ANALYTICAL CHEMISTRY
846 the distribution of a compound within a more insoluble solid as exemplified by the present study. By taking as the most representative values, those distribution coefficients obtained from runs in which 0.6 or 0.8 ml. of methyl sulfate was used, an average value of X = 0.030 5 0.004 is obtained. A number of authors (8, 21,25) have suggested that a relationship exists between solubilities and distribution coefficients. While solubility data are not available for barium and strontium sulfate in an ionic environment such as the methanol-water medium used here, the aqueous solubilities of these sulfates, as a function of temperature, are available (14, 18). From these data, the solubility products a t 83” C. were calculated to be 2.5 x 10-10 for barium sulfate and 8.8 X lo-’ for strontium sulfate. The solubilities of barium sulfate and strontium sulfate, in moles per 100 ml., are 1.6 X and 9.4 X 10-6, respectively. It is seen that the ratio of the solubility products, 2.8 X differs from the observed heterogeneous distribution coefficients by a factor of approximately 100. The ratio of solubilities, 0.017, on the other hand is a t least of the same order of magnitude as the distribution coefficient experimentally obtained. Actually, with the exception of the radium-barium sulfate system, there is little evidence for the existence of a simple relationship between the solubility ratio of salts and the observed distribution coefficients. CONCLUSIONS
A study of the distribution of strontium between the liquid phase and barium sulfate precipitated by several methods has shown that a homogeneous precipitation method more nearly approaches equilibrium formation of barium sulfate than does the conventional sulfate precipitation method. Strontium sulfate-barium sulfate in 20% methan01-80% water medium was observed to be a derichment system in which the strontium appears to be heterogeneously distributed throughout the solid phase. At 83” C., for initial barium t o strontium concentration ratios varying from 1.3 to 2700, the logarithmic distribution coefficient was found to be 0.030 i 0.004. Derichment systems of this type present many more experimental difficulties than do enrichment systems.
ACKNOWLEDGMENT
The authors are indebted to H. W. Kirby, Mound Laboratory, Monsanto Chemical Co., for his assistance during the radiochemical investigations. LITERATURE CITED
Biedermann, W., and Schwarzenbach, G., C h i m i a , 2, 56 (1948). Bolton, B., and Kelley, W., M.S. thesis, Syracuse University, 1947.
Banner, H. A . , and Kahn, M., Nucleonics, 8, No. 2 , 4 6 (1951). Chlopin, V., Trav. inst. &at r a d i u m (U.S.S.R.),4 , 3 4 (1938). Downer, H., and Hoskins, W., J . Am. Chem. SOC.,47, 662 (1925).
Elving, P., and Van A t t s , R., ANAL.CHEM.,2 2 , 1 3 7 5 (1950). Fischer, R., Ibid., 23, 1667 (1951). Flood, H., 2. anorg. u. allgem. Chem., 2 2 9 , 7 6 (1936). Gordon, L., . ~ N A L .CHEM.,2 4 , 4 5 9 (1952). Grahmann, W., Y e u e s Jahrb. Mineral. Geol., 1, 1 (1920). Hahn, O., “Applied Radiochemistry,” Ithaca, Cornell University Press, 1936. Hahn, O., Ber., 59B, 2014 (1926). Henderson, L., and Kracek, F., J . Am. Chem. Soc., 49, 738 (1927).
International Critical Tables, Val. 6, p. 256, New York, McGraw-Hill Book Co., 1929. Kirby, H., A N ~ LCHEY., . 24, 1678 (1952). LaLIer, V., and Dinegar, R., J . Am. Chem. Soc., 73, 380 (1951). RIanns, T., et aZ., ANAL.CHEM.,24, 908 (1952). Mellor, “Comprehensive Treatise on Inorganic and Theoretical Chemistry,” Vol. 3, London, Longmans, Green and Co., 1923. JIerkulova, XI.,Tral;. inst. itat r a d i u m (U.S.S.R.), 3, 141 (1937).
ilIumbrauer, R., 2. physik. Chem., A156, 113 (1931). Ratner. A, J . Chem. P h y s . , 1 , 7 8 9 (1933). Riehl, K.,2. p h y s i k . Chem., A177, 224 (1936). Riehl, N., and Kading, H., Ibmd., A149, 180 (1930). Salutsky, lI.,Stites, J. G., and Martin, A. W., ANAL.CHEM., 25, 1677 (1953).
Shulyatikov, B., J . P h y s . Chem. (U.S.S.R.),2 1 , 9 7 5 (1947). Willard, E., ASAL CHEM., 22, 1372 (1950). RECEIVED for review August 22. 1953. Accepted February 19, 1954. Abstracted from the thesis submitted by Carl C. Reimor to the Graduate School of Syracuse University in partial fulfillment of the requirements for the degree of doctor of philosophy, July 1953. Study supported in part by g r a n a from the Btomic Energy Commission and from the Research Corp. One of the authors (C.C.R.) was the holder of 8 Charles A. Coffin Fellowship (General Electric Co.) in 1951-52 and the Procter and Gamble Fellowship in 1952-53.
Recrystallization of Barium Sulfate from Fused Salts MORRIS GALLANT, GEORGE J. SCHMITT, and JOSEPH STEIGMAN Department o f Chemistry, Polytechnic Institute of Brooklyn, Brooklyn, N. Y. The usual method for sulfate determination consists of the precipitation of the barium salt. The latter strongly coprecipitates many ions, and most variations of the method rely upon either counterbalancing errors or empirical weight calibrations. This investigation has shown that when appropriate precautions are taken, coprecipitated sodium and chloride ions can be quantitatively removed by one fusion of the original precipitate with alkali chlorides, allowing it to resolidify, leaching with barium chloride solution, washing, filtering, and igniting. The necessary precautions include the purification of the alkali chlorides and barium chloride and the use of an excess of barium chloride in the solution used for washing the recrystallized precipitate. The procedure is shown to be applicable to the determination of sulfate in the presence of a large excess of sodium chloride. On the other hand, nitrate is much
more difficult to remove by this procedure. The extremely strong tendency of this ion to form solid solutions in barium sulfate is demonstrated by introducing it into the precipitate from an alkali chloride melt.
0
NE of the important inorganic analyses is the estimation of soluble sulfates. The most widely used method is the
precipitation of the insoluble barium salt and its gravimetric determination. In many situations the accuracy is poor. Other gravimetric methods have been proposed, such as the precipitation of the lead, benzidine, or complex cobalt sulfates. These compounds, while free from certain disadvantages of the barium salt, are in general too soluble for many applications. The overwhelming bulk of sulfate analyses is still performed by precipitating, igniting, and weighing the barium salt. There are many researches in the literature, too numerous t o
V O L U M E 26, NO. 5, M A Y 1 9 5 4 cite here, which describe the well-known tendency of barium sulfate to bring down other ions during precipitation-the so-called coprecipitation phenomenon. Among the anions which coprecipitate are iodide, bromide, chloride, chlorate, nitrite, and nitrate, in increasing order ( 7 ) . Cation coprecipitation has been found for lithium, sodium, potassium, copper, cadmium, manganese, ferric iron, and aluminum (6). The actual proportion of contaminant depends in a complex fashion upon concentration, acidity, temperature, speed and order of reagent addition. digestion, heating, and other variables. The precipitation of barium sulfate from aqueous solution is an extremely complicated phenomenon. Among other things, the precipitation in the presence of sodium chloride results in the appearance of sodium sulfate in the solid (6). In addition, oxonium ion is known t o be formed in the lattice ( 1 ) and is balanced by bisulfate, chloride, or some other anion. The careful work done by Walden and his collaborators (1, 11, 12) with both anions and cations showed, by means of x-ray diffraction, that solid solutions were formed by the incorporation of ions like nitrate, permanganate, perchlorate, and sodium. A reasonable conclusion which can be drawn from their work is that barium sulfate which is precipitated from water solution is not a pure compound slightly contaminated by accidental impurities, but is rather one or more of a large number of inevitable solid solutions. The methods which have been suggested to minimize or eliminate coprecipitation have, in the main, operated on the basis of the empirical weight relationships obtained from controlled experiments in which known masses of pure sulfates have yielded weighed precipitates of unanalyzed barium sulfate. The customary procedure for the sulfate determination involves adding dilute barium chloride solution dropwise to a hot, stirred, diluted acid solution of the sulfate unknown, digesting the precipitate overnight, filtering, washing, and igniting to constant weight. The variants which have been proposed include: The method of Hintz and Weber ( 4 ) , in which a slight excess of barium chloride solution is added rapidly to the unknown, followed by digestion on a steam bath for a t least an hour. The method of I'opoff and Neuman ( g ) , in which the sulfate solution is added slowly to a slightly acidified barium chloride solution. The method of Fales and Thompson ( S ) , in which the customary determination is carried out, except that the precipitate is ignited a t 110" C., rather than a t a higher temperature.
For the particular case of alkali chloride contamination the method of Hintz and Weber yields excellent empirical results. However, it is not known whether the method will prove satisfactory for other ions-e.g., nitrate-and, in addition, the chemical purity of the precipitate has not been established. I n this sense it may represent a more successful application of the principle involved in the other methods-counterbalancing errors. The background for the present research is in the observation of Kolthoff and hlacxevin (6) that when a precipitate of barium sulfate (formed from sodium chloride solution) is heated above 800" C., the sodium chloride and sodium sulfate are driven from the crystal lattice and are either volatilized or can be leached with water. There exists a parallel in some old industrial processes for the purification of barytes in which the impure ore is fused in a flux of molten sodium chloride and the impurities are then washed out with water ( 2 ) . What appears to be involved in both situations is the destruction of a solid solution which is unstable a t higher temperatures and which also has a chance to break up a t a reasonable rate. In short, what is proposed is the purification of precipitated barium sulfate by dissolving it in a molten salt bath, and allowing it to re-form a t a rather high temperature. Preliminary work showed that coprecipitated chloride was quantitatively removed by one recrystallization from a fused sodium chloride-potassium chloride eutectic (9).
847 EXPERIMENTAL
The first problem to arise was the selection of a suitable crucible for the fusion process. Silica and Alundum crucibles cracked, porcelain was attacked, and nickel crucibles were oxidized and flaked off into the melt. Early work with platinum showed that it was n e t by the barium sulfate melt (unlike its behavior with chlorides alone); upon cooling, the crucibles contracted badly, and cracks appeared after several cycles. Accordingly, equal quantities of sodium chloride and potassium chloride were taken. These form a eutectic, melting a t 660" C.-considerably lower than the fusion point of either component. The mixture was fused a t as low a temperature as possible to effect clear solution. -4fter each fusion the crucible was re-formed on a wooden mold. In this way a platinum crucible was usable even after more than 200 fusions. Sodium, potassium, and barium chlorides were purified by precipitation from water solution by hydrogen chloride gas. They were filtered and washed with purified hydrochloric acid. This precaution was necessary because impurities in the reagent grade chemicals consistently added to the weight of barium sulfate precipitate. Favorable conditions were determined for carrying out the fusions and subsequent operations. The purified mixed flux was melted and allowed to cool. A known weight of ignited commercial barium sulfate was added. Usually 1 gram of sulfate was added to 10 grams of flux. The mixture \\-as heated until the sulfate dissolved. After cooling, it was leached with about 100 to 200 ml. of water containing 0.5% hydrochloric acid and 2% barium chloride by weight. This measure was taken because the dilute acid solution alone resulted in a continuing loss in weight, beyond the expected loss due to solubility. The filtrate gave positive tests for both barium and sulfate ions. I t was hypothesized that there was a slight reversal of the precipitation reaction in the melt-Le., that small quantities of barium chloride and sodium sulfate were formed. These reacted in the leaching solution to form barium sulfate, but the precipitation was very slow. Von Weimarn has shown that when the concentrations of barium and sulfate salts in a solution are about 0.0002N, the precipitate takes a month to appear. The loss in weight which was observed corresponded to approximately this concentration of barium sulfate. Accordingly, the leaching solution was prepared ivith barium chloride. The precipitate was collected on a filter, ashed free of chlorides, dried, and ignited to constant weight. It was observed that the precipitate obtained on leaching was very coarse, settled rapidly, and could be filtered and washed with ease, in contrast t o freshly precipitated barium sulfate. RESULTS
When purified reagents Tere used and barium chloride was added to the leaching solution, samples of commercial barium sulfate showed constancy in weight to 1 mg. or less (this was the arbitrary goal chosen for reproducibility, although most results were decidedly better) after a number of cycles of fusion, precipitation, and ignition Table I shows some typical results:
Table I.
Fused-Salt Recrystallization of Commercial Barium Sulfate
Original Weight, Grams 1.0182 1.0108 1.0137 0.9989
1st 1.0217 1.0103 1.0118 0.9982
After Purification, Grams 2nd 3rd 1.0223 1.0224 1.0100 1.0101 1.0116 1.0116 0,9979 0,9979
4th 1,0216
....
.... ....
In all four cases, constant weight t o within 1 mg. was obtained after one purification. Spectrographic analysis showed sodium and potassium to be present to the extent of 0.001 % in samples after one purification. KO other metal except barium could be detected. The addition of sulfuric acid to the precipitate followed by fuming to dryness resulted in no significant change in weight, indicating the absence of chloride. It can be concluded that a pure barium sulfate is obtained after one fusion with a sodiumpotassium chloride flux. Quantitative work was then begun on known soluble sulfate systems, examining the weight of purified barium sulfate which was obtained from the known sulfate.
848
ANALYTICAL CHEMISTRY
Reagent grade sodium sulfate was recrystallized twice from water, dried a t 110" C., and heated to constant weight a t 600" C. immediately before use. The sample was dissolved in 175 ml. of water containing 1 ml. of concentrated hydrochloric acid and heated t o boiling, and 5 grams of barium chloride in 25 ml. of water were added slowly. The preci itate was aged overnight, collected on a filter, washed free of chtride, dried, and ignited to constant weight. It was then dissolved in flux, leached with water containing 2% barium chloride and O.5y0hydrochloric acid, aged overnight, and collected in the same way as before. This general precipitation and purification procedure was used in all experiments, unless specific mention is made of another method. The results are shown in Table 11.
Table 11. Fused-Salt Recrystallization of Barium Sulfate Obtained from Sodium Sulfate iya1s04 Taken, Gram 0.7252 0.6273 0.3150 0.3136
0.3060 0.3021
BaSO4, Grams Theoretical Found 1.1916 1.1892 1.0308 1.0282 o 5176 0.5168 0.5156 0 5161 0.5024 0.5028 0.4904 0.4967
After Purification, Grams 1st 2nd 3rd 4th 1.1918 1.1931 1.1912 1 1918 .... 1.0322 1.0333 1.0336 0.5192 0 . 5 ~ 0 1 0.5205 . .., 0.5187 0.5199 0 . 5 2 0 5 0 5200 0.5059 0.5067 0.5069 .... 0.4995 0.5003 0.5015 0 5025
Table 111. Ignition of Sodium Sulfate NaiSOI, Grain Aftei heating a t 600' C. After fuming with Hz804 and then heating a t 600' C. After being ignited above 884' C.
0 6106 0.6160
O.G14C,
In five of the six samples the actual quantit- of harium sulfate recovered was higher than the theoretical. This was probably causcd by small amounts of sodium bisulfate in the sodium sulfate, because samples of the latter, after being hented to constant weight at 600" C., shoned further weight locses when heated above the melting point (884" (2.). Table 111 illuqtrates this point. The precipitation of barium sulfate from sodium sulfate was repeated, with the latter undergoing ignition to rotistnnt weight above its melting point. Table IV describes thew experiments. It is evident that, the theoretical weight of barium sulfate from a known quantity of sodium sulfate is obtained after one recrystallization from the alkali chloride flux. Table IV. Fused-Salt Recrystallization of Barium Sulfate from Purified Sodium Sulfate NanSOi Taken, Gram 0.6170 0.6167
BaSOa, Grams Theoretical Found 1.0138 1.0104 1.0133 0.1094
After Purifiration, Grama 1st 2nd 1.0114 1.0140 1.0125 1.0127
Table V. Recrystallization of Barium Sulfate Formed from Solutions Containing Sodium Chloride Na2SO4 Taken, Gram 0.6207 0.6237 0,6328
BaSO4, G r a m s Theoretical Found 1,0084 1.0199 1.0249
1.0398
1 0110
1.0279
After Purification, Grams 1st 2nd 1.0195 1.0197 1.0245 1 ,0243 1.0397 1.0390
Sodium chloride was then introduced into the sulfate solution before precipitation, as this is a common analytical problem. T r n grams of sodium chloride were added to each sample. The precipitates were brought down and purified in the usual manner. Table V shows the results of these operations. The initial weight of barium sulfate was low by about 12 mg. because of the coprecipitation of sodium sulfate ( 3 , 6). One recrystallization yielded the theot etlcal quantity of precipitate, demonstrating that the intei fercnce caused by Podium chlot ide can be successfully eliminated by this one step. It ha- been reported ( I . ? ) that in barium sulfate contaminated
Table VI.
Effect of Roasting on Hemoval of Impurities from Barium Sulfate
Contaminant. 10 Grams NaCl NaCl KNOs KNOI
Naz304, Gram 0.6140 0.6295 0.6248 0.6260
Theoretical 1.0089 1.0344 1.0267 1.0286
BaSO4, Grams After After roasting roasting and leaching 0.9919 0.9980 1.0210 1.0272 1.0502 1.0432 1,0420 1.0538 1.0488 1,0439
1:ound 0.9923 1.0214
hy alkali metals, the latter could be quantitatively removed by roasting the precipitate a t 900" t o 1000" C. and then leaching with hot dilute sulfuric acid. It was decided to test this method in order to compare it with that using the alkali chloride flux. In addition, since nitrate ion is one of the most serious contaminants of barium sulfate, samplee of sodium sulfate were prepared to which were added 10 grams of recrystallized sodium chloride or potassium nitrate. The precipitate was washed, ignited to constant weight, roasted in a muffle furnace at 975" C. for 1 hour, reweighed, and leached \r.ith 2% barium chloride solution. iifter overnight aging, each preripitate was collected in the usual manner, washed, and ignited to constant weight. The results are given in Table VI. The results sholv that roasting under these conditions removes or restores only part of the sodium sulfate arising from sodium chloride addition. This method therefore has limited analytical value, unless different heating cycles offer better results. Nitrate ion has long been regarded as a most troublesome source of ei't'or in the sulfate determination because of heavy coprecipitation. This, in turn, has been shown to be due to extensive solid solution formation with barium sulfate (11). Because roasting appeared to have little effect on the removal of nitrate, experiments wwe conducted to pee whether the recrystallization from fused salts could be applied to precipitates containing nitrate. Sample? of sodium sulfate were mixed with 10 grams of recrystallized potassium nitrate; barium sulfate ivas then precipitated, and purified in the usual manner. The results appear in Table VII.
Table VII.
Effect of Recrystallization on Nitrate Content of Barium Sulfate"
After Purification, Graniv T&en. Na2604 ___ BaSOI. Grams Ratio of Flux 1st 2ndGram Tlieoretiral Found to BaS04 1.0400 1.0370 , 1.0288 0.6248 1 02iiIi 5 to 1 1.0315 1.0261 1.0138 1.0406 5 to 1 0.6182 1.0460 1.0345 1.0284 1.0537 10 to 1 0.6259 1.0438 1 0320 1.0576 0.6251 1.0271 10 t o 1 a Potassium nitrate added, potassium chloride-sodium chloride flux.
Table VIII. Effect on Barium Sulfate of Potassium Nitrate Introduced into a Fused-Salt Melt Flux Composition, Grams KC1 NaCl KNOI 4.75 0.: 4.75 4.5 1.0 4.5 2.0 4.0 4.0 3.5 3.0 3.5
BaSOi, Grams Added Recovered 1.0194 1.0029 1,0392 0.9930 1.0002 1.0428 1.0377 1.0000
Gain, Grain 0.0165 0.0482 0.0426 0,0377
It is evident that one recrystallization is insufficient to remove nitrate quantitatively. Pure sodium chloride and pure barium chloride were also tried as fluxes, because it was hoped that the higher temperature a t which these crystallized from the melt would decrease the region of -table existence of solid solutions and hence yield a more pure precipitate. There mas a slight improvement over the results obtained with the eutectic, but not enough t o justify this approach for nitrate removal. Since nitrate appears t o form a stable solid solution Kith barium sulfate which persists at high temperature, it seemed appropriate to attempt to introduce it directly into the precipitate from the melt. -4ccordingly, known quantities of precipi-
V O L U M E 26 NO
5, M A Y 1 9 5 4
tate were treated in fluxes containing varying percentages of potassium nitrate. After the melts cooled, they were treated in the usual manner, and finally ignited to constant n-eight. Table VI11 describes these experiments. The increase in weight is marked, reaching approximately 46 nig. per gram of barium sulfate at a higher ratio of nitrate to flux. The weight decrease beyond this point is probably due to the partial decomposition of nitrate a t the higher conceiitrations, sinrc nitrogen oxides were liberated when these samples were leached. I t is tentatively suggested that nitrate be removed before precipitation by repeated tieatmcnt n i t h hot, concentrated hydrochloric acid. COhCLCSION
Sodium and chloride ions can br quantitatively ienioved f i om barium sulfate in one purification by a flux containing equal amounts of potassium chloride arid sodium chloride. Leaching with slightly acidified barium chloi ide solution yields the theoretical quantity of pure barium sulfaic. Sitrate cannot be removed in one purification. ACKNOW LEDGXIEhT
One of the authors (J. S.) wishes to thank Roland Ward and
849 Joseph Greeuspan for many helpful discussions. bcknowledgnient is a1.o made to Isidore hdler for spectrographic analyses. LITERATURE CITED .Irerell, P. R., a n d W a l d e n . G. H., J . .4m. Chem. Soc., 59, 906 (1937). B o o t h , H. S., Pollard, E. F., a n d Rentschler, M. J., Ind. Eng. Chem., 40, 1981 (1948). a n d T h o m p s o n , W.S., ISD. ENG.C H E W . ,ANAL. ED.,1 1 , 206 (1939). H i n t s . E., a n d W e b e r , H., 2. anal. Chem., 45, 31 (1906). J o h w t o n . .J.. a n d A d a m s , L., J . A m . Chem. Soc., 33, 829 (1911). Kolthoff, I. AI., a n d Rlaciievin, W.II.,J . Phys. Chem., 44, 921 (1040). Kolthoff, I. AI., a,nd Sandell, E. H., “ T e x t b o o k of Q u a n t i t a t i v e I n o r g a n i c .Iiialysis,” 3 r d ed., p. 112, N e w Y o r k , Macmillan
Co.. 1952.
S.,a n d S e u m a n , E. W., ISD. ENG.CHmi., ANAL.ED.,2, 45 (1930). S c h m i t t , G. ,J., B.S. thesis, Polytechnic I n s t i t u t e of B r o o k l y n , J u n e 19.50. Schneider, F., a n d R i e m a n , W.,J . -4711. Chem. Soc., 59, 354 (1937). W a l d e n , G. 11.. a n d C o h e n , 11. U., Ibid., 57, 2691 (1935). W a l t o n . G.. a n d W a l d e n , G. R . , I b i d . , 68, 1742 (1946). Popoff,
RECEIVED for review August 24, 1953. A c c e i ~ t r dFebruary 10, 1954.
Determination of Residual Crag Herbicide 1 and Its Hydrolysis Products on Food Crops J. N. HOGSETT and G. L. FUNK Carbide and Carbon Chemicals Co., Division o f Union Carbide and Carbon Corp., South Charleston, W. V a .
The increasing use of sodium 2-(2,4-dichlorophenox~)ethyl sulfate, Crag herbicide 1, for the weed control of various agricultural crops has necessitated the development of analytical procedures for the determination of spray residues. A procedure is described for the quantitative estimation of as little as 0.018 mg. of Crag herbicide 1 on food crops by measuring the intensity of the colored complex formed with methylene blue chloride. Application of the procedure to the determination of 2-(2,4-dichlorophenoxy)ethanol and 2,4-dichlorophenol, hydrolysis products of the herbicide, after a preliminary sulfation is also described. Data are presented to illustrate the applicability of the procedure to many food cropg. In addition to the determination of Crag herbicide l, the method of sulfation and colorimetric measurement of the resulting sulfate, suggests a sensitive method for the determination of many long-chain alcohols.
S
ODIUM 2-(2,4-dichlorophenoxy)ethylsulfate.
c1 c i a - 0 - c
H?CH~-o--so~N~
Crag herbicide 1 (6, 7’). an effective new chemical for weed control, has the unique property of being noninjurious to plants when sprayed or dusted directly on the foliage a t the concentrations that will kill weed seedlings in the soil. I t is activated upon contact with the soil and is very effective for the control of shallow-rooted, broad-leaved plants and grass roots. Microorganisms in nonsterile soil convert the herbicide to 2-(2.4-dichlorophenosy)ethanol and 2-(2,4-dichlorophenoxy)acetic acid, which are active plantgrowth regulators ( I O ) . Its application and
effectiveness are facilitated by virtue of its water solubility and its surfactant properties. Crag herbicide 1 is commercially available as a powder with a minimum purity of 90.0% by weight. The compound is used in dilute concentrations, 0.4 to 0.5’% in water, which is sprayed upon the soil. Because of the widespread use of synthetic herbicides, fungicides, and insecticides, i t has become necessary to develop sensitive methods of analxsis for their detection on food crops. To ensure adequate protection for the public from possible health hazards arising from the use of these chemicals, adequate procedures for the determination of spray residues in trace amounts must be developed so that reliable data can be obtained for toxicological studies. During an extensive herbicide-spray program, accurate methods of analysis for low concentrations of Crag herbicide 1 and ita hydrolysis products, 2,4-dichlorophenol and 2-(2,4-dichlorophenoxy)ethanol, were needed. Their determination by measuring total aromatic chlorides was rejected because of a lack of selectivity of the method. -4few methods of analysis for chlorinated agricultural chemicals have been reported. None of the three procedures for DDT as outlined by Fahey and Rusk (S) proved applicable to the analysis of Crag herbicide 1 or its hvdrolysis products The residue control of 2.4-D has been successfully developed 15 ith the use of Swanson’s bioassay method ( 9 ) , Bandurski’e spectrophotometric method ( 1 ). and two colorimetric methods-the method of Freed (Q), involving the substitution of chlorine in 2,4-D by an amino group \’i ith the subsequent detection of the amino group with sodium 1.2-naphthoquinone-4-sulfonate. and the method of Marquardt and Luce ( 8 ) ,involving the reaction of 2,4-D with chromotropic acid (l$-dihydroxynaphthalene-3.6-disulfonic acid) in concentrated sulfuric acid. The methods of Swanson, Bandurski. and Freed could be applied to the residue analysis of Crag herbicide 1 and its hydrolysis products, hut all three lack sensitivity and ease of application,
