198
INDUSTRIAL AND ENGIKEERIKG CHEMISTRY
through the barn. I n the drying unit used in this investigation 20 feet per minute was found to be the minimum satisfactory air velocity.
Conclusions The application of air conditioning to the curing and drying of bright leaf tobacco improves the process in the follo\ving ways : 1. The time of curing and drying is reduced ap roximately 50 per cent. The yellowing period is not materially s f h e n e d , since this rate depends upon a physiolo ical proress. The fixing period is reduced considerably, and thetilling (9) and ordering periods are shortened to a large extent. 2. Optimum atmospheric conditions favorable to excellent curing are closely controlled, and air may be distributed uniformly throughout the barn. There is no loss of tobacco due to excessive temperature or uncontrolled humidity, and a more uniformly high quality may be obtained. 3. Labor and fuel requirements are considerably reduced. Automatic control leave the farmer free for other duties. Present methods require considerable day and night observation. A large portion of the fuel is consumed in the barn during the killing period, where high temperatures are maintained for considerable time. This consumption of fuel should be greatly lowered as a result of the shortening of this period. 4. Increase of the capacity of the barn, as a result of the shortened process time, will reduce the investment and maintenance cost. This method indicates possibilities of community operation in contrast to the large number of individually owned barns. 5 . The fire hazard of hot flues within the barn is reduced.
VOL. 32, NO, 2
Bibliography (1) Badger, W. L., and McCabe, W. L., “Elements of Chemical Engineering”, 2nd ed., p . 298, New York, McGraw-Hill Book Co., 1936. ( 2 ) Darkis, F. R., Dixon, L. F., Wolf, F. A., and Gross, P. M., IND. ENG.CHEM.,28, 1214-23 (1936). (3) Delamar, C. D., thesis, Va. Polytechnic I n s t . , 1938. (4) Garner, W. W., U. S. Dept. Agr., Farmer’s Bull. 523 (July, 192.1). ~ . - _ Garner, W. W., Bacon, C. W., and Bowling, J. D., Jr., ISD.ENQ. CHEM.,26, 970-4 (1934). Killebrew. J. B., and Myrick, H., “Tobacco Leaf”, pp. 209-32, New York, Orange J u d d Co., 1920. McCready. D. W., and McCabe, W.L., Trans. Am. Inst. Chem. Engrs.. 29, 131-59 (1933). Maxirnov, N . A., “Textbook of P l a n t Physiology”, New York, McGraw-Hill Book Co , 1930. Sherwood, T. K., Chemical Engineers’ Handbook, p. 1260, New York, McGraw-Hill Book Co., 1934. Sherwood, T. K., Trans. Am. Inst. Chem. Engrs., 23, 28-44 (1929). Sherwood, T. K., Ibid., 32, 150-68 (1936). Sherwood, T. K., and Comings, E. W., IXD.E m . C H m f . , 26, 1096-8 (1934). Smith, H. B., thesis, Va. Polytechnic Inst., 1939. Stillwell, S. T. C , Trans. Inst. Chem. Engrs (London), 6 , 91-101 (1928). \----. Thatcher, R. W . ,, “Chemistry of P l a n t Life”, New York, McGraw-Hill Rook Co., 1921. Walker, W.H., Lewis, W.K., MoA4dams,W. H., and Gilliland, E. R., “Principles of Chemical Engineering”, 3rd ed., p. 642, Kew York, McGraw-Hill Book Co., 1937. PRESENTED before the Division of Industrial and Engineering Chemistry a4 the 97th Meeting of the American Chemical Society, Baltimore, Md.
The System Trisodium PhosphateSodium Carbonate- Water KENNETH A. KOBE AND ALEXANDER LEIPPER University of Washington, Seattle, Wash.
HE solubility data for trisodium phosphate in water disagree badly, and no data exist for the system trisodium phosphate-sodium carbonate-water. The purpose of this paper is to present the data for this ternary system, which will include the binary system trisodium phosphate-water, from the cryohydric point to 100’ C. The reasons for disagreement among the various solubility determinations will be given. It is well known that in commercial trisodium phosphate a certain amount of excess sodium hydroxide is present which will give a formula approximately represented by Na8PO4.l/?NaOH with 10 or 12 molecules of water of crystallization. I n their study of the system Na20-P206-H20, D’Ans and Schreiner (3) state that one of the solid phases is Na3P04.12HzO, but their data are not clear. They report that the crystals were analyzed after being dried between porous plates. The analytical method is of importance so that the excess alkali may be determined. If the crystals are Na3P04.12H20, the P206content is 18.7 per cent, whereas Na,P04.1/7 NaOH.12 HzO has 18.4 per cent PzOs. This difference is so small that it could readily be attributed to imperfect
T
drying of the crystals. Since it is known that commercial trisodium phosphate (t. s. p.) does crystallize with an excess of alkali, any analysis of the solid phase must consist of a determination of both the PzOsand KazO content to determine their ratio in the crystal. Extensive experiments were performed by Smith (16) to prepare pure trisodium phosphate free from excess alkali. Using theoretical proportions of phosphoric acid and sodium hydroxide, he found that the first crop of crystals representing about 25 per cent of the total P20bhad the composition 2Na3P04.Sa2HP04. K i t h 4 per cent excess sodium hydroxide crystals were obtained which had the composition 171/2Na3P04.Ka20. As a result of his work, Smith concluded that either trisodium phosphate cannot be crystallized from its components or else it does not exist under norma1 conditions. hlenzel and von Sahr (9) studied in detail the composition of the solid phase in equilibrium with the saturated solutions in the system Na20-P2OrH20. Beginning a t the transition point of the dibasic to tribasic phosphate where the ratio of Na20 to P20sin the solution is 2.67, the solid phase shows
FEBRUARY, 1940
INDUSTRIAL AND ENGINEERING CHEMISTRY
199
formula lYa3P04.'/~NaOH (without considering the water of crystallization). Upon recrystallization from hot water, a product with the formula KasPOl.l/aNaOH was obtained. Phosphoric acid with three equivalents of sodium hydroxide gave crystals analyzing T\l'a3P04.1/6NaOH. The presence of sodium carbonate in the solution did not cause carbonate to appear in the crystals. The method of Booth and Gerber ( 2 ) gave crystals which contained excess alkali, t,hough less than Baker's c. P. salt.
Method of Analysis Analysis was made of a solution containing trisodium phosphate, sodium hydroxide, and sodium carbonate. An acidimetric method devised by Smith (15) depends upon the reactions :
+ ++ NaaPO4 + HC1+ NalHP04 + HC1+
+
NaOH HC1+ NaCl HzO (1) NaHCOa NaCl (2) NaZCO3 HC1+ NaHCOa HC1+ NaCl COz Hs??,
Courtesy, Victor Chemical Works
CONTROL PANELIN
A
SODIUM PHOSPHATE PLANT
3.09. As the ratio in solution increases, the ratio in the solid phase increases rapidly to 3.22 and then slowly to 3.24. The solid phase was called "pseudotribasic sodium phosphate dodecahydrate", since the stoichiometric trisodium phosphate could not be prepared. Conditions for transition t o other hydrates were determined. Their data (9, Table I) show the minimum Na2O-P2OS ratio for a solution from which t. s. p. can be crystallized. A ratio of 2.70 in the original solution gives 3.09 in the solid; 2.90 gives 3.14 and 3.00 gives 3.18; above this the ratio 3.22 in the solid is approached rapidly. Thus, a neutral t. s. p. cannot be crptallized from solution. Since usual plant practice appears to start with a ratio of 3.0 in the solution and crystallize, the crystals should show about 3.18Na20-1P206. As crystallization occurs, the solid, because of its higher ratio of NaPO t o P20j, depletes Na20 from the solution; this, in turn, decreases the ratio in the crystal. Starting with a solution on the acid side will reduce the excess alkali somewhat, but not to a 3.00 ratio for the solid. The so-called neutral cleaners on the market contained an amount of disodium hydrogen phosphate equivalent to the excess alkali in the t. s. p. The literature describes two processes for the manufacture of alkali-free t. s. p. A patent issued t o Booth and Gerber in 1929 describes a process for the manufacture of neutral t. s. p. ( 2 ) . The specific gravity and excess of phosphoric acid is controlled within specified limits. Another patent issued to Westbrook (19) claims that an alkali-free t. s. p. may be obtained by crystallizing out the decahydrate between 49 " and 62.8" C. instead of the usual dodecahydrate. Richards and Churchill (11) give the transition temperature of the dodecahydrate to the decahydrate as 73.4"C. Below 73.4" C., according to most investigators, the stable phase is the dodecahydrate, and this is the form which would ordinarily be in equilibrium with a saturated solution. If formed below 73.4" C., decahydrate crystals are obtained as a result of a metastable condition of equilibrium. Baker's c. P. tertiary sodium phosphate dodecahydrate was analyzed, and the actual analysis found to correspond to the
++ +
++
ti'
Na2HP04 NaCl NaHZPO, NaCl (5)
Equations 2 and 4 are complete a t the phenolphthalein end point; Equations 3 and 5, a t the methyl orange end point (6-8, 15). The titration error in Equation 5 is about 0.5 per cent (7). The solution containing the three salts is first titrated with acid to the phenolphthalein end point, represent1/2Na2C03 I/&a~P04. Continued titration ing NaOH to the methyl orange end point represents '/ZKa&O3 '/aXa3P04. The solution is then boiled to expel carbon dioxide
+
+
+
The solubility data for trisodium phosphate in water disagree badly, and no data exist for the system trisodium phosphatesodium carbonate-water. Commercial trisodium phosphate (t. s. p.) has the approximate formula Na$04. l/,NaOH.12H20, and the alkali content of the t. s. p. varies with the composition of the solution from which the t. s. p. is prepared. It is shown that the crystals forming from a solution do not have the same Naz0-P20b ratio as that of the solution. The solubility of t. s. p. in water was determined at O", 25', 40°, 60', 80°, and 100' C., and the results were compared with those in the literature. Disagreement is due to the various Na,O-P,O, ratios that exist in various kinds of t. s. p. The system Na20-PzO5-H20was studied at 25' C. to show the changing solubility. The system t. s. p.-NazCOa-HzOwas studied at the above temperatures. No double salt is formed in this system.
200
INDUSTRIAL AND ENGINEERING CHEMISTRY
VOL. 32, NO. 2
(Right) FEED END OF A ROTARY COOLER FOR COOLINQ CRYSTAM OF SODIUMPHOSPHATE (Below) FILLING AND SEWINQ BAGSOF SODIUM PHOSPHATE Courtsay, Victor Chemical Works
found for the t. s. p. used in this work. Thesodium carbonatet. s. p. sample was analyzed similarly. From the amount of phosphate ion, the equivalent alkali in Na8P04.*/,r\’aOH could be calculated. This amount subtracted from total alkali gave the sodium carbonate present. Baker’s c. P. tribasic sodium phosphate, lot No. 52835, bottle 1,analyzed as follows: A random weight was dissolved and diluted to 250 ml. : Moles P = 0.02024 Titrated alkali to MO end point = 0.04332 mole 0.04332 - (2) (0.02024) = 0.00284 mole NaOH in t. s. p. 0.02024/0.00284 = 7.13 ratio Na3P04/NaOH Bottle 2 showed a ratio of 6.91. The average of all results closely approximates 7.0; analysis of this lot of t. s. p. is taken as Na3PO4.l/,NaOH. Calculation of a point on the 25’ C. isotherm of the t. s. p.h’a2COrHz0 system was as follows: a weighed sample (12.334 grams) was diluted to 250 ml. and analyzed:
and titrated with alkali to the phenolphthalein end point, representing */sNa3P04. From the three titrations the amounts of trisodium phosphate, sodium carbonate, and sodium hydroxide present can be calculated. Use of this method showed the phenolphthalein end point to be well buffered and lingering. The method is not accurate enough for these solubility determinations but may be applicable as a control method in the crystallization of the two components, t. s. p. and sodium carbonate. Mixed indicators equivalent to phenophthalein might improve the method. The method of Gerber and Miles (4) for the determination of phosphate mixtures has appeared since the completion of this work. I n the solubility determinations the following method was used: A sample of the t. s. p. was titrated to the methyl orange end point (Equations 1, 4, and 5), and this solution was then analyzed for phosphate ion by the alkalimetric ammonium phosphomolybdate method (1). From these two determinations the constant ratio Na3PO4.l/7NaOH was
Mole P = 0.004024 = 0.6836 g. NasPOd.l/,NaOH Titrated alkali t o MO end point = 0.04434 mole (0.004024) (21/1) = 0.00862 mole acid for t. 8. (0.04434 - 0.00862) (0.5) = 0.01786 mole Na2&03 = 1.893 g.
Na.CO.”
12.334 - 0.6836 - 1.893 = 9.7574 g. HzO (0.6836) (100)/9.7574 = 7.01 g. NaIP04.1/rNaOH/100 g. H20 (1.893) (100)/9.7574 = 19.40 g. NazCOs/lOO g. HzO
Experimental Procedure
If we conclude that there is no simple system trisodium phosphate-water, then i t would be of little use to prepare neutral trisodium phosphate, for if such a substance were dissolved in water to form a saturated solution, the solid phase would change to a composition different from the composition of the solids in solution. Such a system must be considered as one of three components-namely, NazOPzOrHzO-with the restriction that the composition of the whole mass, both solid and solution, have the same acid-base ratio as in pure trisodium phosphate. If to such a solution sodium carbonate is added, the system becomes one of four The analysis of this components, Naz0-P2OsCOt-H20.
FEBRUARY, 1940
INDUSTRIAL AND ENGINEERING CHEMISTRY
quaternary system would require determinations for Na20, P206,and COP,with the HzO being obtained by difference. I n order to simplify the analytical procedure, it was decided t o consider the material I\'aaPOr.l/&aOH as a simple salt (without any inference that such is actually the case), thereby making the system under consideration a ternary This could be done one: iYasPO~.1/7NaOH-P\'a~C03-HP0. by adding the Na3P04.1/7NaOHto the solutions until only a small crystal remained in excess as the solid phase. No assurance could be had that the composition of the crystal would remain as NaaPOo.l/~NaOH, but if it did change, the amount involved would be so small that the composition of the solution would not be altered. Since the trisodium phosphate and sodium hydroxide are present in solution in the definite molar ratio of 7 to 1, the method of analysis previously described could be used to analyze the solutions. The question now arises as to how the composition of the solid phases are to be found. Unfortunately the foregoing procedure does not permit an exact determination of the character of the solid phases. However, the most important information desired about the composition of the solid phase is whether or not t. s. p. and sodium carbonate form a double salt, and if so, over what range of temperature and composition it exists. The presence of a double salt in a ternary system may be readily ascertained by an examination of the shape of the solubility isotherms of mixtures of the salts. If the curve has but two branches, the solid phase represented by each branch is the pure salt whose solubility is given by the terminal end of the curve. If a double salt is formed, the solubility curve at a given temperature mill have three branches, the new branch having the double salt as its solid phase. The determination of the solubility isotherms consisted of finding the solubility of the hypothetical salt l\'asPOh.l/?NaOH in water and sodium carbonate solutions of varying concentration. Baker's c. P. tertiary sodium phosphate was used; analysis showed that it had the constant ratio Ka3P04.1/7NaOH. The sodium carbonate used was Baker's c. P. anhydrous. For determinations a t 0" and 25" C. the decahydrate was prepared; a t 40" and above, the anhydrous salt was added which immediately formed the monohydrate.
20 1
The apparatus used t o attain equilibrium consisted of a disk wheel to which could be attached tangentially eight 8-inch (20.3cm.) Pyrex glass test tubes. The wheel was rotated at about 3 r. p. m. by a geared-head motor through a belt drive the wheel being immersed in a thermostat. Neoprene stoppers were found to be satisfactory a t 60" and 80" C., but at 100" they became plastic, and it was necessary to use glass-stoppered bottles and break the neck off if the stopper stuck. The run of 0" C. was made with the thermostat well filled with ice, and the temgerature was constant to within O.0jo, as it was at 25" C . At 40 and 60" C. the resulation was 0.1",and at 80' and 100" C . the regulation was 0.3 .
0
2
4
PERCENT
6
50, I N Liauio
8
10
I2
PHASE
FIGURE2. THE SYSTEMNa20-P20a-HzO 25" C. IX THE REGIONNasP04
AT
The experimental procedure for determining the various solubility isotherms was as follows: For runs at a given temperature, sodium carbonate solutions of varying concentrations were made up and rotated in the thermostat. T. s. p. was added until only a few small crystals were in excess, and the rotation was continued for B longer period. A sample was withdrawn into a weighed 10-ml. pipet through a coarse filter paper, and the increase in weight obtained. The sample was then diluted in a volumetric flask, and aliquot portions were analyzed by the method previously described. The solutions were then rotated for an additional period, usually a day, and reanalyzed to make certain that equilibrium had been reached. The cryohydric points of the system were found by adding the solid salts to finely ground ice and stirring the whole mass slowly in a Thermos bottle. A Beckmann thermometer showed that the temperature dropped rapidly to a minimum value at which it stayed constant as long as the solid phases were present and the stirring was continued. Samples were withdrawn and analyzed in the usual manner.
Solubility of Trisodium Phosphate Several workers have determined the solubility curve of
t. s. p., but the lack of agreement is startling. The earliest
0
20
40
60
BO
100
I20
T E M P E R A T U R E , 'C.
FIGURE 1. SOLUBILITY OF TRISODIUM PHOSPHATE
work is that of Mulder in 1894, quoted by Seidell (13), over the temperature range 0" t o 100" C. The solubility values of International Critical Tables (6) are based almost entirely on his results, though his work is nnt mentioned in their references. I. C. T. give 73.4" as the transition temperature from Na3P04.12Hz0 to Na3P04.XH20, as determined by Richards and Churchill (11) who did not investigate the composition of the solid phases. The work of Apfel in 1911 over the temperature range 0" to 75" C. is quoted by Seidell (14). The solid phases are stated to be Na3P04.12Hz0from 0" to 40" C., NasPOJOHzO from 50" to 60°, and Na3POa.8H20 from 70" to 75". Schroeder, Berk, and Gabriel ( l a ) determined the solubility of t. s. p. over the range 83" to 350" C.
INDUSTRIAL AND ENGINEERING CHEMISTRY
202
VOL. 32, NO. 2
Their solubility curve intersects that of Apfel a t about 7 5 " , and they interpreted this as a change in phase, possibly from 10 t o 8 molecules of water. They definitely showed that the solid phase from 120" to 215" C. was Na3P04.H20and above 215",P\;a3P04anhydrous. Solubility values a t 25" and 105" C. were determined by Obukhov and Mikhailova (IO) as part of a study of the system t. s. p.-sodium chloride-water. One solubility value a t 20" is reported by Teeple ( 1 7 ) . The data of these various workers are included in Table I for comparison with the results of this work. The data are given as moles t. s. p. per 1000 grams of water; this permits comparison by eliminating the difference in molecular weight between S a 3 POc and Na3P04.1/,KaOH. (See Figure 1.) OF TRISODIUM PHOSPHATE TABLE I. SOLUBILITY
Temp., C.
Moles T. 9. P. per 1000 Grams Hz0
c
0
This work 0.270
Mulder
Apfel
2.46
0.091 0.28 0.67 0.944 1.22 1.89 2.62 3.35
3.75
...
0.27 0 501 0.737 0.835 0,992 1.23 1.79 2.42 2.96 3.30
4.93
...
...
...
10
20 25 30 40 50 60 70 75
0:ioo
...
1.22
...
... ...
SO
...
... 5.29 ...
83 100 101 105 115 121 129 139
... ... ... ... .
.
6.57 .
a
... ... ... ... ...
...
Obukhov
... ...
0 : 944
...
... ... ... ... ...
...
...
... ...
... ... ... ... ...
... ...
6.03
...
... ...
...
...
Teeple or Schroeder
... 0'698 ( 1 7 )
... ... ... ... ... ... ...
3 : 7 i (it)
4 : + 3 (12) 8:4b (it) 5.68 (It) 5 . 5 0 (22) 5.40 (ft)
The values obtained by ilpfel are in fairly close agreement with those of the present work, which indicates that he also was dealing with a basic salt. Schroeder, Berk, and Gabriel (13) used Merck's c. P. trisodium phosphate in which the phosphate was determined by colorimetric analysis and the results calculated to Na3P04. The curve of Mulder and International Critical Tables crosses the other curves a t 20"C. and lies a remarkable distance above them a t the higher temperature. Mulder, Apfel, and Teeple all agree a t 20". It is known that sodium hydroxide has a depressing influence on the solubility of t. s. p. (19). It was therefore decided to determine its solubility in a solution in which the ratio of NazO t o PZOs was 3 to 1, the theoretical ratio in pure trisodium phosphate. Various saturated solutions of Xa3P04.l/7SaOH were made up; sodium hydroxide was added to some, to increase the alkalinity; to others, disodium hydrogen phosphate was added to decrease the alkalinity. The solutions were rotated as before a t 25" C.; the amount of solid phase was kept t o a minimum. Table I1 and Figure 2 show the data obtained, with a comparison with that of D'Ans and Schreiner (3). TABLE 11. THESYSTEM Na20-PzOs-H20 AT 25" C. % ?azo in Liquid Phase
% FzOn in Liquid Phase
YoNag0 in Liquid Phase
% PzO~ in Liquid Phase
4.85 4.91 5.06
1.79 2.38 1.29
5.11 7.80 8.56
2.80 6.30 7.12
The straight inclined line in Figure 2 is the locus of all compositions corresponding t o the theoretical 3-1 Na20-P205 ratio of pure trisodium phosphate. All points to the right of the line represent solutions containing more than the theoretical amount of acid, and all points t o the left represent solutions containing an excess of base. The curve of the present work crosses that of D'Ans and Schreiner exactly on the 3 to 1 ratio
3 GRAMS Ne2C4 PER IO0 GRAMS H P
FIGCRE3. THE SYSTEMNa3POa.l/7NaOHNa2C03-H20 FROM 0" TO 100" C. line, which means that the solubility in an exactly equivalent solution is the only point where agreement exists. At this point the solubility is 0.851 mole of Na3P04per 1000 grams of water as compared with 0.700 mole of iVa3P04.1/7KaOH per 1000 grams of water. The former value checks the work of Apfel exactly, which tends to show that a t 25O, a t least, his solutions must have been iYa3P04. The solubility found by Mulder a t 25" was 0.943 mole of Na3P04per 1000 grams of water, a much higher value. Higher solubilities are found when there is an excess of acid present, which might account for his higher results.
Trisodium Phosphate-Sodium Carbonate-Water Data for the system Na3P04.1/7NaOH-Xa2CO~-H~0 are given in Table I11 and Figure 3. Each isotherm has but two branches; thus the solid phase represented by each branch must be the same as the solid salt on whose coordinate axis the curve branch ends. The solid phases of the system sodium carbonate-water are well known; and since the solubilities found for pure sodium carbonate check those in the literature ( I @ , there is no doubt concerning these solid phases. I n the t. s. p.-water system all investigators call the solid phase . trisodium phosphate but disagree as to the hydrate existing a t the various temperatures. I n the industrial preparation of t. s. p. an excess of sodium hydroxide usually is added, and the crystals have a variable composition Ka3P04.1/7NaOH to Ka3P04.1/5NaOH. When this is recrystallized from its saturated solution, a more basic salt is formed. For this reason the designation " t. s. p." represents the salt used and is justifiable in view of the lack of information as to the true status of the system Xa20-P20j-H20at various temperatures.
FEBRUARY, 1940
INDUSTRIAL AND ENGINEERING CHEMISTRY
TABLE 111. THE SYSTEM Na3P04.1/7NaOH-Ka2C03-H20 Temp.,
c.
-2.48 -2.10 -1.21 0
40
Grams per 1000 Grams H20 IiazPOa.L/;NaOH xa2co3 1.8 5.5 0.0 6.1 4.2 0.0 4.58 0.00 2.58 6.43 0.00 6.93
20.8 15.1
11.6
11.1 0.0
60
41.8 36.6 31 .o 28.0 11.4 0.0
80 100
63.8 52.3
0.0 90.0
88.0
0.0
15.1 35.0 43.1 49.2 0.0
11.1 23.5 31.2 40.0 46.3 0.0 20.0 45.1 0.0
67.1
10.9 17.4
0.0
27.7 35.2 44.8
62.1 39.0 23.0
18.7
Solid Phase t. s. p. NazC03.10HzO KazC03.10Hz0 t. 8 . p. t. s. p. t. s. p. NazC03.10HzO
+ +
PiazC03.10Hz0
t.
9.
p.
t. 5 . p. t. s. p. t. s. p.
+
NazC03.HzO NanCOs.HI0 t. 8 . p. t. 8 . p. t. s. p. t. 8 . p. NasC03.Hz0 NazC03.Hz0 lia~C03.H~O t. 9. p. t. 8 . p. NazCOa.HZ0 NazCOs.HI0 t. 8 . p. t. 8 . p. f N a ~ C 0 3H20 KazCOS.Hz0 NazC03.HgO xa2C03.H~O KazC03.HzO NaKXhH20
+ +
Separation of T. S. P. and Sodium Carbonate
203
Literature Cited Assoc. Official Agr. Chem., Official and T e n t a t i v e Methods of Analysis, pp. 19-21 (1936). Booth, C . F., and Gerber, rl. B., U. S. P a t e n t 1,700,972 (Feb. 5, 1929). D’Ans, J., and Schreiner, O., 2. physik. Chem., 7 5 , 95-107 (191 1). ESG. CHmr., Anal. E d . , 10, Gerber, A. B., and Miles, F. T., IKD. 519-24 (1938). International Critical Tables, Vol. IVY, p. 237, New York, Mc.. Graw-Hill Book Co., 1926. Kolthoff, I. M., “Indicators”, p p . 120, 132, New York, J o h n Wiley &- Sons, 1926. Kolthoff, I. M., “Volumetric Analysis”, Vol. I, p. 129, Vol. 11, p. 142, New York, John Wiley &. Sons, 1928. Mellor, J. W., “Comprehensive Treatise on Inorganic a n d Theoretical Chemistry”, Vol. V I I I , p . 958, London, Longmans, Green and Co., 1923. Menzel, H., and Sahr, E. yon, 2. Electrochem., 43, 104-19 (1937). Obukhov, A. P., and Mikhailova, M . E.,J . Applied Chem. (U. S. S. R.), 8, 1149-57 (1935). Richards, T. W., and Churchill, J. B., 2. physik. Chem., 28, 31316 (1899). Schroeder, W. C., Berk, A. A., and Gabriel, A,, J . A m . Chem, SOC.,59, 1783-90 (1937). Seidell, A , , “Solubilities of Inorganic and Organic Compounds”,. Vol. I, p. 662, New York, Van Nostrand and Co., 1919. Ibid.,Vol. 11, p . 1432 (1928). Smith, J. H . , J . SOC.Chem. Znd., 36, 415-19 (1917). Ibid., 36, 420-4 (1917). Teeple, J. E., “Industrial Development of Searles Lake Brines”, A. C . S.Monograph 49, p. 164, New York, Chemical Catalog co.., ~1929. - - Wildeck, W.F., Lynn, G., and Hill, A. E., J . A m . Chem. S O ~ . , 54, 928-36 (1932). Westbrook, L. R., U. S.P a t e n t 1,711,707 ( M a y 7 , 1029).
PRESENTED before the Division of Industrial and Engineering Chemistry ati the 96th Meeting of t h e American Chemical Society, Milwaukee, \Tis.
The shape of the solubility isotherms indicates that the system is well adapted to the separation of sodium carbonate and t. s. p. Consider a solution whose composition is represented by some point inside the 25” C. isotherm. It is decided to evaporate the solution a t 100” C. As the water evaporates, the locus of the composition will be a straight line extended through the origin and the original composition When the composition reaches the 100” C. isotherm, the solution will be saturated and the solid phase represented by that particular branch of the curve will begin to crystallize out. At 100” C. the ;l\a2C03.H20branch of the isotherm is by far the longer, and unless the solution has a great excess of t. s. p. over sodium carbonate, the first solid phase t o appear would be Na2C03.H20. As evaporation continues, Na2CO3.H20will crystallize out and the composition will travel along the isotherm, finally stopping a t the invariant point, the intersection of the two branches of the curve. A t this point t. s. p. would begin to crystallize out along with the sodium carbonate. But advantage may now be taken of the inverted solubility curve of sodium carbonate. If the solution is cooled, the solubility of sodium carbonate will increase slightly while that of t. s. p. phosphate will decrease enormously, producing crystals of t. s. p.
Summary Kew data are given for the solubility of trisodium phosphate in water from the ice point to 100” C. They are compared with the earl’er data. A study of the system Na20-P206-H20 a t 25” C. showed the great influence of excess alkali or acid on the solubility, which explains the lack of agreement between the various workers. It was not possible to produce trisodium phosphate as the salt which crystallizes from a saturated solution is not of the same composition as the solution. A salt of composition NaSP04.’/7NaOH was used for this investigation. The solubility isotherms for the system Na,P04.1/,NaOH-Ka2C03-H20 were determined a t 0 O, 25”, 40”,60”, 80”,and 100” C. The separation of the two salts is discussed and a control method given.
Courtesv, Victor Chemical Works
DCSTCOLLECTOR AT THE ENDOF THE CONTINUOUS PRODUCTION SYSTEMFOR SODIUM PHOSPHATE
