SOLUBILITY OF AMMONIUM PERSULFATE
*
645,
results are another proof that although iron salts are absolutely necessary for the studied reactions, their concentration is only of secondary importance. SUMMARY
The catalytic oxidation of formic acid by natural mineral waters and hydrogen peroxide was studied. The oxidation is started by the ferrous ions present in the mineral water and is continued by a chain mechanism. The main factors influencing this chain mechanism are the concentration of hydrogen peroxide and its dissociation as indicated by the pH of the solution. The addition of cupric salts greatly increases the rate of the reaction. REFERENCES (1) (2) (3) (4) (5)
CRONHEIM, G . : J. Phys. Chem. 46, 328 (1941). KRAUSE,A., AND KANIOWSKA, D.: Ber. 6% 1982 (1936). KUHN,R . , AND WASSERMANN, A.: Ann. 609, 203 (1933). SIMON, A., AND HAUFE,W.: 2.anorg. Chem. 230,148,160 (1936). WEISR,J.: J. Phys. Chem. 41, 1107 (1937).
SOLUBILITY OF AMMONIUM PERSULFATE IN WATER AND IN SOLUTIONS OF SULFURIC ACID AND AMMONIUM SULFATE J . F . GALL, G. L. CHURCH, AND RALPH L. BROWN Research Division, Pennsylvania Salt Manufacturing Company, Philadelphia, Pennsylvania Received July 66, 194.9 INTRODUCTION
In spite of the growing importance of per compounds, much of the physical data relative to even the inorganic members is lacking. Thus we have found only two references to the solubility of ammonium persulfate (1, 2), and these report results of single measurements only. We have had occasion to measure the solubility of this salt in water and in solutions of sulfuric acid and of ammonium sulfate; the results obtained are presented in this paper. EXPERIMENTAL
Solutions of sulfuric acid and ammonium sulfate were prepared from C.P. reagents. Ammonium persulfate of the same quality was finely pulverized and added in considerable excess to the solutions in brown-glass rubber-stoppered bottles. Continuous agitation at constant temperature was obtained by tumbling in a water thermostat held within &0.5OC. of the desired temperaturs. In most cases two saturation bottles were used at each composition, and samples were withdrawn a t intervals of 1 hr. or 4 hr. Analysis was made for persulfate,
646
J. F. GALL, G. L. CHURCH, A N D RALPH L. BROWN
Caro’s acid, and hydrogen peroxide.’ The densities of the solutions before an& during saturation with persulfate were determined by means of the Westphal density balance. EQUILIBRIUM
Solubility measurements on persulfates are complicated by hydrolysis t o permonosulfuric (Caro’s) acid and to hydrogen peroxide :
+
+
(NHa)zSzOs €LO -+ HzS06 (NHa)zSOd ILS06 H20 + HzSO4 H202 and by decomposition of the hydrogen peroxide: HzOz -+ HzO -k $ 0 2 First attempts to approach saturation equilibrium showed a continued rise, lasting many days, in concentration of oxidizing value, which Tvas more rapid in the solutions of higher sulfuric acid content. Closer examination showed that saturation equilibrium with respect to ammonium persulfate was actually achieved n-ithin a few hours’ time, but that slow hydrolysis to Caro’s acid, catalyzed by higher acid concentrations, caused the slom rise in oxidizing value. Figure 1 shows typical variation of composition and density with time for one of the compositions studied. As will be seen, readings from duplicate sample bottles were in good agreement. Saturation vith respect to ammonium persulfate was established within an hour. The total oxidizing value rose slowly and continuously, oiving to development of Caro’s acid by hydrolysis of persulfate. The persulfate content diminished slowly and continuously, owing to decreasing solubility of ammonium persulfate as Caro’s acid and ammonium sulfate concentrations increased through hydrolysis of the persulfate. No hydrogen peroxide as detected in any of the solutions. The density increased slowly with time, oiving to the increasing total concentration of substances in solution. At each composition the curves for persulfate concentration and total oxidizing value (each expressed as peroxide equivalent) in all cases were found to extrapolate to the same point at zero time. This was taken as a further indication that true saturation equilibrium mas reached. The final values for the solubility of ammonium persulfate were therefore taken as the common point of extrapolation of these two curves. Similarly, the density of the saturated solution was taken as the value found by extrapolation of density versua time to zero time. Analysis of the solid phase was not made; the equilibrium form was assumed to be the anhydrous salt, (NIL)zSzO2.
+
+
1 Unpublished procedure (developed by F. J. Frere and employed in this laboratory): Hydrogen peroxide is titrated directly with ceric sulfate, using orthophenanthroline ferrous coniplex indicator. Caro’s acid is determined in the solution after peroxide determination, by reduction with vansdyl sulfate and back-titration with permanganate. The total available oxygen is dcterrnined by iodimetric titration on a separate sample. Persulfate is calculated by difference. 2 The existence of a dihydrate, ( X H 4 ) & 0 8 . 2 H 2 0 ,in equilibrium with saturated solution below 36°C. has been reuorted (Heidt: J. Chem. Pilys. 10,298 (1942)). We have been unable to prepare crystals by evaporation from the saturated solution a t 7”, 25’, or 35°C. which did
647
SOLUBILITY OF AMMONIUM PERSULFATE RESULTS
Results of the solubility and density measurements are given in table 1, which shows compositions and densities before and after saturation with ammonium persulfate. Concentrations of sulfuric acid and ammonium sulfate in the saturated solutions were not determined analytically, but were calculated by the expression
22
18
-=.= e-
14
1.34
F% 0-0
DeneitP
o-o-o-
where A and A' are concentrations of sulfuric acid or ammonium sulfate before and after, respectively, saturation with ammonium persulfate, dl and dz the densities before saturation and after saturation, respectively, and S is the concentration in grams per liter of ammonium persulfate in the saturated solution. not assay above 98 per cent (NHI)zSZO~. X-ray diffraction patterns of crystals so prepared were identical with patterns of the anhydrous salt. We thus fail t o confirm the reported existence of a dihydrate.
TABLE 1 Density and composition of HzSO,-(NH&SOI miztures saturated with (NH&SZO~ OPIGIXAL COMPOSITION
HdOd
1
(NH4rSOd
I
INITIAL DENSITY
~
FINAL DENSITY
FINAL CO1(POSITION ~
-__ BSO,
1
I
("4dO'
(N&)dlOa
15°C. grams per liler
0
0 0
gram per lifer 1 grams per cm.: gram2 per cm.8
grams
per liter
0 50 100
0 I999 1.028 1.056
1.254 1.254 1.254
0
0
1.258 1.260 1.263
166.9 169.3 171.6
1.304 1.311
362.7 371.4 376.3
200 200 200
100
1.125 1.147 1.169
400 400 400
0 50 100
1.242 1.260 1.279
50
0 0
grams per liter
gram$ per Icier
0 37.5 76.2
522 483 449
0
319
0 37.0 75.1
543 505 474
20°C. 0 0 0 200
200 200 400 400 400
50 100
0
0.998 1.027 1.055
1.266 1.266
0 50 100
1.123 1.145 1.166
1.267 1.268 1.269
161.1 163.2 164.8
0 40.8 82.4
362 334 308
0
1.238 1.257 1.275
1.302 1.310 1.320
358.3 366.4 375.9
0 45.8 93.9
193 159 122
50
100
TABLE 2
SoZubiEity* of (NH4)dhOs i n solutions of HzSOIand (NH4)2SO~ CONCEhTRA'IION
3W
15°C. 0 25 50 75 100
334
428
~
369
288
1
I I 350
4W
I
247
318
20°C. 0 25 50 75 100
542 518 496 475 452
481
2:;
421 401
1
426
1
371 352
:
374 356 340 324 307
,
326 310 295 278 263
281 265 250 235 219
239 224 208 194 178
199 183 169 155 139
162 146 131 116 100
* All concentrations are expressed as grams of constituent per liter of solution at the temperature indicated. 648
SOLUBILITY OF AMMONIUM PERSULFATE
649
The solubility values were obtained for odd values of ammonium sulfate and sulfuric acid concentrations, so that for convenient use of the data a double interpolation to rounded values of these concentrations is necessary. The following graphical procedure was developed: On a plot of ammonium sulfate concentration vs. sulfuric acid concentration, the compositions studied will be represented by an array of points. Normals, erected from each such point and proportional in length to the solubility of ammonium persulfate in the corresponding composition, will generate a surface (which need not be constructed) representing the solubility of ammonium persulfate in all compositions of sulfuric acid and ammonium sulfate within the range of compositions investigated. Lines drawn in any convenient way joining measured points on the solubility surface will be projected onto the ammonium sulfate-sulfuric acid plane as lines joining the corresponding compositions. Solubilities may now be plotted against units measured along the projected connecting lines and interpolated solubility values for rounded concentrations of, e.g., ammonium sulfate, may be read off with reference to the field plot of ammonium sulfate us. sulfuric acid. The so rounded solubility values may now be plotted as a family of curves, for constant ammonium sulfate, against concentration of sulfuric acid. This plot may now be used to obtain solubility values rounded with respect to both ammonium sulfate and sulfuric acid. As a check, the process may be repeated, rounding first with respect to sulfuric acid and then with respect to ammonium sulfate, Table 2 presents the final solubility data for rounded values of ammonium sulfate concentration at intervals of 25 g. per liter, and for rounded values of sulfuric acid concentration at intervals of 50 g. per liter. The solubility is given in grams of (NH4)tSzOs per liter of solution at the temperature indicated. ACCURACY
Inspection of the experimental data leads to an estimate of error in solubility as less than 2 4 per cent; the density data are undoubtedly correct to better than 1per cent. The uncertainty in both quantities arises largely in the extrapolation to zero time. Linear interpolation between the rounded values in table 2 will introduce error no greater than that of the experimental data. SUMMARY
1. A method for accurate determination of solubility in slowly hydrolyBing and decomposing solutions of per compounds was developed. 2. The solubility of ammonium persulfate in water and in solutions of sulfuric acid and ammonium sulfate, at 15°C. and 2OoC., was measured.
Acknowledgment is made to Pennsylvania Salt Manufacturing Company for permission to publish the data of this paper. REFERENCES (1) ELBS: J. prakt. Chem. [2] 48, 185 (1893). (2) MARSHALL: J. Chem. SOC. 69, 771 (1891).
