INDUSTRIAL AND ENGINEERIXG CHEMISTRY
568
R = gas constant, B. t. u./(lb. mole)(" F.) T = temperature elevation of fuel bed above surroundings, O F. T , = temperature elevation of air and combustion gases, F. Ti = ignition temperature of fuel, F. T o = temperature of fuel at 5ottom of bed, F. T = absolute temperature, Rankine Tl5,T7j = temperatures at which rate of increase of fuel temperature in Coal Research Laboratory reactivity test reaches 27" F. (l!' C.) and 135" F. (75" C.) per min., respectively, F. U = rate of ignition, lb./(sq. ft.)(hr.) a = coefficient of heat transfer between solids of fuel bed and combustion gases, B. t. u./(cu. ft.)(hr.)(' F.) 8 = coefficient of heat transfer from bottom of fuel bed to surroundings,B. t. u./(sq. ft.)(hr.)(' F.) O
O
++
-
rate of reaction C 0 2 C O , lb. OZ/(CU.ft.)(hr.) rate of reaction C 0 2 + CO2, a t the plane of ignition, lb. 02/(cu. ft.)(hr.) = specific heat of fuel, B. t. u./(lb.)(" F.) = =
pi pi
p
This nomenclature is the same as that used previously (7, 8) except that P here is equivalent to Vi in the earlier papers.
VOL. 32, NO. 4
Literature Cited (1) Allcock, H. J., and Jones, J. R., "The Somogram", London, Sir Isaac Pitman & Sons, 1936. ( 2 ) Furnas, C. C., Trans. Am. Inst. Chem. Engrs., 24, 142-66 (1930). (3) Hewes, L. I., and Seward, H. L., "Design of Diagrams for
Engineering Formulas", New York, McGraw-Hill Book Co., 1923.
(4) Hottel, H . C., and Smith, 1 '. C., Trans. Am. SOC.Mech. Engrs., 57. 463-70 (1935).
( 5 ) Mayers, M. A , , Chem. Rev., 11, 31-53 (1934). (6) Mayers, M. A . , J . A m . Chern. SOC.,6 1 , 2053 (1939). (7) Mayers, M.A,, Trans. Am. Inst. M i n i n g Met. Engrs., 130, 40823 (1938). (8) Mayers, M. -4.,Trans. Am. SOC.Mech. Engrs., 59, 279-88 (1937). (9) Ibid., 59, 260, Equation 6 (1937). (IO) I b i d . , 59, 282, Equation 21 (1937). (11) Nicholls, P.. U. S. Bur. Mines, Bull. 378 (1934). (12) Ibid., Bull. 378, 9-11 (1934). (13) Parker, A. S.,and Hottel, H. C., IXD. ENG.CHEM.,28, 1334-41 (1936). (14) Sebastian, J. J. S., and Mayers, M. A., Ibid., 29, 1118-24 (1937). (15) Smith, D. F., and Gudmundsen, A,, Ibid., 23, 277 (1931). (16) Terres, E. et al., Angew. Chem., 48, 17-21 (1935). (17) Thiele, E. W., IND.ESG. CHEM.,31, 916-20 (1939). (18) T u , C. M., Davis, H., and Hottel, H. C., Ibid., 26, 749-57 (1934). PRESENTED in the Symposium on t h e Combustion of Solid Fuels before the Division of Gas and Fuel Chemistry a t the 98th Meeting of t h e Amerioan Chemical Society, Boston, Mass.
Barium Chloride from Barite
and Calcium Chloride Factors Affecting Production R. NORRIS SHREVE AND R. K. TONER1 Purdue University, Lafayette, Ind.
HIS investigation was a continuation of experiments in
T
these laboratories concerning the utilization of organic solvents as a n aid in inorganic reactions; it was an extension of the published work of Shreve and Pritchard (6, 7 ) and was carried out simultaneously with that of Shreve and Watkins (8). The reaction Bas04
+ CaCI2e BaCL + Cas04
may be carried out in aqueous solution or in the fused state. This is in contrast to the reaction as it norm all^ occurs, since the reverse reaction is ordinarily assumed to be quantit'ative in the presence of water. Under the former condition i t was thought desirable to study the effects of reaction time, molar ratio of reactants, and quantity of water present, changing one variable a t a time. The effect of ternperature is not reported here since it mas adequately discussed in a previous publication ( 6 ) . The effect of varying time, temperature,and concentration of calcium &loride concurTenth' was next studied. Finally the effects of time, temperature, and molar ratio of reactants under conditions of fusion were investigated. These factors plus a study of the reaction from the opposite side (is e,, the formation of barium sulfate 1
Present address, Lehiph University, Bethlehem, Penna.
from barium chloride and calcium sulfate) gave the answer to the questions: How far can the reaction proceed under given conditions? How far does i t actually go? How fast is it? It has already been shown (6) that ethylene glycol-methano1 (1 to 3) is a good solvent for both barium and calcium chlorides, and that these salts can be removed quantitatively from the corresponding sulfates by its use. Attempts were made to substitute aqueous methanol for this more expensive solvent, the purpose being to increase the solvent power of methanol without incurring the reversion of the reaction which was caused by water alone.
Procedure All reactions in aqueous solution involving but one variable were carried out in a de Khotinsky constant-temperature Oven at 1750 * 10 C. (Other temperatures are given by Shreve and Pritchard, 6.) The reactants, consisting of carefully dehydrated c . P. chemicals, were weighed out into heavy-walled Pyrex glass \
~
;
f
$lltg ,t$a$i:dwt?
~,',",~~a~~ifi,",: ~ b ~ ~ ~
means of microvolumetric apparatus, and then the tube was sealed. A t the conclusion of the reaction time, during which the tubes constantly revolved, the tubes were cooled and broken open, and their contents quantitatively transferred to 125-ml. tincture bottles by the desired organic solvent. The reaction mass was agitated Tvith the organic solvent for 3 hours on a shaker. It was then filtered and the extract analyzed for barium chloride. Aqueous reactions involving variation of more than one factor were carried out on a pilot-plant scale in an electrically heated ball mill. A suspension of barium sulfate in a calcium chloride solution was charged to the mill, and heat xyas applied. As the
APRIL. 1940
IKDLSTRIAL AND ENGINEERING CHEMISTRY
The reaction BaSOa
+ CaCL e BaC12 + Cas04
was studied in concentrated aqueous solution at 173' C. and under conditions of fusion. The reaction was shown to have a true equilibrium state in each instance by studying it also as a reversion-i. e., from right to left as written. Under the former conditions a reaction time of 6-12 hours in the presence of sufficient calcium chloride to give a saturated solution gave large, practical conversions to barium chloride. Time and temperature had no appreciable effect on conversion in the fused state. Aqueous methanol was shown to be a satisfactory extraction solvent if the insoluble type of anhydrite was formed in the reaction. Otherwise anhydrous organic solvents had to be used to extract the barium chloride from the reaction mass.
of maximum conversion. This undoubtedly corresponds to a saturated solution of calcium chloride. Increasing the water content decreased the concentration and hence the conversion. Decreasing the water below a certain value left insufficient reaction solution and increased the mechanical difficulties of stirring. At 175 O C., the temperature a t which all single-variable aqueous reactions occurred, a saturated solution contains 75 grams of calcium chloride per 25 grams of water. This is a ratio of 0.678 mole of calcium chloride to 1.39 moles of water or 1 mole of calcium chloride to 2.05 moles of water, disregarding any bound water of crystallization. For ratios more than this the solution is unsaturated; for ratios less than this there will be undissolved calcium chloride. At least one molecule of water will be bound by the calcium chloride and probably t\vo (the transition point of CaCl2.2H2Oto CaC12.H20 is a t 175.5' C. in a solution of calcium chloride only), so that the solution of water must be in excess of this. This means that saturation may be reached when the ratio of calcium chloride to water is as low as 1 to 4. Any water in excess of this will certainly result in dilution of the solution. Theoretically when this ratio increases to approximately 1 to 2, all the water will be bound and reaction hindered. Under conditions of the experiment i t was probable that some solution was present even a t very high ratios of salt t o water, but this may have been small and thus slowed u p the reaction materially.
temperature increased, the water escaped. The reaction continued until the water was entirely removed. Extraction and analysis followed in the same manner as before. The fusions were carried out in a Hump furnace equipped with a Leeds & Northrup rectrder-controller which maintained the temperature within 1 1 0 F. (5.56' C.) of that desired. The chemicals used were of c . P. quality, and the actual fusion took place in platinum crucibles. The fusion masses were ground in a mortar before extraction. Inasmuch as no strontium salts \?-ere present and as calcium chromate is soluble in a dilute acetic acid solution while barium chromate is not, a simplification of the analysis previously reported (6) was used. In these experiments the filtered extract was diluted t o exactly 250 ml. Ten-milliliter aliquots were taken; each was diluted to 250 ml., acidified with 6-8 drops of glacial acetic acid, and heated to boiling. A 10 per cent neutral ammonium chromate solution was then added dropwise with stirring until precipitation was complete. After settling, the precipitate was filtered through a prepared Gooch crucible, dried over a direct flame, and weighed. This method gave results of quantitative accuracv. The curves shown here are probably accurate to *1-3 per cent. Those in Figure 1A are doubtless more accurate than those in Figure lC, for example, since the former are representative of reactions containing much water. Because of the hygroscopic nature of the salts used, traces of moisture were of more importance, the smaller the water content of the reaction mabs. Secondly, mechanical mixing became increasingly difficult as the water content decreased.
569
AMOLAR R A
I
w (r 0 I
1
I
6
I2
18
6
12
18
24
M
I
36
I
I
30
36
42
R-20
Reaction in -Aqueous Solution Figure 1 shows the effects of varying reaction time for a variety of reaction ratios. The major portion of reaction occurred in the first 2-3 hours, and after 6-12 hours the increment of increased yield was small. Figure 2 indicates the desirability of using excess calcium chloride when much water is present. As the water content was reduced, less excess calcium chloride was required for the same conversion. Figure 3 illustrates the eflect of varying the proportion of water. Each set of curves has an approximately flat portion
24 TIME
IN
48
HOURS
FIGURE 1. EFFECTOF TIMEo s CONVERSION A T 175" C
INDUSTRIAL AND ENGINEERING CHEMISTRY
570
100
VOL. 32, KO. 4
1 MOLAR RATIO
MOLAR RATIO CLL,
: B&O,
MOLAR RATIO
FIGURE 2. EFFECT OF CALCIUM CHLORIDE o s CONVERSION AT 175" C.
Inspection of the experimental results confirms the above conclusions. For ratios of calcium chloride to water less than 1 to 4, conversion decreased. It was approximately constant a t a maximum while saturated solution was present; but when the water content became too small, conversion dropped Off.
If water was allowed to evaporate as the reaction proceeded, good yields were obtained with only a slight excess of calcium chloride. Thorough mixing TYas attained while considerable water was present, and if the mater was driven off slowly (6-12 hours), conditions for maximum conversions were obtained. It was difficult to duplicate these experiments since there were many variables, but conversions of 83-92 per cent were achieved with but 10 per cent molar excess calcium chloride (Table I). If conversion computations are based on 1 mole of barium sulfate, if n is the number of moles of calcium chloride (n is to be greater than 1 in all these experiments) used to react n-ith this quantity of barium sulfate, and if the reaction takes place in the presence of z moles of mater, then the equivalent reversion reaction will consist of 1 mole of barium chloride, 1 mole of calcium sulfate, (n - 1) moles of calcium chloride, and 1: moles of water. If calcium sulfate dihydrate is used, (a: - 2) moles of water will be required. To put it another way, if two runs (one a conversion and one a reversion) are
C&:
B~sO,=l.l : I
HLO : b S O +
FIGURE 3. EFFECT OF WATER ON CONVERSION AT 176' C.
made up according t o the specifications just listed and then carried out under identical conditions for an infinite length of time, they will shon- identical analyses if the reaction is a true equilibrium. It is impossible and unnecessary to wait an "infinite" length of time in the case a t hand. There is evidence from the time-temperature relations already determined that, as far as conversions are concerned, 48 hours suffice t o obtain essentially equilibrium yields. It is well known that reverse reactions do not always proceed at the same rate as the forward ones; they are slower in some cases and more rapid in others. However, it is reasonable to suppose that at the end of 48 hours there would be adequate indications as to whether or not equilibrium conditions existed.
TABLE
Run
NO.^
2ad 3c
I. BALLM I L L RUNSON
-Compn., GramsBas04 CaClz HzO 2334 1220 1290 1125 467 244 467 244 1125 488 410 934
A PILOT-PL.4ST SCALE
Time, HI. Overnight 4
7 8
All runs had 10 per cent molar excess CaCls. Commercial Barite and Dowflake used. C c. P. chemicals used. d First 4 hours same as run 2. 0
Bas01 Final Converted t o Temp.,' C. BaCI1, 7% 260 85 105 83 310 90 300 92
APRIL, 1940
INDUSTRIAL AND ENGlNEERIKG CHEMISTRY
gardless of how it is carried out. This equilibrium is achieved immediately upon fusing the reactants and after several hours in aqueous solutions.
REVERSION
"A
n
n
571
L
Extraction of Barium Chloride Figures 7 and 8 correlate the effects of aqueous methanol as a solvent for barium chloride and on reversion of the reaction
W
b~ 1 j * T
I
12
I
REVERSION
I
1
,CONVERSION I 24 36
mass during extraction. The solubility effect of the increased water was marked. This increase had but little influence on reversion if the reaction had been carried out, by fusion. The effect on the evaporated ball-mill reaction mass was pronounced and thus rendered the solvent unsuitable for use in this case.
I
48
TIME IN HOURS
IS AQUEOES FIGURE 4. APPROACHTO EQUILIBRIUM SOLUTION
CACL. : bso4=l.l:I lo
Much difficulty was experienced in getting check results for the reversions studied in aqueous media. The first runs were made with calcium sulfate that had been thoroughly dehydrated over a direct flame. Practically no reversion occurred unless much water was present, and even then reversion was less than was expected. Calcium sulfate dihydrate was substituted. Reversion occurred and the results of these experiments indicated equilibrium, but no smooth curve could be drawn through the results when they were plotted. Frequently more reversion occurred for a short time of reaction than for a long time, and it was almost impossible to duplicate the results with the same degree of accuracy t h a t was obtained for the conversion studies. Finally, calcium sulfate which had been carefully dehydrated below 200 O C. was used; the results are shown in Figure 4. The discontinuity in the reversion curve of Figure 4 was caused by the fact that the reaction of barium chloride with calcium sulfate to give barium sulfate was more pronounced in cold solutions than hot, for reasons soon to be given. Undoubtedly, this reaction proceeded to a point beyond the equilibrium conditions that exist a t 175" C. soon after the tube was put in the oven. As the temperature increased, the barium sulfate reacted with the calcium chloride; i. e., the reaction actually became a conversion.
Reaction i n the Fused State Figure 5 and Table I1 show the results obtained by fusing the reactants. Time and temperature had practically no effect on conversion, and a slight excess of calcium chloride was all that was needed for Iarge conversion. Reversion was also studied under conditions of fusion. Since the resulting graphs are almost identical straight lines, they are plotted separately in Figure 6. The results lead to the conclusion that the reaction BaS04 CaClz e BaCll CaS04 has a true equilibrium state, re-
+
+
100
Effect of Time: l o 7 Molar Excess CaCIr: 1.6000' F. i
.
-___.
150
175
2OO
FIGURE 5 . EFFECT OF VARIABLESo\i CONVERSION DCRING FUSIOX The differences in behavior of the two reaction masses toward aqueous methanol during extraction can be traced to two different forms of anhydrous calcium sulfate and possibly also to the fineness of subdivision of the reaction mass. Calcium sulfate is known to exist in two anhydrous forms, anhydrite and soluble anhydrite (2, 4). The former is the natural mineral anhydrite and is also the result of heating gypsum or plaster of Paris to a red heat. This calcium sulfate scarcely reacts with water a t all. Accelerators are always added to this product in order to permit it to set-i. e., hydrate-in a reasonable length of time as in the production of Keene's cement. I t crystallizes in the rhombic system. Soluble anhydrite is supposedly formed when the dihydrate or
TABLE 11. EFFECT OF FUSIOS ON REACTIOS BaSOl Effect of Temp.: lOy0 BIolar Excess CaC1:: 30 Min.
125
+ CaCll e BaCI? + Ca30r
EffecG of CaC12:
l6OO0 F.: 30 hlin.
% Bas04 % Bas04 % molar Tzmp., converted to excess converted to Time BaCh F. BaCir CaClz T o clear fusion" 90.6 1400b 92.5 0 1500 92.8 30 min. 93.0 10 1600 93.0 1 hr. 91.0 25 1700 93.5 2 hr. 91.3 60 1800 92.5 4 hr. 91.0 100 a Called "zero" time; actually requires about 5 minutes. Sintered b u t not completely fused; all others are clear fusions,
Reversion: Molar Rnrio of Reactants, BaCh:CaSOl:CaCl? = 1:l:O.l: 1600' F. d _
% Bas04 oonverted to BaClz 82.9 93.0 97.6 98.0 98.2
Time T o clear fusiona 30 min. 1 hr. 2 hr. 4 hr.
100% BaClz reverted t o Bas04 91.8 91.6 91.8 91.0 92.3
INDUSTRIAL AND ENGINEERING CHEMISTRY
572
hemihydrate is heated below 200" C. This takes up water quickly. It has been reported to exist as triclinic needles. This has been denied, and it has even been suggested that the supposed modifications of anhydrite owe their differences largely to a difference in the fineness of their particles. Soluble anhydrite appears to be a metastable form. There are possible cases of supersaturated forms, the degree of each depending on the size of the particles. The interrelations of gypsum and of soluble and insoluble anhydrite account for the fact that soluble anhydrite is not formed in nature. It is unstable with respect to the other forms (2, 4 ) . ? Ullmann (9)states that complete transformation to the insoluble form of anhydrite takes place a t 600" C. Mellor (4)reports that "B 1 1 w E > z TIME IN HOURS when monoclinic gypsum, the usual type, is heated, i t is probably converted into a second or rhombic form and that this happens before dehydration occurs. If this is true, all the products TIME IN HOURS of the dehydration of gypsum (hemihydrate, FIGURE 6. APPROACHTO EQUIsoluble anhydrite, and LIBRIUM BY FUSION ordinary anhydrite) have rhombic symmetry. Gypsum may be changed to soluble anhydrite a t 93" C. in the presence of water. However, the general conclusions are that gypsum and natural anhydrite are the only forms of calcium sulfate which are stable in the presence of any solution. This does not preclude the presence of the metastable form as a n intermediate which may persist for some time ('J 4)'
-
PER CENT H O ,
FIGURE 7.
\
VOL. 32, NO. 4
tures and concentrations employed are not available. However, the data which are available, taken in conjunction with those obtained in these experiments, make it possible to suggest a mechanism. First, the reaction is probably ionic in character. It took place readily in the fused state, occurred in aqueous solution, but did not occur below the fusion temperature if all water v a s excluded. I n dilute aqueous solution one may write the qualitative statement
,
I N AQUEOUS METHANOL
EFFECTOF WATERIN AQUEOUSMETHANOL ON REVERSION DURING EXTRACTION
Keeping these facts in mind, it seems likely that under conditions of fusion a form of anhydrous calcium sulfate is obtained which is very resistant t o the effects of water, while in the ball-mill reactor conditions are favorable not only for the formation of soluble anhydrite but also for a finely subdivided reaction mass.
Theoretical Considerations
It was not within the scope of this investigation t o determine rigorously the mechanism of the reaction. Thermodynamic data such as activity coefficients a t the high tempera-
if both sulfates are assumed sufficiently insoluble to have true solubility product constants (K5.p.). From the values in the literature ( I ) for such constants, a value of K = 2 X is obtained which shows that at equilibrium very few barium ions will be present. This merely means that in dilute solutions the reaction is essentially quantitative to the left. The following data are indicative of the effects which may be anticipated in concentrated aqueous solutions a t high temperatures: 1. The solubility of calcium sulfate decreases with rise in temperatull. For example, a t 100" C., 0.850 gram anhydrite is dissolved per liter of saturated solution; at 200" C. this value is 0.075 ( 5 ) . 2 . The solubility of barium sulfate increases with rise in temperatures. For example, at 10' C., 0.002 gram barium sulfate is dissolved per liter saturated solution; at 100" C. this value is 0.004 ( 5 ) .
3. The presence of calcium chloride materially reduces the solubility of calcium sulfate. For example, at 25" C. a saturated solution of calcium sulfate contains 2.05 grams calcium sulfate per liter if no calcium chloride is present, 1.02 grams if 51.33 grams are present, and only 0.03 gram if 367.85 grams of calcium chloride are present (5). 4. Barium chloride is virtually insoluble in concentrated calcium chloride solutions. 5. Under conditions of these experiments (175" C.) highly concentrated solutions of calcium chloride are obtained. At 175" C. approximately 75 grams of calcium chloride dissolve per 100 grams of solution (3). 6. The activity of calcium chloride in high concentrations increases tremendously with increase in concentration. The activity coefficient of calcium chloride at these temperatures is not known but would be high. For example, at 25' C. the activity coefficient of calcium chloride is 0.519 if the molarity of the solution is 0.5 and 3.41 if the molarity is 3.0 ( 3 ) . Under the conditions of experiment the molality (moles per 1000 grams water) of a saturated solution of calcium chloride is approximately 27. Values of the activity coefficient have not been determined at this concentration or at this elevated temperature, but it may be inferred from data on less concentrated solutions at low temperatures that the activity coefficient of calcium chloride under the conditions employed must probably be large. Hence, the beneficial effect of temperature.
I n light of the above facts i t is evident that under conditions of high concentration of calcium chloride and a t high temperatures solubility relations must favor the right side of the equation. Under these conditions, two insoluble substances barium chloride and calcium sulfate, are formed. The driving force of the reaction is therefore t o the right. Temperature is believed to affect the equilibrium in so far as i t affects the concentration, solubility, and activity of the salts in question. Otherwise its role is that of a rate increaser (6)* The tremendous change in activity of the calcium chloride with change in concentration explains the marked effect of water content observed. When saturation is reached or approached, no further effect should be expected. This is qualitatively proved by the conversion us. water curves. Too little water may cause difficulty from two sources. The first is the
APRIL, 1940
ISDUSTRIAL AND ENGINEERIKG CHEhlISTRY
mechanical difficulty of getting thorough mixing, the second is the absence of appreciable “solutionJ’ water. Theoretically if the ratio of calcium chloride to water is kept constant, as the ratio of calcium chloride to barium sulfate increases, a maximum and constant value for the yield should be obtained. The evidence presented here shoms this to be qualitatively true although there seems to be a tendency foraslight decrease in the yield of barium chloride upon increasing the calcium chloride-barium sulfate ration. This is possibly due to the errors inherent in the experiments.
Commercial Possibilities for Making Barium Chloride Although i t would be possible to carry out this reaction on a commercial scale using stirred autoclaves in place of sealed tubes, such a procedure would necessitate employing a considerable excess of calcium chloride in order to get practical yields. On the basis of data presented here, it is possible to suggest two different commercial methods.
PERCENT WATER BY VOLUME I N AQUEOUS METHANOL
FIGURE 8. EFFECTOF WATERON THE SOLUBILITY OF BARIUM CHLORIDE IN AQUEOVShIETH.4SOL AT 30”
c.
BALL-MILLPROCESS. A saturated solution of calcium chloride (10 per cent molar excess) is added to a heated ball mill containing finely ground barite. The temperature is kept low until the solution and barite are thoroughly mixed. Then the temperature is raised slowly, which causes t4e water to evaporate so that a t the end of 6-12 hours the mass is completely dry. The ball mill is cooled, and glycol-methanol d i - e n t added. The use of aqueous methanol is prohibited because of the large reversion which would result from the solubilized form of anhydrous calcium sulfate which is present. Extraction proceeds in the mill for about 3 hours, after which the suspension is fed to filter presses. Here the calcium sulfate produced in the reaction and the barium sulfate left unreacted are removed. The chlorides of barium and calcium are retained in solution. This solution is distilled under atmospheric pressure to recover the methanol and then under vacuum or by steam to obtain the glycol. The residue crystals are dissolved in water, and from this solution BaC12.2H20 is crystallized. All aqueous washes and liquors are returned to the reaction step to avoid loss of dissolved barium chloride.
573
This process gives a yield of approximately 90 per cent barium chloride based on the barium sulfate content of the barite. X modification of this process is t o heat the dried reaction mass above 600” C., say in a rotating kiln, to insolublize the calcium sulfate. [Ullmann (9) specifies 600” C. as the temperature necessary to convert the calcium sulfate to the insoluble form of anhydrite.] Aqueous methanol might then be used as the extraction process. Although this is referred to as the ball-mill process, presumably any type of stirred heated vessel might be employed. For example, one might use large kettles, heated by gas and equipped with efficient stirrers, such as are utilized in the manufacture of gypsum. The reactants, in the form of a paste or sludge, could be heated to dryness in such a vessel; or if considerable calcium chloride was present, it would be necessary to heat only to a thick consistency. Organic solvents n ould then be needed to perform the extraction. FUSION PROCESS.I n this process the reactants, barite and 10 per cent molar excess calcium chloride, are fused a t 900-1000° C. in appropriate vessels. As soon as i t is liquid, the fused mass is poured into shallow pans to solidify. The solid is broken u p and ground fine. Dilute aqueous methanol can serve as an extraction solvent. This step is followed by filtration to remove the calcium and barium sulfates. The methanol is easily recovered by distillation. KO crystallization occurs a t this point because the chlorides are more soluble in water than in the aqueous methanol. The water solution of the chlorides is then subjected to crystallization in open pans to obtain the desired product. This process gives a net yield (conversion minus reversion and mechanical losses) of about 83 per cent based on the barium sulfate content of the barite. Both suggested processes have their advantages and disadvantages. The fusion process permits the use of a cheaper solvent which is also more readily recovered. The reaction is rapid as compared to that occurring in the ball mill. But the cost of maintaining the high temperatures required, the corrosion of equipment by the fused chlorides, and the reduction in net yield of product due to reversion during the extraction step may offset the advantages of this method. Only semiplant operation can decide which process is the most economical. This phase of the problem is being further investigated.
Literature Cited (1) Curtman, L. J., “Qualitative Chemical Analysis”, appendix, New Y o r k , MacMillan Co., 1933. (2) Friend, J. N., “Text-Book of Inorganic Chemistry”, Vol. 111, P a r t 1, p p , 69-73, London, Griffin and Co., 1925. (3) Landolt-Bornstein, “Physikalisch-chemische Tabellen”, Band I, p. 642 (1923); Erganzungsband I I b , p . 1114 (1931); Berlin, Julius Springer. (4) Mellor, J. W., “Comprehensive Treatise on Inorganic and Theoretical Chemistry”, Vol. 111, pp. 760-77, London, Longmans, Green and Co., 1923. (5) Seidell, A , . “Solubilities of Inorganic and Organic Compounds”, 2nd ed., Vol. 11, pp. 120, 196, 215, 1147, h-ew York, D. Van Nostrand Co., 1928. (6) Shreve, R. K.,and Pritchard, I T . N., IXD.E m . CHEM.,27, 1488 (1935). (7) Shreve, R. 3.. Pritchard, W. S . , and Watkins, C. H., Purdue Univ., Eng. Bull. 22, KO.l a , Feb., 193s. ( 8 ) Shreve, R. K., and Watkins, C. H., IKD. ENG.CHEY.,31, 1173 (1939) (9) Ullmann, F., Enzyklopsdie der technischen Chemie, 2nd e d . , Vol. 111, p. 56, Berlin, Urban & Schwaraenberg, 1929. PRESEITED before the Division of Industrial and Engineering Chemistry a t the 98th Meeting of the Amerioan Chemical Society, Boston, Mass. Based upon a thesis submitted by R. K. Toner to t h e faoulty of Purdue University in partial fulfillment of the requirements for the degree of doctor of philosophy.
