SOLUBILITY O F CALCIUM S U L P H A T E I N AQUEOUS SOLUTIONS O F SULPHURIC ACID I
BY E'. K. CAMERON AND J. F. BREAZEALE
In the course of an investigation now in progress, we have had occasion to study certain solutions containing various amounts of sulphuric acid and saturated with respect to calcium sulphate. Some apparent abnormalities which appeared in this work led us to believe that the solubility of calcium sulphate in solutions of the corresponding acid was not in accordance with prevailing views as to the effect of one electrolyte upon another when both are supposed to yield a common ion; and for the ;proper iiiterpretation of the data we were obtaining, a more precise knowledge of the solution curve for this pair of electrolytes would be necessary. Ostwald" has published some results of work by Bantisch on the solubility of calcium sulphate in solutions of hydrochloric, nitric, monochloracetic and formic acids, respectively, from which it appears that, qualitatively at least, the indications of the electrolytic dissociation hypothesis are followed, No systematic study, however, seems to have been made of the solubility in sulphuric acid solutions, and it was therefore determined to make the investigation, the results of which are here recorded. The calcium sulphate used was prepared by triturating very pure precipitated calcium carbonate with successive portions of pure sulphuric acid until no trace of carbonic acid remained, and then washing the product free of sulphuric acid with distilled water. The product thus obtained was carefully examined, found to be a very pure salt, and was most satisfactory for solubility work, as it was in a finely precipitated form, offering a large surface for the action .of the solvent, and shortening much 1 2
Published by permission of the Secretary of Agriculture. Jour. prakt. Chem. 29, 52 (1884).
572
R K Cameron ana‘]. R Breazeale
the time required to produce conditions of definite equilibrium between solvent and solute. T h e sulphuric acid used was obtained from Baker & Adanison and labeled ‘‘ chemically pure.” It was taken from a bottle, freshly opened for the purpose, and appeared to be entirely satisfactory as regards its freedom from other substances. A series of solutions of sulphuric acid of the concentrations indicated in the table was prepared, and an excess of the dried calcium sulphate introduced into each in the proportion of one gram of the salt to every IOO cc of the solution. T h e solutions were then immersed in a bath kept at a temperature of 25” C, the containing bottles being shaken occasionally. Portions of these solutions were withdrawn from time to time and analyzed for calcium, until successive analyses showed that final equilibrium had been established in all cases. This was effected in about 70-80 hours, but the solutions were retained in the constant temperature bath for six days to establish final equilibrium beyond any possible question. Portions of 50 cc were then withdrawn and analyzed. The calcium was precipitated from a hot solution by adding ammonia to excess and then ammonium oxalate. T h e precipitate was carefully washed, dried, ignited, and the calcium weighed as oxide. T h e filtrate and washings, containing considerable quantities of ammonium sulphate and oxalate, were in all cases brought to dryness in porcelain dishes, which were then carefully ignited to remove ammonium salts, the residue taken up with small quantities of hydrochloric acid, washed into cylindrical bottles and the small quantities of calcium present estimated by a modification of the well-known ‘‘soap” method originally suggested by Clarke. T h e small amounts of calcium found were added as corrections to the amounts determined gravimetrically. By this procedure the error due to the solubility of the precipitated calcium oxalate in rather concentrated solution of ammonium sulphate was avbided, since the modification of the soap method in use in this laboratory has been found to be accurate to about one part of calcium per million of solution.
SolubiliLy of Galciurn Su@hate
573
The amounts of free sulphuric acid present were checked by titration against a standard solution of potassium hydrate, using both phenolphthaleine and methyl orange as indicators, before and after the introduction of the solid calcium sulphate into the solutions of the acid. After making the determinations for 25" C, the flasks were brought to 43" C, being held at this temperature for six days before being analyzed. They were then brought to 2.5" C again and held at this temperature for several days and again analyzed, yielding results in such good accord with those formerly obtained at this temperature, that it is not thought worth while to state them here. A few determinations were made at 35" C, in which, however, the calcium sulphate used was finely powdered gypsum from a sample employed in some earlier solubility work already described in this Journal. The weights of solutions at 2.5" C were carefully determined, the volume of the pycnometer used being standardized by determining the weight of boiled, air-free, distilled water it would contain and finding the corresponding volume from Gray's tables. These data will enable any one who may desire it, to compute the solubilities on the basis of the mass of solvent ; rather than the volume of solution. But conceding the latter form to be t h e less rational, we have retained it here in order to show the comparison of the results obtained at the different temperatures. There does not appear to be any particular advantage in stating the results in terms of reacting weights, and they are therefore only given as grams per liter of solution. The significance of the results is more readily seen when they are plotted to sonie convenient scale, as in the accompanying chart. It will be observed that increasing the temperature increases the solubility of the calcium sulphate in the presence of the sulphuric acid over the range covered by our experiments, and the maximum point on the solubility curve for calcium sulphate in pure water at various temperatures' disappears at least 37.5O C according toMarignac. Ann. Chim. Phys. (5) I, 274 (1874), and 40' C according to the recent work of Hulett and Allen. Jour. Am, Chem. SOC.24, 667 ( I~OZ),
F. K. Cameron and]. P. Breazeale
574
within the range'here considered. In this respect, the system seems to be quite different from those composed of calcium sulphate and solutions of other electrolytes, such as sodium or magnesium chloride, etc.,' formerly studied in this laboratory. Solubility of Calci&n Sulphate in Solutions of Sulphuric Acid Grams CaSO, per liter at Weight of 1000cc of >rams H,SO, per liter solution at 2 5 O C Grams Grams
999.1067 X002.493 1002.553 1005.og I I009 * 787 1030.I 5 I 1 043.470 1075.6I 3 1113.392 II4I.755 1168.143
-
2.126 2.128 2.144 2.203 2.382 2.727 2.841 2,779 2.571 2.313 1.901 1.541
0.00
0.48 4.87 8.11
16.22 48.67 75.00 97.35 146.01 194.70 243.35 292.02
Fig.
43 Grams
35 Grams
25
Grams
2. I45 2.236 2.456 2.760 3. I 16 3.843 4.146
2.209 2.451
-
3.397
-
-
3.606 3.150
4,139 3.551 2.959 2.481
-
I
At the three temperatures studied, the curves show maxima points, which by interpolation appear to correspond to the following concentrations : At 25' C, 75 gms. H,SO, and 2.84 grns. CaSO, per liter. "
"
35O ( I 85 430 ( ' 105
( I I i
" 'I
'I ( I
3-70 4.26
I' ( (
.
' I
' I
I I
i(
I (
ii
Jour. Phys. Chem. 5, 556, 643 (1901). Bulletin No. 18,Division of Soils, U. S. Department of Agriculture, 1901.
SolzibiZity of Calcium Sudphate
575
A remarkable feature presented by this system is that the solubility of the calcium sulphate increases with increasing amounts of sulphuric acid, up to comparatively high concentrations with respect to the latter. We seem to have here another of the increasing number of known cases where the indications of the hypothesis of electrolytic dissociation in aqueous solutions do not hold. It is difficult to see how calcium sulphate could dissociate any way other than that indicated thus :
CaSO,
Ca
+ SO, ;
but on the basis of conductivity, freezing-point and similar measurements, we are accustomed to regard sulphuric acid as dissociating thus : H,SO, '*", H HSO,,
+
and, then a further dissociation of one of the ions, thus : HSO,
WV-b
.l-e
H
4-so,,
and that at lower concentrations this latter dissociation is approximately complete. This would suggest in the system we are considering, the presence of two electrolytes which yield a cotnnion ion, and that we should obtain a decrease in the solubility of the calcium sulphate with increasing amounts of sulphuric acid. That the reverse was actually found at all concentrations suggests several hypotheses : I. Sulphuric acid, possibly influenced by the presence of dissolved calcium sulphate, may dissociate with the formation mainly of HS04 ions. Opposed to this view is the work to which reference has just been made. 2. Possibly the calcium sulphate, in dissolving, more or less completely unites with the sulphuric acid to form complexes. yieldingsuch ions, for instance, as Ca(SO,), or HCa(S04)2. No evidence in support of this view could be obtained from a critical examination of the solid phases and the determination of the acidity of the solutions before and after adding the calcium sulphate, yielding practically identical results, failed to give any in. formation on this point. 3. Possibly the two dissolved substances, or the substances
K. Cameron and]. R Breazeale
576
and the water may have other effects, one upon another of such magnitude as to mask those normally to be expected from the ions which they might yield. There does not seem to be any known reason why even strong electrolytes niight not exert such effects which are well known though not well understood in the case of other and especially not dissociating solutes. No satisfactory methods of testing these several hypotheses at present suggest themselves, and it only seems worth while to bring these views forward to show that while the facts established by this investigation are not in harmony with the indications of the dissociation hypothesis as we know it, it would be going too far to say that they are in actual opposition to the hypothesis, and to emphasize the desirability of further study along this line. Another point brought out by this work as in our former studies' is that there seems to be a condensation, not alone of the solution, but of the solvent itself, brought about by the introduction of these solutes. For instance, a liter of solution at 2 5 O C and containing 0.49 gram of sulphuric acid with 2.128 grams of calcium sulphate weighs 1002.493 grams, so that the weight of the water present is actually 999.875 grams, while the weight of an equal volume of pure water at the same temperature is given as 997.12 grams2 Finally, it seems worth while to call attention to the figures given for the solubility in pure water. T h e figure corresponding to 2 5 O C especially, was determined with unusual care and is the average of eight determinations, six of which were practically identical. This matter is of some moment in view of the careful work on the solubility of gypsum recently done by Hulett and Allen.3 T h e figures given by these authors represent what they term a mwmaZZy saturated solution, that is, a solution in equilibrium with a solid presenting only a plane surface to solvent action, whereas Hulett4 has shown that by secur-
4
Loc. cit. Gray. Smithsonian Physical Tables, 1896. Jour. Am,Chem. SOC.24,667 (1902). Zeit. phys. Chem. 37,388 (1901).Seealso Ostwald. Ibid. 34,495(19.0).
SoZulliZity of Calcium Su@hafe
577
ing sufficiently convex surfaces, through powdering or otherwise reducing the size of the particles, he could obtain an apparent increase in the solubility, of about 20 percent. But the solubility of calcium sulphate or gypsum with which one most commonly has to deal, and probably is most interesting on that account, is not that which would be obtained either by bringing the solution into equilibrium with plates of selenite or gypsum free from powder or with particles less than 0.3 micron in diameter. And the fact that this figure 2. I 26 has been obtained so often in this laboratory,' working with calcium sulphate of widely different origin, leads us to think that it represents with considerable accuracy the solubility, under ordinary conditions, to be expected at 25' C. Hulett and Allen's suggestion that former investigators have had present varying amounts of material consisting of particles less than 0.3 micron in diameter, does not seem very probable, since no matter what the physical character of the calcium sulphate, whether precipitated from solution or powdered mineral, we have obtained such concordant figures. Bureau of Soils, U. S. Department of Agriculture, Washington, D. C. LOC.cit.
x
