T H E SOLUBLE BASIC SU1,PHATES OF BERYI,T,IUM BY CHAS. L. PARSONS, W. 0. ROBINSON AND C. T. FULLER
Normal salts of beryllium in aqueous solution have the property of dissolving large amounts of beryllium hydroxid or basic beryllium carbonate. This property is much more marked with this element than with the normal salts of iron, aluminum and zinc, all of which have been shown to dissolve considerable amounts of their own hydroxids. With beryllium the excess of base dissolved varies with the anion of the salt, the temperature, and the concentration. With the sulphate, in concentrated solutions, it is as great as three equivalents of beryllium oxide t o one of sulphur trioxide and even greater in case of the nitrate, chloride, and acetate. When these concentrated solutions are diluted with water, precipitates of a highly basic nature are thrown out. These general reactions were pointed out by one of us,' and the precipitate in the case of the sulphate shown to be the hydroxid containing small variable amounts of occluded qr dissolved normal salt. The same was also found to be true in the case of the oxalate.2 The object of this research was to determine the character of the soluble basic beryllium sulphates. A number of authors have shown that beryllium hydroxid dissolves readily in solutions of the hydroxids of sodium and potassium and it has been assumed that beryllonates are formed. This is confirmed by experiments to be cited later. The fact that the normal and basic sulphates are strongly acid in reaction to indicators, readily set free carbon dioxid from carbonates, and give off hydrogen when treated with zinc, probably indicate a high concentration of hydrogen ions; although Stewart3 has proven that a very small excess of base completely destroys the power to invert cane sugar. Parsons: Jour. Am. Chem. SOC.,26, 1347 (1904). Parsons and Robinson: Ibid., 28, 555 (1906). a Jour. Am. Chem. SOC., 26, 1432 (1904).
a
C. L. Parsom, W. '0.Ro6irzso?z arzd C. T.Fzder
652
Further, Leysl and Brunner,' have shown that the salts of beryllium are less hydrolyzed than the corresponding salts of iron and aluminum. It was hoped that light might be thrown on the nature of solutions of basic beryllium salts by freezing-point, dialysis, transference, and conductivity experiments with the sulphate. Preparation of Material The normal sulphate was prepared in the usual mannerS from a pure oxide and a very pure product obtained by several crystallizations from alcohol and finally from distilled water in platinum. A very pure beryllium hydroxid containing a slight amount of carbon dioxid was made by dissolving in ammonium carbonate and precipitating the basic carbonate with steam. After thorough washing with hot water the basic carbonate was boiled with frequently renewed water and air bubbled continuously through the mass. Another unpublished investigation has shown that the ammonia could not be nearly all removed by washing alone but by the treatment outlined above the hydroxid contained only the merest trace of ammonia and a very little carbon dioxid which was immediately given off when the hydroxid was dissolved in the hot solution of the sulphate. The hydroxid thus obtained contained 5 2 . 5 9 percent BeO. Freezing-point Determinations , Cryoscopic determinations were carried out, using from 5 0 to 60 cc for the concentrated solutions and larger apparatus and 500 cc for the dilute solution measurements. As the normal tetrahydrated sulphate was stable in air, small successive portions were dissolved in a weighed amount of water. The usual precautions were carried out and three determinations made for each point. The data are given in Table I and are graphically represented in Fig. I. ~
Zeit. phys. Chem., 30, 1 1 , 213 (~Sgg). Ibid., 32, 133 (1900). 8 Jour. Am. Chem. Soc., 26, 1434 (1904)
2
SoZzrbZe Basic Szi@haL'es of BeryZZizinz
1 per Grams BeS0,.4H,O loo grams H,O I 2
,
3 4 5 6
0.0000
0.0000
0.6680 1.4950 2.3680 3.0210
0.3018 0.6754 I .0700 I . 3650 I . 8300 3.1170 4.0170 5.1050 7.7470 9.0800 10.09 1 2 * 39 14.27
4.0500
6.9000 8.8960 I I .3000 17.1500
7 8
8 9 IO I1 I2
20.1000
22.33 27.43 31.59
I3 14
Grams
653
Lowering
0.772 1.033 I . 283 2.000
2.390 2.780 3.525 4.285
4
rr
5
15
10
Fig.
I
To determine the effect of the dissolved hydroxid upon the freezing point, a concentrated solution was prepared by dissolving an amount of beryllium hydroxid in 'an analyzed solution of the normal salt so that $he resulting solution contained 12.00 percent BeO, 25.51 percent SO, and 62.49 percent H,O. This solution was added in small successive portions to 5 0 cc of water and the determinations made as with the normal sulphate.
The ratio
Be0
so, = 3 : 2 ~
was chosen
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C. L. Parsons, W 0. Robinsoit and C. T. F2cZZer
because it is the highest basicity that is soluble in all concentrations. The basic salt was weighed out in the form of a concentrated solution already in equilibrium with the atmosphere, because the solid residue was so hygroscopic that i t could not be weighed. The results are given in Table I1 and plotted in Fig. I as curve ( 2 ) .
TABLEI1 Initial weight of water = 49.842 grams Basic solution contains 25.51 percent SO,, 1 2 . 0 0 percent B e 0 and 62.49 percent H,O
1
Wt. sol.
1
Wt. SO,
I . 132
6 7 8 9
8 ..6690 11.1685 14.2060 16.5955
1 * 593 2.212
2.847 3.624 4.234
1,
Wt. H,O
Total H,O
2.768 3.904 5.544 6.977 8.880 10.37
50.34 50 * 87 51174 52.61 54.75 55.39 56.86 58.12 60.21
>rams SOB ,owering per IOO g.
0.407 0.821 1.496 2.152
2.911 3.992 5.012 6 . I73 7.031
0.118' 0.203 0.329 0.451 0.603 0.787 0.991 I . 228 1,431
It is apparent that, in concentrated solutions at least, the solution of beryllium hydroxid in beryllium sulphate diminishes the osmotic effect and consequently raises the freezing-point. It is also interesting, though of doubtful application, that the curve for the normal sulphate starts out like any dissociated salt, showing a decrease in ionization with the concentration. After a certain point the curve changes its direction and becomes convex toward the axis and exactly coincides with a theoretical curve formed by the tetrahydrate withdrawing water from the solution t o form a decahydrate molecule. The freezing-point lowering in dilute solution was also measured, using 500 grams water and the same basic solution as was employed in Table 11. The results obtained are given in Table 111. Each determination is the mean of from three to five observations.
Soluble Basic S u Q h a t e s of Beryllizriiz
655
TABLE 111 -
1
~
Grams solution
-
-
,
I I 2
I
3 4 5
Lowering
I
0.4028 0.8451 I. 5014 2.1619 2.7852
o 2056 0.4312 0.7658 I . 1028
0.0107~ 0.0210
0.0313 0.0423
The results,are plotted in Fig. 2, curve ( 2 ) , together with the curve for the normal salt taken under the same conditions. Here, again, it is plain t h a t there is a distinct rise in the freezing-point due to the dissolved hydroxid.
Fig.
2
The results from the freezing-point determinations alone would seem to indicate that (I) there is an actual compound formed between the normal sulphate and the hydroxid giving a less number of effective osmotic parts; or (2) there is in some way an aggregation of the dissolved molecules; or (3) that the excess of hydroxid is colloidal in its nature and interferes with the free movement of the crystalloid particles.
Dialysis Trials t o precipitate the soluble basic sulphate by electrolytes gave negative results. Dialysis through parchment tubes and gold beaters ’ skin, although unsatisfactory from a quantitative standpoint, showed conclusively that the membranes tended to separate the excess of base from the normal salt when the basic sulphate was dialyzed, though large amounts of the basic material came through.
-
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C. L. Pamom, W. 0. Robinson and C. T.FzrZZer
The qualitative results entirely confirmed the conclusion drawn from the experiments with the’ electrolytes, that we were not dealing with a true colloid. Two characteristic trials may be cited to advantage. (I) A solution of basic beryllium sulphate, Be0 : SO, = 3 : I , was put inside a carefully tested membrane and dialyzed in a beaker of pure water. The dialyzed solution varied in composition from 1.66 B e 0 : I SO, to 1.81 Be0 : I SO,, and a considerable portion of Be(OH), was found on the inside of the tube. (2) Another solution Be0 : SO, = I .5 : I was dialyzed through gold beaters ’ skin. Analyzed portions of the outside liquid after twentyfour to forty-eight hours showed ratios varying from 1.rBe0 : I SO, to 1.2 BeO: I SO, and again Be(OH), was thrown out on the inside of the tube. When normal beryllium sulphate was dialyzed, no such separation of the hydroxid took place. These conditions are perfectly analogous to the results obtained when a solution of camphor in aqueous acetic acid is dialyzed as will be shown in the following paper.
Ion Transference If beryllium sulphate formed a complex ion with beryllium hydroxid, it would be difficult to conceive of its existing in any other form than an anion. If the excess of beryllium Over the amount present in the normal salt does not exist, it can be easily shown by qualitative transference experiments. To see whether complex ions were formed or not, the following experiments were carried out in the usual form of apparatus used for that purpose, the cut of which is shown in Fig. 3. A coil of platinuni wire was used as the cathode and a piece of pure silver as the anode. The voltage between the terminals varied between sixty and seventy-five. The resistance of the solution was such that less than one-tenth ampere passed during the first stages of the electrolysis. The wTlole apparatus was immersed in a thermostat regulated at 25’. It was first entirely filled with half normal sulphuric acid and when the temperatures had been equalized, half nornial sulphuric acid containing normal beryllium sulphate
SoZfiBZe Basic Su@/zntes of Beryllium
657
was drawn into portion B through the ttibe T. The current was passed for four hours. At the end of this time portions EE were carefully removed with a pipette and were found t o contain no beryllium.
-
I
Fig. 3
The experiment .was repeated, using half normal potassium nitrate. In this case as before the beryllium formed no anion. The basic solution, Be0 : SO, --- 3 : 2, was drawn into section B and electrolyzed for four hours in a solution of half normal potassium nitrate. No anion of beryllium was formed. Beryllium sulphate was dissolved in potassiuni hydroxid drawn into section B and electrolyzed in a half normal solution of potassium hydroxid as before. At the end of the experiment beryllium was found in both the anode and cathode portions. These experiments show that the basic sulphate solutions contain no complex anion,of which beryllium is a constituent. Conductivity The conductivity was determined in the usual manner at 25' upon the normal sulphate and upon the basic sulphate solution of the ratio 1.5 BeO: I SO,, with the following results :
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C. L. Paysoas, W. 0. Robinson and C. T. FuZZer
I
I
h-ormal sulphate Normality
N/5o N/IOO
F”
j
67.1 77.2
. Basic solution F”
1
56.9 67.0
These figures show that beryllium hydroxid diminishes the conductivity of normal beryllium sulphate solutions. To Summarize We have shown that dissolving beryllium hydroxid in normal berylliuni sulphate solutions raises its freezing-point and diminishes its conductivity; that the solutions so obtained are not true colloids; and that they contain no beryllium in the anion. When this paper was presented a t the Ithaca meeting of the American Chemical Society in June, 1906, i t was suggested by Dr. Wilder D. Bancroft, in the discussion that followed, that the facts presented were perfectly analogous t o such cases as the solution of water in mixed solvents of benzene and alcohol and that the whole phenomena might be a similar case of simple solution. This is the point of view we now hold, and our reasbns are discussed by one of us under the broader questions of solutions in binary mixed solvents in the following paper. New Hambshire College, June, 1907.
