ANALYTICAL CHEMISTRY
1504 Principal Lines I/Il
d
5.03
4.09
3.75 3.63 3.53
Weak 2 2 5 Weak
d
2.05 1,992 1.955 1.920 1.840
AIolecular Refraction ( R )(5893 h.; 25" C). = 33.1; R (obsd.) = 38.8.
I/Ii
1 1 2 Weak 2
q z
d 3.45 3.35 3.24 3.04 2.99
I/II
2.79 2.74 2.60 2.50 2.46
1 1 4
9 10 7 4 3
3
3
1.606 1.581 1.540 1.606 1.440
1 Weak 2 1 Weak
1,735.
On warming, the crystal size and growth rate both increase. Bisectrix figures are possible with 2V about 90".
FUSION DATA.4-Aminosalicylic acid decomposes on melting (220" C.) to give a pure decomposition product melting, in turn, a t 122" C. The latter crystallizes spontaneously from the melt after considerable supercooling. The melt is fairly viscous, the crystallization velocity is very slow, and the crystals are small.
Irenecorvin is responsible for powder x-ray diffraction data.
R (calcd.)
ACKNOWLEDGMENT
CONTRIBUTIOKS of crystallographic d a t a for this section should be sent to Walter C. McCrone. Analytical Research. Arinour Research Foundation of Illinois Institute of techno log^-, Chicago 16, 111.
CORRESPONDENCE
Conductometric Titrations with Dimethy IgIyoxi me SIR: This note reports the conditions under which conductometric methods, including high-frkquency methods, may be used to determine the end point for the titration of solutions of nickel(11) ion and dimethylglyoxime and shows that for some recently reported titrations of dimethylglyoxime with cobalt, nickel, lead, and manganese (3),no changes that occur under the conditions specified permit detection of the claimed end points by two recognized types of high-frequency apparatus or by a conventional conductance method. Sakano, Hara, and Yashiro ( 3j report high-frequency titration curves having very abrupt changes of slope a t molar ratios 2 to 1 and 1 to 1 for the addition of O.OO5M nickel sulfate, cobalt nitrate, lead nitrate, or manganese chloride to 0.001M aqueous dimethylglyoxime in the absence of any buffer. It is the opinion of the present authors that their results must be related to factors other than the reported reactions and would not be substantiated over any considerable range of concentration and volume. The high-frequency titrations, using concentrations and conditions specified by Nakano, Hara, and Yashiro, have been repeated in this laboratory using the apparatus based on the General Radio Twin-T impedance measuring circuit (1j a t a number of frequencies from 2 to 10 Mc., and using a crystal oscillator apparatus (2) a t 5 Mc. A thorough study has also been made of these reactions, a t the same and other concentrations, with and without added ammonia and buffers. These titrations were followed with the highly precise 1000-cycle conductance bridge huilt around the Leeds and Northrup Co. No. 1553 ratio box, as well as with the radio-frequency apparatus just listed.
duced a t any point during the titration by addition of a few drops of ammonia. Similar conductance curves and color changes were produced over concentrations ranging from 0.001M to 0.005M. Similar color changes and converse conductance changes were produced by addition of dimethylglyoxime to nickel sulfate solution. More concentrated solutions of dimethylglyoxime were made by using a small amount of ethyl alcohol in the water, but this did not affect the nature of the results.
DIMETHY LGLYOXIME-NICKEL
Figure 1. Conductometric titration of dimethyl glyoxime with nickel ion in presence of ammonia
For all unbuffered solutions, a plot of instrument response against volume of nickel sulfate added to the aqueous dimethylglyoxime gave a smooth curve, indicating a continuous increase i n conductance with no abrupt changes. For the addition of O.OO5M nickel sulfate to 100 ml. of 0.001M aqueous dimethylglyoxime (for which the 2 to 1 and 1 to 1 ratios would occur a t 10 and 20 ml. of nickel sulfate, respectively), the first 1 or 2 ml. of salt solution resulted in a faint yellow color. More nickel sulfate gradually changed the color to a faint pink, with a slight turbidity increasing to a just visible precipitate at 15 to 25 ml. of added solution. Greatly increased precipitation may be pro-
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These observations suggest that in these unbuffered solutions a small amount of the usual 2 to 1 complex is formed and precipitated but that no reaction is completed. This conclusion was further supported by p H measurement as dimethylglyoxime was added to nickel sulfate solution. Upon addition to a 0.001M salt solution the pH falls continuously to well past the 2 to 1ratio, hydrogen ion being released as the 2 to 1 complex is formed. Prior addition of ammonia to the dimethylglyoxime solution
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 pH from the initial pH 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 . W I L K T N S ~ X
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 0.0LVI 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
