for our purpose, and we can analyze more than 20 samples and standards per hour.
ACKNOWLEDGMENT The authors thank R. L. Berger, NHLI, and E. J. Prosen, NBS, for their advice on calorimetry and the thermochemistry of the reaction of LiOH with C02.
LITERATURE CITED (1) M. Burg and N. Green, Kidney Int., 4, 301 (1973). (2) C. R. Caflisch and N. W. Carter, Anal. Biochem., 60, 252 (1974). (3) B. Karlmark and M. Sohtell, Anal. Biochem., 53, 1 (1973).
k,:,
",";tl~,n~'n~~~~h,*~~~:95629),968),
(6) F. J. A. Prop, f x p . Cell Res., 7,303 (1954). (7) F. W. Williams, F. J. Woods, and M. E. Umstead. J. Chromatogr. Sci., 10, 570 (1972). (8) F. Hevert, Pflugers Arch., 344, 271 (1973). (9) w. J. Jones, US. Patent 3716337 (1973). (10) D. A. Borfla and A. J. Mass, hd. Eng. Chem., Process Des. Dev., IO, 489 (1971). (11) S. Dushman,
"Scientific Foundations of Vacuum Technique", John
Wiley and Sons,New 'fork, N.Y.. 1949, p 84. (12) D. D. Williams and R. R . Miller, lnd. f n g . Chem., fundam., 9, 454 (1970).
RECEIVEDfor review September 6, 1974. Accepted December 19,1974.
Convenient Preparation of Standard Thiosulfate Solutions Edwin H. Funk and Albert W. Herlinger Department of Chemistry, Loyola University of Chicago, Chicago, IL 60626
The preparation of standard thiosulfate solutions of sufficiently high concentration to be of practical utility in iodometric titrations has relied almost exclusively upon sodium thiosulfate pentahydrate. Although standard thiosulfate solutions may be prepared by direct weighing of Na2S203.5H20 or the anhydrous sodium salt (1,2),special reagents and care are necessary to ensure a definite composition of these materials. Additionally, solutions of sodium thiosulfate slowly change in titer with time and require periodic restandardization. In actual practice, approximate solutions using ordinary reagents are prepared and standardized against primary standard materials. To obviate the problems associated with Na2S203 solutions, we initiated a search for a primary standard thiosulfate salt which would exhibit good stability characteristics. A convenient method for preparing standard thiosulfate solutions utilizing barium thiosulfate monohydrate as a primary standard has been devised. Plimpton and Chorley, as early as 1895, suggested the use of this salt as a primary iodometric standard ( 3 ) . However, it was not until 1953 that there were any quantitative data ( 4 ) to support this suggestion. Barium thiosulfate monohydrate, which may be obtained commercially from several sources, has a high equivalent weight (267.51 amu) and can be dried to constant weight without special procedure at 40 "C. Above about 50 OC, the hydrate slowly loses water, and the effects 'of high drying temperatures have been reported previously ( 4 ) . The solubility of the salt, 0.283 gram of Bas203 per 100 cm3 of saturated solution at 25 "C ( 5 ) , is insufficient to allow preparation of a 0.1Nsolution. Consequently, its application has been restricted almost exclusively for use as a solid in the standardization of 12 solutions. We have utilized the complexing ability of EDTA, disodium (ethylenedinitri1o)tetraacetic acid dihydrate, to increase the solubility of BaS203.H20. The conditional stability constant for the EDTA-Ba2+ complex (log K = 6.4) a t pH 9 (6), the optimum pH region for stabilizing thiosulfate solutions, indicates that the EDTA concentration should be approximately 0.1F if the desired concentration of S 2 0 3 2 - is to be achieved. Thiosulfate solutions prepared by direct weighing of BaS203. H20 and addition of EDTA are shown to be suitable for several applications of the iodometric method; however, the presence of EDTA may interfere in determinations in which the redox potential or reversibility of the
.
reaction are adversely affected. The barium thiosulfate solutions appear to have the additional advantage of stability characteristics which are superior to the more conventionally prepared thiosulfate solutions.
EXPERIMENTAL Apparatus. The p H of the solutions was determined with a Corning Model 12 research p H meter. Reagents. Primary standard grade barium thiosulfate monohydrate was purchased from Sargent-Welch Scientific and was used without further purification. Primary standard grade KI03, KzCrz07, and copper were used. Hydrogen peroxide was 3?6 Fisher certified reagent grade Hz02 and commercial Clorox was used as an unknown hypochlorite solution. The EDTA and all other chemicals were analytical reagent grade materials and were used without further purification. Procedure. Finely ground BaSz03. HzO was dried at 40 "C for 2 hours and weighed exactly to prepare 0 . 1 N thiosulfate solutions (26.751 g k ) . Two methods were used in preparing the standard thiosulfate solutions with equally good results. In both cases, triple distilled water which had been recently boiled was used. Method A . An amount of EDTA (37.224 g/L) equivalent to the previously weighed Bas203 H20 (26.751 g/l.) was dissolved in exactly twice the equivalent amount of 1 N NaOH (188 ml of 1.065N base). This solution was used to dissolve the BaS203. HzO which had been placed in a 400-ml beaker. The resulting mixture was stirred with a magnetic stirrer to facilitate rapid and complete dissolution and then transferred quantitatively to a volumetric flask. After diluting to the mark and mixing thoroughly, the solution exhibited a p H of 9.5. Method B. The procedure outlined in Method A was followed except that the requisite amounts of EDTA and NaOH calculated from the weight of BaS203. H z 0 to be dissolved were weighed to within 0.1 gram and added to 150 ml of H20. Solutions prepared in this manner have a p H in the range of 8 to 10 and are as applicable as the more carefully prepared solutions in Method A. Large excess of base, i.e., 2 or 3 grams more than that necessary to neutralize the EDTA, causes the titer of the solution to be high and this is to be avoided. Standardization, Stability, and Applicability of Bas203 Solutions. Standardization of the thiosulfate solutions was accomplished by titrating vs. potassium iodate using the method described by Kolthoff (7). These solutions were then restandardized at approximately weekly intervals and found to be stable for a t least six weeks when excess base was avoided. The applicability of the Bas203 solutions was tested in the iodometric determinations of dichromate, iodate, hydrogen peroxide, hypochlorite, and copper using standard procedures (8). Generally, acidifications of the solutions to be titrated were accomplished using hydrochloric acid in preference to sulfuric acid since it was felt that the precipitation A N A L Y T I C A L C H E M I S T R Y . VOL. 4 7 , N O . 4 , APRIL 1 9 7 5
767
~~~~~~~
Table I. Thiosulfate Normality at Various EDTA and Base Concentrations Sample
EDTA
KaOH
No. 1
(formality)
(formality)
Theoreticala
Observedb
0.09994 0.1002 0.1004 0.1005 0.1005 0.1010 0.1080 0.1104 0.1000 0.1002
0.1840 0.2002 0.2000 0.2025 0.2026 0.2117 0.2802 0.2798 0.3600 0.4004
0.09969 0.09982 0.1025 0.09977 0.09975 0.09969 0.1025 0.1025 0.1025 0.1025
0.09975 (0.6) 0.09984 (1.3) 0.1026 (1.5) 0.09984 (5.1) 0.09975 (0.5) 0.09982 (3.7) 0.1023 (2.2) 0.1025 (0.7) 0.1033 (3.9)c 0.1091 (2.9)c
2 3
4 5 6 7 8 9 10
Standard
Calculated from weight of barium thiosulfate monohydrate. Determined by titration against potassium iodate and unchanged after six weeks. Standard deviation X 104 given in parentheses. c Observed normality after 4 weeks.
of Bas04 might be objectionable. However, for one 1 0 7 - determination, sulfuric acid was used and the end point was not obscured by the precipitate which formed. In f a k , good agreement between the theoretical and observed values was obtained in this instance.
RESULTS A N D DISCUSSION These studies have shown that 0. IN thiosulfate solutions may be prepared by direct weighing of primary standard Bas203 H20 when the solubility of the salt is increased by addition of EDTA. Agreement of better than 1 part in 1000 between the calculated normality of the thiosulfate solution based on a weighed amount of BaS203 H20 and that determined by titration vs. primary standard KI03 was observed (Table I). The effect of variations in the formal concentration of EDTA and NaOH on the observed normality of the thiosulfate solutions was investigated and found to be insensitive to small changes in these reagents, Table I. These findings led to the development of the rapid procedure described in Method B for preparing standard thiosulfate solutions which are suitable for iodometric titrations without prior standardization. A large excess of base, as demonstrated by samples 9 and 10 in Table I, caused the titer to be considerably above (ca. 10%) the theoretical value, reminiscent of the errors which occur when titrating iodine in weakly alkaline solutions. Consequently, use of this alkaline buffered standard solution when titrating iodine in unbuffered or very slightly buffered solutions of p H 6 to 4 could result in errors due to partial oxidation of thiosulfate to sulfate. The stability of these solutions is good; 0.1N and 0.01N barium thiosulfate solutions remained unchanged for a t least 6 weeks without added preservatives. The stability of the solutions appears to be independent of small variation in the formal concentration of EDTA or base. Additionally, solutions prepared using deionized water which was not boiled have also remained unchanged for six weeks. Table I1 shows the applicability of Bas203 solutions with added EDTA to determinations which use the iodometric method. Theoretical values for IO,3-, Cr*O.;"-, and copper determinations were provided by weighing these materials as primary standards. For the H202 and hypochlorite determinations, commercial solutions of these reagents were diluted to approximately 0.1N and the theoretical concentrations were provided by titrating the 12, liberated after reaction with KI with a standardized Na2S203 solution. In general, excellent agreement between observed and theoretical values was obtained and suggests that these solutions should also be applicable in determinations of organic
-
768
Table 11. Iodometric Determinations Utilizing Barium Thiosulfate Solutions
A N A L Y T I C A L CHEMISTRY, VOL. 47, N O . 4, APRIL 1975
Determination
Io3-
cr,o,?Cr,07'-
Theoretical
Observed
(noma1ity)a
(norma1ity)b
0.1025 0.09982 0.09982
0.1025 0.0992 1 0.09979
lo4) 0.7 0.9 0.8
0.1071 0.1643
0.1071 0.1645
0.4 4.6
deviation (x
[ 3 m l additional i o n i d HC1)
W? oc1-
Potassium iodate and potassium dichromate were weighed as primary standards. Peroxide and hypochlorite values determined iodometrically using Na2S203 solutions. Reported values are averages of at least three determinations. (1
peroxides, free halogens, and the other oxyhalides. In one instance (the dichromate determination), it was necessary to deviate slightly from the normal procedure (8), an additional 3 ml of concd HCl was necessary for'the reaction to proceed quantitatively. The reaction of KI with Cr2072does not take place instantaneously but the speed increases with increasing hydrogen ion concentration. Consequently, the reaction must be carried out in a highly acidic solution. The results for this determination using standard procedures were 0.6% low indicating that the reaction had not proceeded to completion (see data in Table 11). Apparently, the addition of a thiosulfate solution containing EDTA with all acid groups neutralized raised the pH sufficiently to prevent quantitative reaction between KI and K2Cr207. When the acidity of the dichromate solution was maintained by additional HC1, excellent agreement between the observed and theoretical values was obtained. These titrations were carried out under a blanket of CO2 to prevent air oxidation of iodide which may occur in highly acidic solution. However, later work showed these precautions to be unnecessary. These results imply that BaS203 solutions may also be useful in determinations of barium and lead which are precipitated as the chromate salts in neutral medium. The determination of copper by the iodometric method using barium thiosulfate solutions with added EDTA is hindered by the formation of a stable CuZ+-EDTA complex. Controlling this effect by variation of the pH was not feasible as even a t p H 1.0 the complex has an apparent stability constant of ca. 2.5 X lo3 (6). Consequently, these solutions cannot be used in this application of the iodometric method. Similar difficulties are anticipated in the iodometric determination of iron using thiosulfate solutions prepared as described above. LITERATURE CITED (1) S. W. Young, J. Am. Chem. SOC.,26, 1028 (1904). (2) H. M. Tomlinson and F. G. Ciapetta, Ind. Eng. Chem., Anal. Ed., 13, 539 (1941). (3) R. T. Plimpton and J. C. Chorley, J. Chem. SOC.,67,314 (1895). (4) W. M. MacNevin and 0. H. Kriege, Anal. Chem., 25, 767 (1953). (5) T. 0.Denny and C. B. Monk, Trans. faraday SOC.,47,992 (1951). (6)A. E. Ringbom, "Complexation in Analytical Chemistry," Wiley-lnterscience. New York, N.Y., 1963, p 360. (7) I. M. Kolthoff, E. 6.Sandell, E. J. Meehan, and S. Bruckenstein, "Quantitative Chemical Analysis," 4th ed., Macmillan Co., London, 1969, p 849. (8) I. M. Kolthoff and R . Belcher. "Volumetric Analysis." Vol. Ill. 2nd ed., Wiiey-lnterscience, New York, N.Y., 1957.
RECEIVEDfor review August 29, 1974. Accepted December 31, 1974. Acknowledgement is made to the Donors of the Petroleum Research Fund, administered by the American Chemical Society, for partial support of this research.
