1505
V O L U M E 2 7 , N O . 9, S E P T E M B E R 1 9 5 5 resulted in immediate precipitation as each increment of nickel sulfate was added, and a distinct break occurred in the conductance or high-frequency plot a t the 2 to 1 ratio. S o break was detected a t the 1 to 1 ratio under any of the conditions tried. One or two drops of 3 5 ammonia for each 100 nil. of O.0OlL4J dimethylglyoxime pioduced distinct 2 to 1 end-point breaks with considerable curvature of the lines leading to them. As the ammonia concentration was increased, the linearity of the plots on the two sides of the end point improved. Addition of ammonia beyond 6 ml. or addition of acetate buffer decreased the change of slope a t the end point. The conditions selected as beet for the conductometric titration of dimeth) lglyoxinie with nickel viere as follows. An 0.005M stock solution was made by dissolving a weighed amount of Eastman Kodak dimethylglyoxime in the minimum amount of alcohol and diluting to volume with water. -4n 0.005M stock solution was made by dissolving a weighed amount of Baker & Adamson nickel sulfate hexahydrate reagent and diluting to volume. This solution was found to be 0.00496M by gravimetric dimethylglyoxime determination. Twenty millimeters of the dimethylglyoxime solution and 5 to 6 ml. of approximately 31T ammonia were mixed and diluted to approximately 100 ml. The resistance was measured with the Leeds and Northrup Co. bridge folloving each addition of nickel sulfate solution. The end points, as determined from graphical plots, agreed within 0.1 ml.
il sample curve is shown in Figure 1. The cell constant for the cell used in this titration v . a ~0.0962. An average of three determinations gave the value 10.14 ml. Using the gravimetric value for the concentration of the nickel, the concentration of the dimethylglyoxime was calculated to be 0.00503, in good agreement with the value from the weight used. This method should permit awuratr standardization of dimethylglyoxime solutions.
OTHER IONS
The high-frequency and conventional conductance proretlures described above for nickel were repeated for cobalt(II), nitrate, lead(II), nitrate, and manganese(I1) chloride. As for nickel, the conditions of Nakano, Hara, and Yashiro were first duplicated as nearly as possible, then other conditions were tried. I n no case, with or without the addition of ammonia, was there any indication of end points a t any ratio. The cobalt solutions gave a yellow to amber color and whenever ammonia v a s present the lead solution yielded a small amount of a-hite precipitate. ilddition of the first 5 ml. of dimethylglyoxime to the 0.001M cobalt solution produced no change of p H from the initial p H of 1.8. h gradual drop t o pH 4.6 occurred between 5 and 25 ml., possibly indicating some slight displacement of hydrogen ion from the dimethylglyoxime by coordination. For both the lead and manganese salts the pH rises continuously on addition of dimethylglyoxime, indicating no coordination involving displacement of h),drogen ion, ACKKOWLEDGMEYT
The n-ork upon which thir report is based vias a joint undertaking of the Department of Chemistry of Kest Yirginia University and the Office of Ordnance Research, U. S . Army. LITERATURE CITED
Hall, J. L., and Gibson, J. A , , Jr., h.41,. CHEM.,23, 966 (1951). Hall, J. L., Gibson, J. A., J r . , Phillips, H. O., and Critchfield, F. E., J . Chem. Educ., 31,54 (1954).
Xakano, K., Hara, K . , and Tashiro, K., AXAL.CHEM.,26, 636 (1 954).
JAMESL. H ~ L L K e s t Virginia University Slorgantown, W. T’a.
JOHN A. GIBEOS.JR. HAROLD 0. PHILLIPS P.~.ur. R . WILKTNSOX
Effects of Impurities upon Rate of Precipitation and Particle Size SIR: It has been observed ( 1 , a ) that crystals of barium sulfate, precipitated with a freshly prepared barium chloride solution, are distinctly s n i d e r than those obtained with a reagent solution which has been Ptanding for some time. If “barium chloride” is understood to mean pure barium chloride, the explanation of the “aging” of its solution invites speculations involving equilibria between hydrated and complex ions and considerations of kinetics including that of the growth of an anhydrous lattice of barium sulfate. With such possibilities in mind, a quick experimental survey was carried out, which soon led to the suspicion that an impurity of the reagent might be responsible for the phenomena observed. I n the final series of experiments, about 15-ml. portions of 0.02~21barium chloride solution and of a solution 0.01M with sulfuric acid and 0 . M in hydrochloric acid were briefly heated in test tubes by inserting them into a steam bath. The hot solutions were simultaneously poured into an empty beaker which was being heated upon the steam bath. The mixture was kept hot for 5 minutes with frequent mixing by swirling. I t was then allowed to cool to room temperature during 10 minutes, after which the average diameter of the crystals of barium sulfate vias determined under the microscope. Batch A. Barium chloride dihydrate, maximum impurities listed “heavy metals 0.0.” When a freshly prepared solution of the salt was used within 10 minutes from the time of dissolution, the crystals of the precipitated barium sulfate had a diameter of approximately 3 microns. This diameter increased to 5 microns when the freshly prepared solution of the barium chlo-
ride was filtered through S. & S. Blue Ribbon, ash-free filter paper and used within 10 minutes from the time of dissolution; average particle diameters of 6 microns were obtained with solution of the salt, filtered or not filtered, which had been standing in glass containers for one to several days. The results were not affected if the solution was 0 . 2 M during the aging period and was diluted with water to 0.02M just previous t o use in precipitating barium sulfate. Batch B. Some of batch *4 was recrystallized from hot water. The hot saturated solution of about 50 grams of batch A assumed, for a moment, a dark gray color and this was immediately followed by flocculation of a very small amount of partly tawny and partly brownish black preci itate which was removed by filtration and gave a positive test g r iron, .4 freshly prepared solution of batch B used within 10 minutes from the dissolution of the salt gave barium sulfate with average grain diameters of 6 microns. Aging or filtration of the solution had no effect on the particle size of the barium sulfate. Batch C. Barium chloride dihydrate meeting ACS specifications and furnished by a different firm. Solutions of this pure salt behaved exactly like solutions of batch B-Le., the aging of solutions had no effect on the particle size of precipitated barium sulfate. The conclusion seems unavoidable that solutions of pure barium chloride do not age and that in the instance of my brand A, impure barium chloride, the increase of the particle size of precipitated barium sulfate is caused by a removal of the impurity. Adsorption on the walls of storage bottles or on the surface of filter mats could bring this about, since the amount of impurity is very small. Freshly prepared solutions of batch A gave a turbidity of barium sulfate immediately upon addition of sulfate solutions
ANALYTICAL CHEMISTRY
1506 With aged solutions of batch A and with pure barium chloride, the turbidity is perceived after a slight delay. The time lapse between addition of the reagent and perception of the first turbidity is much more pronounced with strontium sulfate and was studied with 0.01dl solutions of strontium salts and sulfuric acid. The turbidity became visible 20 to 30 minutes after combining the reagents if a solution of our strontium chloride hexahydrate was used. The time lapse increased to 1 hour if our strontium chloride hexahydrate was first rendered anhydrous by drying i n a n oven and then dissolved in water. Our strontium nitrate gave the visible turbidity within 2 to 5 minutes, but the time lapse was extended to 2 hours if the solution as acidified with hydrochloric acid and allowed to stand 2 hours before adding t h r sulfuric acid. These strontium salts retained the described brhavior also after attempted purification by recrystallization from water. After the experience with the barium chloride, it does not seem proper t o attach any importance to the odd behavior of our strontium salts unless it mere proved that they are entirely free from impurities. Traces of barium could be responsible for the observed phenomena, and they would not be removed b j recrystallization from m t e r . All these observations show that the rate of precipitation and the particle size of precipitates may be strongly affected by tr-ic’e impurities in solutions. I n the instance of the barium sulfate, the particle size decreased by more than 50% without any noticeable change in the habit of the crystals of barium sulfate. Considering this remarkable sensitivity of the particle size, the analyst will do well to consider the origin and history of his solutions when trying to obtain easily filterable precipitates. I n addition, fine-grained precipitates might be contaminated bv the trace constituents which are responsible for the small particle size.
paper. The recrystallized salt, used in a 5 % solution, required 102 minutes for 28 mg. of barium sulfate t o settle. When this same solution was used 48 hours later, the barium sulfate took 49 minutes for 28 mg. to settle. These sedimentation results cannot be compared with those above, for they were carried out a t n lower temperature. It \vas suggested during the examination over my thesis that air was serving as nuclei. This was checked with a 5% solution: Standing for 1 hour a t 757.5-mm. pressure gave barium sulfate that required 47.4 minutes for 30 mg. to settle, standing over the same period of time but under 27.5-mm. pressure, with vigorous agitation every 5 minutes, gave barium sulfate that required 26.6 minutes for 30 mg. to settle. I have no doubt that impure barium chloride will cause fine precipitates of barium sulfate. I believe, too, that no one single item will explain all the interesting facts ahout this precipitation LITERATURE CITED
(1) Bogan, E. J.. thesis, Ohio State University, 1947. University of hIaine EDGAR J. BOQAN. Orono, Maine
Separation of Orthonhosnhates from Organic Phisphates
SIR: I n attempting to follow the procedure of Martin and Doty [lIartin, J. B., and Dotv, J. H., ASAL. C H m f . , 21, 965 (1949)] for the separation of orthophosphates from organic phosphates, n e found that the aqueous phase developed a blue cwlor of greater intensity than the isobutyl alcohol-benzene phase. I n an effort to eliminate this trouble, we prepared a new set of reagents, using only redistilled watpr and chemic& LITERATURE CITED of the highest purity. (1) Bogan, E. J., thesis, Ohio State Universlty, 1947. The persistence of a deep blue color in the aqueous phase (2) Fischer, R. B., and Rhmehammer, T. B., -4s.t~.CHEM.25, 1544 suggested that a contaminant in the ethyl alcohol used to rinse (1953); 26, 244 (1954). Queen’s College A. 8 . BENEDETTI-PICHLERthe pipet was responsible. I n fact, addition of 1 ml. of absolde Flushing, N. Y ethyl alcohol to a mixture of 2 ml. of 4y0 ammonium molybdate in 4-V sulfuric acid and 0.50 ml. of stannous chloride caused a blue color to develop aithin a few seconds. The addition of 1 SIR: Benedetti-Pichler’s work on the “Effects of Impurities ml. of acetone gave a very intense blue color (absorbance > 2 upon Rate of Precipitation and Particle Size” is interesting. a t 730 mp in a Beckman Model D U spectrophotometer with 1-cin. His conclusions on the aging and filtering of a solution of recryscuvettes), while 0.3 ml. of acetone gave a pale blue (absorbance tallized barium chloride dihydrate do not agree with mine. The 0 05). One milliliter of acetone-free methanol (as checked differences could well be explained bv the differences in the preby a mass spectrometer) had an absorbance of about 0.05 a t 200 cipitation. His conditions are: 0.02M barium chloride added seconds and about 0.15 a t 800 seconds. t o a solution 0.12147 in hydrochloric and sulfuric acids after a 5Methyl ethyl ketone, isopropvl alcohol, and isobutyl alcohol minute heating period compared to mine: 5% barium chloride also caused some blue color to develop. Formnldehyde and added to sodium sulfate in O . O O 4 N hydrochloric acid a t once a t ether did not cause the formation of any color. Heptyl and room temperature. Our filtering of the solutions was also difhexyl alcohol caused the formation of a bright yellow color in the ferent, his being filtered through S. & S. Blue Ribbon while mine aqueous phase, xhile cyclohexanone caused a yellow color in the was filtered through Pyrex F filters Tyith suction. cyclohexanone phase. I n my work ( I ) , a sedimentation method \vas used t o comptre I n addition to suggesting a rather sensitive method for acetone particle sizes. Barium chloride dihydrate, without previouh and possibly for some alcohols, these results explain the residual treatment, was used to make a 5% solution. After precipitation blue color in the aqiieous phase even after extraction of all the with this solution, 66 minutes were required for 28 mg. of barium orthophosphate. sulfate to settle. This same barium chloride dihydrate wap disJACOBJ. BLUM solved in hot water, filtered through a Pyrex F filter with suction, Navy Medical Research Institute and recrvstdlized. This salt \vas used to make a 5% solution. Bethesda, ’Id. The barium sulfate obtained required 48 minutes for 28 mg. to settle. This same 5 % solution of recrystallized salt was filtemd through a Pyrex F filter with suction. It was used within 5 minutes of the solution above and the barium sulfate took only 5.2 minutes for 28 mg. to settle. This last was not included in my thesis, for I felt that recrystallization had little to do with the effects on particle size. Also not included in my thesis was an experiment on recrystallization where the hot concentrated solution \vas filtered through
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