Coulometric Titration of Thiosulfate with Iodine Application to the Determination of Oxidizing Agents KEITH R O W L E Y and ERNEST H. SWIFT Gates and Crellin Laboratories of Chemistry, California Institute of Technology, Pasadena, Calif. When carbon dioxide was used for the titrations under oxygenfree conditions, the gas was bubbled through the water and the acid for 5 minutes to sweep out the dissolved oxygen. Then the thiosulfate and iodide solutions were pipetted into the titration cell, the cell was attached to the titration apparatus, and the titration was conducted under a stream of carbon dioxide. The methods of determining blank time and titration time, of adjusting the generator current, and of maintaining the sensitivity of the indicator electrodes were the same as those described by Ramsey, Farrington, and Swift (8). A potential difference of 135 mv. was impressed across the indicator electrodes. The generation rates used were 1.037 X and 1.036 X 10-8 equivalent per second.
There is a need for reducing intermediates for secondary coulometric titrations which will react completely, rapidly, and stoichiometrically with a large number of oxidizing constituents. In conventional volumetric titrations these requirements are frequently met by the addition of excess iodide with subsequent titration of the liberated iodine by thiosulfate. If thiosulfate could be titrated with electrolytically generated iodine, the coulometric process could be extended to many of the conventional iodometric determinations of oxidizingagents. This investigation was undertaken to determine the feasibility of the coulometric titration of thiosulfate with iodine, and to establish the limitingponditions. The results of two applications of this titration to the determination o f oxidizing agents are reported.
DISCUSSION
The data obtained from coubmetric titrations of 10 nil. of millinormal thiosulfate (approximately 1 mg. of thiosulfate) with iodine a t various pH values are presented in Table I, and show, in agreement with Bradbury and Hambly (2),that betvieen a pH of 1 and 8 titrations can be made with errors of less than 2 parts per thousand without exclusion of atmospheric oxygen. Titrations of 0.1-mg. quantities show the same absolute error, :tpproximately 1 y. The trend from a small positive error a t pH 8 to a small negative error a t a pH 1 is of doubtful experimental significance. At higher pH values positive errors are obtained. The data of Tables I, 11, and I11 show that in 0.6 tu 1.2VF hydrochloric acid solutions values within 2 parts per thousand ran be obtained in the absence of oxygen; larger negative errors rew l t in the presence of oxygen. These effects are discussed below.
A
LTHOUGH the titration of millinormal thiosulfate with millinormal iodine, with either the starch end point (6,9, IO) or the amperometric end point ( 5 4 , has found analytical applications, the optimum conditions have not been studied thoroughly until recently. Since the completion of experimental xork, two papers pertinent to this titration have been published. Tutzundic and Mladenovic ( 1 2 ) have titrated thiosulfate with elertrolytically generated iodine and have used the starch end point. They did not extend their investigation to solutions as dilute as those in this investigation nor did they investigate the permissible p H limits for the titration. The effect' of various p H values on the volumetric titration of millinormal thiosulfate solutions with millinormal iodine, with both the starch and the amperometric end points, has been studied by Bradbury and Hambly ( 2 ) . I n this work the effect of higher concentrations of arid has been investigated.
-
._.___
..
Table 1. Titrations of Thiosulfate at Various pH Valucsa Thiosulfate, BH
'3.7 9.0 8.0 7.0
EXPERIMENTAL
Chemicals. Reagent-grade chemicals were used. A standard solution of 0.1VF (volume formal) sodium thiosulfate was prepared by dissolving 25 grams of sodium thiosulfate pentahydrate and 0.1 gram of sodium carbonate in 1 liter of freshly boiled distilled water. The resulting solution was standardized against a standard solution of potassium iodate. The millinormal thiosulfate solutions used in the titrations were prepared by approriately diluting this standard solution with boiled distilled water. it was found that the iodine titer of a millinormal thiosulfatt, solution decreased significantly within 2 days, a fresh dilute thiosulfat,e solution was prepared each day that titrations were made. A 1.0VF potassium iodide solution was prepared by dissolving reagent grade potassium iodide in boiled distilled water. The solution was made 0.001VF in sodium carbonate in order to retard air oxidation of the iodide. The buffers listed by Ramsey, Farringt,on, and Swift ( 8 ) were used to control the p.H values of the solutions. Hydrochloric acid was used in order to obt'ain hydrogen ion concentrations of 0.1VF or greater. ~ ~ i distilled l ~ d water was used exclusively in making ul, solutions; boiling removed a small anlount of oxidizing agent from the distilled water and gave smaller blank corrections. Apparatus, Adjustment, and Procedure* The apparatus used was that described by XIeier, Myers, and Swift ( 7 ) with the ~~ ~ ~ and ~ SJvift ~~(8). The ~ ~pre-, modifications of R liminary adjustment,s and the titration procedure were essentially those described by Ramsey, Farrington, and Swift (8)with the following changes: I n place of the 25 ml. of arsenic (111) solution, 10 ml. of millinormal thiosulfate and 15 ml. of boiled distil!ed mater were pipetted into the titration cell, ~h~ titrated solutlon was 0.1ViF in potassium iodide and had a volume of 50 ml.
,j.0
i:;
'Ic1, 0 . 1V F 0.6
knee
a .
y
Number of Titrations
Taken
Found
2 3 5 6 3
1071 1138 1138 1071 107.2
1192 1158 1138 1071 106.5
2 4 3
1071 1138 1071
1070 1137 1070
-1
0.0 0 .!? 0.9
5 2
1138 107.2
1136 106.3
-2 -0.9
0.3
5
1138
1127
Error +121 +2;
0 -0.7 -1
-1
-11
Standard deviation 1 6 I 7 1.1 0.7
0.7
0.6
3.6
Oxygen not excluded. ~
-
Titrations at High p~ values, A t p~ values of 9 9.7 too much iodine was required and the indicator currents were unstable. At these pH values, the indicator current obtained with a blank solution in which iodine had been generated decreased slolvly on continued (without further generation iodine) thus indicating that hydrolysis of iodine was becoming significant (8); however, the decrease of indicator current was so slow that reproducible blanks could be obtained. \?%en a thiosulfate titration was made a t these pI-1 values, excess of iodine before was obtained i ~ was required ~ ~ an end point ~ ~ and the resultant indicator current then decreased much more rapidly than with the blanks. This indicates that during the to a higher oxiditof thethiosu]fakis being tion st#atethan tetrathionate, and that a t least one of the produrtu formed during the titration is being further oxidized by the excew
373
~
,
ANALYTICAL CHEMISTRY
374 Table 11.
Titrations of Thiosulfate in Hydrochloric Acid Solutions" Thiosulfate,
Number of HC1, V F Titrations Taken
Found
7
Error
Standard deviation
a All solutions except thiosulfate oxygen-free, titrations made under carbon dioxide.
Table 111. Titrations of Thiosulfate in Hydrochloric Acid Solutions" Thiosulfate, y Xumber of Standard HCl, V F Titrations Taken Found Error Deviation -2 0.4 1073 1071 0.6 3 -2 1.3 1073 1071 1.2 3 -8 0.8 1073 1065 2.4 2 0 All solutions oxygen-free, titrations made under carbon dioxide.
of iodine. The formation of sulfate during the oxidation of thiosulfate by iodine in basic solutions is known (11), and could account for the excess iodine required a t these high p H values (2). Titrations at Low pH Values. Excess iodine was required when solutions 0.05N in thiosulfate and having a p H of 1 were titrated volumetrically, presumably because of the following reactions:
SzOa" + 2H+ S + HpSOa + HZS03 + 21- + HSOn- + 3 H + +
HZ0
12
+
a possible explanation of the consistentlj- negative errors in these oxygen-free solutions.
(1)
(2)
However, the kinetics of reaction 1 apparently are such as t o cause smaller percentage errors in the titration of more dilute thiosulfate solutions (1). Analytical procedures have been proposed ( 3 , 5) involving the titration of an excess of millinormal thiosulfate in O.IN acid solutions, but it is doubtful if discrepancies of less than a few per cent could have been detected in the confirmatory work reported. About 2% excess of iodine was required when 10 ml. of millinormal thiosulfate were added t o 50 ml. of 0.5VF hydrochloric acid, the solution was allowed to stand for 25 minutes under carbon dioxide, then buffered to p H 7 , iodide added, and the solution titrated coulometrically. However, when the thiosulfate was added to the acid, iodide added a t once, and the titration made immediately without buffering the solution, less than the equivalent quantity of iodine was required a t the above and all other acid concentrations investigated. With these dilute solutions the air oxidation of sulfite and/or of iodide in acid solutions apparently has a much greater effect than the acid decomposition of thiosulfate. In obtaining the values shown in Tables I, 11, and I11 for titrations in hydrochloric acid solutions, the titration was made immediately after acidification. I n O.1VF hydrochloric acid titrations with errors of less than 2 parts per thousand can be made without the use of carbon dioxide. When a solution which was O.1VF in potassium iodide and O.1VF in hydrochloric acid was placed in a titration cell in the apparatus and stirred, the indicator current rose very slowly from air oxidation, but the magnitude of this effect did not cause significant errors in the subsequent titrations made in this concentration of acid. In 0.6V'F hydrochloric acid the error is approximately - 1% if oxygen is not excluded. If all of the solutions except the thiosulfate are swept with carbon dioxide for 5 minutes and the titration is done under carbon dioxide, the error decreases to 0.2 to 0.3%. When all solutions, including the dilute thiosulfate solution, were made up with water through which nitrogen had been passed, the error decreased slightly to less than -0.2'%. In 1.2VF hydrochloric acid, under the same conditions, the error was about the same, but rose to -0.8% when the titration was done in 2 . 4 V F hydrochloric acid. Loss of sulfur dioxide is
DETERMINATION OF OXIDIZING AGENTS
The determinations on a macroscale of both iron(II1) and chromium(V1) by adding an excess of potassium iodide to an acid solution and then titrating the liberated iodine with standard thiosulfate have long been standard analytical procedures; therefore these substances were selected as typical oxidizing agents with which to test the analytical application of the coulometric thiosulfate titration. The determinations were carried out by adding the solutions of iron(II1) or chromium(V1) t o an iodide solution which was O.1VF in hydrochloric acid, then adding an excess of standard thiosulfate solution and coulometrically titrating the excess with iodine. EXPERIMENT 4 L
Chemicals. A solution of O.IVF ferric chloride was prepared by dissolving the salt in O.1VF hydrochloric acid. This solution was standardized iodometrically. +4standard solution of 0. I S potassium dichromate was prepared by diluting to a known volume a weighed quantity of the salt which had been dried for 1 hour a t 150" C. The more dilute standard solutions were prepared by making appropriate dilutions of these standard solutions. At first the dilute iron solutions were prepared ith ordinary distilled water, but it was found necessary to use water through which carbon dioxide had been passed for 20 minutes and then to keep these solutions under carbon dioxide in order to avoid large oxygen errors. Titration Procedure. TI?enty milliliters of distilled water, 5 ml. of 1VF hydrochloric acid, and 5 ml. of 1VF potassium iodide Tvere placed in a titration cell and carbon dioxide mas passed through this solution for 5 minutes. Then 10 ml. of the iron(II1) or chromium(V1) solution were added. Ten milliliters of the dilute thiosulfate solution were added immediately, the solution was placed on the apparatus, and the excess thiosulfate was titrated a t once under carbon dioxide in the same manner as described above. Two methods were used t o obtain a correction for reagent impurities, etc. I n the first method, a blank was run under the same conditions as was the titration with 5 ml. of l V F potassium iodide, 5 ml. of 1 V F hydrochloric acid, and 40 ml. of xater. In the second method, which gave results similar to those obtained bv the first method, a thiosulfate titration was run as a blank. The blank solution was the same as the titration solution except that the iron(II1) or chromium(V1) solution was replaced by distilled water. The titration time was then subtracted from the blank time to give the net titration time. This method is preferred since, in addition to a correction for impurities, a standardization of the thiosulfate is also made under the conditions of the titration. DISCUSSION
Preliminary experiments showed that the titration of iron(II1) is satisfactory only if oxygen is carefully removed from all solutions. Oxygen dissolved in the solutions caused positive errors of about 6% in the titration of 500-mg. quantities of iron even though the titrations were done under carbon dioxide; a similar titration which was not done under carbon dioxide resulted in a positive error of 13%. The dilute standard solutions used for the titrations shown in Table IV were prepared by diluting the
Table I\'.
Confirmatory Determinations of Iron(II1) and Chromium(VI)a Iron, y
Number of Titrations 5 5 4
5
Found 5758 578.3 287.0 58.0
Error 3 3.0
Standard deviation 2.2
0.8 0.5
Chromium, y 158.3 158.5 0.2 Titrations done under carbon dioxide, solutions oxygen-free 5
a
Taken 5755 575.3 286.2 57.5
0.6
0.2 0.7
0.1
V O L U M E 26, NO. 2, F E B R U A R Y 1 9 5 4
375 Evans, D. P.. and Simmons, N. T., J . Soc. Chem. I n d . (London),
stock standard solutions with distilled water through which carbon dioxide had been passed for 20 minutes. Carbon dioxide was also passed through the other solutions used in these titrations and the titrations were done under carbon dioxide. These titration values are slightly high, and show an absolute error of 3 y with quantities of iron greater than 500 y and of less than I y with smaller quantities. The data from the dichromate titrations indicate that under the conditions of these titrations the errors are within those of the experimental measurements.
63. 29 f19441,
Harris, E: D., and Lindsey, d.J., Analust, 76, 647 (1951). Hewson, G. W., and Rees, R. L., J . Soc. Chem. Ind. (London), 54, 2 5 4 T (1935).
Kolthoff, I. XI., Z . anal. Chem., 6 0 , 341 (1921). Meier, D. J., Myers, R. *J., and Swift, E. H., J . A m . Cheni. Soc., 71, 2340 (1949). Ramsey, W. J., Farrington, P. S., and Swift, E. H., AXAL. CHEM.,22, 332 (1950).
Reith. J. F.. Biochena. Z.. 216. 249 11929). Sadusk, J. F., and Bell, E. G., IND.ENG CHEW.,SAL. ED.,5 ,
ACKNOWLEDGMENT
3 5 6 (1933).
Topf, G.. Z . anal. Chem., 26, 137 (1887). Tutrundic, P. S., and Mladenovic. S.,A d . Chzm. Acta, 8 , 184
The authors are indebted to Jerry C. Mitchell for aid in carrying out the confirmatory titrations of thiosulfate.
(1953).
LITER4TURE CITED
R E C E I V Sfor D reriew August 2-1, 1953. Accepted October 2 3 , 1953. PreCHEMICAL SOCIETY,Los s m t d a t the 123rd Meeting of the AMERICAN
(1) Bassett, H., and Durant, R. G., J . Chena. S o c , 1927, 1458. (2) Bradbury, J. H., and Hambly, A. K.,Australian J . Sci. Research. A5, 541 (1952).
Angelen. Calif., hlarch 1953. Contribution KO.1841, Gates and Crellin Laboratories of Cheniistry, California Institute of Technology, Pasadena, Calif.
Acetylacetone as an Analytical Extraction Agent Extraction of Aluminum, Gallium, and Indium JOHN
F. S T E I N B A C H and H E N R Y FREISER
D e p a r t m e n t o f Chemistry, University o f Pittsburgh, Pittsburgh 73, P a .
-is acetylacetone has already proved to be a useful and interesting extraction agent, the extraction characteristics of some trivalent metals, such as aluminum, gallium, and indium were studied. Equilibrium extraction curves for these metals by acetylacetone have been determined. The aqueous solubility of acetylacetone as a function of pH and ionic strength has been determined. The relation between the shape and position of the extraction
K
OLTHOFF and Sandell ( 6 ) and Irving and Williams ( 6 ) have derived the following equation for the distribution of metal chelates between organic and aqueous phases:
n here p , and p , are the pal tition coefficients of the chelate and ieagent hehveen the tn-o phases; K f is the over-all formation constant of the chelate, JIR,, and K , is the acid dissociation constant of the reagent, HIt. The subscripts o and zu refer to the organic and water phases, respectively. Usuallj- the solubility of the chelate is so much greater in the organic phase than it is in water that the term ( p c - D )reduces to p,. From this equation it can be seen that the distribution is dependent on two variables, ( H +) and (HR),. Consequently, two methods of verification may be employed. The p H may be maintained constant through use of a buffer and the distribution Ftudied as a function of reagent concentration. However, as pointed out by Irving ( 5 ) , more useful analytical data may be obtained by maintaining the reagent concentration constant and studying the distribution as a function of pH. Irving has shown that when the per cent extracted is plotted against the p H when the reagent concentration is constant, a sigmoid extraction curve whose shape is independent of its position along the p H axis is obtained. The slope of the curve depends only on the valence of the metal; the greater the valence, the greater the slope. The p H of 50% extraction, ( D = I ) , has been called the PHI/*value and the spread of these values for two metals is an indication of their separability in extraction processes.
curves and properties of the acetylacetone chelates is examined and the usefulness of such curves in predicting the feasibility of analytical separations and determinations is indicated. The factors which must be controlled in obtaining equilibrium extraction curves are discussed. The results indicate the possible separations of aluminum, gallium, and indium and have a more general bearing on the development of solvent extraction as an analytical tool.
However, it is rather difficult t o maintain a constant reagent concentration in the organic phase when a reagent such as dithizone is used in, for example, a chloroform-water system, because the distribution of the reagent itself is a function of p H and also the amount of reagent consumed in the formation of the chelate must be considered if the ratio of metal to reagent is large. If acetylacetone is used as both the solvent and reagent, the concentration of acetylacetone in both organir and aqueous phase is of necessity constant, providing the solubility of acetylacetone in water does not vary appreciably n i t h pH. Hence, equilibrium extraction curves are more easily obtained. Other advantages of acetylacetone as a solvent have been discussed ( I O ) . I n this paper the extraction behavior of aluminum, gallium, and indium is considered, along with the solubility of acetylacetone in water as a function of pH which must be knoxn to evaluate the term, (HR),. EXPERIMENTAL
Purification of Acetylacetone. The purification of the acetylacetone has been discussed (IO). Solubility of Acetylacetone in Water as a Function of pH. Since existing methods ( 8 , 9 ) of determining acetylacetone in Rater were found unsatisfactory when excess acid was present, a method of analysis was developed in which the amount of chelate formed on adding an excess of metal ion to the buffered sample would indicate the amount of acetylacetone originally present.
An excess of ferric ion was added to a I-ml. sample of dilute acid saturated with acetylacetone. Sufficient strong sulfuric acid-sodium sulfate buffer was added to bring the p H of all
