Conductometric Titration of Perchlorate with Tetraphenylarsonium Chloride R. J. Baczuk and W. T. Bolleter Bacchus Works, Hercules, Znc., Magna, Utah
MANYVOLUMETRIC METHODS have been reported for the determination of perchlorate. Most of these are indirect, requiring reduction to chloride. Either the chloride is subsequently measured (1-5) or the amount of reductant is determined by back-titration (4, 6, 7). These methods are all subject to interference by other halide salts or reducible substances. Several direct methods have also been reported. A microgravimetric technique using 3,5,6,8-tetramethyl-l ,lO-phenanthroline to precipitate perchlorate was reported by Brandt and Smith (8). Perchlorate has been determined colorimetrically by Bodenheimer and W d e s (9)in pyridine as the cupric tetrapyridine salt and by Frii:z, Abbink, and Campbell (IO) as ferrous 1,lo-phenanthroline perchlorate in n-butyronitrile. Tetraphenylphosphonium chloride has been used by Nezu (1I) as an amperometric titrant for perchlorate. Recently, Morris (12) described the amperometric determination of perchlorate by titration with tetraphenylstibonium sulfate. The use of tetraphenylarsonium ion for perchlorate analysis was reported as early as 1939 by Willard and Smith (I.?), who proposed a potentiometric back-titration with triiodide. Recently, the reagent was used for the gravimetric determination of ammonium perchlorate by Glover and Rosen (14). Another approach to the use of this reagent for perchlorate determination, herein proposed, incorporates the use of conductance for the titration of perchlorate with tetraphenylarsonium chloride.
EXPERIMENTAL Reagents. Tetraphenylarsonium chloride (J. T. Baker Chemical Co.) and teixaphenylarsonium chloride hydrochloride (Eastman Organic Chemical Co. and Hach Chemical
(1) N. L. Coump and N. C. Johnson, ANAL.CHEM.,27, 1007 (1955). (2) A. de Sousa, Chemist Aizalyst, 49, 18 (1960). (3) A. de Sousa, Anal. Chinr. Acta, 24, 424 (1961). (4) G . P. Haight, Jr., ANALCHEM., 25, 642 (1953). (5) E. Kruz, G. Kober, and M. Berl, ANAL.CHEM., 30,1983 (1958). (6) E. A. Burns and R. F. IvIuraca, Ibid.,32, 1316 (1960). (7) D. C. Eagles, Chem. Ind. (London),33, 1002 (1954). 21, 1313 (1949). (8) W. W. Brandt and G. F. Smith, ANAL.CHEM., (9) W. Bodenheimer and H. Weiles, Ibid.,27, 1293 (1955). (IO) J. S. Fritz, J. E. Abbink, and P. A. Campbell, Ibid., 36, 2123 (1964). (11) H. Nezu, Bunsaki Kqaku, 10, 561 (i961); Chem. Abstr., 56, 26f (1962). (12) M. D. Morris, ANAL.CHEM., 37, 977 (1965). (13) H. H. Willard and G. 111. Smith, IND.ENG.CHEM., ANAL.ED., 11, 186 (1939). (14) D. J. Glover and J, M. Rosen, ANAL.CHEM., 37, 306 (1965).
Co.) were used to prepare the titrant solutions. Ammonium perchlorate, Fisher certified reagent, was used for the conductometric standardization of the titrant. Ammonium perchlorate samples were obtained from the American Potash Co. All other chemicals were of analytical reagent grade, and were used without purification. Apparatus. A Metrohm Potentiograph, Type E-336, with conductance assembly, Type E-165, was used for recording conductance curves and performing titrations. A dip-type, platinized, platinum electrode having a cell constant of 0.7 cm-1 was used for conductance measurements. Procedure. Tetraphenylarsonium chloride, O.O5M, was prepared by dissolving tetraphenylarsonium chloride in an appropriate amount of distilled water. The pH of the solution was adjusted to 7 with dilute HCl or NaOH. The titrant was also prepared from tetraphenylarsonium chloride hydrochloride by passing a concentrated aqueous solution of the reagent through an ion exchange column, containing Dowex 1-X4 resin activated to the hydroxide with 2 5 x (w./v.) NaOH. The eluted tetraphenylarsonium hydroxide was converted to the chloride upon neutralization of the solution with dilute HC1. The solution was then diluted to an appropriate volume with distilled water. The perchlorate sample (or standard perchlorate salt) to be titrated is dissolved in 200 ml of distilled water. The sample should contain about 1 to 1.5 mmoles of perchlorate. If necessary, the pH of the solution is adjusted to about 7 with dilute HC1 or NaOH. The sample is then titrated with 0.05M (CGHj)IAsClat a rate of about 2 ml per minute. For manual titrations there is a waiting period of about 30 seconds between incremental additions of titrant for stabilization of conductance readings before recording data. The indicated times will, of course, vary with stirring rate.
RESULTS AND DISCUSSION Conductance titrations were initially performed in acidic solutions, in keeping with the potentiometric back-titration method of Willard and Smith (13) and the gravimetric method of Glover and Rosen (14). However, the titration curves, prior to the equivalence point, were found to be nonlinear, making it difficultto determine the end point readily. Curves were also of decreasing conductance in this area, which was not expected. Both the decrease in conductance and curvature were attributed to a dilution effect. Preliminary titrations at a neutral pH gave curves that were linear with a slight increase in conductance prior to the end point, and thus appeared to be satisfactory for quantitative work. Neutral and acidic titration curves are compared in Figure 1. By proper adjustment of concentration of reagent and sample solutions, it was possible to keep the conductance of the solution fairly constant before the end point. These adjustments were not critical and were easily satisfied by using a 0 . 0 5 ~ titrant 4 and a sample solution that was 0.005 to 0.0075M VOL. 39, NO. 1, JANUARY 1967
93
in perchlorate. Extrapolations to intersection of the linear segments of curves thus obtained included angles of less than 120°, which made the end point intersection readily discernible. Standardization. Two methods were tried for reagent standardization. One employed a gravimetric approach, while the other was volumetric. The gravimetric method yielded a lower molarity and was considerably less precise than the volumetric technique. The coefficient of variation was 0.30 and 0.06, respectively, for the two methods. The gravimetric standardization was essentially the method of Glover and Rosen (14). However, the precautions required in handling the precipitate and the relatively poor precision of the results suggested problems with solubility of the precipi-
I
I
IO
I
I
20
30
T I T R ~ N TY D L U M ~ I ~ I J
Table 1. Recovery of Ammonium Perchlorate in 1 to 1 Salt Mixtures Ammonium Ammonium Salt perchlorate perchlorate Salt added, gram added, gram recovered, gram 0.1495 0.1501 0.1506 KCl 0.1525 0.1501 0.1501 KC103 0.1499 0.1501 0.1509 KBr 0.1516 0.1501 0.1505 KBrOa 0.1500 0.1501 0.1506 KNOI 0.1520 0.1511 0.1514 KzCrOa
Table 11. Precision Data for Determination of Purity of Perchlorate Salts Mallinckrodt American Potash NHK104 reagent grade Lot 90-2 Lot 9015-00006 KC104, Lot 7053 99.41 99.69 99.41 99.75 99.83 99.41 99.62 99.89 99.27 99.32 99.95 Av. 99.37z 99.72z Over-all standard deviation = 0.074 Over-all 95% c.1. = A0.16z
99.89Z
Table 111. Comparison with Other Methods in Area of Recovery Z NHaClOd Conductometric titration with Titanous Nonaqueous tetraphenylchloride arsonium titration of reduction (6) ammonium ion chloride Ammonium 100.3 100.4 perchlorate 99.98 100.0 100.6 Sample A 100. 1 100.2 100.8 99.93 100.2 loo. 8 99.86
x= &0.07 99.97 S =
Ammonium perchlorate Sample B
100.6 k0.2
100.2 100.0 100. 1 100.1
99.91 99.64 99.68 99.55
x= 100.1 s *0.1
100.2 100.2 100.0 100.1 -
99.69 +O. 15
100. 1
=
94
100.2 f O .2
ANALYTICAL CHEMISTRY
Figure 1. Comparison of titration curves obtained from acidified and neutral ammonium perchlorate solutions with 0.044M tetraphenylarsonium chloride A . Titration of 200 ml of 0.006M ammonium perchlorate, 0.01M in HCl, illustrating dilution effect B. Titration performed on 200 ml of 0.006M ammonium perchlorate in absence of other ions
tate. To determine the solubility of the precipitate, an ultraviolet spectrophotometric method was developed, based on the absorbance of the phenyl rings. The absorbance of a saturated solution of tetraphenylarsonium perchlorate at 25 O C was measured at 270 mp. Precipitated tetraphenylarsonium perchlorate was collected on a medium-porosity sintered glass crucible, washed with distilled water, and added to distilled water at 80" C. The mixture was then stirred for one hour and allowed to stand at 25' C overnight. The absorbance of this solution was compared to those of solutions of known concentration in tetraphenylarsonium chloride in the range of 2.48 X to 1.98 x 10-4M. The K,, for the perchlorate salt from these results was approximated at 1.2 X lo+. This is comparable to the K., of 3.5 X 10+ reported by Morris (12) for the analogous stibonium compound. The volumetric standardization was performed by the procedure for the sample titration, using highly pure ammonium perchlorate as the primary standard. Because of the better precision and the fact that this method more closely approximated sample titration conditions, it was given preference. Interferences. Previous investigators (13, 14) had reported a lack of specificity of the tetraphenylarsonium ion for perchlorate. Interferences from a variety of ions were noted. However, these studies were conducted in an acidic medium. The present investigation was carried out in a neutral solution and no interference resulted from the presence of the several ions studied except iodide and permanganate. Table I shows data on the recovery of perchlorate from approximately 1 to 1 mixtures of ammonium perchlorate and various salts. In addition, semiquantitative work showed no interferences from iron(III), sulfate, iodate, fluoride, molybdate, and phosphate. Precision. Titration data are given in Table I1 for replicate analysis of three samples: ammonium perchlorate from two lots of material purchased to a federal specification (15) and certified to have a minimum ammonium perchlorate content of 99a/o, and reagent grade potassium perchlorate. The precision of the method expressed as the 95% confidence ljrnit calculated from the data was *0.16x:.
zko. 1
(15) MIL-A-192A.
A direct comparison of this method with several other techniques for the analysis of ammonium perchlorate in given in Table 111. The titanium trichloride determination is the method of Burns and Muraca (6). The third method, limited to the determination of ammonium perchlorate, is based on the volumetric measurement of the ammonium ion with alcoholic potassium hydroxide in a nonaqueous solvent system. From these data it can be seen that the proposed conductometric titration is competitive with the other two methods in recovery of ammonium perchlorate. Other Applications of Reagent. The interference from permanganate in the determination of perchlorate was found
to be quantitative. Thus, the procedure given above can be used for the determination of permanganate in the presence of a variety of ions. This should be particularly advantageous for permanganate solutions which contain ions that interfere with spectrophotometric and/or redox methods of determining this ion. Iodide also interfered with the perchlorate determination. However, the precipitation of iodide with tetraphenylarsonium chloride was not quantitative in either neutral or slightly acidic solutions. No further attempts to make the reaction quantitative were made. RECEIVED for review June 9,1966. Accepted October 10,1966.
Amperometriic Titration of Lead, Cadmium, and Zinc in 2-PropanoI with (1,2-Cycl ohexa nedinitriIo)tet raacetic Acid Paul Arthur and Brenda R. Hunt1 Department of Chemistry, Oklahoma State Unicersity, Stillwater, Okla.
IN AQUEOUS SOLUTIONS, chelometric titrations have been employed with great success and chelating agents, particularly ethylenediamine-tetraacetic acid (EDTA) and its analogs, have been so extensively and commonly used that even elementary texts in analytical chemistry include applications of one or more of these reagents. The usefulness of such titrations would be greatly extended if they could be performed in suitable completely nonaqueous solutions ; for there are many situations in industries dealing with petroleum, petroleum products, paints, fats arid oils, etc., as well as in research where need arises for the quantitative determination of metals in substances which are immiscible with water. Apparently, however, very few chelonietric titrations in nonaqueous solutions have been described. Brummet and Hollweg ( I ) determined nickel(II), copper(l;I), and cobalt(I1) in 20 methanolbenzene by potentiometrically titrating the hydrogen ion freed by the reaction of these metals with dimethylglyoxime, with 8-quinolinol, and with 1-nitroso-2-naphthol. A photometric titration with dithizone was used by Marple, Matsuyama, and Burdett (2) for the determination of zinc in lubricating oils, the sample being dissolved in an 8 methanol-benzene mixture, while spectropliotometric titrations of nickel(I1) and cobalt(I1) with dimt.thylglyoxime in a 1 :1 mixture of chloroform and 2-propanol were described by Behm and Robinson (3). Spectrophotometric methods were used also for titrations in dimethylformamide, by Takahashi and Robinson ( 4 ) for the determination of copper(I1) and nickel(I1) with 1-nitroso-2-napthol and by Boyle and Robinson (5) for the determination of zinc, copper(II), cadmium, lead, and nickel(I1) with 8-quinolinol. EDTA-type chelating agents apparently have not been employed in nonaqueous titrations. Of the reasons for this, two Present address, Dow Chemical Co., Freeport, Texas.
are undoubtedly outstanding: the solubilities of these reagents are extremely low in most (if not all) common organic solvents, and little is known of either the stabilities or the solubilities of their metal chelates in organic solvents. In the work reported here, 2-propanol was selected as the principal solvent because it has been found in our laboratory (6) that this solvent, either alone or mixed with equal volumes of petroleum or any of several petroleum products tested, can be used readily for polarographic studies and amperometric titrations. EXPERIMENTAL
Reagents. All reagents used were reagent grade with exceptions noted below. Unless otherwise specified, all were used without further purification. 2-Aminoethanol, Eastrnan White Label, was distilled in nitrogen and stored under nitrogen. (1,2-Cyclohexanenedinitrilo)tetraacetic acid was Matheson, Coleman and Bell practical grade. Diethylenetriaminepentaacetic acid was K. and K. Laboratories, Inc. reagent. 2-Propanol, Fisher Scientific Reagent grade, was distilled before use. Apparatus. A Sargent Model XXI polarograph equipped with a Model A IR-Compensator was used in this research. The cell was like that described by Arthur and VanderKam (7) except that the bridging tubes connecting the DME compartment to the reference electrodes had asbestos fibers at each end and had filter tubes so any desired electrolyte solution could be used in the bridges. With this arrangement, if aqueous calomel electrodes were used, clogging of the asbestos fibers with potassium chloride precipitated by organic solvents could be avoided by using aqueous lithium chloride as the bridging solution; also diffusion of substances from the reference electrodes into the DME compartment was essentially prevented. In the work reported here, the bridging tubes were filled with 0.1M lithium chloride in 2-propanol. Until, as was learned early, it was found that small amounts of water did not apparently bother, this solu-
(1) B. D. Brummet and R. M. Hollweg, ANAL.CHEM.,28, 448 (1956).
( 2 ) T. L. Marple, G. Matsuyama, andL. W. Burdett, ANAL. CHEM.,
(6) P. E. Moran, “Amperometric Titrations of Selected Petroleum
30,937(1958). (3) R . K. Behm and R. J. Robinson, Ibid., 35, 1010 (1963). (4) I. T. Takahashi and R. J. Robinson, Ibid., 32, 1250 (1960). ( 5 ) W. G. Boyle, Jr., and R. J. Robinson, Ibid., 30, 958 (1958).
Additives and Sulfonic Acids in Nonaqueous Solutions,” Ph.D. Thesis, Oklahoma State University, 1959. (7) Paul Arthur and R. H. VanderKam, ANAL.CHEM.,33, 765 (1961). VOL. 39, NO. 1, JANUARY 1967 e
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