The System Sodium OxidePhosphorus Pentoxide-Water J
BERNARD WENDROW AND KENNETH A. KOBE University of Texas, Austin, Tex.
T
H E sodium orthophosphates are important industrial chemicals, yet data pertaining t o their solubilities and t h e sodium phosphate-water system are not plentiful and also are in disagreement. It is the purpose of this paper to present data for the system sodium oxide-phosphorus pentoxide-water. These d a t a will necessarily include t h e binary systems trisodium phosphate-water, disodium phosphate-water, and monosodium phosphate-water.
hot water, Kobe and Leipper obtained a product with the formula Na3P04.12H~0.1/5NaOH. Using 3 equivalents of sodium hydroxide with 1 of phosphoric acid the same product was obtained.
PREVIOUS WORK
Results of previous investigations of the chemistry of the sodium orthophosphates are not consistent with one another, and they have led t o considerable confusion. When trisodium phosphate is crystallized from water solutions, the product contains sodium oxide in excess of t h a t required to combine in stoichiometric proportions with the phosphorus pentoxide present. It is this property t h a t has led t o t h e discrepancies in t h e reported solubilities of trisodium phosphate (TSP). The investigation of t h e behavior of these solutions has been the subject of a number of studies and patents.
COURTESY VIOTOR CHEMICAL WORKS
Storage Tanks for Phosphate Liquors Used in Production of Sodium Phosphates In his attempts t o prepare trisodium phosphate, Smith (29) found t h a t upon adding theoretical quantities of phosphoric acid and sodium hydroxide, crystals were obtained having the approximate composition: 2Na3P04 Na2HPOd. The mother liquor, which upon slight concentration solidified t o a mass of crystals, possessed the composition: 18NaaPO4 2XaOH. Using a 4% excess of sodium hydroxide, crystals were obtained 2NaOH. As a result of having a composition of 17.5 NaaPOa further experiments, Smith found t h a t he could not obtain a pure trisodium phosphate from phosphoric acid and sodium hydroxide, and concluded t h a t trisodium phosphate either cannot be crystallized from its components or else i t does not exist under normal conditions. A complete investigation of trisodium phosphate solutions at 20" C. was made by Menzel and von Sahr (16). They found that as the mole ratio of sodium oxide t o phosphorus pentoxide in t h e solution was varied from 2.69 t o 3.68, the same ratio in t h e equilibrium solid phase quickly increased from 3.11 t o 3.22. Thereafter, as the sodium oxide-phosphorus pentoxide ratio in the solution was increased t o 144.9, the ratio i n t h e solid phase increased slowly from 3.22 t o 3.24. They concluded that the solid phase in equilibrium with solutions of trisodium phosphate is not Na3P04,which has a sodium oxide-phosphorus pentoxide ratio of 3.00; rather, t h e solid phase appears t o be a series of solid solutions in which the sodium oxide-phosphorus pentoxide mole ratio continually changes from a value of 3.11 t o 3.24. The former value conforms t o a formula having the composition Na3P04.12Hz0.1/9NaOH, and t h e latter t o the formula Na3P04.12Hz0.1/4NaOH. However, x-ray patterns of four members of the series did not reveal any differences among them.
+
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COURTESY VICTOR CHEMICAL WORKS
K i l n and Cooler in Production of Trisodium Phosphate Material goes from cooler to hopper for packing into Multiwall paper bags
It is an established fact t h a t both commercial and reagent grade trisodium phosphate dodecahydrate contain some free alkali, which is apparently held in the crystal lattice. Kobe and Leipper (11)found t h a t a number of samples of reagent grade salt corresponded to the formula Na3PO4.12HzO.l/7NaOH, and other investigators have reported t h a t the fraction of sodium hydroxide per molecule of trisodium phosphate varies from 1/11 t o 2/9 (21). Upon r&crystallizationof the reagent grade salt from
1439
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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
Vol. 44, No. 6
Figure 1. Comparison of Previous and Present Work on the System Sodium Oxide-Phosphorus Pentoxide-& ater at 20" a n d 25" C. Prior t o the work of Menzel and von Sahr, there were several patents issued describing processes for the manufacture of alkalifree trisodium phosphate (5, 32). Sodium fluoride is sometimes used in the crystallization of trisodium phosphate in order t o reduce caking of the crystals formed. The composition of the crystals has been in question, and they have been reported both as trisodium phosphate decahydrate and as a double salt of trisodium phosphate and sodium fluoride. The latter is a definite possibility, since sodium fluoride is a persistent impurity in phosphates. Mason and Ashcraft ( 1 4 ) report t h a t the refractive indices of the supposed decahydrate and the double salt are the same. They undertook the determination of the composition of the compound in question, first seeding a solution of the composition given in Westbrook's patent (58) with the double salt, which they had prepared earlier. KO crystals other than the double salt were obtained, indicating absence of the decahydrate. They then investigated the system trisodium phosphate-sodium fluoride-water a t 25 O C. and found that the only solid phases existing are trisodium phosphate dodecahydrate, the double salt 2NaaPOa.NaF.19Hs0, and sodium fluoride. They concluded that in all cases in which trisodium phosphate decahydrate has been reported, the substance was actually the double salt. Bell (5)studied the hydrates of true trisodium phosphate and prepared several of the complex hydrates. He states that there are three hydrates of true trisodium phosphate, a hemihydrate, a hexahydrate, and an octahydrate, and t h a t any hydrate of trisodium phosphate higher than the octahydrate is actually a complex containing a monobasic anion. These higher hydrated complexes are divided into two groups, one given by the formula nNaaPO~.NaY.zH~O where n is 1 or 2, Y is a monobasic anion, and x is 18 or 19. The second type is given by the formula NasP04.zH20.l/nNaYl where n is 4 to 7, x is 11 or 12, and Y is a monobasic anion. The ordinary trisodium phosphate dodecahydrate of commerce is a complex belonging to the second type, and it is this one that Bell studied in detail. He varied the mole ratio of sodium oxide t o phosphorus pentoxide in the crystallizing solution from 2.82 to 3.75, and obtained complexes having ratios of either 3.21 or 3.12. This shows t h a t there are two hydroxide coinplexes having definite formulas, rather than a series of solid
solutions in which the sodium oxide-phosphorus pentoxide mole ratio varies. To the ratio 3.21 he assigned the formula Sad'O4.12H20.1/5NaOH and t o the ratio 3.12 the formula Na3P04.12H,o.~/~N~oH. The amount of previous work on the system sodium oxidephosphorus pentoxide-water is small, and the ranges of concentration and temperature are entirely inadequate. The work of D'Ans and Schreiner (7) a t 25" C. is the only isotherm recorded in the literature, and it is not complete. I n the tertiary region, they list the solid phase as trisodium phosphate dodecahydrate, as determined by the wet residue method. The wet residue from only one of their samples was analyzed, which is not an acceptable procedure when using this method. KOmention is made of the fact that the dodecahydrate is actually a complex. The isotherm was not carried far enough into the highly alkalinc rrgion and thus the solubilities of trisodium phosphate hexahydrate and trisodium phosphate hemihydrate were not found. In the secondary, or disodium phosphate (DSP) region, the solid phases found were the dodecahydrate and the heptahydrate. Here again the data were insufficient t o characterize the solid phases. In the monosodium phosphate (MSP) region the solid phase is given as the dihydratp. The hemisodium phosphate region wtt8 not investigated. Kobe and Leipper (11)studied the system at 25" C. over a short range in the trisodium phosphate region, the phosphorus pentoxide concentration varying from 1.79 to 71.2%. The solid phase was recognized to be the trisodium phosphate dodecahydrate complex. The work of Menzel and von Sahr (16)at 20" C., and of Kobe and Leipper ( 1 1 ) and D'Ans and Schreiner ( 7 )at 25' C. are shown in Figure 1 in comparison with the results of the present work a t 25O
c.
EXPERIMENTAL
SAMPLE PREPARATION. A standard-type constant temperature bath fitted with automatic controls was used throughout this work. For the work a t 0' C. a vapor compression refrigeration system provided the cooling required. Ethylene glycol was added to the water in order to prevent ice formation. Water was used as the bath liquid at 25', 40°,and 60' C., and white mineral oil a t 80' and 100' C. To provide agitation for the samples, a 12-
June 1952
INDUSTRIAL AND ENGINEERING CHEMISTRY
inch diameter brass disk wheel and stand were constructed, the wheel having clips bolted t o it, so t h a t 16 samples could be accommodated at one time. The wheel rotated in a vertical plane and was driven at a rate of 4 r.p.m. by a sprocket and chain drive through a gear reducer. Borosilicate glass-stoppered tubes were used for the samples a t 40" C., but for all other temperatures, sealed borosilicate glass tubes were used. At temperatures of 60" C. and above it was found that solutions in the highly alkaline region attacted the glass tubes. Under these conditions, sample tubes made of pure nickel were em loyed. All chemicai used were of the C.P. reagent grade. Orthophosphoric acid, disodium phosphate, and monosodium phosphate from J. T. Baker and Co., and disodium phosphate and hemisodium phosphate from Monsanto were used. The sodium phosphates were all checked for purity by determining the sodium oxide-phosphorus pentoxide mole ratio. Merck's sodium hydroxide pellets were used in the preparation of both t h e samples and the standard sodium hydroxide solutions. For t h e latter, a 50% solution of sodium hydroxide in distilled water was made and kept in a wax-lined bottle. Immediately before preparing a fresh 0.1 N solution, t h e proper amount of 50% solution was filtered through a sintered-glass crucible t o remove precipitated sodium carbonate and was poured into distilled water which had previously been aerated for about '/e hour. When prepared in this manner, t h e solution kept its titer for about a month and did not deteriorate. The samples were prepared using a modification of the method of complexes described by Purdon and Slater (20). It is known t h a t the compositions of t h e liquid phase, of the solid in equilibrium with the saturated liquid phase, and of the total mixture will lie on a straight line. Usually, it can be predicted what solid phase will be in equilibrium along a branch of the solubility curve A straight line is drawn between the points representing the composition of this solid phase and the desired liquid phase, and the total mixture of solid and liquid can then be made up t o correspond to a point on this line. After mixing t h e components, the entire sample was heated until the solid was all dissolved, and the sample tube was then placed in the constant temperature bath in order that t h e solid phase would crystallize a t t h e proper temperature. In some cases i t was found necessary t o seed the solution or t o stir i t vigorously with a stainless steel rod in order t o promote crystallization. After reaching equilibrium, the sample tube was removed from the agitating wheel and placed in a special rack within t h e bath. Liquid phase samples were removed after the solid phase had completely settled out. The liquid was drawn into a pipet, the tip of which was wrapped with filter paper in order t o exclude any crystals. The sample was immediately transferred t o a tared weighing bottle and weighed. This procedure worked very well at teniperaturas u p t o and including 40' C., but at the higher temperatures i t was Eound necessary to heat the pipet to a temperature slightly higher than t h a t of the sample, in order t o prevent crystallization. Solid phase samples were usually taken by quickly filtering the sample through a coarse sintered-glass filter crucible. A portion of the solid, with any adhering mother liquor, was then transferred t o a weighing bottle and weighed. Filtration through the sintered-glass crucible is not safe a t higher temperaturee because the sample immediately begins t o cool and crystallization from the mother liquor results. T o accompliah removal of the solid phase, a small stainless steel spoon was inserted in the sample tube, and a portion of the wet crystals was removed and immediately placed in a tared weighing bottle. ANALYTICAL METHOD I n the analysis of both the liquid and solid phases, it is necessary t o determine the amount of sodium oxide and phosphorus pentoxide. Approximately 800 samples were analyzed in this present work, so it is obvious t h a t a method of analysis which is both accurate and rapid was necessary. This suggested a volumetric method, confirmed by periodic checks using an established gravimetric method. The volumetric method involving precipitation of ammonium phosphomolybdate was first tried, but it was found undesirable for this work because of the time involved in washing the precipitate free of nitric acid, and of the uncertainty of the composi-
1441
tion of the precipitate. Other methods in which the phosphate ion is precipitated as silver phosphate, and t h e residual silver determined by t h e Volhard method, were also considered but abandoned. The method of Gerber and Miles (8) for the analysis of phosphates was found t o incorporate t h e desired characteristics of accuracy and speed, and was used throughout this work except for highly alkaline solutions in which the phosphorus pentoxide content was 0.8'% or less. For these solutions, the magnesium ammonium phosphate gravimetric method according t o Kolthoff and Sandell (1W) was used. The gravimetric method was used as a check on the volumetric method of Gerber and Miles and also in the analysis of t h e starting materials. Comparison of t h e results of each method shows t h a t t h e volumetric method is completely reliable and has t h e same precision as the gravimetric method. I n addition, during the work on t h e isotherms, many of the samples varying in phosphorus pentoxide concentration from 0.139 t o 60.95% were analyzed by both methods. The results, shown in Table I, are in excellent agreement, except for concentrations below 0.8yo. However, as mentioned earlier, the gravimetric method was used in all the latter cases.
TABLEI. COMPARISON OF VOLUMETRIC ANALYSES WITH GRAVIMETRIC ANALYSES P206, %
Gravimetric
Volumetric
Gravimetric
Volumetric
The composition of the equilibrium solid phase was determined by the use of Schreinemakers' wet residue method. The wet residue consists of the crystals, plus some amount of adhering mother liquor. For increased accuracy in t h e use of Schreinemakers' method, the mother liquor should be removed as completely as possible. A weighed sample of t h e wet residue w&s taken along with each liquid phase sample, dissolved in water, and analyzed by the method of Gerber and Miles. I n addition t o a chemical analysis of the wet residue, the crystals from each sample were examined wider t h e microscope. This microscopic examination proved t o be a n extremely valuable aid in comparing t h e various solid phases or in differentiating among them. For instance, both t h e crystals of t h e alkaline complex of trisodium phosphate dodecahydrate and octahydrate appear t o be long needles, and no difference can be detected between them in ordinary light. However, in polarized light, the crystals of the dodecahydrate exhibit parallel extinction, whereas those of the octahydrate have extinction angles varying from 7 t o 23 degrees (14). This fact was used t o prove t h a t there is no transition from t h e dodecahydrate t o the octahydrate a t 25" C. as there is Let 40" and 60" C. SODIUM OXIDE IN PHOSPHORUS PENTOXIDE-WATER SYSTEM
The system is shown in Figure 2 and the d a t a are given in Table 11. The highly alkaline region of t h e system is shown in Figure 3, in which the abscissa has been greatly expanded t o show clearly t h e isotherms under conditions of low phosphorus pentoxide content and high sodium oxide content of the liquid phase. It can be seen that t h e system is not a simple one, there being as many as 10 different solid phases appearing a t 25' C., and a comparable number at the other temperatures. The solid phases found in t h e tertiary region of the isotherms are the anhydrous salt, the hemihydrate, hexahydrate, octshydrate, and the alkaline complex salt. T h e anhydrous salt is
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
1442
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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
Vol. 44, No. 6
found only in highly alkaline solutions a t the lower temperatures but a t 100" C. it exists as the solid phase over a relatively large concentraAs the tion range. s o d i u m oxide-phosphorus pentoxide inole ratio in the solution is decreased, the hemihj-drate of trisodium phosphate is e n c o u n t e r e d . Further decrease in this ratio causes the appearance of the hexahydrate, which is seen t o be the major solid phabe existing in the tertiary region a t 80" C. At 25", 40", and 60" C. there is a transition from the hexahydrate to the alkaline complexsalt. At 25"C., two of the complcsea were found to have a sodium oxide-phosphorus pentoxide mole ratio of 3.14, representing 1/7 mole of exce.5s sodium hydroxide pc:r mole of trisodium pho5phate, and the remairidcr of the complexes had mole ratios of 3.25, representing 1/4 mole of excess sodium hydroxide per mole of trisodium phosphate, This indicates that there are two different alkaline complexes represented by the f o r mu1 a s NaJ'04.12H20.1/7NaOH and K a n P O * ,1 2 H 2 0 . 1 / 4 XaOH, r e s p e c t i v e l y , substantiating the viork of Bell (S), although he found ratios of 3.12 and 3.21. Since solid phases in this p r e s e n t w o r k were found by the wet residue method, the ratios 3.14 and 3.25 are claimed to be accurate only t o the limits of this graphical method. A t 40" and 60" C. all of the solid phases in the region in which the complex c r y s t a l l i z e s w e r e found to have a sodium oxide-phosphorus pento x i d e m o l e r a t i o of 3.25. Examination of the tertiary region shows t h a t there is no alkalifree dodecahydrate of trisodium p h o s p h a t e ,
I
INDUSTRIAL AND ENGINEERING CHEMISTRY
June 1952
and 40" C. there are transitions from higher t o lower hydrates in the acid side of the region. Points of minimum solubility are encountered on the line representing the locus of all the points having sodium oxide-phosphorus pentoxide mole ratios of 1.00. At 40" C., the highest hydrate of monosodium phosphate found is the monohydrate, although Imadsu (IO)in 1911 reported t h a t the dihydrate existed up t o 40.8" C. From the results of identification by Schreinemakers' method and by microscopic examination of crystals of both t h e mono- and dihydrate, carried out in this present work, it is believed that the temperature reported by Iinadsu is in error. As the phosphorus pentoxide content of the solution is further increased, a region is encountered in which the solid phase is hemisodium phosphate, NaHa(PO&. The solubility branch of this salt is quite similar in shape in each isotherm. T h e effect of an increase of temperature is to decrease the size of the region of crvstallisation and t o displace it toward Ihigher co'ncentrations of sodium oxide and phosphorus pentoxide.
WEIGHT PERCENT P 2 0 5
Figure 3.
Highly Alkaline Region of the System Sodium OxidePhosphorus Pentoxide-Water from 25" to 100" C.
but t h a t whenever trisodium phosphate appears with 12 molecules of water of hydration it also contains some free alkali. The octahydrate of trisodium phosphate was found only at 40" and 60" C. Figure 1 shows the system sodium oxide-phosphorus pentoxidewater including previous work and this present work. The data of Kobe and Leipper a t 25" C. agree well with those of this present work, and the data of Menzel and von Sahr at 20" C. also show a similar trend. It is concluded t h a t the data of D'Ans and Schreiner are incorrect in t h e tertiary region. I n contrast to the tertiary region, the secondary (disodium phosphate) region shows a marked similarity at each temperature. The decahydrate, oetahydrate, heptahydrate, and dihydrate of disodium phosphate, and t h e anhydrous salt were found as solid phases. I n the disodium phosphate region, including the transition point from the trisodium phosphate region, the data of D'Ans and Schreiner check those of this present work very well, until a concentration of about 14% phosphorus pentoxide is reached. The two isotherms begin to diverge here, and from this point on they differ markedly. I n this present work, a transition was found from the dodecahydrate of disodium phosphate to the octahydrate, while D'Ans and Schreiner report the lower hydrate as the heptahydrate. Again, as in the trisodium phosphate region, their result is based on only one wet residue analysis. To show that the solid phase was not the heptahydrate, solutions were made and seeded with the dihydrate, the heptahydrate, and the anhydrous salt. At equilibrium it was found that all had changed over t o the octahydrate, which must be the stable phase in this region. The secondary and primary regions of the isotherms are eonnected by short branches along which the solid phases are double salts. At 25 and 40 C. the double salt is NazHP04.2NaHzP04.21320 and at 60" and 100"C. it, is NazHPOa.NaHzPO1. A characteristic of each isotherm in the secondary region is the point of minimum solubility which coincides exactly with the line representing the locus of all sodium oxide-phosphorus pentoxide mole ratios of 2.00. These minimum points appear as sharp breaks in the isotherms, yet do not mark a change in the solid phases. Similar points appear in the systems potassium oxide-phosphorus pentoxide-water and ammonia-phosphorus pentoxide-water as found by Berg ( 4 ) , Muromtzev and Nazarova ( I ? ) , Parker (19)' and Ravich (33). D'Ans and Schreiner (7) thought the singular points may be rounded minimums, but the data of Muromtzev and Nazarova prove this is not so. I n the monosodium phosphate region the solid phases found are the dihydrate, the monohydrate, and the anhydrous salt. The isotherms are similar in shape a t each temperature, but at 25'
1445
SOLUBILITY OF TRISODIUM PHOSPHATE
The points at which the isotherms cross the line of constant sodium oxide-phosphorus pentoxide mole ratios of 3.00 represent t h e solubilities of trisodium phosphate in water. The solubilities of t h e sodium phosphates were found at 0' C. as well as a t 2 5 " , 40 ', 60 80 and 100 C. There is quite a lack of agreement in the published data on the solubility of trisodium phosphate, and the discrepancies are believed t o be due to the starting material used, since if reagent grade trisodium phosphate is dissolved in water the sodium oxide-phosphorus pentoxide mole ratio in t h e solution will be about 3.12, instead of the desired value of 3.00. From the &ta of this present work i t is possible t o pick a t each temperature the point a t which the solution has a sodium oxidephosphorus pentoxide mole ratio of 3.00, and thus calculate t h e solubility of trisodium phosphate. These calculated solubilities will therefore differ from those reported in the literature. These data are shown in Figure 4 and Table 111. O,
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TEM'PERATURE, " C . Solubility of Trisodium Phosphate i n Water
INDUSTRIAL A N D E N G I N E E R I N G C H E M I S T R Y
1446
TABLE 111. SOLCBILITY OF TRISODIL-x PHOSPHATE Temp., ' C. 0 10
20 25 30 40 50 60 70 75 80 83 100 101 105 115 121 129 139
(Moles of trisodium phosphate per 1000 grams of water) Kobe and Aluider Apfel Obukhov This Work Leipper (11) (26) (27) (18) 0.328 0.270 0.091 0.27 ... ... ... ... 0.25 0.501 0.67 0.737 0:887 0:ioo 0.944 0,835 0:044 ... ... 1.22 0.992 1.42 1.22 1.89 1.23 ... ... ... 2.62 1.79 3.31 2.46 3.35 2.42 ... ... ... ... 2.96 ... ... ... ... 3.30 ... 3.75 4.15 4.93 ... ...
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present work agree with Menzies and Humphrey at 40" C., but are consistently lower from 50 to 95" C.; a t 100' C. the value is almost coincident with t h a t of Shiomi. Transition points were not determined directly in this work, but can be approximated from the extrapolations of the data, as shown in Figure 5. They are as follows:
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Vol. 44, No. 6
35.50 4 8 . 0 ' C. 9 5 . 0 ' C.
These data agree rather well with the ones reported by the other investigators. I n addition, Hammick, Goadby, and Booth (9) report the existence of two forms of NazHP04.12H20, with a transition point a t 29.6" C. Because of the nature of the present work, this was not checked. SOLUBILITY O F MONOSODIUM PHOSPHATE
Below about 55" C. a solid phase of lia3PO4.1/4NaOH.12HZO is in equilibrium with a solution having a sodium oxide-phosphorus pentoxide ratio of 3.00. Therefore it is a matter of d e h i tion whether t h e solubility of trisodium phosphate is defined as the composition of a saturated solution having a sodium oxidephosphorus pentoxide ratio of 3.00 or 3.25. Here the former viewpoint has been chosen: The solubility is the amount of trisodium phosphate t h a t can be dissolved in a solution before a solid phase will separate. If the solubility is desired in terms of a sodium oxide-phosphorus pentoxide ratio of 3.25 or 3.14 (11) these values can be read from Figure 2. T h e heavy line in Figure 4 connects the data obtained in this work. The dotted portions of the curve are extrapolations, and their intersections indicate transition points. The complex salt is the soild phase u p t o about 55" C., then NapPOa.8H20 t o about 65" C., and finally NaaP0,.6H20 up t o 100" C. Apfel ( 2 7 )repcrted the solid phases as dodecahydrate from 0 to 40" C., the decahydrate from 50" t o 60" C., and t h e octahydrate from 70" t o 75 O C. The decahydrate does not exist, as has been shown, and his temperature range for the octahydrate is higher than that of this work. Mulder, whose data are given in the first edition of Seidell (86), determined solubilities over the temperature range 0' to 100" C. but did not investigate the composition of the solid phases. The transition from the decahydrate t o a lower hydrate at 73.4" C. reported by Richards and Churchill ( 2 3 ) appears t o be in error. Kobe and Leipper ( 1 1 ) determined the solubility of NatPOa.12Hz0.1/7KaOH over the range 0" to 100" C. Their results agreed fairly well with those of Apfel and with a single determination at 20' C. made by Teeple (30). The work reported here at 25" C. agrees fairly well with t h a t of Apfel ( 2 7 ) , Mulder (26), and Obukhov and Mikhailova ( 1 8 ) but is somewhat higher than t h a t of Kobe and Leipper. For the other temperatures, except 60 C., the values of this work are in between those of Mulder and the other investigators. At 60" C. there is almost exact agreement with Mulder. Schroeder, Berk, and Gabriel ( 2 5 ) extended the temperature range from 75 ' to 350 " C. hut they did not identify solid phases below 120" C. Their values are considerably below those reported here. O
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SOLUBILITY OF DISODIUM PHOSPHATE
D a t a for the solubility of disodium phosphate are given in Figure 5 and Table IV. The values of all the investigators agree very well from 0" t o 30" C.; a t 35" C., the d a t a of Hammick, Goadby, and Booth (9) and Shiomi ( 2 8 ) agree with the extrapolated data of this present work, but are higher than t h a t of Menzies and Humphrey ( 1 6 ) . Shiomi and Menzies and Humphrey agree well in the range from 45" t o about 75' C., but their values differ above the latter temperature. The data of this
The solubility data for monosodium phosphate found in this work agree well with existing data a t O", 40°,and 100" C. but are 6.
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SOLUBILITY OF DISODIUM PHOSPHATE
(Moles of disodium phosphate per 1000 grams of water) Temp., This Hammick Menzies and Shiomi Menzel and 0 c. Work et al. ( 9 ) Humphrey (16) (18) Gabler ( 1 4 A ) -0.47 0.00 0.05 6 10.26 15.11 18 19.95 20.0 25.00 25.75 29.5 30.1 30.21 30.76 32.5 33.04 34 34.7 36.5 37.27 40.0 45.0 47.23 50 55.17 60.0 70.26 80.0 89.74 90.2 96.2 99.77
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