318
T. D. FARR, GRADYTARBUTTON AND H. T. LEWIS,JR.
TVA data for the heats of solution of the respective calcium compounds and for the heats of formation and Of phosphoric acid; (b) published data for the heats Of formation4 Of 'Oh3- and F-, for the entropies4of Ca++, F-, Ca(s), F2(g) and
SYSTEM CaO-PzOb-HF-HzO:
02(g), for the entropy6 of Po43- (-52 and for the entropy? of fluorapatite.
Vol. 66 f
2 e.u.),
(6) C. C , Stephenson, J. Am. flhem, Sot., 66, 1436 (1944). (7) E. P. Egan, Jr., Z. T. Wakefield and K. L. Elmore, ibid., 75, 5581 (1951).
EQUILIBRIUN AT 25 AND $0''
BYTHAD D. FARR, GRADYTARBUTTON AND HARRY T. LEWIS,JR. Division of Chemical Development, Tennessee Valley Authority, Wilson Dam, Alabama Received September 91,1081
Phase equilibria in the system CaO-Pz06-HF-HaO a t 25 and 50" were determined for the region represented by liquid P206, 0.7 to 5,5y0 CaO and less than o.o7yO F. The stable solid phases in equilibrium with the phases containing 4 to saturated solutions were calcium fluoride and fluorapatite or calcium fluoride and monocalcium phosphate monohydrate. The invariant points representing solutions saturated with all three compounds were located a t both temperatures. Measurements on the saturated solutions included pH, density and vapor pressure.
The presence of fluorapatite, Ca,o(PO&Fz, as a major component in rock phosphate focuses interest on the system CaO-PzOs-HF-HzO in its relation t o the manufacture and use of phosphatic fertilizers. Phase equilibrium in a portion of the system was studied with three major objectives: To determine whether dicalcium phosphate fertilizer can be made directly from rock phosphate and the stoichiometric proportion of phosphoric acid at temperatures in the range 25 to 50°, to determine the relative rates of equilibration from the directions of supersaturation and of undersaturation, and t o determine some of the thermodynamic properties of the saturated solutions. Materials and Methods Monocalcium phosphate monohydrate, dicalcium phosphate and phosphoric acid were crystallized twice from the reagent materials. Hydrofluoric acid was redistilled in platinum and the middle fraction was retained. Calcium fluoride was prepared from twice-recrystallized calcium nitrate and the purified hydrofluoric acid. Tricalcium phosphate and fluorapatite were prepared by methods described previously.2 About half the equilibration mixtures were prepared from phosphoric acid and fluorapatite to approach equilibrium from the direction of undersaturation. The other mixtures were prepared from phosphoric acid and various combinations of monocalcium phosphate monohydrate, dicalcium phosphate, tricalcium phosphate, fluorapatite, calcium fluoride and hydrofluoric acid. The mixtures were equilibrated in 500-ml. hard rubber or polyethylene screw-cap bottles that were rotated end over end in water-baths a t 25 f 0.08" or 50 rt 0.04". The wet solids were identified petrographically, with some confirmations by X-ray diffraction. Slow settling of solids complicated sampling of the liquid phases, which were analyzed in duplicate. The phosphorus content generally was determined by double precipitation as magnesium ammonium phosphate, with ignition t o magnesium pyrophosphate. A few phosphorus determinations were made by a differential spectrophotometric method .a Calcium was determined by double precipitation as the oxalate, generally with ignition to calcium oxide. Some of the precipitates were ignited a t 475 to 500" and weighed as calcium carbonate. Fluorine was determined by a method that has been de~cribed.~ (I) Presented before the Division of Phyeioal and Inorganic Chemistry, 132nd National Meeting of the American Chemical Society, New Y o r k , N. Y., September 1957. ( 2 ) E. P. Egan. Jr., 2. T. Wakefield and X. L. Elmore, J . Am. Chem. Soe., 1 8 , 5581 (1951). (3) A. Gee and V. R. Deitz, Anal. Chem., 26, 1320 (1053).
Densities of the solutions were measured in 25-ml. pycnometers calibrated a t 25 f 0.005' or 50 f 0.005". The pH's of the solutions were measured by means of a Beckman Model H-2 meter, with glzkss and saturated calomel electrodes calibrated at 25 and 50". A few check determinations of p H were made with a hydrogen half-cell in a different system. In measuring the vapor pressure of a saturated solution, a technique of alternate stirring and freezing was used for removal of dissolved extraneous gases without changing the composition of the solution. The solution was stirred in a 25-ml. bulb by means of a perforated platinum disk, lifted magnetically 8 times per min. After 5 to 10 min. of stirring, the liquid was frozen quickly at -78" and evacuated. The solid was melted a t room temperature with the bulb isolated from the vacuum system and the cycle was repeated until the ressure over the solid was 10-4 mm. on a McLeod gage usually 8 to 10 cycles). The bulb with degassed test solution and stirring mechanism, a mercury isoteniscope, and the critical connecting lines were immersed in a water-bath a t 25 f 0.005" or 50 f 0.005". The nitrogen pressure needed to balance the isoteniscope, as indicated by an electronic relay system, was measured with a manometer containing Monsanto Arochlor No. 1242, which gave a magnification factor of 9.78. The manometer was read with a cathetometer. The apparatus was tested by measuring the vapor pressure of conductivity water. The results of replicate determinations, 23.76 f 0.01 mm. a t 25" and 92.58 i 0.01 mm. a t 50°, are compared with the value 0.032287 kg./cm.2 (23.75 mm.) for 25" recommended by Osborne, Stimson and Ginningss from their compilation of published vapor pressures and the value 92.56 mm. for 50" given by Keyes.*
P
Equilibration at 25".--?dost of the complexes at 25' were sampled at 12, 16, 21 and 24 months. Some of the more viscous liquids were sampIed only twice because of slow settling and the difficulty of getting clear samples. Keither the compositions of the liquid phases nor the forms of the solid phases changed significantly after 18 months. Properties of the system at equilibrium are summarized in Table I. The low fluorine content of the liquid phases (C + C F FA + C F 11 MC + C F DC 4.71 4.95 1.24 1.06 DC + F A + C F FA CF CF CF FA 12 MC CF DC HI? 26.91 26.93 5.36 5.30 DC C F FA CF DC CF DC FA 13 hlC CF TC 4.11 4.21 1.11 1.13 T C CF CF CE' 14 MC CF TC IIE' 26.56 26.55 5.61 5.56 Aqueous H#04 was part of each charge; FA = fluorapatite, MC = monocalcium phosphate monohydrate, CF = calcium fluoride, DC = dicalcium phosphate, T C = a-tricalcium phosphate, TIP = hydrofluoric acid. * Identified petrographically in wet solids; first symbol represent,s major phase. 1
2 3 4 5 6 7
+ + + + +
+ + +
+
+ + + +
+ + +
+ +
+
major phasc, and fluorapatite was observed as a precipitated phasc in various degrees of crystal growth (acicular to well-formed hexagonal prisms) on nuclei of dicalcium phosphate. At 11 months, dicalcium phosphate had decreased to a minor phase comprising skeletal prisms that had grown on tricalcium phosphate cores no longer present; fluorapatite, the major phasc, was present as rod crystals in radial clusters around dicalcium phosphate seeds. The change in complex 14 was extremely slow; calcium fluoride w n s the only solid phase that could be identified petrographically at 11 months. Most of the fluorine in complex 14 was added as hydrofluoric acid.
Conclusions E'luorapatite and calcium fluoride are the stable solid phases in equilibrium with saturated liquid phases whose compositions cover the respective ranges PD6, 70 CaO, % A t 25" At 50"
4 to 25.8 4 to 29.9
0 . 9 to 5.47 0.7 to 5.51
In this rcgion, tri- and dicalcium phosphates are inctastablc phases; tricalcium phosphate, when added to the system, was converted to dicalcium phosphate and, in turn, to fluorapatite. Monocalcium phosphate moriohydratc and calcium fluoride are thc stable solid phases in equilib-
+
+ + + + + + +
+
+ +
+
+
rium with saturated liquid phase whose compositions cover the respective ranges p2061 % CaO, % At 2.5-
At 50"
2
