1256
ANALYTICAL CHEMISTRY
than 0.003 p.p.m.) is predicated on the assumption that D D T is quantitatively sorbed by the solids. In exploratory tests (6) of the assumption, quantitative sorption of D D T by suspended solids was not obtained. The quantity sorbed was influenced by the amount and the nature of the solids-i.e., different physical properties resulting from different amounts of clay, silt, and very fine sand, different sedimentation rates, etc. Ability of sui+ pended solids to adsorb microquantities of D D T from river water will no doubt vary with the individual case. Although the experiments on synthetic samples failed to show quantitative sorption by the silt, the procedure described may nevertheless be valuable in obtaining some approximate data for D D T concentrations in samples taken a t long distances-e.g., 68 miles ( 6 ) from the point of application of the DDT. In such cases the solvent extraction procedure is not practicable because of the very minute amounts of D D T in the water, and the photometric error becomes large.
ACKNOWLEDGMENT
The author wishes to thank C. R. Twinn, H. E. Gray, and -4. P. Arnason, Division of Entomology, for cooperation; B. N. Smallman, now of the Science Service Laboratory, London, Ont., for completing the arrangements whereby this aspect of the black fly problem was explored; J. W. T. Spinks, Department of Chemistry, University of Saskatchewan, and members of his staff, for advice and excellent working conditions in the laboratory supplied for the investigation; and W. A. E. McBryde, Department of Chemistry, University of Toronto, for very helpful criticism of the manuscript. REFERENCES
(1) Archibald, R. h l . , ANAL.CHEW,22,639 (1950). (2) Arnason, -1.P., Brown, A. W. d.,Fredeen, F. J. H., Hopewell, W. W., and Rempel, J. G., Sci. Agr., 29, 527 (1949). (3) ;lyres, G. H . ,AK.\L. CHEM.,21, 652 (1949). (4) Clifford, P. A . , J . Assoc. Ofic.Agr. Chemists, 28,152 (1945). (5) Davidorv, E., Ibid., 33, 130 (1950). (6) Fredeen, F. J. H., Arnason, A. P., Berck, B., and Rempel,
SUMMARY
J. G., Can. J . Agr. Sci., in press. D D T in river water (as low as 0.003 p.p.m. of D D T ) and in suspended solids wag determined, using solvent extraction and the Schechter-Haller procedure. Interferences in the extract were removed by Davidow's chromatographic method. I t was found that solids suspended in river water (clay, silt, very fine sand) adsorbed DDT. In samples containing suboptimal amounts of DDT, for whirh the regular solvent extraction of water was not feasible, D D T was determined directly from the suspended solids fraction. Larger samples of water may thus be processed with gain in sensitivity.
(7) Herriott, R. M., Science, 104, 228 (1946). (8) Lowry, 0. H., and Bessey, 0. S., J . Bid. Chem., 163,633 (1946). (9) MoBryde, W. A. E., . ~ N A L . CHEM.,24, 1639 (1952). (IO) Ringbom, A., Z . anal. Chem., 115, 332 (1939). (11) Schechter, M. S..Pogorelskin, M. A,, and Haller, H. L., ANAL. CHEX, 19,51 (1947). (12) Schechter, M. S., Soloway, S. B., Hayes, R. A., and Haller, H. L., IND.E m . CHEN.,ANAL.E D . , 17, 704 (1945). (13) Stiff, H. a., and Castillo, J. C., Science, 101, 440 (1945). RECEIVED for review October 17, 1952. hccepted March 9, 1953. Contribution 3021, Dix-ision of Entomology, Science Service, Department of Agriculture, Ottawa, Canada.
Quantitative Insolubility of Thorium Oxalate HAROLD L. KALL
AND
LOUIS GORDON
Departriient of C h e m i s t r y , Syracuse Cnicersity, SyraczLse 10, V. I-
u
THE
course of a study of coprecipitation with thorium
I-oxalate, some doubt arose as to the quantitative insolubility (2, 6, 9) of this salt, since the thorium content of a solution analyzed by an oxalate procedure ( I O ) was found to be consistently less than the value obtained by the hexamine method (3). This apparent solubility loss of thorium oxalate in dilute acid solutions was previously reported by Rider and Mellon in a paper describing a colorimetric method for thorium ( 5 ) . Although these authors did not completely describe the conditions under which the solubility loss was obtained, they reported that 0.1 mg. of thorium remained in 20 ml. of solution a t pH 0.7. They further reported that quantitative precipitation of thorium oxalate could be obtained in solutions of pH 0.7 to 3.0 without the use of a large excess of oxalate. Since oxalate is widely used for the determination of thorium, an investigation of the solubility losses of thorium oxalate in dilute acid media was undertaken. A recent sensitive and accurate method for the colorimetric determination of microgram quantities of thorium (8) was utilized in the analysis of filtrates. APPARATUS AND REAGENTS
Thorium Nitrate Solution Solutions were prepared by dissolving pure thorium nitrate (Lindsay Chemical c~.,. Code 103) in distilled water, adding 100 ml. of concentrated nitric acid, and then diluting to 1 liter. These solutions contained approximately 68 mg. of thorium oxide per 50 ml. l-(o-Arsonophenylazo)-2-naphthol-3,6-disulfonicacid. A 0.1 % reagent solution was prepared by dissolving 0.1 gram of this reagent (marketed as Naphtharson by the SmithaNew York Co.) in 100 ml. of distilled water. Lanthanum Perchlorate Solution. A solution was prepared by dissolving lanthanum oxalate (Lindsay Chemical Co., Code
518) in a mixture of nitric and perchloric acids, evaporating to fumes of perchlorate, and then diluting to 1 liter. The solution contained 20 mg. of lanthanum oxide per milliliter. Dimethyl Oxalate. The Matheson Go. product was used. Hexamine. Hexamethylenetetramine, The Matheson Co. product, was used. All other chemicals were C.P. .4 Beckman Model B spectrophotometer with 1-cm. Corex cells was used to measure the transmittancies of the thorium solutions. ANALYSIS O F FILTRATE FOR THORIUM CONTENT
I n each case, except where otherwise noted, 68 mg. of thorium oxalate were precipitated as oxalate in a 250-ml. volume as described, and the precipitate was then filtered with S o . 42 Whatman filter paper. The filtrates were evaporated to dryness with nitric acid, the residue was dissolved with nitric acid, a suitable aliquot was evaporated to dryness with perchloric acid, and the residue was analyzed by the procedure of Thomason, Perry, and Byerly (8). RESULTS
Effect of pH on Solubility Loss of Thorium Oxalate Precipitated from Homogeneous Solution. Solutions containing 68 mg. of thorium oxide were diluted to approximately 225 ml., then adjusted to various pH values using ammonia or hydrochloric acid. The solutions xere diluted to 250 ml., and 1 gram of dimethyl oxalate was added. These solutions were heated for 45 minutes a t 70" C. Subsequently, 7 grams of oxalic acid dihydrate were added, and the solutions were allowed to digest at room temperature for 12 hours. The results of these experiments are summarized in Table I. Effect of Variation of Oxalate and Thorium Concentrations on Solubility Loss. Precipitations were carried out, except for
1257
V O L U M E 2 5 , NO. 8, A U G U S T 1 9 5 3 Table I. Solubility Loss of Thorium Oxalate at Various pH Levels Initial p H
Final p H
ThOn in Filtrate, hIg.
0 .2 j a
0.32 0.250 0.30 0.50 0.65 0.50 0.60 1.0 0.85 1 .o 0.90 2.0 0.95 3 0 0.90 4.0 0.90 a Hydrochloric acid added to adjust p H ; all other ammonium hydroxlde for adjustment to initial p H .
0.49 0.50 0.31 0.35 0.86 1.1 >1.5 >l.5 >1.5 solutions required
Table 11. Effect of Variation of Oxalate and Thorium Concentrations” ThOs Present, hlg. 68 68 68 68 68 68 68 68
Volume, M1. 260 250 125 125 125 125 250 250
Methyl Oxalate, Gram 1 0 1 0 0.5 0.5 1.0 1.0 1.0 1.0 1.0 1.0
Oxalic Acid, Granls 7 0 7 0 3.5 3.5 7.0 7.0 0.0 0.0 7.0 7.0
250 34 b 346 250 a At a n initial p H of 0.5. b No ammonia or hydrochloric acid was added.
Final pH 06.5
060 0.48 0.45 0 50 0.50 0.65 0.70 0.60 0.60
ThOz in Filtrate, Jlg. 0 31 0 35 0.17 0.12 0.37 0.42 0.74 1.1 >1.7 >1.4
homogeneous procedure was repeated with filter paper pulp present, the solubility loss of thorium oxide was reduced to 0.19 mg. which is about the same as in the heterogeneous method. The initial and final p H values of the solutions were approximately 0.90. Precipitation of Thorium in the Presence of Rare Earths. An oxalate procedure is often used as one step in the determination of thorium in monazite sand. In many of these procedures oxalate is used t o precipitate thorium and the rare earths. To investigate the effect of rare earths on the solubility loss of thorium oxalate, several precipitations were made with 20 mg. of lanthanum oxide present in addition to the thorium. The results of these experiments are summarized in Table IV. Lanthanum alone, under the conditions used, was itself precipitated a t p H 0.60 or higher. Hexamine Procedure. The analysis of filtrates obtained a ith the hexamine procedure (3) showed no more than 0.02 mg. of thorium oxide in the 250-ml. volume of filtrate. In these experiments precipitation was effected a t 70” C. rather than a t 30’ C. as recommended. Precipitation is more rapid a t the higher temperature and the precipitate still retains its desirable properties. DISCUS SIOB
Thorium oxalate is not as quantitatively insoluble as one would infer from the literature. One source of solubility loss occurs in
Table 111. Solubility Losses of Thorium in Some Standard Oxalate Procedures modifications in concentrations of constituents, in much the =me manner as described in the previous paragraph. These data are shown in Table 11. Effect of Varying Amounts of Ammonium Salts. Varying amounts of ammonium chloride were added to solutions at different pH values. The results followed an erratic pattern, but the tendency seemed to favor greater solubility at large concentrations of ammonium salts even a t p H values of 0.25 and 0.50. However, the presence of a certain quantity of ammonium salts materially decreased the total solubility. For example, in one case where 68 mg. of thorium oxide were precipitated in the presence of 2 grams of ammonium chloride, less than 0.1 mg. of thorium oxide was found in the filtrate; the initial and final pH values were both 0.25. However, even a slight variation either in pH or quantity of ammonium chloride produced a very sharp increase in the total thorium loss. Effect of Digestion Time. Prolonged digestion of thorium oxalate a t room temperature resulted in a decrease in the solubility loss. For example, under one set of precipitation conditions, 0.34 mg. of thorium oxide was found in the filtrate after a 12-hour digestion period; this loss was reduced to 0.20 mg. after a 36-hour digestion period. Other Methods of Precipitation of Thorium Oxalate. Results of the solubility losses encountered with some other methods for the determination of thorium are summarized in Table 111. Of the methods investigated, the procedure by Eniiig and Banks ( 1 ) resulted in the most complete precipitation. I n this method a thorium solution containing 5 ml. of perchloric acid i8 diluted to 300 ml., filter paper pulp is added, and 5 grams of oxalic acid dihydrate are then added to the hot solution. After a digestion period of 15 minutes a t 95’ C., the solution is cooled, and the precipitate is filtered. I n order to improve the characteristics of the precipitate obtained in the Ewing and Banks procedure, thorium was precipitated with I gram of dimethyl oxalate followed by the subrequent addition of 4 grams of oxalic acid dihydrate. The total thorium oxide loss found in this procedure was 0.50 mg. as compared with 0.14 in the Ewing and Banks method. When the
Method
T h o , in Filtrate, Mg,
Bureau of Standardsa
0.19 0.17 Bureau of Standardsb 0.24 0.22 Heterogeneous C 0.90 >l.5 Mayer and B r a d s h a d 1.2 1.2 Ewing a n d Banks8 0.14 0.12 a This method ( 7 ) involves solution of a thorium hydroxide precipitate in hydrochloric acid, evaporation t o dryness, a n d addition of 20 t o 25 ml. of 10% oxalic acid t o the residue. The solution is stirred, diluted t o 100 ml. gently boiled for a few minutes, a n d allowed t o stand overnight a t r o o d temperature before filtration through S o . 42 Whatman filter paper. b Same procedure as in but with a final volume of 250 ml. containing 50 to 60 ml. of 10’7 oxalic acid solution added to effect precipitation. c Procedure Jescribedjn “Scott’s Standard Methods of Analysis’’ (2). d Thorium hydroxide is precipitated a n d then dissolved in 5 ml. of concentrated hydrochloric acid. This solution is diluted to I50 ml., heated, 10 grains of oxalic acid dihydrate in 50 ml. of water are added, a n d the precipi. tate is filtered after a n overnight digestion at room temperature (4). e Previously described; 200 mg. of thorium oxide taken.
Table IV. Solubility Losses of Thorium Oxalate Precipitated in Presence of Lanthanuma Homogeneous procedure b No digestion period: filtered hot Initial plI 0.25 0.50 1
.o
1.5 2.0
T h o , in Filtrate. Mg. >1.6 >1.2
0.13 0.07 0.07
Twelve-hour digestion period; filtered cold Initial p H T h o , in Filtrate, M g . 0.53 0.25 0.14 0.50 0.02 1.0 0.02 2.0 Heterogeneous procedurec Initial p H 0.25 0.50
Final p H 0.45 0.60 0.90 1.0 a 20 mg. of lanthanum oxide present. b As described in Results. As described in Scott ( 2 ) .
ThOa in Filtrate, Y g , >1.8 0.80 0.66
1258
ANALYTICAL CHEMISTRY
the presence of ammonium salts. Although it appears that a certain quantity of an ammonium salt, as is shown in Tables I and 11, can be beneficial in that a minimum solubility loss is obtained, erratic results due to the solvent effect of ammonium oxalate would require too rigorous control over precipitation conditions. It appears, therefore, desirable to effect precipitation of thorium oxalate in the absence of ammonium salts. Of the procedures tested, those by the Bureau of Standards (7) and Ewing and Banks, and the modification of the latter to include a homogeneous step as described in this paper, approach quantitative precipitation. The small solubility loss of 0.1 to 0.2 mg. of thorium oxide requires the use of a sample of suitable size to offset the resultant loss in accuracy. The use of filter paper pulp seems mandatory. The data in Tables I1 and 111 confirm the statement by Rider and Mellon (6) that the solubility of thorium oxalate depends on the total volume of solution and not on the amount of thorium present. When a rare earth is present which may also be precipitated by oxalate and thus serve as a carrier for traces of thorium normally remaining in solution, the solvent effect of moderate quantities of ammonium salts becomes negligible. I n this case quantitative precipitation of thorium may be effected in the p H range from 1.0 to 2.0. Precipitation with hexamine can then serve to separate thorium from the rare earths. A slight variation in the concentration of oxalate does not seem to affect appreciably the solubility of thorium oxalate, although a very large excess tends to increase the loss. However, as is generally recognized, the mineral acid concentration critically affects the solubility of thorium oxalate. This is also evidenced by the data in the last three methods in Table 111.
On the whole it would seem desirable to avoid, whenever possible, the use of an oxalate procedure for the determination of thorium. If it is used, for example, to separate thorium from titanium, zirconium, or phosphate, ammonium salts should be absent or else rare earths should be present as a carrier. Hexamine can be used to precipitate thorium quantitatively in the presence of ammonium salts or the rare earths but titanium, zirconium, and phosphate would interfere. ACKNOWLEDGMENT
The authors wish to acknowledge a grant-in-aid from the Research Corp. in support of this investigation. LITERATURE CITED
(1) Ewing, R. E., and Banks, C. V., AKAL.CHEM., 20, 233 (1948). (2) Furman, N. H., ed., “Scott’s Standard Methods of Analysis,” 5th ed., Vol. I, p. 953, New York, D. Van Nostrand Co., 1948. (3) Ismail, A. M., and Harwood, H. F., Analyst, 62, 185 (1937). (4) Mayer, A,, and Bradshaw. G., I b i d . , 77, 154 (1952). (5) Rider, B. F., and Mellon, hl. G., Anal. Chim. Acta, 2, 370 (1948). (6) Rodden, C. J., ed., “Analytical Chemistry of the Manhattan Project,” 1st ed., p. 169, New York, McGraw-Hill Book Co., 1950. (7) Rodden, C. J., “Thorium Standard Samples,” United States Atomic Energy Commission, MDDC-1220 (1947). (8) Thomason, P. F., Perry, M. A, and Byerly, M. M.,ASAL. CHEM., 21, 1239 (1949). (9) Willard, H. H , and Diehl, H., “Advanced Quantitative 4nalysis,” p. 323, New York, D. Van Nostrand Co., 1943. (10) Willard, H. H.,and Gordon, Louis, ANAL.CHEM., 20,165 (1948). RECEIVED for review March 26, 1953. Accepted June 1, 1953.
Two Techniques for Determination of Organic Sulfides WM. H. HOUFF AND ROBERT D. SCHUETZ Kedaie Chemical Laboratory, Michigan S t a t e College, East Lansing, Mich.
s
and Edsberg (9) have described a method for the determination of dialkyl sulfides which involves the oxidation of the sulfide to the sulfoxide by means of standard bromide-bromate solution. The end point is detected by the observance of the first yellow color due to excess bromine. This procedure suffers from the disadvantage of having an end point difficult to distinguish and is thus inapplicable to even slightly colored com. pounds. A procedure surmounting these difficulties was sought by the study of a back-titration technique and by the application of the dead-stop end-point method. Furthermore, it was desirable to investigate whether bromine oxidation was applicable to the determination of aromatic and heterocyclic sulfides. The oxidation of organic sulfides by bromine is described by the following equations ( 5 , 6): IGGIA
+ Br2 +(RpSBr)+Br+ H2O +R2SO + 2HBr RzSO + Br2 + H,O +R,SO2 + 2HBr R2S
(R2SBr)+Br-
(1) (2) (3)
Reaction 1, addition of the bromine to the sulfur atom of sulfides, occurs more or less readily and is influenced somewhat by the nature of the organic groups attached to the sulfur. This was observed experimentally in the use of the dead-stop technique which allowed a qualitative comparison of the rates of addition of bromine to the sulfides. In the case of the dialkyl sulfides the rate of addition was rapid in comparison to the rates for compounds having either an aryl or heterocyclic group attached to the sulfur, presumably because of steric hindrance of the sulfide linkage by such groups. In general, the addition compounds, (R2SBr) +Br-, tend to
lose the halogen and by hydrolysis (Reaction 2) yield sulfoxides. The stability of the dibromides has been found to be determined to some extent by the groups attached to the sulfur ( 3 , 4). This \vas confirmed in the present study since it was found necessary to carry out the bromine oxidation of the aryl and heterocyclic sulfides a t slightly elevated temperatures to obtain the sulfoxides. I n several cases in the determination of aryl and heterocyclic sulfides, employing the dead-stop technique, the end point, which depends upon the first excess of bromine in the presence of bromide, was not particularly apparent unless bromide was added in excess of the stoichiometric amount required in the bromidebromate reaction. This presumably was necessary because of the stability to hydrolysis of some of the bromine addition compounds of the aryl and heterocyclic sulfides n-hich prevented the formation of bromide ion. In addition, it was found that the sulfoxides of aryl and heterocyclic sulfides were resistant to further oxidation to sulfones by bromine a t temperatures a little above room temperature (Reaction 3). This was in contrast to the easy oxidation of dialkyl sulfides to the corresponding sulfones. These observations allowed the addition of excess bromidebromate and the determination of excess reagent iodometrically. As stated by Siggia and Edsberg (9) and observed in the present work, this procedure could not be applied to dialkyl sulfides without obtaining high results since dialkyl sulfides are easily oxidized to the sulfones by excess bromine. In the case of aromatic and heterocyclic sulfides a slight excess of bromide-bromate solution is added to an acetic acid solution of the sulfide along with 3 ml. of concentrated hydrochloric acid. Since both the bromine addition and the hydrolysis steps of the reaction proceed slowly, the solution is stoppered and heated to 45’ C., and then allowed to
