V O L U M E 25, NO. 5, M A Y 1 9 5 3 Table I V .
767
Acid Effect on I n d u c e d Oxidation of U r a n i u m (IV)
[Titration of uranium(1V) solution with 0.01173 4‘ ferric sulfate] Total Acidity Uranium Uranium (HzSOd of Taken, Found, Error, Reduced Solution. Mg.
6.270 6.270 6.270 6.270 6.270
3Ig. 6.001
3Ig.
.v
0.269
0.14
6.057 6.111 6.158
0.213 0.15g 0.112
6.004
0.266
0.14
0.19
0 35
2.0
uranium with dissolved oxygen in the solution, but also allowed the oxidation of uranium(1V) to take place in a more acid and more concentrated solution where the conditions are less favorable for induced oxidation. RESULTS
The results given in Table I are representative of niany titrations of uranium solutions with 0.009374 S cerium(1V) sulfate as the titrant, using a spectrophotometric end point a t a wave length of 360 mH. Some results of the titration of iron(I1) with 0.009374 A’ cerium(1V) sulfate are given in Table IT. These results indicate that 0.6- to 8-mg. quantities of uranium and 0.3- to 3-mg. quantities of iron in a volume of 90 to 100 mi. can be titrated with an average accuracy of 3.4 and 2.9 parts per thousand, respectively, KO attempt was made to extend the range of the uranium titration below 0.6 mg. or of the iron below 0.3 mg. The sensitivity of these methods should certainly make possible the titration of much smaller quantities of either of these substances. The analyses of mixtures of uranium and iron are given in Table 111. Some samples in this table were titrated without adding cerium(1V) solution to the titration cell before the reduction of the uranium-iron solution. All other titrations in Table I11 were made by first adding ea. 95% of the theoretical amount of cerium(1T’) solution required for the titration of uranium(1V).
This latter procedure gave sharp uranium end points with an average error of less than 0.6% even when 25 times as much iron(I1) was present as uranium(1V). Although no reason could be found for this error, the iron(I1) end points had a tendency to be a little high. The spectrophotometric method for the simultaneous determination of uranium and iron is not only more convenient than the usual potentiometric method but it is applicable to very diluteuraniumandironsolutions where other methods for detecting the end points would be insensitive. The spectrophotometric method can probably be applied to the simultaneous determination of other metals in the presence of each other. ACKNOW LEDGMEVT
The authors wish to thank James ‘4.Wright I11 for doing some preliminarp work on this problem. A portion of this researrh was supported by Contract ,4T(30-1)-937, Scope I of the U. S. Atomic Energy Commission. LITERATURE CITED
(1) Bricker, C. E., and Sweetser, P. B., ANAL. CHEW,24, 409 (1952). (2) Crouthamel, C. E., and Johnson, C. E., Ibid., 24, 1780 (1952). (3) Ewing, D. T., and Eldridge, E. F., J . Am. Chem. Soc., 44, 1484 (1922). (4) Furman, X. H., Bricker, C. E., and Dilts, R. V., ANAL.CHEM., 25, 482 (1953). (5) Rodden. C. J., Editor-in-chief, “Analytical Chemistry of the Manhattan Project.” Sational Kuclear Enerav Series, Division VIII, Vol. 1, Chap. 1, Kew York, McGra%-Hill Book Co., 1950. (6) Sill, C. W., and Peterson, H. E., A N a L . CHEM., 24, 1175 (1952). (7) Smith, G. F., “Cerate Oxidimetry,” pp, 39-42, Columbus, Ohio, G. F. Smith Chemical Co., 1942. . 24, 1107 (8) Sweetser, P. B., and Bricker, C. E., - 4 s . i ~CHEY., (1952). (9) Sweetser, P. B., and Bricker, C. E., Ibid.,25, 253 (1953). RECEIVEDfor review September 5 , 1952. Accepted February 5 , 1963. Based upon a thesis to he submitted by P. B. Sweetser in partial fulfillment of the requirements for the degree .of doctor of philosophy at Princeton University.
Barium Thiosulfate Monohydrate, Standard for Thiosulfate Iodometry WILLIAM M. MAcNEVIN AND OWEN H. KRIKGE McPherson Chemical Laboratory, Ohio State University, Columbus, Ohio
T
HE first suggestion that barium thiosulfate monohydrate be used as a standard for iodine solutions was made hy Plimpton and Chorley (14) in 1895. They showed that this material was easy to prepare, was stable, and had a large equivalent weight. They also noted that it had a low solubility. Their anttlytical results, hom ever. showed that barium thiosulfate monohydrate could be used as a standard for iodine solutions, although their data indicating accuracy and precision are sparse; several qurstions relating to its use as a standard went unansnered. llutnianski ( I d ) , in 1897, reported an improvement in Plimpton and Chorley’s method of preparing the monohgdrate. I n 1944, Gaspar and Santos (6)made an application of the monohydrate and its reaction with iodine by using it as a piecipitating form for the determination of barium. Brief reference only is made to the Plimpton and Chorley work in books by Mellor ( I I ) , Gmelin ( 6 ) , Ahegg and Auerbach ( I ) , Friend (4),and Iiolthoff and Furman ( 7 ) . The use of the reagent has not become popular. I11 view of studies made in the authors’ laboratory, and reported
in this paper, it now appears that its general use as a standard is justified. The objection to its low solubility has also been overcome. Many of its properties make it an ideal standard substance. PREPARATION OF BARIUM THIOSULFATE MOUOHYDRATE
Barium thiosulfate monohydrate is soluble in water to the extent of about 2.66 grams per liter a t 25” C. (9). I t is less soluble in most organic solvents. I t is prepared by mixing warm water solutions of barium chloride and sodium thiosulfate. The white, needlelike crystals of barium thiosulfate are easily filterable with suction, and are washed with alcohol and ether. The ether is removed b y suction and the monohydrate crystals are stable a t ordinary humidities. Procedure for Preparation. Dissolve 40 grams of barium chloride dihydrate and 50 grams of sodium thiosulfate pentahydrate each in 300 ml. of Tvater. Filter each solution if it is not perfectly clear. Warm the two solutions to between 50” and 60” C.
ANALYTICAL CHEMISTRY
2'68
Although the use of barium thiosulfate monohydrate as a standard in iodometry was first suggested in 1895, no data indicating the advantages, limitations, and accuracy have been published. Barium thiosulfate monohydrate of 99.85% purity was prepared. Using this solid material as a standard for determining the titer of dilute iodine solutions, a precision of 5 parts per 10,000 was easily obtained. Dilute solutions of barium thiosulfate monohydrate were also prepared and found satisfactory as standard solutions. .A study of the thermal stability of barium thiosulfate monohydrate showed t h a t the monohydrate is stable only up to 40" C. I t may be dried satisfactorily a t this temperature. Anhydrous barium thiosulfate was
and pour the barium chloride into the thiosulfate solution slowly with stirring. White, silky crystals of barium thiosulfate monohydrate form. Filter the solution on a medium-porosity, sinteredglass funnel with suction. R a s h the precipitate with 3 liters of distilled watSer,300 ml. of 95% ethyl alcohol, and 300 ml. of diethyl ether. Store the crystals under ordinary atmospheric humidity. These crystals have the composition, BaS?03.H20.
I
0
20
,
I
!-.
40
60
,
I
,
I
.
100 120
I
140 TEMPERATURE. DEGREES C
80
,
I
160
,
I
,
entirely unsatisfactory because of its slow rate of dissolving in water. The major advantages in the use of barium thiosulfate monohydrate as a standard are: The equivalent w-eight is large, 267.50. No additional reagents are needed for preparation for the titration reaction. The reaction between thiosulfate ion and iodine is single, rapid, and well defined. The monohydrate is a substance of definite composition, BaS203.HzO. The preparation of the salt is rapid and easy. The salt is commercially available and per cent purity is easily determined. An accuracy better than 1part per thousand is easily obtained in the standardization of iodine solutions. The chief limitations are: its low solubility and its thermal instability above 40" C.
The results, shown in Figure 1, indicate that the monohydrate is stable up to about 50' C., that a large loss in weight occurs between 50" and 100' C., and that further losses occur between 100" and 200' C. When the loss in weight at a single temperature was studied as a function of time, it was found that a steady state is reached for each temperature. Figure 2 shows the behavior of the monohydrate when heated a t 170' to 176' C. The weight loss corresponds exactly to the theoretical for barium thiosulfate monohydrate. At higher temperatures additional loss in weight occurs owing to decomposition of the thiosulfate radical. The anhydrous barium thiosulfate prepared in this way dissolves in water very slowly and although its composition might be suitable as a standard, the slowness of dissolving makes it useless.
I
180 xx)
Figure 1. Decomposition of Barium Thiosulfate Monohydrate a t Various Temperatures Heating time, 3 hours
Thermal Stability of Barium Thiosulfate Monohydrate. Thermal stability data for barium thiosulfate monohydrate in the literature are both vague and contradictory. bianchec (IO) stated that the monohydrate loses water a t an elevated temperature to form the anhydrous salt. Rose (16) reported that the monohydrate loses a molecule of water a t 100' C. or when stored over sulfuric acid. Letts (8) and Curtius ( 2 ) found similar results but did not state how long a heating period mas necessary. Rammelsberg (16) reported that a temperature of 170" C. was needed. Pape (IS) found a temperature of 215" C. was necessary to produce the anhydrous salt. Plimpton and Chorley (14) reported a slow loss in weight a t 56' C. but were able to remove 88% of the water by heating for a "long time." They obtained the anhydrous salt by drying a t 120' C. I n view of these reports, experiments were performed to determine the thermal stability of crystals whose preparation was described above. A thermobalance was used to simplify the work. A platinum crucible mas suspended on the end of a fine chain in an electric oven. The platinum wire extended vertically through the top of the oven and was attached to the end of a balance beam. With this apparatus, weight changes could be followed continuously. I n the first experiment, samples of the monohvdrate were heated for 3-hour periods a t gradually higher temperatures.
0
5
10
I5
TIME IN HOURS
Figure 2. Decomposition of Barium Thiosulfate Monohydrate a t 173' C.
The conclusion is therefore reached that the monohydrate is stable only up to 50" C. and that the anhydrous form is not suitable as a standard. PURITY OF BARIUM THIOSULFATE MONOHYDRATE
I n order to determine the purity of the monohydrate crystals, solutions of the monohydrate were titrated against the iodine liberated by solutions of potassium iodate. Weight burets were used and the percentage purity ( % BaSoOs HzO) was calculated. The iodate solution was prepared by dissolving 0.062787 gram of dried { l l O " c.) potassium iodate in 957.485 grams of water. The thiosulfate solution was prepared by dissolving 0.520503 gram of air-dried barium thiosulfate monohydrate in 236.697 gram of water. Typical data for the titration are shown in Table I.
V O L U M E 25, NO. 5, M A Y 1 9 5 3 Table I.
769
Determination of Purity of Barium Thiosulfate Monohydrate
Wt. of Iodate Solution, Granls
Wt. of Thiosulfate Solution, Grams
99.570 99.951 99.778 99.999 99.669
22.436 22.438 22.388 22.450 22.364
BaS203.Hz0,
Av.
Tahle 11.
Yo
99.86 99.83 99.88 99,83 99,87 99.85
Standardization of 0.1 IN Iodine Solutions
w t . of BaS?Oa.H?O, Grams
Yol. of I2 Solutlon, X I .
1.1710 1,3472 1,4853 0,9028 1.2182
41.63 47.96
of
52.85
32.05 43.36
AV.
Normality I2 Solution 0.1050 0.1049 0.1049 0.1051 0.1049 0.105c-
It should be analyzable with accuracy. The reaction it undergoes should be single, well defined, rapid, and complete.
5.a.
Barium thiosulfate monohydrate meets all of these requirements excellently except the second. While it cannot be dried at 100” C., it nevertheless is stable toward air drying when prepared as described. From the experience in this laboratory, it is admirably suited to the determination of the normality of 0.1 N solutions of iodine where an accuracy of 1 part per thousand is sought. A precision of 1 part in 10,000 has also been obtained M ith weight burets. Barium thiosulfate monohydrate is available commercially. lIowever, samples analyzed in this laboratory showed purity of 99.5%. In view of the simplicity of preparation, a sufficiently pure form of the reagent for standardization purposes could easily I)e prepared and distributed. LITERATURE CITED
Repetition of the preparation gave a product of the same purity. USE OF MOROHYDRATE FOR STANDARDIZING 0.1 N IODINE SOLUTION
The low solubility of barium thiosulfate monohydrate in n-ater (0.01 111) might be considered a serious limitation t,o its use. However, the monohydrate dissolves so rapidly that there is no delay during the course of a titration. Indeed the disappearance of the solid salt serves as a crude indication of the approach of the end point. X o decomposition of the undissolved thiosulfate seems to occur. Several titrations were run to determine the suitability of the monohydrate for use as a standard for 0.1 AV iodine solutions. Air-dried barium thiosulfate monohydrate samples were weighed into 250-ml. iodine flasks and 100 ml. of water were added to each. The solutions were titrated with an iodine-iodide solution whose normality against arsenic oxide (.%s20j)was known to he 0.1049+ S. Table I1 shows typical data. Recently Farr, Butler, and Tuthill ( 3 ) listed the following five desirable characteristics for a primary standard: 1. The substance should be stable and of definite composition. 2. I t should be stable toward drying (preferably a t 100” C.) Lvithout decomposition. 3. It should have a large equivalent weight in order to minimize the weighing error.
Abegg, R., a n d .luerbach, F., “ H a n d b u c h der anorganischen Chemie,” Vol. 11, P a r t 11, p . 271, Leipsig, P. Hireel, 1905. Curtius, T., J . prakt. Chem., 24 ( 2 ) , 233 (1881). F a r r , H. V., Butler, A. Q . , a n d Tuthill, S.XI., ANAL.CHEW., 23, 1534 (1951). Friend, J. S . , “Textbook of Inorganic C h e m i s t r y , ” Vol. 111, P a r t I. D. 227. London. Charles Griffin a n d Co.. L t d . . 1925. G a s p a r , ‘f.,a n d Santos,‘XI., Analcs fis. y guim ( M a d ~ i d )40, , 660-77 (1944). Gmelin, “ H a n d b u c h der anorganischen Chemie,” 1‘01. XXX, p 283, Berlin, Verlag Chemie, 1932. Kolthoff, I. M., a n d F u r m a n , hT. H., “Volumetric Analysis,” Yo]. 11, p. 373, New Y o r k , J o h n TTiley & Sons, 1929. Letts. E. -A,. J . Chem. Soc.. 23. 427 11870). Luchinski, G. P., a n d Suadaleva, V. S . , J . ken. Chem., U.R.S.R., 10,2047-51 (1940). AIanchec, .I., J . pharm. chim., 21, 481 (1935). Xlellor, J. W., “Comprehensive Treatise on Inorganic a n d Theoretical C h e m i s t r y , ” Vol. 5,p. 544, London, Longmans, Green a n d Co., 1930. llutiiianski, h i . , 2. anal. Chem., 36, 220-1 (1897). P a p e , C . , Pogg. Ann., 122, 414 (1864). P l i m p t o n , R. T., a n d Chorley, J. C., J . Chem. Soc., 67, 314 (1895). Rammelsberg, C., P o g g . Ann., 56, 300 (1842). Rose, H , Ibzd., 21, 441 (1831).
RECEIVED for review August 28, 1952, Accepted February 16, 1953. Based on a thesis presented by Owen H. Kriege to the Graduate School of Ohio State University in partial fulfillment of the rerjriirenients for the degree of master of science.
Determination of Disulfides in the Presence of Thiols T. E. EARLE Research Department, Standard Oil Co. (Indiana), Whiting, Znd.
0
S E of the methods for the desulfurization of gasoline is the extraction of thiols by aqueous alkaline solutions of cresols. Spent cresol solutions are regenerated by air oxidation of the alkali mercaptides to disulfides, which are sparingly soluble in aqueous alkali and can be readily separated. If the disulfides are not completely removed from the alkaline solution, they ill contaminate the next gasoline treated. Disulfides are undesirable in gasoline because they reduce the ability of tetraethyllead to improve octane number. Disulfides are usually determined by an indirect method that measures the increase in thiol content after reduction of the disulfides to thiols ( 5 ) . This determination by difference introduces large errors when the amount of thiols present is large in comparison with the amount of disulfides. In seeking a method by which disulfides could be determined directly, Consideration was given to methods utilizing precipitation
or extraction to eliminate the interfering thiols. Removal of thiols as silver salts ( 5 ) ,or as other insoluble heavy-metal mercaptides, was unattractive because of the gummy nature of the precipitate formed by the mixed thiols present in petroleum distillates. Removal of thiols by extraction with alkali is not quantitative, and air oxidation of thiols to disulfides is accelerated in alkaline media. A specific reaction of the thiols was therefore sought, the reagents and products of which would not interfere in the subsequent reduction of disulfide to thiols or in the titration of thiols with silver nitrate. A4crylonitrileand other conjugated unsaturated nitriles, esters, and ketones react with thiols in the presence of small amounts of alkaline condensing agents to form thioethers (2-4). This reaction suggested the use of an excess of an unsaturated compound of this type to remove thiols quantitatively from mixtures x i t h disulfides.
