Spectrophotometric Titration of Uranium and iron CLARK E. BRICKER AND PHILIP B. SWEETSER Department of Chemistry, Princeton University, Princeton, N . J . uranium(1V) and iron(III), and the effect of induced air oxidation of uranium(1V) in the presence of iron(I1) ions were studied, iMilligram quantities of uranium and iron individually can be determined in a volume of 90 to 100 ml. with an accuracy of about 3 parts per thousand. With mixtures of these two ions, the accuracy of the iron titration is about 8 parts per thousand, whereas that of the uranium determination is about 6 parts per thousand. These titrations are rapid and accurate, require no modification of a spectrophotometer, and utilize very little special equipment.
Spectrophotometric titrations have been shown recently to be extremely sensitive and accurate in a wide variety of applications. The use of this method for detecting the end point in determinations of uranium and iron when present separately or in mixtures with ceric sulfate has not been studied previously. Procedures which involve the reduction of the uranium and/or iron solutions with a 90% cadmium amalgam reductor followed by titration with a standard ceric sulfate solution are given. The acid concentration, the composition of the cadmium reductor, the rate of the oxidation between
T
traces of the nitrate nere removed. The solution was then standardized m ith the cerium(1V) sulfate stock solution using a Jones reductor and ferrous o-phenanthroline as indicator. The final uranyl sulfate solution (0.01059 IV)was made from the stock solution of uranium and was 2 iV in sulfuric acid.
H E R E have been a considerable number of methods in the literature for the determination of various amounts of uranium (5); however, only a relatively few of these procedures are suitable for the determination of uranium and iron in the presence of each other. Ewing and Eldridge (3) described a method for the simultaneous titration of uranium(1V) ion and ferrous ion using permanganate or dichromate as the titrant. While this potentiometric method is suitable for larger amounts of uranium and iron, it has serious limitations in the titration of very dilute solutions of these ions. Recently, Crouthamel and Johnson (2) have described a spectrophotometric method for the determination of uranium in the presence of iron based on the formation of the thiocyanate complex in an acetone-water solvent. The use of the spectrophotometer for detecting the end point in titrations with dilute cerium(1V) sulfate has been shown to be a very sensitive and accurate method of end point detection (1). A spectrophotometric titration has been described for the simultaneous determination of arsenic and antimony using a bromatebromide solution as titrant (8). In this paper, procedures are given for the determination of milligram quantities of uranium and iron and for the simultaneous determination of both uranium and iron by a spectrophotometric titration with cerium(1V) sulfate.
PROCEDURE
Determination of Uranium. The 90% cadmium reductor was washed before each run with 20 to 30 ml. of 2 S sulfuric acid. If the reductor was not in use for a considerable period of time (overnight or longer), the reductor n'as JT-ashedwith 100 ml. of the dilute sulfuric acid which was usually sufficient so that an additional 30 ml. of the 2 S acid gave no detectable blank titration. Occasionally it is advisable to wash the reductor \Tith 2 A- hydrochloric acid or 2 S perchloric acid, folloaed by 2 -Ysulfuric acid.
APPAR4TUS AND REAGENTS
A Beckman Model B spectrophotometer was used for all titrations. S o modifications to this spectrophotometer are needed other than a titration cell and cover. The T-shaped titration cell with a i-cm. light path, 10-ml. microburet, and an electric stirrer were the same as previously reported (9). The 90% cadmium reductor was prepared from fine cadmium filings (ea. 20 grams of 30- to 60-mesh) which were washed with 1 N hydrochloric acid and then treated with the proper amount of 2% mercuric chloride solution to make the amalgam 90% cadmium. The resulting amalgam was washed thoroughlv n ith 1 h ' hydrochloric acid and then tightly packed into a column of 0.5 em. in diameter which gave a reductor ~ i t ah total height of approximately 20 em. Very dilute sulfuric acid was kept in the column when the reductor was not in use. A standard stock solution of 0.09374 N cerium(1V) sulfate was prepared from Ce(HSO4)a (G. F. Smith Chemical Co.) and standardized against arsenious oxide by the usual procedure ( 7 ) . The 0.009374 S cerium(1V) solution was prepared by taking an appropriate aliquot of the standard stock solution and diluting it to volume with 2 N sulfuric acid. The iron(II1) solution used for making the synthetic samples for analysis and for the titration of the uranium(1V) was prepared by dissolving reagent grade ferric sulfate in concentrated sulfuric acid and diluting the resulting mixture to the proper volume, leaving the mixture 2 N in the acid. This solution was standardized against the cerium(1V) sulfate using a Jones reductor for the reduction of the iron. A stock solution of uranyl sulfate (0.2119 N ) was prepared from reagent grade uranyl nitrate by first evaporating a sulfuric acid solution of the nitrate to fumes of sulfur trioxide until the last
I 2.0
2.5
3.0
3.5
4.0
ML. OF 9.374 X
Figure 1.
4 3
5 0
5.5
I
lo" N clCS0.h
Spectrophotometric Titrations Standard Cerium(1V) Sulfate
with
A . Iron(I1) solution B. UraniumW) srrlution Wave length, 360 mw.
The uranium sample, if not present as the sulfate, is evaporated with a slight excess of sulfuric acid to fumes of sulfur trioxide and is then dissolved in 2 N sulfuric acid. An aliquot of appropriate size of the resulting solution is passed through the reductor a t a rate of about 4 ml. per minute and is then washed with 30 ml. of 2 N acid. The last 20 ml. of the wash acid may be passed through the reductor a t a much greater rate than specified for the uranium solution. The reduced solution is collected directly in the titration cell which has been deaerated with nitrogen. The resulting solution is diluted with distilled wate! to make a total volume of about 90 ml. An atmosphere of nitrogen is kept over the uranium(1V) solution a t all times. The cell is placed in the spectrophotometer and the tip of the buret and the stirrer rod are immersed in the solution. A wave length for the titration is 764
V O L U M E 25, NO. 5, M A Y 1 9 5 3
765
i I 2.0
2.5
30
40 4.5 ML. OF 9,374 x
5.0 5.5 N. C ~ ( S O + J ~
6.0
6.5
Figure 2. Spectrophotometric Titration of Uranium(1V) and Iron(I1) Solution with Standard Cerium(1V) Sulfate A.
Uranium e n d point
E . Iron e n d point
selected on the basis of the normality of the cerium(1V) solution (360 mp was employed with the ca. 0.01 A' titrant). The slit width of the instrument is adjusted so as to give an absorbancy reading of zero for the solution. Cerium(1V) solution is then added from the microburet, taking only one or two absorbancy readings before the end point since the uranium(1V) and uranium(V1) ions have very small extinction coefficients a t this wave length. After the end point has been reached, as indicated by a sudden increase in the absorbancy reading, three or.four 0.10-ml. aliquots of the titrant are then added and the absorbancy is recorded after each addition. The end point is determined bv plotting the absorbancy us. milliliters of reagent added as shown in Figure 1. Determination of Iron. The procedure for the determination of iron with cerium(1T:) is identical with that described for uranium, except that in t,he actual titrat'ion of the iron(I1) with cerium(1V) more points are taken prior to the end point (every 0.50 to 1.00 ml.), since iron(II1) has a considerable ext,inction coefficient a t t h e wave length employed (360 mp). A resulting titration curve is shown in Figure 1. Determination of Uranium and Iron in a Mixture. The eimultaneous determination of uranium and iron is complicated by the fact that when iron is present in a uranium solution there is an induced oxidation of the uranium(1V) eit'her by the dissolved oxygen or 11)- impurities in the solution. The extent of this error is of the order of 2 to 8% depending upon the various factors discusqed later. The following procedure was designed to overcome this error as much as possible while still retaining the convenience of a single titration for the determination of both uranium and iron. .4lt,hough the acidity of the solutions for the individual determinations of uranium and iron is not critical and can be a t least 2 in sulfuric acid, the final reduced solution for the simultaneous titration of uranium and iron should be between 0.2 and 0.4 N in sulfuric acid. The sample, if not present as the sulfate, 18 evaporated with a slight excess of sulfuric acid and then diluted 80 that the aliquot of sample taken for each analysis does not contain over 20 meq. of sulfuric acid. When such a sample is reduced, washed through the reductor with 30 ml. of 0.2 S SUIfuric acid, and diluted to 90 ml. with water, the final solution will be about 0.3 .V in acid. For best results a preliminary titration must be made using a Drocedure similar to that described for the uranium. The SUIfuric acid solution containing uranium and iron is passed through the 90% cadmium reductor a t a rate of 3 to 4 ml. a minute and is then washed x i t h 30 ml. of 0.2 N sulfuric acid. The reduced solution is caught as before in the titration cell which has been deaerated n-ith nitrogen and the solution is diluted to 90 ml. with water. An atmosphere of nitrogen is kept over the solution in the cell a t all times. The titration is carried out as in the uranium titration until the absorbancy reading approaches 0.20 within one half a minute after the addition of the cerium(1V) sulfate solution. The reaction a t this point is slow SO that approximatelv 2 minutes should elapse before the first absorbancy reading beyond the uranium end point is taken. This time interval all ox^ equilibrium to be established between iron(II1) and the last traces of uranium(1V) SO t h a t the absorbancy reading is a true
measure of the iron(II1) present when the uranium(1V) is completely titrated. Then two or three additional 0.20-ml. aliquots of the titrant are added and absorbancy readings are taken immediately after each addition. It is also convenient a t this point to obtain approximately the amount of iron(I1) present. This may be done quickly by merely observing the rate of increase in the absorbancy as the titrant is added s l o ~ l y . A sudden increase in the absorbancy will indicate the end point of the iron(I1) titration. X o plot need be made of this latter titration since only a rough estimate of the end point is required. After the preliminary titration has been completed and the uranium end point has been determined by the usual plotting method, the titration of a similar aliquot of the uranium and iron solution is repeated, but this time 90 to 97% (an accurately measured volume) of the theoretical amount of cerium(1V) required for the uranium end point is added to the titration cell before collecting the reduced uranium and iron. The reduction procedure is otherwise the same as in the preliminary step. The wave length is set a t 340 to 345 mp and the elit is adjusted so that the absorbancy reading is zero. Cerium(1V) solution is added from the microburet as before until the absorbancy reading approaches 0.20 1% ithin approximately one half a minute after the addition of the cerium(1V); 2 or 3 minutes are allowed before recording the absorbancy so that equilibrium is attained. Then three or four 0.20-ml. aliquots of the cerium(1V) are added and, since equilibrium is now reached very quickly, the absorbancy readings are taken immediately after each addition. If the iron(11)end point is not more than 2.0 to 2.5 ml. beyond the uranium end point, the titration with the cerium(1V) may be continued until the iron(I1) end point is reached whereupon three to four readings are taken after the addition of 0.10-ml. portions of the titrant. If the amount of cerium(1V) required to titrate the iron(11) is much greater than 2.0 to 2.5 ml. the absorbancy would soon become prohibitively large. Therefore, a more convenient procedure is to reset the spectrophotometer, after obtaining the uranium end point, to a wave length of 360 mp and to make the absorbancy reading zero again with the slit. The titration is then continued using 0.5- to 1.0-ml. aliquots of the titrant between absorbancy readings until the iron(I1) is all titrated, whereupon the usual 0.10-ml, portions of cerium(I\') are added an 1 absorbancies are taken. 0.61
\
WAVE LENBTH,
MY
Absorption Spectra in 2 N Sulfuric Acid
Figure 3. 1. 2. 3.
5.30 X 10-5 M uranium(V1) sulfate 1.19 X 10-4 .M iron(II1) sulfate 9.37 X 10-5 cerium(1V) sulfate
The resulting plot of the recorded absorbancies will give a plot similar to Figure 2 where the first portion of the curve corresponds to the amount of titrant used in the oyidation of uranium(IV), and the second portion to that used in the oxidation of the iron(11) to iron(II1). DISCUSSION
Spectra of Ions Involved. The absorption spectra of uranium(VI), cerium(1V) sulfate, and ferric sulfate are given in Figure 3. All spectra viere taken in about 2 N sulfuric acid medium using a Beckman ;\lode1 DG spectrophotometer with 1-cm. cells. The choice of a working wave length is dependent mainly upon the concentrations of the iron( 11) solution and the titrant employed. Although the absorption maximum for cerium( IV) sulfate is 320 mp, a wave length of 360 mp was employed in the titration of
ANALYTICAL CHEMISTRY
766 Table I.
Titration of Uranium(1V) Solutions with 0.009374 N Cerium(1V) Sulfate (90% cadmium reductor used) Uranium Uranium Taken, Found, hIg. Mg. 0.631 0.632 0.893 0.883 1.261 1.261 1.260 1.261 2.536 2.522 3.795 3.783 5.382 5.422 6.269 6.258 6.269 6.290 7.568 7.566 8.814 8.827
Error, Mg. 0.001 0.010 0.000 0.001 0.014 0.012 0.040 0.011 0,021 0,002
0.013
uranium with the ca. 0.01 S cerium(1V) since there is less danger of overstepping the end point to such an extent that abnormally large absorbancy readings would result. Likewise, too short wave lengths must be avoided in the titration of larger amounts of iron since iron(II1) solutions show considerable absorption a t the shorter wave lengths. Holvever, too long wave lengths would decrease the sensitivity of the uranium end point in the simultaneous uranium and iron titration because of the decreased absorption of the iron(II1). The procedure of selecting a shorter wave length for the uranium-iron end point with the subsequent increase of the wave length for the titration of larger amounts of the iron(I1) gives much better sensitivity. Table 11. Titration of Iron(I1) Solution with 0.009374 N Cerium(1V) Sulfate Iron Taken, hf g. 0.328 0.983 1.966 3.280
(90% cadmium reductor used) Iron Found, Mg.
0.329 0.985 1.975 3.288
-
Error, Mg. 0.001 0.002 0,009 0.008
Choice of Reductor. The number of metal reductors available for the reduction of uranium(V1) to the quadrivalent state is limited by the tendency of many reductors to give nonreproducible amounts of trivalent uranium, and also by the type of acid required for the reduction. The most desirable reductor for this work would be one that gave quantitative reduction of uranium(V1) to uranium(1V) in a sulfuric acid medium. A cadmium amalgam reductor has been recently reported by Furman, Bricker, and Dilts (4) to have these desirable qualities. This reductor n-as prepared by adding cadmium turnings to a boiling mixture of mercury and dilute sulfuric acid until no more cadmium would dissolve in the mercury. Preliminary studies with this reductor showed a rather troublesome blank when the reductor was washed with dilute sulfuric acid and also the rate of reduction of the uranium was quite s l o ~ . If the wish solution \vas deaerated with nitrogen and extreme measures were taken to evclude oxygen from the reductor, the hlank was reduced to about 0.01 to 0.02 ml. of 0.01 S cerium(1V) for 30 ml. of 2 N sulfuric acid through the column. A study of the effect in the variation of the cadmium concentration in the amalgam revealed that a 90% cadmium reductor (preparation given under Apparatus and Reagents) gave a much faster rate of reduction and little or no blank even with acid solutions which had not been deaerated. When the amount of cadmium \vas increased to 98,99.9, and 100% the blank continued to be negligible, but the uranium end points were from 2 to 3y0 high, indicating formation of trivalent uranium. The use of an 85 to 90% cadmium amalgam therefore seemed to be a more practical reductor and was employed in all of the reported titrations. Induced Oxidation and Rate of Reactions. The rate of oxidation of uranium(1V) and iron(I1) with cerium(1V) solution is quite rapid a t room temperature although the uranium(1V)
oxidation is slightly slow in the region of the end point. Since the end point in photometric titrations is determined by extrapolation rather than by titration to the exact end point, this slowness is scarcely noticeable. On the other hand, in the simultaneous titration of uranium and iron with cerium(1V) the rate of the reaction near the uranium end point is considerably sloRTer than with the uranium and cerium(1V) alone. This is due to the slow reaction between iron(II1) and the uranium(IF'), where some of the iron(I1) is prematurely oxidized with the cerium(1V) before all the uranium has been titrated. The rate of the reaction betn-een iron(II1) and uranium(1V) is very dependent upon the acid concentration. In 6 S sulfuric acid the titration of uranium(1V) with ca. 0.01 N ferric sulfate requires over 15 minutes to reach equilibrium even when less than one half of the uranium has been titrated. In 2 S acid a t least 10 minutes are required for equilibrium when 90% of the uranium has been oxidized. However, if the final acid concentration is reduced to 0.2 the uranium-iron(II1) reaction will reach equilihrium in only 2 minutes even in the region of the end point. I t is therefore essential that the acid concentration be kept low in all titrations involving uranium and iron. Table 111. Simultaneous Determination of Uranium and Iron Mixtures Uranium Taken,
Uranium Iron Iron Error, Taken, Found, Found, Mg. Mg. Mg. Alg. Mg. 1.261 1.250 -0.011 1.359 1.379 1.261 1.263 f0.002 2.667 2.638 -0.015 4.005 1.261 1.246 4.012 2.523 2.524 t0.001 0.680 0.693 -0,003 2.003 2.521 2.518 2.029 -0.006 3.338 2.521 2.515 3.354 1.352 5,019 -0,018 5.037 1.352 -0.032 2.000 5.046 5.014 2.027 -0.033 3.982 5.042 5.009 3.989 7.569 7.574 +0.005 0.677 0.684 2.533 1-0.014 33.45 ... 2.519 +0.002 66.75 ... 2.522 2.524 33.43 ... 3.747 -0.037 3.784 2.001 -0.163 2.523a 2.360 2.008 1.354 -0.321 1.332 3.783' 3.462 -0,151 1.345 5.04P 4.895 1.352 a Ceric(1V) sulfate solution not added t o titration cell before uranium-iron solution.
Error, hIg. 1-0.020 -0,029 +0.007 4-0.013 +0.026 $0.016 0.000 4-0.027 +0.007 +0.007
..... .....
..... +0.007 -0.022 +0.007 reduction of
By far the most serious difficulty encountered in the titrations of uranium and iron with cerium(1V) was the induced oxidation of uranium(1V) in the presence of iron(I1). A recent paper hy Sill and Peterson (6) has shown that many ions cause an induced oxidation of uranium(1V). With the conditions employed in that paper there was little induced oxidation caused by iron; however, the results in our study of the simultaneous titrations were invariably 2 to 8% Ion for the uranium(1V) end point (see Table 111). Titrations of uranium(1V) with ferric sulfate also shoxed similar low results, moreover, if the resulting iron(I1) formed from these titrations was subsequently titrated u ith cerium(IV), the amount of iron(I1) found was again low as compared with the uranium originally present. The amount of induced oxidation was quite constant provided the acidity of the solution remained the same. The error was absolute rather than relative. il change in the acid concentration of the solution not only affected the rate of oxidation of the uranium(1V) by the iron(III), but also influenced the amount of induced oxidation of the uranium(1V). Table IF' shows the effect of varying the acidity in the titration of uranium(1V) with standard ferric sulfate solution. These titrations xere made using a procedure similar to that already described for uranium with rerium(IY), except that a wave length of 340 mp was employed. I n order to eliminate the induced oxidation of the uranium(IV), a procedure was adopted whereby 90 to 97% of the theorrtical amount of cerium(1V) required for the titration of the uranium(1V) was added to the titration cell prior to the reduction of the uranium-iron solution. This not only allowed less contact of the
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] Uranium Taken, Mg.
6.270 6.270 6.270 6.270 6.270
Uranium Found, 3Ig. 6.001 6.004 6.057 6.111 6.158
Error, 3Ig.
0.269 0.266 0.213 0.15g 0.112
Total Acidity (HzSOd of Reduced Solution.
.v
0.14 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. (8) Sweetser, P. B., and Bricker, C. E., - 4 s . i ~CHEY., . 24, 1107 (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 t o he submitted by P. B. Sweetser in partial fulfillment of t h e 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.
