Quantitative Spectrographic Studies of Co- precipit ation 11. Group I1 Elements with Barium Sulfate LOUISWALDBAUER AND E. ST. CLAIRGANTZ
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Division of Analytical Chemistry, State University of Iowa, Iowa City, Iowa HE study of co-precipitation with barium sulfate has long been of interest and has occupied the attention of a great number of workers. Owing to the difficulty of precise determination of some of the con-
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The co-precipitation of beryllium, magnesium, calcium, zinc, strontium, and cadmium with barium sulfate was studied spectrographically. All precipitates were prepared by the method of Popoff and Neuman, which consists essentially of the addition of the solution of the sulfate ions to the solution of the barium ions. It was found that beryllium, magnesium, and zinc do not coprecipitate, whereas calcium, strontium, and cadmium do. The type of co-precipitation has been suggested for each element.
taminating substances by ordinary analytical m e t h o d s , the authors turned to the spectrograph as a possible means of overcoming this difficulty. P o p o f f , W a l d b a u e r , and McCann (9)applied quantitative spectrographic methods to the study of magnesium in the quantitative determination of calcium as calcium oxalate, and their method was applied to the study of co-precipitation of group 11 elements in the determination of sulfate as barium sulfate.
MATERIALS Water and sulfuric acid were purified by distillation. Sodium hydroxide, carbonate- and chloride-free, was prepared by the Cornog ( I ) method. Hydrochloric acid was prepared by dropping 12 M c. P. hydrochloric acid upon 18 M c. P. sulfuric acid. The hydrogen chloride gas formed was washed and collected in redistilled water. All salts used were recrystallized from redistilled water to remove the very small amounts of the other alkaline earths that were initially present.
SOLUTIONS
rectly to the stock sulfuric acid solution when the presence of beryllium or cadmium was desired. Stock solutions of these salts as cation sources were not used becauseof their i n s t a b i l i t y over a period of time. The zinc chloride solution was made up using the special hydrochloric acid already described and spectroscopically pure zinc ribbon which was obtained from the New Jersey Zinc Co. Ten milliliters of this solution contained enough zinc exactly to replace all the barium in 0.8 gram of barium sulfate. The spectographic s t a n d a r d s were prepared by weight, using the purified salts already described. A range of standards containing from 0.05 to 1 per cent of the cation was prepared for each element. Each solution was of the same barium sulfate and concentrated sulfuric acid content as those prepared from the precipitates under examination. Some difficulty was encountered in the preparation of the cadmium standards. After the solution of the necessary amount of barium sulfate in the concentrated sulfuric acid used as a solvent, it was impossible to get the cadmium iodide into solution. It was necessary therefore to use a small amount of Kahlbaum’s cadmium carbonate in order to introduce the amount of cadmium needed for the several standards. The precipitates for spectrographic analysis were obtained from a third sample prepared under the same conditions as those for the gravimetric determinations. The precipitate was caught on a small Hirsch funnel, using a filter paper. This method was satisfactory, and eliminated the possibility of contamination from the asbestos in the Gooches, which might have occurred if the gravimetric precipitates had been used. Solutions for spectrographic analysis were obtained by dissolving 0.4 gram of the precipitate in 5 ml. of the concentrated sulfuric acid already described.
Sulfuric acid solutions used as a source of sulfate in all the work were made up by weight from the purified stock solution which was stored in silica bottles. These solutions were made up APPARATUS so that 40 grams of solution were equivalent to 0.8 gram of barium sulfate, and were standardized, using weight burets, with The apparatus for the preparation and study of the prethe standard sodium hydroxide. cipitates is essentially that used d Phenolphthalein was used as an % by Popoff and N e u m a n (8), indicator. The s t a n d a r d b a s e sr and by Popoff, Waldbauer, and had been analyzed using Bureau of Standards potassium acid phthal1. McCann (9). ate, all necessary p r e c a u t i o n s 0.5 having been taken for the excluGENERAL METHODS OB sion of carbon d i o x i d e d u r i n g 0.1 ANALYSIS titration. Vacuum c o r r e c t i o n s were made whenever necessary. 0.05 Except w h e r e o t h e r w i s e Popoff and N e u m a n (8) f o u n d stated, the following m e t h o d that standardization against parecommended b y Popoff a n d tassium acid phthalate agreed to within 0.05 per cent with standNeuman (9) w a s u s e d in the tL preparation of the various preardization against benzoic acid and against c o n s t a n t - b o i 1i ng h cipitates both for gravimetric 0 * hydrochloric acid. and spectrographic determinaBarium chloride solution conFIGURE1 tion: t a i n i n g 21 grams per liter was made up from the purified crystals. Approximately 40 grams of the sulfuric acid were added dropThe magnesium chloride, calcium chloride, and strontium chlo- wise to an acidified (4ml. of 6 N hydrochloric acid) solution conride solutions used as a source of the separate cations were made taining a slight excess (5 ml.) of barium chloride, 10 ml. of the u from the purified salts using redistilled water, so that 10 ml. cation solution, and enough redistilled water to bring the final ofthe stock-cation solution would contain enough cation exactly volume of the mixture to 350 ml. During the addition of the sulfuric acid with constant stirring, the barium solution was kept to re lace all the barium in 0.8 gram of barium sulfate. SoEd beryllium chloride and cadmium iodide were added di- just below the boiling point by means of a hot plate. The pre311
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cipitates were digested for 1 hour just below the boiling point. The mother liquor was decanted and the precipitates were washed with four 15-ml. portions of hot water and washed into the crucibles. The precipitates were heated for 1 hour at 800” C. and weighed. For spectrographic determination, electrodes and solutions were prepared as previously described. Six drops of the solution t o be analyzed, or of the standard, were placed in each electrode. The electrodes so prepared were dried in an oven at least 6 hours at a temperature of 240’ t o 260’ C. The high temperature was necessary in order to get the electrodes completely dried, for when not entirely dry they do not burn well.
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70 Sr 1.
0.5
01
0 05 1
ment conclusively proves the absence of that particular element.” Mellor (5) and Johnston and Adams (4) say that salts of the alkaline earths and magnesium are carried down by barium sulfate. But neither beryllium, magnesium, nor zinc was found to eo-precipitate with barium sulfate. It would seem probable that because of the long digestion (18 hours) used by Johnston and Adams, magnesium and zinc contamination was due not to co-precipitation but to postprecipitation. Precipitates were also prepared from solutions containing double equivalents of the cation as well as by the regular method, but no eo-precipitation of beryllium, magnesium, or zinc was observed. Since beryllium, magnesium, and zinc do not co-precipitate with barium sulfate, the quantitatiie study was confined to calcium, strontium, and cadmium, Precipitates were all prepared by the reverse method, having equivalent amounts of the cation present in the barium solution a t the start. These precipitates were studied both gravimetrically and spectrographically and the results are recorded in Table 1. TABLEI. PRECIPITATION WITH BARIUM SULFATE ELEMENT STUDIED
0.5
0.1
0 06
r?
0” dl
FIGURE2 After exposure and development, h a 1 analysis of the plates was made by visual examination, using a ground-glass light-box to give a white background. Although this method is none too precise, it was felt that it gave values as reliable as the method warranted. With visual estimation, no attempt was made to extend precision beyond one figure. This method is essentially the same as that used by Nitchie ( 6 ) . In comparison with standards, the standards were photographed at the top of the plate followed by the spectra for analysis (Figure 1). This method is as reliable as that used by Nitchie (6) in which the spectra under examination were photographed between several standards (Figure 2). The lines used in examination were as follows: Magnesium, lines 2852, 2798, and 2790 A. Zinc, lines 3345.5, 3345; and 3303, 3302.6 A. Calcium, lines 3968.5, and 3933.7 A. Strontium, line 4077.8 A. Cadmium, lines 2288 and 2265 A. Beryllium, lines 2348, 3130, 3321 b.
EXPERIMENTAL The object of this study was to determine spectrographically the amount of co-precipitation of group I1 elements with barium sulfate. Of the elements in group 11, radium compoundk were not available, and owing to the inherent difficulties in their use no mercury compounds were studied. Beryllium, magnesium, zinc, cadmium, calcium, and strontium were studied. The first course was to determine qualitatively whether these elements did or did not co-precipitate with barium sulfate. The precipitates were prepared from solutions containing amounts of the elements equivalent to 0.8 gram of barium sulfate. Spectrograms were taken and the presence of the element in question was determined by the presence or absence of its raies ultimes or most persistent lines. Pollock ( 7 ) states that “the absence of the residuary lines of an ele-
Vol. 5, No. 5
Calcium Calcium Strontium Strontium Cadmium Cadmium
DIFFERENCE FROM BARIUM SULFATE THEORETICAL ELEMENT Calcd. Found Actual Calcd. IN BaSOa Gram Gram Mg. Mg. % 0.8065 0.8036 2.9 2.23 0.2 2.23 0.2 0.8188 0.8167 2.1 0.7913 0.7764 14 9 14.64 1.0 1.0 0.8066 0.7926 14.0 14.64 0.8215 0.8201 1.4 0.84 0.3 0.7997 0.7970 2.7 0.84 0.3
In column 2, the calculated weight of barium sulfate was found by the usual method from the weight of standard sulfuric acid taken; the weight in column 3 was that found experimentally by gravimetric analysis. In column 4, the actual difference was obtained by subtracting the value in column 3 from that in column 2. The value in column 5 was the change in the weight of the precipitate expected, if the per cent of the element found spectrographically were substituted for an equivalent amount of barium. DISCUSSION The amount of co-precipitation observed with strontium was to be expected, inasmuch as strontium sulfate and barium sulfate are isomorphous salts having the same lattice-type structure. One might also expect a small amount of calcium to be coprecipitated according to the modified Paneth (2) adsorption rule: “On a polar crystal those ions experience good adsorption that form with the oppositely charged components of the lattice a compound that is poorly soluble or slightly dissociating in the solvent employed.” I n the case of cadmium, a search of the literature failed to show whether or not cadmipm sulfate has the same latticetype structure as barium sulfate. Cadmium sulfate crystals belong to the orthorhombic system, as do both barium and strontium sulfates. However, a comparison of the co-precipitation of cadmium with that of strontium would lead one to believe that cadmium sulfate is not isomorphous with barium sulfate. The low gravimetric results (compared with the amount of cadmium present), obtained when cadmium iodide was studied, are probably due to the preferential adsorption of the iodide ion and also of the Cd14-- ion. According to Heym (3),Cd14-- is formed when the slightly dissociated cadmium iodide is dissolved in water. From a study of crystal-structure data one would not necessarily expect beryllium, magnesium, or zinc to eo-precipitate with barium sulfate, and the authors were unable to detect any eo-precipitation, using either the reverse or the regular methods of precipitation.
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INDUSTRIAL AND ENGINEERING
CONCLUSION Calcium, strontium, and cadmium are all co-precipitated with barium sulfate. Beryllium, magnesium, and zinc are not co-precipitated with barium sulfate. The data for cadmium indicate the preferential adsorption of iodide ion or of Cd14--.
LITERATURE CITED (1) Cornog, J . Am. Chem. Soc., 43, 2573 (1921). (2) Fajans, “Radio Elements and Isotopes,” P. 95, McGraw-Hill, 1931.
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(3) Heym, Ann. Phys., (9) 12, 443 (1919). (4) Johnston and Adams, J. Am. Chem. Soc., 33, 829 (1911). (5) Mellor, “Comprehensive Treatise on Inorganic and Theoretical Chemistry,” Vol. 3, p. 766, Longmans, 1923. (6) Nitchie, (3. ( 3 . 3 I N D . EXQ.CHEM.,21, 1 (1929). (7) Pollock, Sci. Proc. Roy. Dublin Soc., 11 (N.S.), 184 (1907). (8) Popoff and Neuman, IXD.ENG.CHEW,Anal. Ed., 2, 45 (1930). (9) Popoff, Waldbauer, and McCann., Ibid., 4 , 4 3 (1932). RECEIVED April 16, 1933. Presented before the Division of Physical and Inorganic Chemistry at the 85th Meeting of the American Chemical Society, Washington, D. C., March 26 to 31, 1933. From the thesis of E. St. Clair Gants for the M.8. degree, State University of Iowa.
Determination of Calcium in Lead-Calcium Alloys of Low Calcium Content BEVERLY L. CLARKEAND LELAND A. WOOTEN,Bell Telephone Laboratories, Inc., New York, N. Y.
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H E purpose of this work was the development of a rapid, high-precision method for determining calcium in lead-calcium alloys containing from 0.02 to 0.06 per cent calcium. Small amounts of several metallic impurities were present in the alloy. After the removal of the lead as sulfate from the nitric acid solution of the alloy, calcium and the impurities remain as the nitrates and sulfates. It is generally agreed that the best form in which to separate calcium for determination is the oxalate, and where speed is desired the volumetric evaluation of the calcium oxal’ate is obviously preferable to the gravimetric. Our problem was therefore that of modifying the classical method for determining calcium by the permanganate titration of the sulfuric acid solution of the oxalate, so that our requirements of precision and rapidity would be met. It was desired to run two concurrent determinations in an hour, with a precision of +0.002 per cent calcium. Shaw, Whitternore, and Westby (2), of the Western Electric Company, have published a method for this determination. They dissolve a 20-gram sample of the alloy in fuming nitric acid, add sulfuric acid, filter off the lead sulfate, and add successively ammonium hydroxide in excess, ethyl alcohol, and ammonium oxalate. The mixture is boiled for 2 minutes, and the precipitate allowed to settle for 5 minutes. The mixture is filtered through asbestos, the precipitate dissolved in sulfuric acid, and titrated hot with permanganate in the presence of the asbestos. A careful trial of this method led us to suspect compensating errors. Repeated determinations on the same sample were generally consistent when performed on the same day and by the same analyst; but were generally inconsistent when either of these conditions did not obtain. This observation pointed to the existence of uncontrolled variables. A careful search for such variables led to the discovery of the following sources of error and experimental difficulties: 1. Incompleteness of the calcium precipitation under the prescribed conditions. 2. Difficulty in selecting end point in presence of asbestos, and possible reducing action of the asbestos. 3. Failure to make an ammonium hydroxide separation of the Second and Third Group Metals before precipitating calcium as oxalate. 4. Solubility of calcium oxalate in the unneutralized hot wash water.
5 . Incomplete removal by washing of the excess ammonium oxalate in the calcium oxalate precipitate. The experimental basis for so characterizing the above follows : 1. The Western Electric Company method specified a cooling period of 5 minutes a t room temperature before filtration of the calcium oxalate precipitate. Since the completeness of the precipitation is a function of the final temperature of the solution, results obtained in summer will show a negative error from this cause measurably higher than in the case of those obtained in winter. Experiment confirmed this; microchemical examination of the filtrates from the calcium oxalate precipitation showed that cooling in an ice bath reduced the error by approximately 50 per cent. 2. The presence of asbestos in the solution being titrated was found definitely to obscure the end point, and to produce erratic results in check determinations. A positive error was introduced, caused either by a reducing action of the asbestos or by the tendency to overtitrate in turbid media, or by both. A certain alloy gave successive values of 0.051, 0.056, 0.053, and 0.055 per cent calcium by the Western Electric Company method. When the same alloy was analyzed by the method modified by the substitution of a Frittig glass filtering crucible for the asbestos Gooch, the values obtained by the same analyst were 0.040, 0.041, 0.042, and 0.042 per cent. 3. It was originally suspected that a source of error in the method lay in the failure to separate the Second and Third Group Metals as hydroxides before precipitating the calcium as oxalate. Upon investigation, however, it was proved that no measurable error was produced from this cause, at least in the case of the alloy in which we were primarily interested. This alloy contained on the average the following impurities: copper, 0.07 per cent; tin, 0.002 per cent; bismuth, 0.001 per cent; antimony, 0.001 per cent; arsenic, 0.002 per cent; and iron, 0.002 per cent. We found, however, that when the Frittig crucible was substituted for the asbestos mat, the hydroxides of these metallic impurities, if not previously removed, tended t o clog the Frittig crucible and render filtration prohibitively slow. For this reason, and also to guard against samples with abnormal amounts of impurities, it seemed desirable to filter off the ammonium hydroxide precipitate before precipitating the calcium. 4. Hahn and Weiler ( 1 ) state that calcium oxalate is appreciably soluble in hot water, but much less so in dilute ammonia solution. I n corroboration of this statement, on a
