystals of oxalic acid were added in excess; ammonia to decided alkalinity was then added. The granular precipitate thus formed was filtered, washed, and dissolved in aqueous hydrochloric acid and again precipitated by ammonia. The precipitate w&s converted to oxide in the muffle, sulfuric acid added, the excess acid volatilized, and calcium oxide calculated from the weighed residue of calcium sulfate. SULFURDIoxIDE-This was determined by treating a sample of the product under investigation with sulfuric acid, and absorbing the liberated sulfur dioxide in concentrated aqueous potassium hydroxide. Vapor Tensions in the System Sulfur Dioxide-Water
Regnault (23) using the static method, determined the vapor pressure of water between -30" and 209" C., and of sulfur dioxide between -30" and 65" C. The vapor tensions of water as found by Regnault have been verified by Scheel and Heuse (Sb),Holborn and Baumann (25), and Henning
1263
(21). Hofbauer (24) found the calculated vapor pressure of water over a wide range of temperatures to be in accord with the observed results of the previously named investigators. Burrell and Robertson (9), Henning and Stock (E?),and Bergstrom (4) determined the vapor pressure of su!fur dioxide as low as -94.4" C. Briner ( Y ) , working with tubes of small bore, determined the vapor pressures of sulfur dioxide to its critical temperature, a t which point the pressure is 77.95 atmospheres. The vapor pressures of the system sulfur dioxide-oxygen have been determined by Briner and Wroczynski (8). A slight cloudiness observed was attributed to the formation of sulfur trioxide. The system sulfur dioxide-water has been studied by Roozeboom (33) and eo-workers between the temperatures -10" and 17.1" C. A compound to which they ascribe the formula SO2.7H2Ois a solid phase between the temperatures -2.6" and 12.1" C. Bailey (2) has shown that su!furous acid may exist as a compound represented by the formula S02.nH20. Baume and Tykociner (3) studied the system from -74" to 12.1' C., and Freeze (17') determined the solubility of sulfur dioxide in water from 0 " to 40" C. In order to know the pressures to be encountered in the study of the system a t higher temperatures than those reached by Roozeboom, it was necessary to extend the work. The mixture sulfur d i o x i d e-w a t e r was placed in a thick-walled Pyrex tube 1, Figure 1, with a bore of 1 cm. This tube was welded directly t o a compound manometer 4 capable of measuring eight to nine atmospheres pressure. A condenser surr o u n d i n g t h e upper portion of the tube 1 was filled with ice water for the purpose of condensing v a p o r s that might escape from the mixture. Efficient agitation of the mixture was obtained by means of a s o f t i r o n l u g 2 sealed within a small glass tube, operated by a solenoid 3 t h r o u g h which a pulsating curr e n t passed. This portion Of the apparaFigure 2-Vapor Pressure-Temperature tus was immersed in an Curves for Sulfur Dioxide, Water, a n d Mixe lec t r i c a1 1y heated tures thermostat i n s u l a t e d with saw dust. In reading the pressure registered by the manometer, the temperature of the thermostat was taken as the temperature of the mixture sulfur dioxide-water. Ample time was allowed for equilibrium to take place on changing from one temperature to another. First, the pressure was determined from 0' to 90' C. by a gradual heating of the thermostat. Then another set of readings was recorded as the system cooled. A number of series of determinations through this range was made and concordant results obtained. The averaged results are assembled in Table I.
The authors' results agree with those of Roozeboom but cover a much Fvider range. At 0" C. the pressure is 117.65 em. of mercury and a solid phase is present-namely, SO2i"20. The slope of the vapor-pressure curve increases gradually with a rise in temperature until 12.2" C. is reached, when an abrupt break takes place in the temperature-pressure curve and the solid phase disappears. The second portion of the curve tends toward a mean of the pressures of sulfur dioxide and water as determined by Regnault. By violent agitation only was it possible to obtain a quick precipitation of the solid hydrate, metastable conditions being easily ob-
1264
IIVDGSTRIAL 9iYD ESGISEERIhTG' CHEMISTRY
Table I-Temperatures
a n d Pressures for S y s t e m Sulfur IXoxideWater PRESSURES TEMP. Temp. risinz TemD. fal!inx AI E R A L E Cm. H g Cm. Hg 117.50 123.01 125.39 132.53 138,13 165.72 167.25 168.59 150.14 171.36 172.60 174.92 177.27 179.25 199.58 222.14 244.00 265.97 286.83 300.75 340.05 410.22 460.25 491.72 527. SO 559.52 597.16 a Transition Point.
tained. The hydrate sometimes persisted as high as 210 c*, while two liquid phases have coexisted as low as I " C. The limitations Of the apparatus did not permit making accurate determinations above c* However, it is quite probable can be made for a higher temperature' Furthermore, it 's certain that the pressures recorded during the time that the temperature was approaching that Of the room are the more accurate* Above the transition temperature a t which the solid hydrate of sulfur dioxide disappears as a solid phase, the curves indicate that some definite form of the distribution law holds. However, the acid (S02~nH20) probably forms complexes or combinations in the liquid phase, and it seems vain to attempt even an empirical generalization. Therefore, no comparison of the authors' results with those of Regnault for the individual components is justified, other than the qualitative ones apparent in Figure 2. This chart shows that, a t a temperature below about 12" C., pure sulfur dioxide has a lower vapor pressure than when water is present. It is desirable, obviously, that measurements should be repeated for the sulfur dioxide. The exigencies of the investigation forbade a repetition at this time.
Vol. 23, No. 11
Table 11-Temperatures a n d Pressures for S y s t e m Sulfur DioxideWater-Tricalcium Phosphate PRESSURES TEMP. Temp. rising Temp. failing AVERAGE O c. Cm. Hg Cm. H g Cm. H a 117.2 117.3 117 2.5 118.8 119.9 118.91 122.1 124 2 122.65 128.7 129.9 129.40 140.8 140.0 140,40 150.2 150 .5 150.35 10.64 160.5 160.4 160.40 12 180.8 179.6 180 40 18 189.2 188.7 188.95 20 19;7.2 195 0 195.15 25 209,2 210.1 209.95 30 225.6 225.0 225.29 35 245.0 245.2 245.10 40 263.5 261.8 262.65 45 284.8 285.1 284.95 50 303.9 305.0 304.45 55 325 2 324.9 325,05 60 352.4 350.8 351.55 65 380.1 378.4 379.25 70 404.8 404.0 404.20 76 433.4 432.5 432.95 80 461.5 461. S 461.65 494.9 85 492.6 493.20 90 525.8 526.5 526.10 95 562.9 562.0 562,55 100 594.1 592.5 593.25 a Transition point.
of the tricalcium phosphate. From experiments which will be described later it was found that, on long standing in contact with water and tricalcium phosphate a t moderately elevated temperatures, sulfur dioxide is more or less decomposed. Sulfur trioxide, free sulfur, sulfides, sulfates, as well as fitesand less basic phosphates were found in the reaction product. They Were not observed in these vapor-tensjon measurements, probably because not anticipated and small in amount. The vapor tensions of sulfur and sulfur trioxide are relatively small in comparison with those of water and sulfur dioxide, and their would not be marked. Saturated Aqueous Solutions of Sulfur Dioxide and Tricalcium Phosphate
Working a t atmospheric pressure, Mebane, Dobbins, and Cameron increased the content of sulfur dioxide in the liquid phase by cooling to 0" C. and found that the solubility of lime (CaO) increased proportionally more than that of phosphoric acid (P205). The reaction was very slow. The experiments just described showed it to be feasible to work with solutions saturated with sulfur dioxide a t temperatures where a greater solvent action on phosphoric acid could be anticipated, with reaction velocities of practical significance and at pressures practicable in plant operations. Solubility e x p e r i m e n t s w e r e carried out as follows:
A Coca Cola bottle was fitted with a rubber stopper held in place by a clamp. In the bottle were placed weighed amounts of water and tricalcium phosphate, and sulfur dioxide was passed through crkwnchmw - m y the solution until saturated a t room temperature. The bottle was then surrounded by a mixture of salt and ice and the passage of sulfur dioxide continued until a liquid layer persisted in the bottle. The rubber stopper and clamp were a d j u s t e d . The stopper was fitted with a narrow, thick-walled glass tube drawn to a capillary seal. The p t t l e was brought to room temperature (about 26 C.), agikrnchm ccnw68 tated vigorously for several hours, and then brought mkrrrum H I D R U I I D I B torest. Sedimentation was very slow, and to hasten Figure 3-Apparatus Used in Analysis of Li u i d Phase for Sulfur Dioxlde, Calcium it 3 grams of ammonium nitrate were customarily Oxide, Water, a n d P%osphoric Acid -4-Sample flask B-Drying flask C-Sulfur dioxide flask added to each 100 grams of water. No evidence of this salt or its components was found in the solid products of the reaction and its solvent action was assumed to Vapor Tensions in the System Sulfur Dioxide-Waterbe negligibly small. Beside the vapor phase and the solids, the Tricalcium Phosphate bottle contained two liquids: one of these being water saturated The presence of tricalcium phosphate produces a noticeable with sulfur dioxide and containing also lime and phosphoric acid; lowering of the vapor tension of a mixture of sulfur dioxide the other, liquid sulfur dioxide saturated with water and containand water, as shown by the results assembled in Table I1 ing negligibly small quantities of the lime and phosphoric acid. a little experience i t was possible to have the liquid sulfur and charted in Figure 2. The transition point a t which the With dioxide present-in quite small amounts. hydrate of sulfur dioxide appears as a stable solid is lowered For analysis of the water solution, the end of the tube with capilby 1.6"C., indicating a considerable amount of solvent action lary seal was connected to a weighed flask A , Figure 3, containing
ISDL'STRIA L d S D EXGINEERIIVG CHEMISTRY
November, 1931
a solution of potassium hydroxide. The connection was made with a thick but soft-walled rubber tube, and, by crushing the capillary seal, any desired amount of the straw-colored liquid was forced out of the reaction chamber into A and determined by reweighing A . The flask A was then connected with the train as indicated in Figure 3, and a slow stream of air drawn through. As soon as the air stream was established, sulfuric acid was introduced a t X t o decompose the sulfites in A , the sulfur dioxide being caught in the weighing flask C, containing potassium hydroxide. The contents of flask A , after elimination of sulfur dioxide, were carefully washed into a beaker and heated, and enough nitric acid added t o dissolve any solid which might be present. The solution was then made t o known volume a t room temperature and an aliquot taken for estimation of the phosphoric acid and lime as previously described. The results of several series of experiments, which need not be recorded in detail, showed t h a t a desirable ratio of tricalcium phosphate t o water would be from one t o three parts of the salt t o eight parts of water; and even with continuous and vigorous shaking, a final state of equilibrium would require several days. Series of bottles were then prepared and continuously agitated in a motor-driven machine. At intervals of 12 hours, or multiples thereof, a bottle would be removed from the machine, the contents permitted t o settle for 2 hours, and a sample of the clear mother liquor withdrawn for analysis. The results are assembled in Table 111.
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water and sulfur dioxide by heating. Unfortunately the total recovery of phosphoric acid would be unsatisfactory, perhaps 65 per cent of the total in the tricalcium phosphate, and the proportion of water to be evaporated would be so large as to make the expense prohibitive. Possibly the mixture obtained when three parts of tricalcium phosphate were added to eight parts of water saturated with sulfur dioxide and agitated for 120 hours represented equilibrium with the realization of a transition point and no degrees of freedom or variance. With four components (water, sulfur dioxide, calcium oxide, and phosphoric acid), there must have been, in addition to the vapor and two liquids, three solid phases-probably calcium sulfite, monocalcium
-1
Ln.
0' , nN
c
Table 111-Effect of T i m e o n Concentration of A u e o u s S o l u t i o n in C o n t a c t w i t h Tricalcium Phosphate and Liqui8 Sulfur Dioxide PzOr cao Hz0 Ca0:PlOa DISSOLVED TIME so2 Ratio Hours % 70 70 5 70
20
RATIO H z O : 3 B. P.L
24 48 72 96 120
20.5 20.7 20.6 23.1 22 5
4.92 5.80 6.50 6.90 6.95
12 24 48 72 96 120
20.1 22 2 21.1 21.5 20.8 22.4
4.20 4.80 5.40 5.60 5.80 6.10
12 24 36 48 60 72 84 96 108 120 132
21.4 22.3 22.2 20.6 20.8 21.1 20.5 22.4 22.6 22.5 21.8
4.37 4.71 4.80 5.02 5.10 4.92 5.80 5.25 5.26 5.25 5.27
4.63 5 26 6.30 6.40 6.52
69.21 68.00 66.30 63.60 62.90
RATIO 8 H20: 2 B. P. L.
4.12 4.63 5.00 5.11 5.14 5.92
RATIO 8 H20
5.50 4.01 4.62 5,40 5.49 5.76 5.85 5.88 5.92 5.93 5.92
0.87 0.88 0.90 0.91 0.92
65.8 69.6 73.2 78.7 79.1
0.86 0.88 0.87 0.91 0.88 0.89
79.2 82.6 87.6 91.4 96.5 96.8
0.79 0.83 0.96 1.07 1.07 1.10 1.12 1.12 1.12 1.12 1.12
82.5 86.27 91.50 93.80 95.60 97.17 97.73 98.67 98.67 98.51 98.50
.
70.50 69.30 68.40 68.50 68.26 67.60 : 6 B. P. L.
69.92 69.83 69.70 66.80 66.53 66.10 65.28 64.83 64.15 63.93 63.95
: 1 B. P. L. 1.65 74.20 0.58 75.16 0.70 2.39 74.11 0.73 2.88 3.25 73.80 0.81 3.50 73.51 0.87 3.70 73,26 0.92 3.10 73.10 0.80 0.97 3.50 71.60 1.02 4.10 71.02 4.19 70.25 1.04 4.30 69.42 1.07 4.35 68.65 1.08 4.36 67.13 1.09 4.41 66,82 1.10 lime, or tricalcium phosphate
RATIO 8 HZ0
22.4 2.81 22.5 3.59 22.1 3.92 20.8 3.97 22.3 3.99 22.5 3.99 20.1 3.90 20.7 4.01 22.1 4.02 21.7 4.01 20.9 4.06 22.8 4.05 22.3 4.06 4.03 22.6 B. P. L., bone phosphate of
12 24 36 48 60 72 84 96 108 120 132 144 156 168 0
69.60 89.50 97.01 98.92 99.17 99.31 99.40 99.42 99.48 99.49 99.50 99.48 99.48 99.52
It appears that the concentration of sulfur dioxide in the aqueous solution is affected but little by the other solutes. It also appears that, a t first, phosphoric acid passes from the solid to the liquid phase more rapidly than does calcium oxide. But the latter continues to go into solution long after the maximum solubility of phosphoric acid has been attained. When the proportion of liquid to solid is sufficiently large, the ultimate solubility of the calcium oxide exceeds considerably that of phosphoric acid, owing no doubt to the continued formation of calcium sulfite as a solute. I n other words, in those cases where the complete extraction of phosphoric acid was produced, solutions with the maximum content of calcium oxide and of water were obtained. They offer no promise of practical application. By inspection of the graphs, Figure 4, it is apparent that with short-time agitation, solutions can be obtained which would yield solid mono- or mixtures of mono- and dicalcium phosphate after removal of
40
90
60
IO0
20
TIMEI N HOURS Figure 4-Time-Concentration Graphs from T r e a t m e n t of Calcium Phosphate w i t h Wet Sulfur Dioxide
phosphate, and dicalcium phosphate. To have realized this transition point would have been, however, a happy accident. More probably a point on a boundary curve was obtained, and the solid phases realized were calcium sulfite and dicalcium phosphate. Microscopic examination of the residue, visually and by photographs, seemed to confirm this conclusion, but not very satisfactorily, because, with continued escape of sulfur dioxide from adhering mother liquor, new crystalline depositions obscured the material being examined. With the mixtures containing relatively larger proportions of water to tricalcium phosphate, points in a field must have been realized (though possibly near a boundary curve), since all the phosphoric acid was in a liquid phase. Consequently, the only solid phase was calcium sulfite. Again, this conclusion could not be established definitely, as dicalcium phosphate was continuously deposited with escape of sulfur dioxide while the solid residues were under examination. A similar set of experiments was carried out where ground rock was used instead of the precipitated tricalcium phosphate. The carbon dioxide in the rock was eliminated in the preliminary treatment with sulfur dioxide. Saturation of the liquid phases was slower when rock was used. The data are assembled in Table IT7. Table I\'-Effect of T i m e on Concentration of Aqueous S o l u t i o n in Contact w i t h Phosphate Rock a n d Liquid Sulfur Dioxide PzOr TIME SOz Pi05 CaO HzO CaO:P,Os DISSOLVED Hours % % % % Ratio 70 RATIO 8 Hz0
24 48 72 144
22.7 22.3 24.9 23.7
2.19 2.30 2.32 4.12
24 48 72
22.1 22 0 20 2
1.90 2.10 2 52
120 144
27.10 23.10
4.38 4.39
. 1 6 ROCK
2.65 1.90 2.77 5.40 RATIO 8 H10
74.40 73.50 69.70 60.90
5.58 5.82
48.2 49.0 50.9 95.80
1.51 1.21 1.08
62.1 67.9 79.2
1.25 1.32
81.4 88.6
: 1 ROCK
2.98 2.56 2.74
R A T I O 8 H20
1.20 0.82 1.19 1.16
74.02 73.40 74 40 :2
ROCK
63.04 64.30
The primary purpose of the experiments a t room temperature was attained on finding that with 5.3 parts of water to one of tricalcium phosphate, the water being saturated with
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INDUSTRIAL A N D ENGINEERING CHEMISTRY
sulfur dioxide a t less than 3 atmospheres, practically all the phosphoric acid went into the aqueous solution. It was disappointing to find that the solutions contained relatively high proportions of lime and water, and that a steady state was attained only very slowly. Hence, experiments a t higher temperatures were tried. Difficulties were encountered. Because pressures up to nine or ten atmospheres were involved, it was necessary to carry out the reactions in narrow tubes. From these it proved to be impracticable to obtain representative samples of the aqueous phase. Because of the high pressures within the tubes, scattering of the contents was unavoidable when the seals were broken, and only a mixture of the solid and liquid phases could be recovered. Furthermore, the rapid evolution of sulfur dioxide was accompanied by a deposition of more solid. It was necessary, therefore, to confine atten-
Figure 5-Distribution of Phosphoric Acid i n Aqueo u s Solutions When C a l c i u m Phosphate Is Treated w i t h Sulfur Dioxide
tion t o the solid mass resulting when the excess sulfur dioxide and water had been removed from the reaction mass by aspiration as quickly as possible. Action of Water and Sulfur Dioxide on Tricalcium Phosp h a t e a t 100" C. A number of tubes were charged each with 30 grams tricalcium phosphate, 15 grams liquid sulfur dioxide, and 10 grams water, and sealed. To insure maximum surface contact, the tubes were placed horizontally in a thermostat. The temperature was raised gradually to, and maintained at, 100" C. A t 12-hour intervals a tube was withdrawn, and the contents discharged and dried a t reduced pressure. The residue was analyzed for water-, citrate-, and acid-soluble phosphoric acid and calcium oxide. The results are assembled in Table V, and charted in Figure 5 . Apparently, a transformation of the phosphates takes place in steps: first dicalcium, then monocalcium phosphate, and finally phosphoric acid. The proportion of acid-soluble or tricalcium phosphate decreased rapidly with time, practically disappearing in 100 hours. The citrate-soluble, with a ratio of calcium oxide to phosphoric acid, practically that for dicalcium phosphate, increased for about 24 hours, then decreased continuously. The water-soluble increased continuously with time, rapidly a t first and then asymptotically, attaining 95 per cent in about 90 hours, The ratio of calcium oxide to phosphoric acid, which was that for monocalcium phosphate to this point, began to decrease rapidly, indicating the presence of free phosphoric acid in the reaction mass at longer time intervals. A practically complete conversion t o available phosphates was attained in 95 hours and to water-soluble form a t a somewhat longer time interval. T a b l e V-Analysis of Solid Residues f r o m Liquid Sulfur Dioxide a n d Water T r e a t m e n t of T r i c a l d u m Phosphate a t loo3 C. WATERSOLUBLE CITRATESOLUBLE ACIDSOLUBLE TIMB CaO : Pi01 CaO : P2Or CaO :Pi06 Hours % Ratio % Ralio % Ratio 1.71 55.59 0.420 26 .-16 0.18 18.81 12 2.07 34.91 0.83 30.20 0.560 35.79 24 3.44 21.16 30.42 0.86 0.430 48.52 36 5.14 14.36 48 0.410 0.79 28.50 57.23 ._ 7.46 9.57 0.80 0,462 20.18 70.31 60 13.31 6.38 15.13 0.85 82.72 0.393 72 15.97 9.57 5.38 0.78 0.351 85.94 84 152.12 0.63 4.37 0.81 95.60 0.210 96 ..... 0.00 0.81 0.223 3.80 108 96.28 0.04 2,677.6 0.157 3.05 0.75 96.97 120 10,440.0 0.01 4.64 0.78 0.126 132 95.83
Vol. 23, No. 11
A striking feature of these residues was the presence of elemental sulfur and calcium sulfate, both in increasing proportions with increasing time intervals. Calcium sulfite, however, was always present until free phosphoric acid appeared. It is necessary to conclude that autoxidation of sulfur dioxide or oxidation of calcium sulfite by sulfur dioxide took place. Similar results have been reported by others. Carter and Butler (11) found ferrous phosphate oxidized to ferric phosphate in the presence of phosphoric acid and sulfur dioxide, free sulfur being formed. Hammick (20) and Schott (33), independently, studied the action of sulfur dioxide on metallic oxides. With calcium oxide, the sulfate and sulfide are formed a t temperatures above 300" C. With the oxides of lead and of manganese they found free sulfur. These results were duplicated by Birnbaum (6),who also found that calcium sulfite was formed a t 150" C. in the presence of water. With dry sulfur dioxide, a t 236" C., a substance with the composition approximating the formula Ca6S& resulted. Veley (37) found that dry sulfur dioxide does not react with calcium oxide below 300" C., but above 380" C. calcium sulfate and calcium sulfide are formed with the liberation of free sulfur. Farnell (16) found that, in aqueous solution, calcium sulfite and bisulfite are readily oxidized to the sulfate, and found sulfur trioxide to be formed above 100" C. Linden (%) found that calcium sulfite is less soluble a t higher temperatures, the dihydrate being the stable solid a t lower temperatures, the anhydrous salt a t the higher. Forester and Kubel (16), working as high as 600" C., found calcium sulfite transformed to the sulfate and sulfide a t the higher temperatures a t atmospheric pressure. Bichowsky (6) found calcium sulfite to be unstable above 160' C., and a reversible reaction between the sulfite and sulfur on the one hand and calcium thiosulfate on the other. In Table VI are assembled the results of a series of experiments similar to those just described, but where ground phosphate rock replaced the precipitated tricalcium phosphate. Except that the reaction rates were much slower, the results were of the same character, as shown by the graphs in Figure 6. Free sulfur was found in the reaction products, sometimes abundantly. JOOl
144
180
216
T ~ MI NEHOURS Figure 6-Distribution
of Phosphoric Acid i n Aque-
ous Solutions When Phosphate Rock Is Treated with Sulfur Dioxide
The adaption of these results to a practical method would involve further and probably expensive research to find a material of construction strong enough to meet safely pressures of 10 atmospheres or more and, a t the same time, resistant to the corrosive action of aqueous sulfur dioxide. Otherwise, no serious engineering considerartions seem involved. If absorption of sulfur dioxide by lime were alone to be considered, a good recovery of the former might be anticipated for its re-use. Unfortunately, a serious destruction and permanent loss of this reagent appears unavoidable. Further investigation in this direction appearing too dubious for a promise of industrial success, attention has been given to the action of gaseous sulfur dioxide on the phosphates of calcium.
INDUSTRIAL AND ENGINEERING CHEMISTRY
November, 1931
T a b l e VI-Analysis of Solid R e s i d u e from T r e a t m e n t of P h o s p h a t e Rock at looo C. b y L i q u i d S u l f u r Dioxide a n d Water WATER-SOLUBLE CITRATE-SOLUBLE TIME Ca0:PzOr CaO: PzOr ACID-SOLUBLE CaO:PzO, r,_ n "UrJ
m
6 12 24 36 48 60 72 84 96 108 120 132 144 156 168 192 204 216 228 240 252 264
2.44 2.66 14.32 21.48 27 91 32 40 36 49 41.02 44.50 49.69 52.89 50.52 59 10 62 55 64 48 70 04 72 14 74.54 76.02 79.98 79.89 60.75
o,m -
D . . . _
na'ru
70
0.54 0.48 0.45 0.42 0.50 0.42 0.41 0.49 0.42 0.40 0.41 0 45 0.42 0.43 0.41 0.41 0.44 0.40 0.42 0.40 0.41 0.43
47 11.72 11.80
i;:;:
20.62 23.29 23.98 23.52
;:g
21.25 21.15 17.01 17.08 15.45 14 75 10.85 11.27 10.98
m
D_l_.^
nullu
D_...-
nur,u
70
82 0.81 0.81
8 7 , o9 85.62 74.15
1,96 2.01 2.21
0.79 0 80
46.94 40.15
3.97 3 61
80
26 33 23 51
5.37 6.01
16.18 14.49
8 24 9.10
:,:' :o.;A
0.82
g:;:
o
80 0.81
0.80 0.82 0.79 0.84 0.78 0.82 0.81
t:;::
2":'
:;
j:;:
:::$,
2;
i::
i:!;
9.98 9,16
12.59 13.55
E:8 ;.:2 5 it::; 16.88
Absorption of Gaseous Sulfur Dioxide by Tricalcium Phosphate
The procedure adopted was to stir continually the solid phosphate in contact with gaseous sulfur dioxide a t a fixed temperature and pressure; to replace, at noted time intervals, the gas absorbed; and, when a steady state was attained, to remove the reaction product and analyze it for its content of available phosphoric acid. Usually, a 5-gram sample of precipitated tricalcium phosphate was subjcrted to treatment. The sample was air-dried and so contained enough water, apparently (since moisture drops were always congealed on the walls of the reaction chamber), to assure the continuous saturation of the gas. A sketch of the apparatus is shown in Figure 7. Two accurately graduated glass cylinders 5 and 5' served t o introduce measured volumes of the gas into the system from the stock supply. For instance, if cocks V4 and V2 be opened, the mercury level in 5 is depressed and gas forFed from 5' through the tube fitted with cock V5 into reservoir 4 . This is connected by tube T4' with the reaction chamber, which in turn is connected t o a second reservoir 4 by tube T4. For reservoirs were employed bottles of 3 liters capacity each. The stoppers were closely fitted, held in place by wiring, and protected by mercury seals. They were insulated by a wooden box fitted with a thermometer. Reservoir 4 was connected t o a manometer and t o 1
-~~ ~~~
~~
___
-____ >z" 6
~
0
Figure 7-Apparatus
mm>rrr
f o r Measuring Absorption of S u l f u r Dioxide b y Calcium P h o s p h a t e
a graduated mercury reservoir which could be raised or lowered at will by a wire passing over a pulley to a spool firmly anchored on the work-table. The reaction chamber or absorption tube, mounted vertically, was made from Pyrex glass and was 3 cm. in internal diameter and 30 cm. in length. The closely fitted stopper carried a pyrometer and leads. The lower 8 cm. of the reaction chamber was surrounded by an especially constructed furnace in which the heating element was made from nickel wire. Heating was accomplished with a 110-volt current from the laboratory lighting circuit and controlled by means of a rheostat. Above the furnace and surrounding the absorption tube was an
1267
asbestos shield and a copper-fin radiator, of which the main purpose was t o protect the glass seals of the connections t o the gas reservoirs. On opposite sides of the reaction chamber, electro marnets 2 and 2' were mounted. A water-driven Pelton wheel conboiled an alternator which permitted a &volt current from a storage battery t o flow through the magnets alternately and actuate a soft iron bar, sealed in glass 3, thus stirring vigorously the contents of the reaction chamber but not a t such a rate as t o produce a dust cloud higher than the top of the furnace. The apparatus being filled with sulfur dioxide, after sweeping out all air, the sample of phosphate t o be treated was weighed into a thin-walled Pyrex-glass thimble and lowered into the reaction chamber. The stirrer was placed, the apparatus closed, and the reaction chamber brought t o the desired temperature and pressure. For the first few minutes both pressure and temperature were altered slowly, after which they were brought t o the desired point as quickly as possible and the time noted. As absorption proceeded, the pressure was maintained by changing the level in the graduated mercury reservoirs. By noting the change in level, the volume of gas absorbed was determined.
These volume determinations were not accurate because of the fact that the gas in the absorption tube was expanded, being a t a higher temperature than in the remaining portions of the apparatus. As long as the temperature of this portion of the apparat,us was kept constant, the error was negligible, especially because the volume of the heated portion of the absorption tube was only 0.01 of the total system. Quite accurate volume measurements were made a t 5 and 5' of the gas necessary to replace that which had been absorbed. The indefinite character of the phosphates treated and the occurrence of complicating side reactions in the absorption chamber did not justify any attempt a t accurate determinations of rates. The general character of the approach to a steady state was the information sought, and only when equilibrium had been attained was an accurate estimation of the volume of gas absorbed a matter of concern. Table VII-Absorption of S u l f u r Dioxide b y T r i c a l c i u m P h o s p h a t e Temp., 24' C . ; pressure, 1.5 atm.; Car(POdz, 5 grams: total SO1 absorbed. 165 cc.; total PzOs, 36.4%; availabl- PzOs, 8.89%; recovery, 34.43% TIME SO2 TIM= so: TIME SOY Hours Cc. Hours cc. Hours cc. 4 50 24 160 48 165 150 36 165 60 165 16 66 165 Temp., 65' C.; pressure, 1.5 atm.; Ca3(POi)i. 5 grams; total SOP absorbed, 825 cc.; total PzO,, 35.90%; available PzOs, 23.04%; recovery, 64.27, 12 330 24 520 40 750 800 16 430 36 700 48 57 825 Temp., 65' C.; pressure. 2 atm.; Caa(POp)E, 5 grams: total SOP absorbed, 860 cc.; total PzOs, 36.01%; available PnOa, 25.01%; recovery, 69.5% 800 7 250 25 625 49 17 490 41 760 65 850 73 860 Temp., 65' C . ; preqsure, 1.5 a t n ; Car(POl)r, 5 grams: total SO? absorbed, 850 cc.; total PnOa, 35.51 %; available PzOr, 23.12%; recovery, 65.10% 24 560 48 790 96 850 100 850 32 660 54 810 105 850 Temp.. 65' C.: pressure, 2 atm.: CadP,Oa)r, 5 pramc; total SO2 absorbed, 805 cc.; total Pz01, 35.75%; available P206, 22.93%; recovery, 64.15% 16 465 24 580 40 750 47 805 Temp., 65" C.; preisure., 1.25 atm.; CadPOi)!, 5 grams; total Sot absorbed, 454 cc.; total PzOs. 25 91 '?&' available PzOa, 12.24%; recovery, 47.25% 19 354 22 414 25 454 Temp.. 200' C.; pressure 1.5 atm: Car(PO4h 5 grams. total SO2 absorbed, 115 cc.; total PiOs. 40.2,5%; availabl; PzOs, 7 Si%.,; recovery, 19.25% 15 74 100 112 110 5 88 105 124 115 17 40 55 98 110 23 Temp., 450' C.: precsure, 1.5 atm.: CadP0i)z. 5 grams: total SO, absorbed, 750 cc.; total PzOs, 35 5 % : available PzOs, 22.07%: recovery, 62.15% 2 40 48 675 72 740 42 640 66 725 90 750 96 750
The results of several series of observations a t 20" C., 65" C., 200" C., and 450" C. are assembled in Table VII. For convenience in comparing the results they are charted in
1268
IXDUSTRIAL AAVD ENGINEERING CHEMISTRY
Figure 8. At 24' C. and under a pressure of 1.5 atmospheres, absorption ceased after 24 hours. The volume of sulfur dioxide absorbed was 165 cc. and its weight approximately 0.65 gram. The available, or water- and ammonium citratesoluble phosphoric acid was 8.93 per cent of the product, while the total phosphoric acid (P205)was 36.4 per cent; hence, the possible recovery was but 24.4 per cent. Addition of strong mineral acid to a sample of the reaction product produced a strong evolution of sulfur dioxide. Microscopic examinations showed the presence of crystals which could be identified as calcium sulfite. No other compound of the mixture, however, was identified positively. I loo0
3 9
2 9 "e M@-
- - -.
2oO.r ~ m /8% k
40 60 TIME OF TRENTMENT / N HOURS
80
Figure 8-Relative Rates of Absorption of Sulfur Dioxide b y Calcium Phosphate a t Different Temperatures
At 65" C. with a pressure of 1.5 atmospheres, a steady state was reached in about 70 hours. The volume of gaseous sulfur dioxide required to replace that which had been absorbed was 825 cc., weighing approximately 3.24 grams. The product gave, on analysis, 23.04 per cent available phosphoric acid, 35.91 per cent total phosphoric acid, hence 64.2 per cent recovery. Microscopic examinations showed gypsum as well as calcium sulfite. The latter was in excess, and the product fumed freely, emitting sulfur dioxide when a sample was touched with concentrated sulfuric or hydrochloric acid. A small amount of free sulfur was observed. A duplicate series a t 65" C. and 1.5 atmospheres came to a steady state in about the same length of time with a slightly greater absorption of gas, 850 cc. or 3.34 grams of sulfur dioxide. Analysis showed the total phosphoric acid to be 35.5 per cent, the available to be 23.12 per cent, making a recovery of 65.1 per cent. Gypsum and sulfur, in small quantities. were noted. A series run a t 65" C. and 2 atmospheres pressure, accidentally interrupted a t 47 hours, showed but little improvement in the absorption by increasing the pressure. The volume absorbed was 805 cc., the total phosphoric acid 35.75 per cent, available 22.93 per cent, hence a recovery of 64.15 per cent. The run was repeated, and again it was found that it required about 70 hours to reach equilibrium. The gas absorption was, finally, 860 cc., or 3.38 grams of sulfur dioxide. The total phosphoric acid was 36.01 per cent, the available 25.01 per cent, making a recovery of 69.5 per cent. Examination of the product showed the presence of gypsum, calcium sulfite, and free sulfur. The latter was removed by leaching with carbon disulfide and recovered by evaporating the solvent, but no quantitative determinations were made. The results from this last series seem to show that increase of pressure does produce an increase in absorption by the solid phase (or phases), as might be anticipated from the principle of Le Chatelier. The effect of the increase from 1.5 to 2 atmospheres pressure was small but real. In support of this conclusion are the results from an uncompleted run a t 1.25 atmospheres pressure and 25" C. which was terminated a t 25 hours. The absorption appeared to be distinctly less than a t 1.5 atmospheres pressure, but the difference was again small.
Vol. 23, No. 11
Runs a t temperatures above 80" C. showed a diminishing absorption of sulfur dioxide and recovery of available phosphoric acid until the temperature was raised t o 200" C. A run a t this temperature (200" C,), cited in Table VII, shows a recovery of but 19.25 per cent. Relatively large proportions of sulfur and sulfates were foundinthe mixture. Above 200" C there was an increase in available phosphoric acid Finally, a run was made a t 450" C. and a pressure of 1.5 atmospheres. A steady state was reached in somewhat more than 70 hours with a disappearance from the gas phase of 750 cc. or 2.95 grams of sulfur dioxide. I n the solid residue was 35.5 per cent total phosphoric acid, available 22.07 per cent, making a total recovery of 62.15 per cent. Very little solid sulfite was present, but much calcium sulfate and relatively much elemental sulfur. It would seem, therefore, that a t higher temperatures sulfur trioxide and not the dioxide is the effective agent in converting the insoluble phosphates to the available form. In no case in these experiments was much more than two-thirds of the total phosphoric acid made available. It became clear that the process as here developed is too extravagant in unrecovered available phosphoric acid and ineffective sulfur dioxide to expect it t o be of use in a commercial development. There does appear to be a definite relation between the amount of sulfur dioxide absorbed or reacted upon by tricalcium phosphate and the phosphoric acid thereby made available. In Table VI11 are assembled the figures found for steady states and, in Figure 9, the corresponding graph. The points apparently lie fairly well on a smooth curve, and the law of absorption would be given by the equation of the curve, which can be determined empirically. It would have little or no value in terms of the parameters chosen, and it cannot be expected to throw any light on the complex group of reactions not known to be involved. By extrapolation it will be seen that it would take 1400 to 1500 cc. sulfur dioxide to transform completely all the phosphoric acid in a 5-gram sample of tricalcium phosphate to the available form. This is an amount quite outside serious consideration for a commercial process. Table VIII-Recovery of Phosphoric Acid by T r e a t m e n t of 5 Grams Tricalcium Phosphate w i t h Gaseous Sulfur Dioxide sot P206 so2 PZOl sot PlOl CC. % CC. 70 cc 70 11.5 19 25 750 62.15 850 65 10 BO5 64 16 860 69 50 165 24 43 454 47 18 825 64 20
Action of Sulfur Dioxide on Tricalcium Phosphate and Phosphate Rock at Higher Temperatures
A procedure was followed which was quite different from that described in the foregoing chapter. A more rugged apparatus was employed. The phosphate to be treated, usually 50 grams, was heated in a tube through which a current of gas was passed slowly for a fixed period of time, usually about 5 hours. Various gas mixtures were employed in which sulfur dioxide was always present. I t was early apparent that sulfur trioxide rather than the dioxide was often the reacting agent. I t was anticipated that metaphosphates and pyrophosphates might be formed, and these would be more or less converted to normal phosphates by hot water or steam. If such reactions did take place, the writers were unable to establish satisfactory proof. Figure 10 is an outline drawing of the apparatus. The reaction chamber or tube in some of the experiments a t atmospheric pressure was a heavy quartz combustion tube of 2.5-cm. diameter. Later, and in all experiments carried out a t pressures greater than one atmosphere, an iron tube, 6 cm. in diameter, was employed. It was thought that the iron surface or the oxides which might be formed on the inner surface would have positive values as catalysts. This tube was surrounded for the greater part by a furnace built of fire brick and asbestos cement. The furnace was 2 feet (0.91 meter) in length and the internal meas-
INDUSTRISL AND ENGINEERING CHEMISTRY
Yovember, 1931
9
4 7
8
$ 5
6
current p a s s i n g t h r o u g h the m a g n e t , the knife-edge was drawn down on the rubber tubing, pinching off the flow of gas. Simu 1t aneousl y the current ceased to feed the motor driving the blower, thus cutting off the blast of air to the furnace. The 110-volt current t h r o u g h the motor driving the blower was controlled by the electro-magnet 14. This in turn carried a 6-volt
1269
mum conversion. Three experiments with tricalcium phosphate point to the same conclusion. At 500" C. the available phosphoric acid was 23.1 per cent, a recovery of 50.22 per cent; a t 600" C. available phosphoric acid was 38.4 per cent, recovery 83.49 per cent; at 700" C. available phosphoric acid was 26 per cent and the percentage recovery was 56.52. The recoveries with tricalcium phosphate were smaller than those obtained with ground phosphate rock treated in the same manner and a t the same temperature. This decrease in recovery is perhaps due to the fact that the tricalcium phosphate contains relatively more phosphoric acid than does the rock and would therefore require a longer period of contact with the sulfur trioxide to complete the conversion. However, in each case sulfur trioxide was issuing from the exit freely when the heating was stopped and there was every appearance that absorption of the oxide had ceased. Table IX-Available Phosphoric Acid f r o m Ground Rock r-d Ti:c a l c i u m Phosphate Treated w i t h Sulfur Trioxide a t Various Temperatures RECOVERY TEMPERATURE AVAILABLE PzO, O
c.
%
%
PHOSPHATE ROCK
375 400 460 500
600 700 800 900 1000 1100
11.04 8.00
26,OO 28.20 28.30
26.50 24 00 15.60 11.20 4.40
35 07 25 81 83 28 90.96 91.01 83.66 77.42 60 32 36.13 14,20
TRICALCIUM P H O S P H A T E
500
Figure 9-phosphoric Acid Re- rupted a connection in the 1.5as ~~~~~i~~ of Sulfur volt circuit as the rod expanded
or contracted in resuonse to the temperature changes with i n the furnace. The furnace was fitted with a pyrometer, 9, read from a wall galvanometer, and over the entire range of experiments here recorded temperatures were maintained within 5" C. Sulfur dioxide, alone or mixed with other gases, passed through an intake tube, 1,fitted with a manometer to a mixing chamber, 3. This chamber consisted of a wide-mouth bottle, and the stopper was held in place by an adjustable metal frame. The stopper carried three tubes: one was the intake tube; the second was a flexible connection with a mercury level, 2; and the third conducted the mixed gases t o the steam generator 4, where they were saturated with water at the boiling point. In experiments where the addition of water was not desired, the steam generator served as an additional mixing chamber in which the gases could be preheated. The gases then passed t o the reaction chamber by way of a conducting tube held in position by a union, 6. The union afforded an easy access to the interior of the reaction chamber. To carry out reactions at pressures above atmospheric, the escaping gases from the reaction chamber were forced to pass a mercury column, 17, before leaving the system. The back pressure thus produced was controlled by the mercury level 2 attached t o the mixing chamber 3. Before making any experiments with sulfur dioxide, the effect of sulfur trioxide upon rock and tricalcium phosphate was tried. It was thought that maxima results might be secured, since it was assured that sulfur trioxide would be more effective than the dioxide. The reactions were carried out in a quartz tube a t atmospheric pressure. The sulfur trioxide was obtained by heating pure concentrated sulfuric acid to vigorous boiling. These fumes, with the accompanying water vapor, were conducted over the heated phosphate. Immediately following the expiration of the 5-hour period, the reaction chamber was opened, and the residue within withdrawn for analysis. The water-soluble phosphoric acid was always small, but the citrate-soluble much greater. The results from phosphate rock are assembled in Table IX; the first column shows the temperature; the second column, the percentage or available phosphoric acid in the reaction product obtained a t that temperature; and the third column, the percentage of the total phosphoric acid converted to the available form. At approximately 600" C. there is a maxiDioxide Absorbed Phosphate
bv Calcium
600 700
23.10 38.40 26.00
50.22 83.49 56.52
The falling off in recovery of available phosphoric acid with an increase of temperature beyond 600" C. may be due to a n increasing decomposition of the sulfur trioxide or the formation of pyro- and metaphosphates resistant to water and citrate solutions. Some support for the latter suggestion was
Figure 10-Apparatus for T r e a t m e n t of Calcium Phosphate or Phosp h a t e Rock w i t h Sulfur Dioxide a t High Temperatures
found in the fact that the reaction products, on long standing in contact with the solvents, did yield a greater percentage recovery of phosphoric acid than in the cases where the standard time limit of the Association of Official Agricultural Chemists was closely followed. As this is a property of the untreated materials also, the explanation is not very convincing. There seems to be no available explanation for the progressively lower results a t temperatures below 600" C. I n all cases sulfur trioxide was issuing freely from the reaction chamber some time before the conclusion of the experiment, and a n inspection of this waste gas failed to furnish clues. Several sets of experiments were made with sulfur dioxide or mixtures of sulfur dioxide and air saturated with water vapor. The results are assembled in Table X, and they prove to be surprisingly low. On the hypothesis that, at the temperature of the experiments, meta- or pyrophosphates might have been formed, a portion of each of the reaction products was re-
INDUSTRIAL A X D ENGIArEERli\TG CHEMISTRY
1270
fluxed with boiling water for 24 hours, the rrater evaporated, and the residue analyzed for available phosphates. The results are stated under the heading, "After Treatment," and the recoveries in the table are based upon them. The results show that an increase in temperature beyond 500" C. decreases the effectiveness of the treatment. A mixture of air designed to oxidize the sulfur dioxide was ineffective, nor did the further addition of water to the sulfur dioxide-air mixture improve the reaction. Likewise an increase of pressure was only slightly effective. Table X-Rock
TEMP. O
c.
Treated w i t h Gaseous Mixtures for 5 Hours in Iron T u b e , Residue Extracted w i t h Water AVAILABLE PzOa Before Alter treatment treatment RECOVERY
%
%
%
SOP a t 1 atmosphere 11.11 5.20 10 20 5.00 5.80 2.10 12.10 5.90 SOz and air at 1 atmosphere 12.00 5.80 11.70 5.60 1.70 4.80 SOa and air a t 1.5 atmospheres 8.10 16.30 7.30 15.30 7.80 12.20 4.20 11.80 2.80 4.20 SOa and Hz0 a t 1.5 atmospheres 5.40 12 10 5.10 10.80
400 600 so0
900 500 600 800 320 400 500 600 800 400 450
TEMP. O
c.
38.7 37.7 15.5
300 350 300 350 400 350 400 450
3.00 4.20 5.03 6.31 8.10 5.30 SO,,air, and Hz0 a t 2 atmospheres 4.20 5.50 SO,, air, and Hz0 at 3 atmospheres 6.10 7.30 16.40 SO9 and air a t 2 atmospheres 3.60 6.80 8.30
3 6 9
60 MESH
0.40 2.80 11.50
28.4 26.2 19.5
4.8 15.2 30.6
5.30 7.00 10.20 11.80 12.70 13.10
30.6 28.5 27.9 27.1 26.5 25.6
17.0 24.5 36.5 43.5 47.9 51.1
3 6
62.1 5R.3
38.3 52.6
100 MESH
24.80 30.70
53.2 49.3 39.3 38.0 13.5 39.0 34.8
%
SOz, air, a n d H10 a t 1.5 atmospheres
300 350 376 400 450 550
Table XII-Effect of Particle Size a n d Regrinding of P h o s p h a t e Rock W h e n Exposed t o Gaseous Sulfur Dioxide a t 450° C. TIXE AVAILABLE PrOl TOTAL PZOS RECOVERY Hours % % %
80 MESH
Treated w i t h Gaseous Mixtures in Iron T u b e for 5 Hours AVAILABLE PZOS RECOVERY
%
drawn from the tube, a sample taken for analysis, and the remainder ground in a porcelain mortar. The ground product was returned to the furnace and the process repeated successively as indicated in Table XI1 where the results are assembled. Regrinding and retreating always increased the recovery of available phosphates. Also, the finer mesh grades gave substantially better recoveries, thus indicating that the surface exposed is a dominant factor in controlling the yield of available phosphates.
35.8 32 9 18 7 39.0
Sulfates were always present in the reaction products and relatively large amounts of sulfur were noted except a t the higher temperatures. Sulfides dominated with the products at higher temperatures, in each case evolving hydrogen sulfide on the addition of mineral acid. The experiments were then repeated in an iron tube, somewhat higher pressure being possible. The results are assembled in Table XI. The residues were not subsequently treated with water, since the increase in available phosphates was too low to justify the labor. A striking feature was a large accumulation of ferrous sulfate (FeS04.7H20) in the cooler portion of the tube outside the furnace. I t appears that much of the sulfur dioxide was oxidized to sulfur trioxide and this latter reacted far more readily with the iron of the tube than with the ground rock. Sulfides were also present as in the residues from the quartz tube. Table XI-Rock
Vol. 23, No. 11
14.0 1-5.2 18.5 23.3 30.2 18.3 14.8 18.7 21.6 26.4 63.3 12.00 24.10 31.01
The best recoveries in either the quartz or the iron tube were realized at temperatures near 450" C.; this temperature was therefore selected for most of the subsequent, work. A probable explanation of the disappointingly low recoveries is that a protective coating formed over the surface of each particle. Consequently, a series of mechanical separations of the rock were made. Each grade was treated a t 450" C . for the time indicated in the table. The reaction product was then with-
A series of experiments were undertaken t,o test the effect of time. Thpy were made with finely ground rock and carried out at 450" C. The results are assembled in Tahle XIII. It is evident that an increase in the time of treatment has a small but real effect on the recovery. Table XIII-Effect of T i m e of Exposure t o Gaseous Sulfur Dioxide a t 450" C. on Percentage Recovery from P h o s p h a t e Rock TIME AVAILABLE PzOr TOTAL PzOa RECOVERY Hours % % % 28: 3 25.1 7.1 27.8 29.8 8.3 24.2 30.1 7.3 7.4 24.3 30 4 8.5 30 6 27.8 6.8 23.4 28.4 7.4 23.6 31.3
Summary
The principal results of the investigation may be summarized as follows: The vapor pressure measurements of mixtures of sulfur dioxide and water have been extended t o 90' C. Below -2.6" C., a solid hydrate of sulfur dioxide is stable. The vapor pressure is always lowered by contact with solid tricalcium phosphate. The transition point, below which the hydrate of sulfur dioxide is a stable solid, is lowered t o -4.2' C. by the presence of the phosphate of calcium. Metastable conditions are frequently realized a t temperatures below 20" C. When tricalcium phosphate is treated a t 25" C. with an aqueous solution of sufficient volume and kept saturated with sulfur dioxide, a complete solution of the phosphoric acid is attained in time. At first the acid is extracted from the solid more quickly than is the base. Ultimately they are present in the aqueous phase in about the proportions corresponding t o tricalcium phosphate. The solid residue is calcium sulfate, calcium sulfite, and free sulfur, whereas the solid obtained from evaporation of the liquid phase is calcium sulfate and a mixture of the mono- and dicalcium phosphates. When tricalcium phosphate or finely ground phosphate rock is treated a t 100" C. with a two-liquid mixture of water and sulfur dioxide, a complicated series of reactions takes place, and equilibrium is attained slowly. Elemental sulfur is formed, showing an uneconomical decomposition of sulfur dioxide. From the reaction product practically all of the phosphoric acid can be extracted with water, together with lime in about the proportion required by the formula for monocalcium phosphate. The solid residue contains much calcium sulfate as well as more or less sulfite and phosphates. The pressure of the vapor in the reaction chamber a t 100" C. is approximately ten atmospheres. When tricalcium phosphate is treated with sulfur dioxide in the presence of small amounts of water, there is a gradual trans-
November, 1931
INDUSTRIAL A N D ENGINEERING CHEMTSTR Y
formation of the phosphoric acid to the available form. The favorable temperature for the transformation is about 65’ C., and a steady state or equilibrium is attained when about twothirds of the phosphoric acid has become available. On increasing the temperature, the amount of transformation falls off, there is much decomposition of the sulfur dioxide, and much more sulfur dioxide is apparently absorbed than corresponds t o the available phosphates produced. Free sulfur, sulfates, sulfides, and sulfites are formed. Probably a t moderate temperatures thiosulfates are formed. When tricalcium phosphate or finely ground phosphate rock is heated in a current of sulfur dioxide, the maximum transformation of the phosphoric acid t o an available form is attained at about 450’ C. Probably it is due t o a maximum formation of sulfur trioxide a t this temperature, the reaction being between the phosphates and the trioxides rather than the dioxide of sulfur. T h e highest possible transformation seems t o be about two-thirds of the total phosphoric acid present in the phosphate. Large amounts of sulfates, sulfides, and elemental sulfur, but no sulfites or thiosulfates, are formed a t high temperatures. Protective coatings form on the reaction mass which can be removed mechanically. With increase in temperature the available phosphate in the reaction product decreases rapidly to a vanishing amount at 1100” C. and 2.5 atmospheres.
(9) (10) (11) (12) (13) (14) (15) (16) (17) (18) (19)
Literature Cited
(31) (32) (33) (34) (35) (36)
(1) (2) (3) (4) (5) (6) (7) (8)
Aufsalzteil, Z . angew. Chem., 34, 272-5 (1921). Bailey, J. Chem. Soc., 121, 1813 (1922). BaumC and Tykociner, J . chim. phys., 12, 270 (1914). Bergstrom, J . Phys. Chem., 26, 358 (1922); Bull. s c i . a t a d . roy. Belg.. 6, 529 (1919). Bichowsky, J . A m . Chem. Soc., 46, 2225 (1924). Birnbaum, Ber., 13, 651 (1880). Briner, J . chim. phys., 4, 476 (1906). Briner and Wronynski, Z . anorg. Chem., 63, 49 (1909).
(20) (21) (22) (23) (24) (25) (26) (27) (28) (29) (30)
(37) (38)
127
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Sorption of Water Vapor b y Paper-Making Materials I--Eff ect of Beating’ C. 0.Seborg a n d A. J . Stamm FORESTPRODUCTS LABORATORY,% FORESTSERVICE, U. S. DEPARTMENT OF AGRICULTURE, MADISON,WIs.
M e a s u r e m e n t s were m a d e of the m o i s t u r e contentrelative h u m i d i t y a n d m o i s t u r e content-electrical conductivity relationships of p u l p a n d stuffs. An a p p a r a t u s was designed for m e a s u r i n g t h e relative h u m i d i t y relationship of materials a t a t m o s p h e r i c pressure, using a m e t h o d in which air, humidified by passi n g t h r o u g h various s a t u r a t e d s a l t solutions t o give the desired relative humidities, was circulated a b o u t the adsorbent m a t e r i a l suspended from a sensitive q u a r t z helix balance. The extensions of this balance were
measured w i t h a cathetometer. T h e hydration (hygroscopicity) of the pulp, which t h e regularity of t h e curves showed to be physical r a t h e r than chemical, was not affected by beating. Electrical conductivity-moisture c o n t e n t relations likewise showed n o differences between beaten a n d u n b e a t e n Pulp. The so-called h y d r a t i o n of t h e paper i n d u s t r y is n o t a t r u e hydration. It is probably a p h e n o m e n o n of fiberfiber bonding r a t h e r than of fiber-water bonding.
. . .. .. . . .. .. .. HE objectives in this series of studies are to determine the hygroscopicity of pulps, stuffs, and the various components of wood under normal atmospheric conditions; to determine the relationship of the sorption hydration (hygroscopicity) to the so-called hydration of the paper industry, that is, the change in properties of fibers produced by beating which in the past has been assumed to result from a change in their ability to retain water; to determine the extent to which this sorption hydration affects the fiberbonding properties; and t o critically study the desorptionadsorption hysteresis. The various materials to be studied are sulfite, kraft,
T
1 Received July 8, 1931. Presented before the Division of Cellulose Chemistry a t the 82nd Meeting of the American Chemical Society, Buffalo, N. Y.,August 31 t o September 4, 1931. Maintained at Madison, Wis., in cooperation with the University of Wisconsin.
soda, and semi-chemical pulps; bleached pulps; beaten pulps; isolated wood components, such as lignin, Cross and Bevan cellulose, and alpha cellulose from both pulp and Cross and Bevan cellulose; and furnishing materials, such as coatings, sizings, and fillers. This paper deals only with the effect of beating on the sorptive hydration of a sulfite pulp. The word “sorption” is the same as used by McBain (8) to include both adsorption and absorption phenomena. Previous Sorption Studies
A number of investigators have published papers on subjects pertaining to the hysteresis of the adsorption and desorption of water vapor by various materials, which ranged from silica gel to textiles and paper ( 1 , W, 6, 7 , 11, 12, 13, 15, 16, 17, 18, 19).
-411 of these investigators, with the exception of Patrick (11)
