INDUSTRIAL AND ENGINEERING CHEMISTRY
July 1948
TABLE
ON
x.
AND
'
EFFECTOF NITRIFICATION B.0.D. CALCULATED CARBoNACEoUs B*o*D*
THEORETICAL
Yo,
Yn,
Nitrite Nitrogen Equiv. to Yn, P.P.M.
Probable Error P.P.M. in Y e , Oxygenb P.P.M. 1.42 0.07 2.39 0.12 3.06 0.15 3.54 0 17 3.89 0.19 4.15 0.20 4.36 0.21 (4.55)C 0.21 (4 74) 0.22 (5 00) 0.22
Y O+
n,
Time, P.P.MSn P.P.M. Days OxweOxygen Nitrogen 1 1.39 0.03 0.009 2 2.35 0.04 0,011 3 3.01 0.05 0.014 4 3.47 0.07 0.02 5 3.79 0.10 0.03 6 4.01 0.14 0.04 7 4.16 0.20 0.06 8 4.26 0 29 0.08 9 4.34 0.40 0.12 10 4.39 0.61 0.18 a For Y c , K = 0.160,L = 4.50. For Y c + n, K = 0.118,L = 5.20. 0 Figures in parentheses include active nitrification B.O.D.
Minus Y* Probable Error in
-
-
7
f + +
0.15. This variation was unpredictable and was apparently independent of the origin of the sewage (military, hospital, or domestic). A statistical comparison of the values found in a series of two dilutions showed t h a t the higher concentration of * sewage on the average gave the higher K value, and that the Coefficientof variation was greater on the lowest concentrations, Careful studies of several series of dilutions of the same sewage showed t h a t t h e data could be divided into two groups, one of which was consistent and the other inconsistent. Examination showed t h a t each inconsistent value was one in which the oxygen demand equivalent of the nitrite nitrogen exceeded the Drobable in the B.0.D. data. In other words, the in K were observed to result from the unpredictable formation of amounts of nitrite after incubation periods of 3 days 01 more in
1295
some series, and dilution was an important factor in the onset of nitrification. The nitrite formation occurred earlier and more frequently in the lower concentrations of sewage. The oxygen demand equivalent of the nitrite formation was so small t h a t it could not be recognized by plotting the observed B.O.D. results. The great effect of the formation of small increments of nitrite on the K value derived from a series of observations was demonstrated. The conclusions which may be based on this study are: 1. The general theory of clear-cut carbonaceous and nitrogenous stages in the B.O.D. reaction needs further study and considerable revision, particularly in its application t o stream pollution problems. 2. Application of blank seed sample demands as correction factors is a n erroneous procedure and should be avoided. 3. Nitrite and dissolved oxygen determinations should be made on all B.O.D. series t h a t are t o be used for the determination Of reaction constants' 4. Much work remains t o be done on other factors besides dilution that may affect the onset of nitrification in sewage LITERATURE CITED
(1) N a t l . Research Council, Sewwe Works J., 18, 791-1028 (1946). (2) Reed, L. J . 9 a n d Theriautt, E. J . 7 J * PhW. Chem.7 35, 950-71 (1931). (3) Ruchhoft, C. c,, worksJ , , 13, 669-80 (1941). (4) T h o m a s , H. A., Jr., Ibid., 9, 425-30 (1937); 12, 504-12 (1940).
sewage
RECEIVED October 18, 1946. Presented before the Division of Water, Sew. age, a n d Sanitation Chemistry a t the 110th Meeting of the AivBRICAN CHEMICAL SOCIETY, Chic-ago, Ill. Published by permission of the Surgeon General's Office, u.8. Public Health Service, Federal Security Agency.
System Calcium OxidePhosphorus Pentoxide-Sulfur TrioxideWater at 15.3" C. A. N. CAMPBELL AND J. W. COUTTS, When sulfuric acid is added to tricalcium phosphate, the number of possible solid phases is large. The three common phosphates of calcium give rise to five or six, depending on the degree of hydration. The occurrence of calcium sulfate a8 a solid phase complicates matters because this substance can exist in at least three, possibly four, forms, some of which, however, are claimed to be metastable. In addition to this, the possibilities of the formation of solid solutions (16) and of double salt (31) have been mooted. The problem of determining what solid phases coexist with solution, under various conditions of acidity and temperature, is one of considerable importance to industry, agriculture, and geology.
T
HE only previous investigation of the complete fourcomponent system, calcium oxide-phosphorus pentoxidesulfur trioxide-water, is that of Cameron and Bell (7), carried out a t 25' C. The present investigation was undertaken at 75.3" C.; a temperature in the region of which the industrial process for the manufacture of phosphoric acid from phosphate rock is operated. The system calcium oxide-phos-
University of Manitoba, Winnipeg, Canada
phorus pentoxide-water has been studied by Elmore and Farr ( I d ) at, among other temperatures, 75 O C. The authors' results are in close agreement with those of Elmore and Farr; in other words, the presence of calcium sulfate has little or no effect on the concentrations of solutions or the nature of the solid phases, except t h a t the number of the latter is increased by the presence of gypsum, anhydrite, or hemihydrate. The literature of this subject is vast and merits a brochure dedicated t o it. The bibliography given with this paper is believed to be the most complete bibliography extant. EXPERlMENTAL
MATERIALS. Analyzed tricalcium phosphate from the J. T Baker Chemical Company was used for the experiments of series 1. Commercial tricalcium phosphate, as has been pointed out by several investigators, seldom has a composition which corresponds exactly to the formula Ca3(P04)2. This was confirmed by the authors when i t was found t h a t a n extract of a sample of tricalcium phosphate supplied by another manufacturer gave a conductivity almost ten times as great as the Baker product. E o foreign radicals, however, were detected; the increased conduc-
INDUSTRIAL AND ENGINEERING CHEMISTRY
Vol. 40, No. 7
late with calcium phosphate (1.9). Therefore, the calcium was first precipitated as a basic phosphate by bringing the sample to a p H of 12 by adding dilute sodium hydroxide. The gelatinous precipitate was allowed to settle overnight, filtered, and dissolved in dilute hydrochloric acid. The solution was then neutralized a i t h ammonia and the calcium precipitated as oxalate, in presence of free acetic acid, according to the standard procedure. Analyses were always done in duplicate and were concordant to
=tnFIX
0 LO fW 30 w XQ 60 G ~ A M IC Ta o P E R /OO GRAMS S O L U T l O N Figure 1. A EImore and Farr ( 1 2 ) for System CaOP20j-Hz0. 0 -4uthors' Investigation for System CaOPzQs-SOa-HzO (Neglecting SOa) tivity was undoubtedly due, as has often been shown, to frec acid, hydroxyapatite, or less basic phosphate. Similarly, dicalcium and moiiocalcium phosphates from British Drug Houses, Ltd., which were used did not give the theoretical ratios of calcium oxide to phosphorus pentoxide on analysis, hut no foreign radicals were detected. It is of little consequence in an investigation of this kind whcther or not materials of fixed calcium oxide to phosphorus pentoxide ratio are used; hence, no purification of materials was carried out, although excellent, but time-consuming methods, are available. British Drug Houses' Analar phosphoric acid was used. APPARATUS.The samples for the solubility determinations were contained in 250-ml. hard-glass reagent bottles. Hard glass was used because Bassett ( 3 ) claimed that calcium phosphate solutions have a corrosive effect on glass. No evidence of this was found in the present investigation, but the Bassett samples were stirred for a much longer time. The bottles were immersed in a thermostat maintained at 75.3 O =t0.1 C. A Wheatstone bridge circuit, containing standard resistances, audio-frequency oscillator, and Wagner ground was connected to a dipping electrode for the measurement of the conductivity of the liquid phase. This was used as a means of determining when equilibrium had been attained. Separation of saturated solution from solid phase was carried out in the usual way, but the solid phase was subsequently washed with anhydrous acetone, so that it was not the wet solid phase of Schreinemaker's method, METHODS OF ANALYSIS
CHEMICAL. Because large concentrations of phosphate and sulfate ions interfere in the analysis of calcium, some of the more convenient methods of analysis were inapplicable. It was, therefore, necessary to use tedious and lengthy procedures for their estimation. In the analysis of solid phase, the calcium was estimated volumetrically with permanganate, after precipitation as oxalate. I n analysis of the liquid phase, which usually had a much higher calcium oxide to phosphorus pentoxide ratio, precautions had t o be taken t o prevent contamination of precipitated calcium oxa-
T i L k e t h o d used for the determination of phosphate radical i b given in some detail, as it is the result of combining what was believed to he the better features of several methods. An aliquot sample, containing 0.07 to 0.1 gram of phosphorus pentoxide, wai neutralized with dilute ammonia, made just acid with a drop or two of dilute nitric acid, and dilutcd to about 60 ml. The solution was then warmed (not above 80 C.) and a mixture of 40 ml of Xoyes' ammonium molybdate reagent and 50 ml. of 5 S nitric acid added slowly from a pipet ( 2 4 ) . The sample was shaken vigorously during and after the addition of the ammonium molybdate. It was then kept for an hour a t a temperature of about 65' C., and again shaken for 3 or 4 minutes. (Complete precipitation was ensured by allowing it to stand overnight.) The solution was then filtered and the precipitate washed aTith 5% ammonium nitrate solution. The precipitate was then dissolved in a fine stream of warm 6 A- ammonia and the washings were collected in the flask in which the original precipitation took place. ( I t is important to use only a small excess of ammonia. The bulk of the solution should not exceed 100 ml.) The solution was carefully. neutralized with 1 to 1 hydrochloric acid, using litmus paper as indicator. (Care must be taken not to add a n excess of acid.) The phosphate was then precipitated with magnesia mixture in the usual way. This procedure gave results which checked within *0.5%. Sulfate radical was estimated gravimetrically in the usual way. OPTICAL~ ~ N A L Y S I S .All solid phases were subjected to optical analysis under the petrographic microscope, using the refractive indices, determined by the Beclre line method, and the interference pattern, as means of identification. The optical analysis of these samples was made difficult by the fineness of the crystals and by the frequent occlusion of almost all of the sulfate phase by the phosphate phase. If a phase 'ryere present in extremely small amount it may be overlooked in any one examination, but this difficulty is overcome by examining many samples of the same solid mixture. The authors' knowledge of the nature of the solid phases rests principally, but with certainty, on the optical examination.
Two methods of M orking were EXPERIMENTAL PROCEDURE. adopted. I n the first, 50 grams of tricalcium phosphate were niixed with 225 ml. of water, the mixture was stirred for 24 hours, and the conductivity taken. An accurately measured volume of concentrated sulfuric acid, sufficient t o produce a predetermined displacement of phosphoric acid, was then run in. Stirring was then recommenced and the conductivity determined a t intervals. Constancy of conductivity was attained after 6 days as is shown by the data of the following typical experiment: CELL RESISTANCE O F S.4XPLE 1B BEFORE ADDITION O F SULFrRIC ACID, 1540 OHMS Date Xov. 13 13 13 14 15 16 17
is
20
21 22 23
Time 10:30 A . M . 1:OoP.M.
5:30 P . X . 9:30A.Y. 9 : 15 A.M. 9 :45 A.M. 4:16 P.M. 3 :20 P.M. 9 :30 A.M. 9 : 15 A.M. 9:3O~.x. 2:30 P.M.
Cell Resistance, Ohms 66.19 71.84 84.73 106.88 113.40 121.95 123.51 124.08 124.13 124.22 124.06 124.14
After the sixth day, the resistance fluctuates by about 0.2 ohm, up and down. If the equivalent conductivity of calcium sulfate is taken as 119.5 a t 18" C., and if the temperature difference is neglected, and the calculation carried out in the usual way, this fluctuation corresponds to a varying in concentration of 0.0088 gram of calcium sulfate per liter, or about 0.0009%. After constant conductivity was attained, the saturated solution and solid phase were separated at the temperature of the thermostat and analyzed. This procedure readily produced
.
July 1948 TABLEI.
Sample
1A B C
D
E 2A B C D E F
G
H I J K L 1M N 0
P
w
8
T
Concn. of Liquid Phase before Addn. of Solid % PaOs % So8 ,.. 2.2 4.4 .. .. 6.7 .. 8.9 11.1' ... 4.0 3.75 3.75 6.5 3.75 10.0 3.75 13.7 8 3.75 20 3.75 16 3.75 21.2 3.75 21.8 3.75 25 3.75 28.4 3.75 20 3.33 30 3.33 27 3.33 50 3.3 40 3.75 41.5 3.75 43 3.75 45 49 54
INDUSTRIAL AND ENGINEERING CHEMISTRY
1297
EQUILIBRIUM CONCENTRATIONS IN THE SYSTEM CALCIUMOXIDE-PHOSPHORUS PENTOXIDESULFUR TRIOXIDE-WATER AT 75.3" C. Solid Added Caa(PO4)z
stirring Time, Days 5 10 9
CaHPOk
5 6 14 14 14 15 14 16 16 16 14 14 15 14 14 15 15 15 15 14 14 14 14
3.75 3.75
,--Equilibrium Concns., G./100 G. ----.hoS Liquid Phase __ Solid Phase CaO PzOa SO8 CaO PzOs SO8 6.69 0.294 0.151 47.61 38.04 0.109 0 , 1 2 2 0.310 0.154 43.86 32.72 14.49 0.118 0.283 0.157 39.25 29.05 19.17 0.281 0.976 0.141 36.50 26.82 22.10 3.07 0.157 35.05 19.63 30.89 0.741 .... 0.125 . . . . 9.49 2.01 4.72 13.35 0.091 40.10 46:50 2.68 .... 0.075 .... 18.09 3.49 .... .... 0,081 . . . . 20.93 3.94 .... . . . . 0.052 . . . . 4 . 9 2 27.51 ,.. . . . .... 30.53 0.248 5.27 . . . . .... 5 . 3 7 31.68 . . . . .... 0,050 32.72 5.45 .... 33.54 0.035 . . . . 5.52 35.51 0.033 40.85 46.15 6.65 5.64 0 . 0 2 8 24.70 51.14 4 . 9 1 5 . 6 8 85.60 0.035 31.62 48.95 5 . 3 0 36.54 5.38 5.65 0.026 23.65 50.15 5 . 1 0 37.60 . . . . .... 4 . 3 5 40.62 0.018 . . . . 8.55 0.017 24.71 48.01 44.63 3.30 2 . 9 6 46.00 Trace 26.85 41 .OO 15.48 .... 2 . 4 0 49.02 . . . . ... 1 . 6 5 54.24 .... 1.275 67.46
Nature of Solid Phase, (Optical Determination) Cas(PO&.HnO; CaHPO4; CaSO4.2HaO CaHPOa, CaS04.2Hz0, CaS04b CaHPOa, CaSOa.2HzO
....
....
CaHPO4, CaSO4.2Hz0, CaSOcl/zHzO CaHPO4 CaSOa.2HzO CaHPO4: CaSO4.1/zHzO CaHPO4, Ca(HzPOa)z.HzO, CaSO4.1/aHO Ca(HaPO4)z. HzO, CaSOa. '/zHnO
...
1.269 57.38 1.150 67.47
.... ....
I
.... .... ....
Ca(HzPOa)z.HzO, Ca(HzPOi)zb, CaSOa. '/zHzO
Ca(HzP04)z.HzOG, Ca(HzPO&b, .... CaSOd. l/zHzO Sample spoiled. Trace. 0 This phase probably not actually present in equilibrium system but formed after separation from liquid phase by hydration of Ca(HzP04)z, which is known to be a very rapid transformation ($2).
u
a b
invariance with the solid phases-gypsum, tricalcium phosphatq, and dicalcium phosphate-but, owing to the high solubility of monocalcium phosphate, i t is not convenient to produce the invariant, for which the solid phases are dicalcium phosphate, monocdcium phosphate, and gypsum by this method. (A very great amount of tricalcium phosphate and sulfuric acid would be required and the vast quantity of gypsum resulting would produce an unworkable paste.) I n the second method, applied to the range where monocalcium phosphate was a desired solid phase, either or both dicalcium or calcium monophosphate was used as starting material. To this phosphoric acid of increasing concentration was added, together with sufficient sulfuric acid to produce an appreciable amount of calcium sulfate in the solid phase. In this second series of measurements, the high conductivities rendered the conductivity method insenditive in determining the attainment of equilibrium. Judging by the results of the first series and the work of Elmore and Farr (12) two weeks of constant stirring were allowed for the attainment of equilibrium. RESULTS
Table I contains the analyses of systems in equilibrium over a range of 0.3 to 57.5% phqsphorus pentoxide. The numbering of the analyses does not indicate the order in which the determinations were made but corresponds to their sequence on the curves. The pre&es 1 and 2 indicate samples of the first and second series, respectively. Optical analyses were made on all solid phases, but chemical analyses only in certain cases; the chemical analysis invariably confirmed the optical one. 75.3' C. ISOTHERMAL OF THE SYSTEM
The graphical representation was simplified by the fact that the concentration of sulfur trioxide in solution is always low. Hence, the liquid phase is effectively a three-component system. I n order t o compare results with those of Elmore and Farr ( l a ) for the system calcium oxide-phosphorus pentoxide-water, the results have been plotted as per cent phosphorus pentoxide against per cent calcium oxide in the liquid phase (Figure 1). The slight difference in temperature between the two investigations is of no consequence.
Figure 1 shows that the compositions of solution in contact with CaHP04, Ca(HzPO&.HzO, and Ca(HzP04)zare only very slightly altered by the presence of some form of calcium sulfate. The curve along which dicalcium phosphate is the stable solid form is slightly lowered by the presence of calcium sulfate. The curve for monocalcium Fhosphate monohydrate is slightly raised up to the point where sulfate is no longer detectable in solution (about 48% phosphorus pentoxide) The composition of the solution a t the invariant point where the solid phases are Ca(HzP04)Z and Ca(H2P04)z.Hz0,obtained by extrapolation from Elmore and F a d s curves, agrees well with experimental determination, where the solid phases were Ca(HzPO&, Ca(HzP04)z.H20, and CaS04.1/2Hz0. Point 2U was also in agreement with Elmore and Farr's curve. Therefore, it was unnecessary to carry the investigation to higher concentratiods of phosphoric acid because the curves for the threeand four-component systems are practically identical after a concentration of about 48% phosphorus pentoxide has been reached. Analyses l A , B , and C represent the invariant system Cas(P04)z.H20-CaHPO~-CaS0~.2H20-solution; analyses 2J and K the invariant system CaHP04-Ca(H2P04)z.Hz0-CaS04. '/zHzOsolution; and 2X and T the invariant system Ca(HzP04)z.H20Ca(H2POa)z-CaS04.l/zHzO-solution, Previous workers have pointed out that addition of calcium sulfate to the system calcium oxide-phosphorus pentoxidewater has only a slight effect at 25" C. (6, as). The effect is appreciable a t 75 C., as Figure 1shows, but it becomes less as the content of sulfate in solution decreases-that is, as the content of phosphoric acid increases. SOLUBILITY OF CALCIUM SULFATE. Table I shows that the solubility of calcium sulfate is always slight and decreases as the phosphoric acid content increases. This is shown in Figure 2, where the sulfur trioxide content is plotted against the content of phosphorus pentoxide. Although the poin!s do not fall exactly on the curve, they indicate the trend from a solubility greater than that in pure water a t low phosphorus pentoxide concentrations to almost zero solubility at high concentrations. NATUREOF THE SOLIDFORM OF CALCIUM SULFATE. I n this investigation, gypsum was the stable form of calcium sulfate up t o a concentration of 30% phosphorus pentoxide. I n a solution containing 30.53% phosphorus pentoxide (point 2F), hemihy-
1298
I N D U S T R I A L AND E N G I N E E R I N G CHEMISTRY
Vol. 40, No. 7
analyses are consistent with a smooth curve also indicates that equilibrium was attained. This relative rapidity with which equilibrium was attained may be due t o the presence of sulfuric acid, the high temperature, or the high rate of stirring. The solutions were, of course, in equilibrium with respect only to calcium phosphate, and not x i t h respect to calcium sulfate, which has been shown to be capable of metastable existence for a period of years. SOLIDSOLUTIONS. It was suspected by Hill and Ilendricks (16) and also by the chemists of the Consolidated Mining and Smelting Company of Canada, that CaS04.2HzO and CaHP04.2H20 might form solid solutions. At the temperature of this investigation hydrated dicalcium phosphate did not form, but ah-ays the anhydrous form. Nevertheless, this does not exclude the possibility of gypsum taking the molecules or ions of dicalcium phosphate into its lattice. Since the refractive indices and x-ray patterns of gypsum and of dicalcium phosphate dihydrate are almost identical, the ordinary methods of attacking this problem are not applicable. Neither does the analysis of the washed solid phase yield anything conclusive because a solid solubility of more than 17, is not claimed. This phase of the problem therefore remains unanswered. SUMMARY
Figure 2. Per Cent Phosphorus Pentoxide Plotted against Per Cent Calcium Oxide in the Liquid Phase
drate appeared along with gypsum. S t first sight it appeared that this concentration of phosphorus pentoxide marked the transition point of the two forms: 2CaS04.2Hz0 ?=? 2CaSOc1/zHzO -I- 3Hz0
but gypsum appeared again, unaccompanied by hemihydrate, a t point 2G (31.68y0 phosphorus pentoxide). The explanation probably is t h a t the initial concentration of phosphorus pentoxide is a more important factor than the final concentration in determining which form of calcium sulfate will appear in the solid phase. The initial concentration of phosphorus pentoxide (and of sulfuric acid) probably determines whether gypsum or hemihydrate is formed by the actionbf sulfuric acid on the calcium phosphate. The initial form of calcium sulfate then persists metastably. Sulfuric acid appears t o exert a greater dehydrating effect than phosphoric acid, for the only samples which contained anhydrite were ID and E , to which 9 and 11% sulfur trioxide had originally been added. On the other hand, solutions containing as much a s 57.57, phosphorus pentoxide deposited no anhydrite. Marshall, Hendricks, and Hill (26)noted, a t a lower temperature, that hemihydrate was the predominant form of calcium sulfate in products prepared from dilute sulfuric acid, whereas anhydrite appeared in products where the acid concentration ranged from 62 to 65% sulfuric acid. It is remarkable t h a t d'.4ns and Hofer ( 1 ) rarely obtained hemihydrate in their study, at 83" C., of the system calcium oxide-phosphorus pentoxide-water. They claim t h a t hemihydrate is very unstable and, t o preserve it, it must be withdrawn from the solution after only a few minutes' contact. The results of the authors' investigation indicate, on the contrary, that calcium sulfate hemihydrate can exist metastably for long periods of time. TIhlE REQUIRED TO REACH EQUILIBRIUM. Earlier workers claim t h a t equilibrium in the system calcium oxide-phosphorus pentoxide-water is attained only after many weeks or even months, I n the present investigation, the conductivity data which were obtained for the lower concentrations indicate that equilibrium is reached in a much shorter time. The f w t that
A comprehensive study of the literature relating to the systems Ca0-PzO6-SO3-H20, Ca0-P205-H20, and Ca0-S03-H20 has been made. The solubility relationships of the system CaOP205-SOa-H20 a t 75.3" C. have been investigated over a range of phosphorus pentoside concentration of from 0.3 to 57.570. The sulfate content of the liquid phase of this system is very slight and decreases steadily with increasing phosphorus pentoxide concentration, finally approaching zero. The kalcium oxide to phosphorus pentoxide ratios in the liquid phase of the system Ca0-P205-H20 are slightly altered by the addition of calcium sulfate. The regions of stability of the solid forms of calcium sulfate and of the calcium phosphates in phosphoric acid solutions were determined. The regions of stability of the calcium phosphates are only slightly altered by the presence of calcium sulfate. Gypsum, CaS04.2H20, is capable of prolonged metastable existence in solutions of low phosphorus pentoxide content, whereas the hemihydrate, CaSOd.1/2HZO, exists metastably in contact with solutions of higher phosphorus pentoxide content. Anhydrite was only obtained in solutions having a relatively high initial sulfuric acid concentration. The following invariant points have been determined experimentally: Caa(P04)z.Hz0, CaHPOa, CaSOa 2HzO-solution CaHPOa, Ca(HzPOdz.Hz0, CaSOi. 1/zHzO-solution Ca(HzPOa)z.HzO, Ca(HzPOi)p, CaSOc'/zHzO-solution
The time required to reach equilibrium was shown to be less than that reported for the system Ca0-P2O6-Hz0. ACKKOWLEDGRIENT
The authors are indebted t o the Consolidated Mining and Smelting Company of Canada, Ltd., Trail, B. C., for the fellowship held by J. IT, Coutts and for a grant for the purchase of apparatus. BIBLIOGRAPHY
Bns, J. d', a n d HRfer, P., 2. angew. Chem., 50, 101 (1937). Armsby, H. P., Am. J . Sci., 12,46 (1878). Bassett, H., J . Chem. Soc., 111, 1620 (1917). Bassett, H., 2 . anorg. Chem., 53, 34, 49 (1907); 59, 1 (1908); Chem. News,95, 21 (1907): J . Chem. Soc., 111, 1620 (1917). Berzelius, J . J., Ann. Chem., Justus Liebigs., 53, 286 (1845). C a m e r o n , F. K., and Bell, J. M., J . Am. Chem. Soc., 27, 1512 (1906). I h i d . 28, 1222 (1BOCj).
I N D U S T R I A L A N D E N G I N E E R I N G CHEMISTRY
July 1948
(8) Cameion, Ei. K., and Huiat, L. -4.,Zbid., 2 6 , 8 8 5 (1904). (9) Cameron, F. K.. and Seidell, A., Ibid., 26, 1454 (1904) : 27,1503 (1905).
(IO) Causse, H., Compl. rend., 1 1 4 , 4 1 4 (1892). (11) Davis, W. A., J . SOC.Chem. Znd., 26, 727 (1907). "
(12) (13)
Elmore, K., andFarr, T., IND. ENG.CHEM.,3 2 , 5 8 0 (1940). Haddon, C. L., and Blown, M . A . W., J . SOC.Chem. Znd., 43, 11 (1924).
(14) Halla, F., 2. Krist., 80, 349 (1931); Z . angrw. Chem., 41, 659 (1931). (15) Hill, A. E., J . Am. Chem. Soc., 59, 2242 (1937). (16) Hill, W. L., and Hendrirks, S . B., IND.ENG.CHEM.,28, 440 (1936). (17) Hill, W. L., andJacob, K. D., J . Assoc. OficiaZAgr. Chemists, 17, 487 (1934). (18) Holt, L. E., La Mer, V. K., arid Chown, H. H . , .J. Biol. Chum., 64, 509, 567 (1925) (19) Jung, H., 2. anorg. Chem., 142, 73 (1925). (20) Kolh, J., Compt. rend., 78, 825 (1874). (21) Le Chatelier, H., "Recherches exp&-imeritales sur la Constitu-
tion des Mot tiers hydroliques," Paris, 1887.
1293
(22) (23) (24)
Linck, G., and Jung, H., 2. anorg. Chem., 137, 407 (1924). Lugg, J. W. H., T r a n s . F a r a d a y Soc., 2 7 , 2 9 7 (1931). McCandless, J. M., and Burton, J. A,, IND.ENG.CHEM.,16,
(25)
Marshall, H. L., Hendrickb, S . B., and Hill, W. I,., Ibid., 32,
(26) (27)
Marshall, H. L., and Hill, W. L., I b i d . , p. 1129. Partridge, E. P., and White, A. H., J . Am. Chem. Sor., 51, 360
(28) (29)
Posnjak, E., Am. J . Sci., 35 A, 247 (1938). Ranisdell, L. S., and Partridge, E. P., Am. M i n e d , 14, 59
(30) (31)
Roller, P. S., J . Phys. Chem.,35, 1132 (1931). Sanfourche,A., and Facet, B., Bull. soc. chim., 53, 1221 (1933). Sanfourche, A,, and Krapivine, A , Zbid., p. 1573. van't Hoff, J. H., 2. physik Chem., 45, 257 (1903). Viard, Compt. rend., 114,414 (1892). Weiser, H. B., Milligan, W.O., and Eckolm, W. E., J . Am. Chem.
1267 (1924).
.
1631 (1940).
( 1 929). (1929).
(32) (33) (34) (35)
Soc., 58, 1261 (1936). RE,CEIVED March 4, 1947
SYSTEM LBUTENE-n-BUTANE Compositions of Coexisting Phases B. H. SAGE AND W. N. LACEY California Institute of Technology, Pasadena, Calif. ,.I
[ h e compositions of coexisting gas and liquid phases in the 1-butenen-butane system were experimentally determined for seventeen equilibrium states at four temperatures ranging from 100' to 280' F. The results, which were not of great accuracy, were correlated graphically by plotting the isobaric difference in the mole fractions of I-butene in the two phases as isothevmal functions of the .mole fraction of 1-butene in the liquid phase. Throughout the interval of temperature experimentally investigated the observed difference between the composition of coexisting phases i s markedly less than that predicted by Haoult's law.
E C E N T developments in petroleum processing have emphasized the need for data on the phase behavior of systems containing both paraffin and olefin hydrocarbons. I n the present investigation equilibrium phase compositions were determined for the 1-butene-n-butane system at loo", 160°, 220°, and 280" F. The volumetric and phase behavior of 1-butene was investigated earlier a t pressures up t o 10,000 pounds per square inch for temperatures between 100" and 340" F. (4). The properties of n-butane at temperatures and pressures of interest in the present study were reported by several investigators ( I , $ , 3 ) . MATERIALS
The 1-butene used in this study was prepared by dehydration of n-butyl alcohol using activated alumina as a catalyst. Details of the process of preparation and purification were described in a n earlier publication (a). A sample of the material used in the investigation varied less than 0.3 pound per square inch in its vapor pressure from bubble point t o dew point at 100'F. The n-butane was obtained from the Phillips Petroleum Company together with a n analysis which indicated the presence of less than 0.003 mole fraction isobutane and negligible amounts of other impurities. This material was fractionated in a 4-foot vacuum jacketed column packed with small helical glass rings. The column was operated at atmospheric pressure with a reflux
ratio of about 50 t o 1. The initial and final portions of the material fractionated were discarded, and the intermediate fraction amounting t o approximately 0.8 of the total overhead product was condensed at liquid-air temperatures at a pressure of about 4 X 10-6 inch (10-4 mm.) of mercury in order t o remove noncondensable gases. A sample of the material prepared in this manner had a vapor pressure at 100' F. which agreed t o within 0.1 pound per square inch with the value of 51.5 pounds per square inch absolute for the vapor pressure of pure n-butane. APPARATUS AND PROCEDURE
Th.: equilibrium chamber used for the study of this system had sample ports at the midpoint and at the top of the cell. After the amount of material in the cell was adjusted so that the lower sample port was in contact with the liquid phase and the upper port with the gas phase, the system was brought t o equilibrium at the desired state with the aid of mechanical agitation which was accomplished by the rotation of a n internal stirrer. Isobaric conditions were maintained in the equilibrium cell during withdrawal of the samples by injecting mercury at, a suitable rate. Duplicate gas and liquid phase samples were withdrawn at each experimentally observed state. The samples were collected in evacuated 500-ml. bulbs provided with glass stopcocks and standard taper ground-glass joints. The bulbs were filled t o slightly more than atmospheric pressure and the samples were then transferred t o the analytical apparatus. The amount of 1-butene in a sample was determined by volumetric, catalytic hydrogenation methods (6). From the measured diminution in volume as a result of the hydrogenation of the 1-butene and from the original volume of the sample and the total volume of the sample and hydrogen, the mole fraction of unsaturated hydrocarbon was computed, taking irito account the deviations of the hydrocarbons from perfect gas and ideal solution behavior. The analytical measurements permitted the calculation of a t least two, and generally three values of the mole fraction of n-butene from each set of data. The average of these calculations was taken as the most probable value for the particular state. t
