Precipitation from Homogeneous Solution A Method for Determining Barium ARNO
H. A. HEY"
and EUGENE SCHUPAK*
Boston University, Boston 75, Mass.
M
Add 3 ml. of the 0.534 ethylenediamine tetraacetate solution for each 100 mg. of barium, but in any case not less than 5 ml. Bemuse calcium and other metals form stable complexes, add sufficient ethylenediamine tetraacetate to complex these metals and the barium. For example, 100 mg. of calcium carbonate will require 2 ml. of ethylenediamine tetraacetate solution in addition to the above amount. Dilute to 190 ml. and adjust to pH 10 to 11 (indicator paper) with concentrated ammonium hydroxide or hydrochloric acid. Add 5.0 grams of ammonium peroxydisulfate and stir until only a few small crystals remain undissolved. Place the beaker on a hot plate a t low heat. After 30 minutes, \Then the temperature reaches about 60" C., precipitation takes place, \Yhen the supernatant liquid becomes clear, test the pH. If the pH is above 4 to 5, cool the solution, add 3 grams of ammonium peroxydisulfate, and continue heating. Generally, the solution will turn yellow because of oxidation products, or, at' times, a hrown fluffy precipitate which looks like hydrous ferric oxide will appear but disappears on ignition. Aft,er 1 to 3 hours on the hot plate, decant the solution through a neighed and ignited filtering crucible, and wash several times with hot' water. Dry the crucible a t 110' C. and ignite a t 900" C. Cool, moisten the precipitate with one drop of dilute sulfuric acid, and reignite at 900" C.
ETHODS for increasing the p H for releasing anions by internal reaction n-ere reviewed by Willard (8). This study presents a procedure for releasing cations by complexing them LT ith ethylenediamine tetraacetate ion (EDTA) a t high pH, and dissociation of the complex to yield the desired cation upon lowering the pH by hydrolysis. A second method of breaking up the complex by oxidative destruction was attempted but abandoned. Piibil and llIariEov6 (6) applied the first principle to the purification of barium sulfate precipitates by redissolving them in hot ammoniacal ethylenediamine tetraacetate solution and reprecipitation of the barium sulfate by acidification. They lowered the p H by direct addition of acid. Subsequently, RlacSevin and Dunton (6) developed the approach of oxidative destruction of the complex. The first method was applied to the determination of barium as barium sulfate in the procedure to be described. Because the success of the method depends on finding a material which. on hydrolysis, will release hydrogen ions a t the desired rate, various compounds were investigated: alkyl halides, halohydrins, and peroxydisulfate ("persulfate"). Limited solubility in water, a rate of hydrolysis which was either too slow or too rapid, or a final p H which was not sufficiently low eliminated most of these materials. Perosydisulfate was found to hydrolyze a t a desirable rate. Because the presence of ammonium salts causes no error in the determination of barium, ammonium salts were used entirely. The first appearance of barium sulfate precipitate occurs when the pH is lowered to about 8 by heating the solution containing the barium complex and ammonium peroxydisulfate. The p H changes rapidly in this region and in a few minutes is lowered to about 2 to 3 and the dissociation of the barium rornples is complete.
__ ____ Table I.
Substance Added Substance Mg.
...... ...... ......
Ca-'G Ca++' Ca + + a Caiia Ca ++a
Mg T +a
Fe+'+b Na +a Na +a
Reagents. Ethylenediaminete traacetate solution, 0.5M. Add 14.6 grams of ethylenediaminetetraacetic acid (Commercially available from several manufacturers) to 80 ml. of water, add C.P. concentrated ammonium hydroxide until the ethylenediamine tetraacetate is dissolved, and dilute to 100 ml. Ammonium peroxydisulfate (ammonium persulfate). Xany batches of this material were found to be unsuitable because of excessive hydrolysis in the reagent bottle. To test, add 2 to 3 grams of the reagent to 100 ml. of water a t room temperature and add concentrated ammonium hydroxide Rith an eye dropper until the solution is slightly alkaline against a universal indicator paper. S o more than 4 to 5 drops of the hydroxide should be required. The presence of a ring in the solid due to moisture generally indicates that extensive hydrolysis has taken place. Barium chloride for known samples. The barium chloride dihydrate content of samples from a bottle of reagent grade material was determined by cation exchange of barium for hydrogen followed by titration of the acid, precipitation of the chloride in the sample as silver chloride by conventional procedures, evaporation of weighed samples with excess dilute sulfuric acid followed by ignition, and weighing of the resulting barium sulfate. All these procedures indicated a purity of the reagent of 100.00% zk O.O50J,, even though some of these tests were performed on the same batch more than one year apart. A standard barium solution containing 1.000 mg. of barium per ml. was prepared by using the theoretical amount of the reagent. Procedure. Beakers, 250-ml., should be cleaned thoroughly and be free from scratches to prevent adhering of the precipitate. Dilute the solution containing up to 200 mg. of barium to 100 ml. T-801 Building, Brookharen Xational Laboratory, Upton, L I , N Y 2 Present address, Allied Chemical & Dye Corp.. Long Island City, N. Y.
Mg.
10 10 100 100
100 100
100 500 100 1000 1000
50 100
+l.8, +2.2 -0.3, -0.3
100
0.0, 0.0
3600
100
IiBrO3
2000
10
KBrOa
2000
2000
+2.7, f2.7
50 10 100
NH4NO3
2000
- 0 . 5 , -0.6
-0.3
10 100 10
1000
NHpBr "08
-0.1, -0.1, -0.3
100 200
100 100 100 100 100 100
&SO4
Error, hlg. Direct precipitation This method
10
10
METHOD
1
Determination of Barium in Presence of Other Substances
100
+0.1, +0.1 -1.7. -1.1, -1.8 0 , 0,0.0 -0.7, -0.3 +7.8
-0.8
-0.7 -0.7
-0.4
-0.5 +1.0, +0.6, f0.7 +0.3, +0.5
- 0.5
-0.7, -1.0 (2000 "02) -2.9. -2.8
c0.9.+ i . n
-4.5.
-n. 1.
(KNOa present) -3.2
(NaBrOlused) -0.2, + 1 . 3 (NaBrOs used)
-0.3
f1.5, f1.7
Hap01 500 100 -1.1, -0.9 a As chloride. b .Is chloride, 0.8 gram of tartaric acid added per 0.1 gram of Fe. cipitates were yellow showing contamination.
Pre-
Comparative results obtained by the given procedure and a conventional procedure ( 3 ) containing various cations and anione in addition to barium and chloride are given in Table I. Four results were excluded because of errors in handling. One result. obtained before the reignition after treatment with dilute sulfuric acid, gave an unexplained low result and was excluded when duplicate redetermination gave consistent results. DISCUSSION
Two methods of determining barium by precipitation from homogeneous solution were described by Elving and Van Atta (I), who obtained the sulfate ion by hydrolysis of methyl sulfate,
1243
17.44
ANALYTICAL CHEMISTRY
and by Wagner and Wuellner (7), who obtained the sulfate ion by hydrolysis of sulfamic acid. When the results of these authors are compared with those ohtained by the given procedure, it is evident that, in presence of large amounts of calcium, better re d t s were obtained by Elving and Van Atta and also by Waguer and WueUner. For small amounts of strontium, the results of the given procedure are intermediate between those of Wagner and Wuellner and Elving and Van Atta, who showed the least error, In the presenoe of sodium or potassium, the given procedure produces satisfactory results, even if the sodium ion is present in large excess, although slightly high results are obtained in the presence of B large excess of potassium ion, I n the presence of nitrate ion, the results obtained by the given procedure are comparable to those of Elving and Van Atta and Wagner and Wuellner up to the maximum amounts used by them, io0 mg:. of nitric acid. Even1 for much larger quantities of nitrate, the errars are smaller then those obtained by the conventional procedure. t ) nr Neit h..m ".I~ _. - ,17) \. ,s~m r o~TO ..suits . on " samples containing bromate. The given procedure yields satisfactory results in the presence of a large excess of bromate. I n samples containing 100 mg. of barium and 2000 mg. of potassium bromate, the high results are probahly due to the presence of potassium, rather than the bromate, as indicated by the direction of the.er ror. Samples Containing up to 0.5 gram of ferric iron give only fair ..x?,,.D f..- tl results, which is probably due to a compensatic.. * f oul.y.l, precipitates were light brown, indicating contamination with iron. Also, samples containing iron precipitated much more slowly and the supernatant liquid remained slightly turbid. ~~
." ".
.".
Sulfate was much retarded even though the proportion of calcium was low. The results, obtained by the conventional procedure in the presence of strontium, were better than those obtained by the recommended procedure or by the sulfamic acid procedure (7). It was observed that errors due to strontium increased with longer times of standing. For samples each containing 10 mg. of strontium and barium and times of 5, 30, and 60 minutes after precipitation, the e i r o r ~were +6.0, f7.6, and +8.5 mg., respretively, whereas solutions containing only harinm showed no such time dependence. This Suggests that strontium sulfate pastprecipitates. Methods of homogeneous precipitation can he expected to favor errors due to postprecipitation because there is R closer approach to equilibrium conditions. Barium sulfate obtained by the given procedure did not creep and was easily filtered. The precipitate consisted of individual crystals of uniform Biae, larger than those obtained,hy the other two methods (1, 7). Figure 1,a shows B photomicrograph of the precipitate obtained by the given procedure wherex b shows a photomicrograph of a barium sulfate precipitate obtained by a conventional procedure (5). Both are taken ttt the mme magnification. Because of the great difference in appearance of tho precipitate under the microscope, it was thought that the p r e cipitate might consist of a barium-organic compound which also contains sulfate. This precipitate was heated a t 110" C. and weighed. Results showed that weights were only 3 to 4 parts per hundred higher than the weight of the same precipitate after ignition to 900' C. This is the loss of weight expected far hasium sulfate. In addition, the cry8tala examined with crossed Nicol prisms were found to be anisotropic, the asme 8s barium sulfate. Growth of regularly shaped crystals of increased size OCCUIS even though the rate of precipitation is fairly rapid its long as concentration changes take place uniformly. In hydrolysis procedures where the pH was lowered more g r s d u d y or to a final p H of 4 to 6, the precipitates were invariably found to be liner and showed a tendency to creep, although precipitation was complete. This is in agreement with the experiments of Levin and Vanee
(4). ACKXOWLEDGMENT
The Bersworth Chemical Co., Framingham, Mass., furnished samples of Versene for preliminary tests. LITER*TURE CITED
(1) Elving. P. J., and Van -4th. R. E., ANAL.C ~ E M 22, . . 1375-12 (1950). (2) Hulett. Z.phgsik. Chem. 37. 385 (1901). (3) Kolthoff. T.M.. and Sandell. E. B.. "Textbook of Quantitative Inorganic Analysis." p. 340. New York. Maomillan Co., 1946. (4) Levin. S. Z.,and Vanoe, J. E., J . An. Chem. Soc.. 74,1433 (1952). (5) MacNevin, W. M., and Dunton, M. L.. ANAL Cmar.. 26, 1246 (1954). (6) Pfibil. R.. and MariEov&. D.. C h a . List?,. 46. 4 0 0 3 (19521. (7) Wagner, W. F., and Wuellner. J. A.. AN 1 (1952). (8) Willard. H. H.. Ibid..22, 13724 (1950). ~~~~~~
Figure 1. Photomicrographs of Barium Sulfate Precipitate a. Ho"ogeneous precipitation pmcedur. b. Customary p m d n r o
Precipitates obtained by the conventional procedure appeared
to be contaminated to nn even greater extent. Phosphate gives a greater error by this procedure than by the sulfamic acid Drocedure. On the whole, the results by the given . .uracedure have slight tendency to be low. If there-is no compensation of errors, results hy this and most other barium procedures would be expected to be slightly low, due to the solubility of b:rrium sulfate in the hot wash water. According to Hulett ( 8 ) at,out 0.4 to 0.7 mg. barium sulfate dissolves per 100 ml. of hot water and this is the maximum loss which may occur. The actual loss is uur smaller because of the failure to establish equilibrium. &I"^ 1L^ absolute error is smaller for smaller quantities of barium because of the relatively small surface area of the precipitate obtained by the given procedure, and the decreased chance for establishment of solubility .equilibria as compared to heavier precipitates. I n the presence of calcium. complete precipitation of the barium
~
~~
~
~
. ~ "....-.. ~,
RECEIVED for review July 24. 1851. Amepted Ap... ~ " , -J before the Division of Analytical Chemistry st the l l Q t h Meeting of the AMERICAN CXEMLICAG SOCIETY. Boston. Mae&. 1851. Abstracted from II thesis aubrnitted by Eugene Eohupak in partial fulfillment of the reouiremen- for the M.A. degree. Boston University. Boston. Mass.
Ddermination of Zinc by Dithizone-Correction T"
I Y
*I ." Ullr
"*&:","
aLYI.lr
""
