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(5) Leisey, F. A., Ibid., 26, 1607 (1951). (6) Lingane, J. J., Ibid., 26, 622 (1954). (7) Meier, D. J., Myers, R. J., and Swift, E. H., /. Am. Chem. Soc.,
(11) Wooster, W. S., Farrington, P. S., and Swift, E. H., Ibid., 21, 1457 (1949).
71, 2340 (1949). (8) Myers, R. J., and Swift, E. H., Ibid., 70, 1047 (1948). (9) Rowley, Keith, and Swift, E. H., Anal. Chem., 26, 373 (1954). (10) Sease, J. W„ Niemann, C., and Swift, E. H., Ibid., 19, 197 (1947).
Received for review April 1, 1955. Accepted June 29, 1955. Presented in part before the Division of Analytical Chemistry at the 123rd Meeting of the American Chemical Society, Los Angeles, Calif., March 1953. Contribution 1981, Gates and Crellin Laboratories of Chemistry, California Institute of Technology, Pasadena, Calif.
Precise Assay of Trichloroacetic Acid by Coulometry at Controlled Potential THELMA MEITES and LOUIS
MEITES1
Sterling Chemistry Laboratory, Yale University, New Ha\
The reduction of trichloroacetate ion from an ammoniacal medium proceeds quantitatively to dichloroacetate ion at a mercury cathode whose potential is maintained at a constant suitable value. Integrating the current which flows during such an electrolysis of trichloroacetate in pure permits the determination samples or in the presence of at least 8 times as much dichloroacetate, with an accuracy and precision of within ±0,2%. and dichloroacetic acids are universal contaminants acid, and no method of assaying trichloroacetic acid (4, 9) has yet been proposed, which is free from interference by the lower chlorinated acids. Such a method is presented in this paper. It is based on the quantitative electroreduction of trichloroacetate ion to dichloroacetate ion in an ammoniacal medium, which proceeds with 100% current efficiency at a mercury cathode at a suitably chosen potential. An integration of the current flowing during this reduction gives the amount of trichloroacetate present in the sample to ±0.2% of trichloroacetic MONO-
better. The reduction of the chloroacetic acids at a mercury cathode has been studied several times by polarographic techniques. Elving and Tang (2. 3) showed that dichloroacetate ion in ammonical media gives a single wave at a fairly negative potential, corresponding to the reaction CLCHCOO- + H20 + 2e ^ ClCH2COO- + Cl~ + OH" or
Trichloroacetate ion under the same conditions gives a double The second wave is identical with the single dichloroacewave: tate wave, while the first wave represents the reduction of trichloroacetate to dichloroacetate. This first wave, which is well defined, was used by Elving and Tang (2) for the polarographic determination of trichloroacetate in the presence of dichloroacetate; of course this is insufficiently accurate for assay purposes. Elving and Tang (3) found that monochloroacetate ion was not reducible from an ammoniacal solution. Neiman, Ryabov, and Sheyanova (8), on the other hand, asserted that monochloroacetate does give a single wave in 0.13/ sodium hydroxide, and that dichloroacetate and trichloroacetate correspondingly give two and three waves, respectively. This the present authors were unable to confirm; the polarographic characteristics of the chloroacetic acids in sodium hydroxide media do not differ in any significant respect from those found in ammoniacal media by Elving and Tang (3). From this information it appeared possible to select conditions under which trichloroacetate ion alone could be reduced, and quantitatively, and to apply this to the coulometric determination of trichloroacetate, either alone or in mixtures with dichloroacetate. Present address, Department of Chemistry, Polytechnic Institute Brooklyn, 99 Livingston St., Brooklyn 1, N. Y. 1
of
i, Conn.
EXPERIMENTAL
The double diaphragm cell, the potentiostat, and the current integrator, which are manufactured by Analytical Instruments, Inc., Bristol, Conn., have been described (6), as has the recording polarograph used (7). All weights and volumetric apparatus were carefully calibrated by conventional techniques. The chloroacetic acids were secured from a commercial supplier. The trichloroacetic acid was dried for a week over anhydrous magnesium perchlorate. Titration of a sample of the dried acid with sodium hydroxide standardized against potassium hydrogen phthalate gave (on the assumption that no lower chlorinated acid was present) an “assay” of 99.98%. Stock solutions of the acid appeared to be stable indefinitely when ammonium and potassium chlorides were present alone, but when ammonia was also present the values secured by coulometric analysis decreased at the rate of about 1% per week. The kinetics of this reaction have been studied by Verhoek (10). Some of the mixtures of di- and trichloroacetic acids were prepared from the commercial dichloroacetic acid. Analysis of this material, either polarographically (2) or by the coulometric procedure described below, indicated the presence of roughly 0.3% trichloroacetic acid. The necessity of correcting for this amount of impurity would have severely limited the accuracy attainable with the mixtures containing much dichloroacetic acid. Consequently, a pure dichloroacetate solution was prepared by the following procedure. A solution of 4.6 grams of trichloroacetic acid in the ammoniacal supporting electrolyte used throughout the work was electrolyzed at a mercury cathode at —0.8 volt vs. S.C.E. until the current had fallen to zero, and the resulting solution was transferred to a 250-ml. volumetric flask and diluted to the mark with the stock supporting electrolyte. Whereas ammoniacal solutions of the commercial dichloroacetic acid rapidly yellowed on standing, the electrolytically prepared solution appeared stable for many months. RECOMMENDED
PROCEDURE
Prepare a stock solution containing 2.53/ ammonia, 13/ ammonium chloride, and 23/ potassium chloride. (The potassium chloride serves primarily to decrease the cell resistance; and none of the concentrations is in any way critical.) Dissolve a sample of trichloroacetic acid weighing 0.03 to 5 grams in a little water and add about 40 ml. of the supporting electrolyte solution. If the weight of the sample exceeds 1 gram, enough additional ammonia should be added to restore that lost by
neutralization. Fill the auxiliary electrode and central compartments of a double diaphragm cell for controlled potential electrolysis in the manner described previously (6), and transfer the trichloroacetate solution to the working electrode compartment. Since the time required to complete the electrolysis is proportional to the volume of the solution (5), not more than an additional 50 ml. of the supporting electrolyte should be used in completing the transfer. Pass a stream of tank nitrogen or hydrogen through an efficient gas washing bottle filled with the supporting electrolyte and containing 1 to 2 grams of hydrazine dihydrochloride to facilitate the removal of oxygen, and thence into the solution in the working electrode compartment of the cell. Since oxygen is reducible at a mercury cathode under the conditions used in this procedure, it is essential to remove all dissolved air before the electrolysis is begun. This normally requires about 5 minutes, provided that a sufficiently rapid stream of gas is used and the solution is well stirred. Add about 30 ml. of pure mercury, read the coulometer register,
ANALYTICAL
1532
and adjust the potentiostat to maintain the potential of the mercury electrode at —0.9 volt vs. S.C.E. Do not disconnect the gas stream until the electrolysis is completed. The electrolysis may now be allowed to proceed unattended. After 60 to 90 minutes the current will fall to zero and the integrator will stop. Subtract the initial from the final register reading to give directly the number of milliequivalents of trichloroacetic acid. (The integrator’s 1-ohm resistor is always used unless the weight of trichloroacetic acid present is less than about 0.05 gram.) Since two faradays are consumed in the reduction of each mole of trichloroacetic acid, the equivalent weight of the acid is 81.70 grams. After the completion of the electrolysis, discard the solution and rinse the cell thoroughly. The mercury may be re-used many times without purification. RESULTS AND DISCUSSION
The half-wave potentials corresponding to the reactions
ClsCCOO- + H,0 +
2e
ChCHCOQ- + OH" + Cl~
-*
the optimum potential at which to conduct the electrolysis is difficult than it would be if the waves were reversible. Normally one would anticipate that a potential of about 1.2 volts vs. S.C.E.. midway between the two half-wave potentials, would effect complete reduction of trichloroacetate without initiating Reaction 2. This is not the case, however, for some dichloroacetate is reduced at this potential during the latter stages of the electrolysis, and the value found is a few tenths of a per cent too high (Table I). At —1.3 volts the reduction of dichloroacetate is even more extensive, and the error becomes about +2%. Even at —1.1 volts one may be on dangerous ground in dealing with solutions containing much dichloroacetate for, as is well known, the potential at which a wave begins becomes more positive as the concentration of the reducible substance increases. more
—
(1)
Table II.
and
ChCHCOO- + H20 +
2e
—
ClCH2COO" + OH" + Cl"
(2)
—0.73 and —1.65 volts, respectively, in the supporting electro(In the absence of the potassium chloride they are —0.93 and —1.62 volts; thus the potassium chloride also serves the unexpected purpose of providing a better separation of the waves.) Since both of the waves are irreversible, and are
lyte recommended.
therefore
cover
a
considerable range of potentials, the selection of
Cathode Potential, Volts S.C.E. -1.00 -0.90 -1.00 -0
90
-1.00 -0.90 -1.00 -0.90
Determination of Trichloroacetic Presence of Dichloroacetic Acid Trichloroacetic
Dichloro-
Taken, Grams
Added, Grams 0.0125 0.0286 0.0250
0.2994 0.3773
0.2994 0.2995 0 3773
0.2996 0.3773 0.3773 0.2991
0.3773 0.3773 0.3773 0.3773
Table I. Cathode Potential, Volts vs. S.C.E. -0.80 -0.90
Assay of Trichloroacetic
Trichloroacetic Acid Taken,
Trichloroacetic Acid Found,
0.03762
0.03763 0.03760 0.03760 0.07364 0.07320 0.07337 0.07315 0.1505
Grams 4.5894
0.07369
0.15096
0.3649
0.3773 0.7575 1.5150 1.5183 5.6928 3.9976
Grams 4.5855
-1.00
0.2993
99.95 99.93 99.33 99.56 99.26
99.70 99.56 99.63 99.56 99.56
99.70 99.62
100.16 99.58 99.92 99.87 100.11 99.67 99.54 99.66 99.72 99.74 v.
0.2993 0.2990 0.2993
99.90
0.2992
99.97
0.3628 0.3638 0.3629 0.3638 0.3633 0.3635 0.3634
99.42
Average of 12 values at —1.00 -1.10 -1.20
-1.30
0.2993 0.2993 0.2993
99.72 ± 0.22 (standard deviation) 100.00 100.00 99.93
0.2991
0.3649
Assay, % 99.91 99.95
0.1503 0.3638 0.3635 0.3655 0.3757 0.3770 0.7565 0.7583 0.7550 1.508) 1.5132 3.6823 3.9870
Average of 24 values at —0.90
Acid
100.03
0.1503 0.1504 0.1503
0.2993
0.3002 0.3056 0.3050
99.70
99.45 99.70 99.56 99.62 99.59 v.
99.74 ± 0.21
(standard deviation) 100.00 100.30 102.1 101.9
CHEMISTRY
0.1501
acetic
0 0376
0.0718 0.0627 0,108 0.145
0.125 0.217
0.362 0.579 0.868 1.16
Trichloroacetic Acid Found, Grams 0.3005 0.3770 0.2986
0 2992
0,3767
0.2988 0.3766
0.3766 0.2976 0.3764
0.3767
0.3763
0.3770 0.1500 0.1500
Acid in
Assay, %
100.37 99.92 99.73 99 90
99.85 99.73 99.81 99.81 99.50
99.76 99.84 99.73 99.92 99.93 99.93
Average 99.85 ± 0.18 (standard deviation)
As the averages of the 24 values secured at —0.90 volt and of the 12 values secured at —1.00 volt differed by only 0.02%, it seems evident that electrolysis at either potential suffices for the complete reduction of trichloroacetate without any danger of reducing dichloroacetate. Table II shows this to be true even when the original solution contains 8 times as much dichloroacetate as trichloroacetate. Though samples poorer in trichloroacetate than this would probably be most efficiently analyzed by the procedure of Elving and Tang (3), there is no reason to believe that a coulometric analysis conducted at, say, —0.90 volt, would not provide a more accurate result if the need arose. Since a considerable excess of dichloroacetate does not measurably affect the results, it is equally clear that the presence of substantial amounts of monochloroacetate would also be without effect. The average of the 36 values at these two potentials, 99.73 ± 0.22% (standard deviation), together with the result of the alkalimetric “assay” previously described, indicates that the dried trichloroacetic acid contained approximately 0.25% lower chlorinated acids. Though this seems to be entirely credible, its confirmation by any other method would be extremely difficult. Of special interest from a theoretical standpoint is the result of the electrolysis at —0.80 volt shown in Table I. This is only 0.07 volt more negative than the half-wave potential of trichloroacetate, and the current measured at this potential with a dropping electrode is barely two thirds of the diffusion current, yet the reduction at the large electrode proceeded to completion. Of course an electrolysis at so low a potential requires considerably The fact that more time than does one at —0.9 or —1.0 volt. quantitative reduction is eventually secured merely reflects the chemical irreversibility of the electrode reaction. At —0.8 volt one is dealing, not with an equilibrium mixture of di- and trichloroacetate at the surface of the electrode, but with a
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reduction which simply proceeds at a lower rate than it would at a higher potential. Though these considerations were outlined by Delahay (1), this is believed to be their first practical application.
(3) Elving, P.
LITERATURE
CITED
(1) Delahay, P., “New Instrumental Methods of Electrochemistry,” p. 277, Interscience, New York-London, 1954. (2) Elving, P. J., and Tang, C.-S., Anal. Chem., 23, 341 (1951).
Chem. Soc., 72, 3244
(4) Lalic, M. A., and Canic, V. D., Bull.
soc. chim. Belgrade, 14, 111 (1949). (5) Lingane, J. J., J. Am. Chem. Soc., 67, 1916 (1945). (6) Meites, L., Anal. Chem., 27, 1116 (1955). (7) Meites, L., “Polarographic Techniques,” p. 16, Interscience,
ACKNOWLEDGMENT
It is a pleasure to thank Analytical Instruments, Inc., of Bristol, Conn., for the grant-in-aid which made this work possible.
j., and Tang, C.-S., J. Am.
(1950).
New York-London, 1955. B., Ryabov, A. V., and Sheyanova, E. M., Doklady Akad. Nauk S.S.S.R., 68, 1065 (1949). (9) Rosin, J., “Reagent Chemicals and Standards,” p. 457, Van Nostrand, New York, 1937. (10) Verhoek, F. H., J. Am. Chem. Soc., 67, 1062 (1945). (8) Neiman, .
Received for review March 29, 1955. Accepted May 3, 1955. Contribution No. 1330 from the Department of Chemistry of Yale University.
Application of Organic Solvent Extraction to Flame Spectrophotometry Determination of Iron in Nonferrous Alloys JOHN A. DEAN
and J.
HAROLD LADY1
Department of Chemistry, University of Tennessee, Knoxville, Tenn.
This investigation describes an application of organic to flame spectrophotometry. Iron solvent extraction was selectively extracted from aqueous solutions with acetylacetone, which served as chelating agent and solvent. The extract was aspirated directly into an oxyaeetylene flame. This procedure circumvents many spectral and radiation interferences encountered in iron determinations by flame spectrophotometry and provides a superior aspirating medium compared to an of aqueous solution containing variable concentrations diverse elements. Acetylacetone contributes greatly to the size of the flame and increases sixfold the luminescence of the 372-mg iron line. No interferences were found when this method was applied to aluminum-, Replicopper-, and nickel-base alloys and limestone. cate samples show a standard deviation of 3%. present investigation resulted from the need for a faster the flame spectrophotometric determination of iron. In a preceding study a considerable number of elements were found either to interfere directly or to affect the luminescence from the 372.0-, 373.7-, 374.7-, or 386.0-mg flame emission lines of iron (5). Manganese emits a series of weak bands in the region 363 to 400 mg, and magnesium possesses three intense band systems whose heads appear at 372, 375, and 385 mg. In addition to these coincidences, interference with the excitation process has been observed for aluminum, zinc, the alkaline earths, lithium, and potassium. Through the addition of cobalt, as an internal standard, interference from aluminum, zinc, and lithium was circumvented. However, difficulties due to magnesium, manganese, and the alkaline earths could be circumvented only if the iron were first separated as the hydrous oxide or by some other chemical method. Although solvent extraction is recognized as a powerful method for accomplishing analytical separations, it has not yet been exploited in flame spectrophotometry. The solvent extraction of iron as the iron acetvlacetonate chelate is a convenient and rapid method for isolating iron from other elements encountered in nonferrous alloys and limestone. Iron is 50% extracted at a pH of zero (8). Thus several extractions with iron acetylacetone (2,4-pentanedi one) quantitatively remove
method for THE
1
Pa.
Present address, Westinghouse Research Laboratories, East Pittsburgh,
Figure
1.
Flame emission spectra of iron
Present 65 p.p.m.
iron.
Slit width
0.030
mm.
from can
a l.V acid solution. As a result, the extraction of iron be accomplished in the presence of metals which hydrolyze
at higher pH values. The use of acetylacetone, as both solvent and reagent for iron, offers several other advantages. Iron acetvlacetonate is very soluble in acetylacetone, soluble to an extent that macroseparations are possible. The solvent is readily available at reasonable cost and is combustible itself, thus contributing to the degree of excitation of the iron lines. EXPERIMENTAL
WORK
Apparatus. A Beckman Model DU spectrophotometer with Model 9220 flame attachment and photomultiplier unit was used. A metal atomizer-burner unit, supplied with the flame attachment, was used as the excitation source. Oxygen and acetylene were the gases used. The wave-length knob on the spectrophotometer was replaced by a 4 to 1 gear-reduction knob to facilitate positioning the wave length dial. The iron lines are sharp, arc-emission lines. One must move slowly back and forth over the peak of each line to ascertain
