Determination of Microgram Quantities of Sulfide Sulfur BOLESLAW L. DUNICZ AND TERKEL ROSENQVIST Institute f o r the Study of Metals, University of Chicago, Chicago, Ill.
A reliable method is needed for the determination of microgram quantities of hydrogen sulfide and sulfur in metal sulfides. A volumetric method based on the oxidation of hydrogen sulfide into sulfate by means of alkali hypochlorite has been investigated. Systematic studies show that this reaction is influenced by the pH and the temperature of the hypochlorite solution. Quantitative oxidation of sulfide into sulfate occurs, at room temperature, in hypo-
F
OR the determination of small amounts of hydrogen sulfide, or sulfide sulfur t h a t can be converted into hydrogen sulfide by acid evolution, various methods may be used. Probably the most common is based on the oxidation of hydrogen sulfide with iodine t o form elementary sulfur, either by titration with a standard iodine solution ( 3 ) or by addition of excess iodine and back-titration with thiosulfate. As this reaction utilizes the exchange of 2 electrons only per atom of sulfur, the amount of iodine consumed is relatively small, and very dilute iodine and thiosulfate solutions are needed when microgram quantities of sulfur are t o be determined.
n
//I/
T Figure 1. Apparatus for Evolution and Absorption of Hydrogen Sulfide A bettcr method would be one where the hydrogen sulfide is oxidized t o sulfate: corresponding t o the exchange of 8 electrons per atom of sulfur. The most suitable oxidizing agent in this respect seems t o be hypochlorite solution. This type of analysis has been mentioned (1, .9), but without description of the exact procedure t o be followed. Recently Pepkowitz ( 5 )presented an account of his study on microdetermination of sulfur by oxidation t o sulfate with calcium hypochlorite solution. Pepkowitz did not obtain a theoretical yield of oxidation, but suggested the use of an empirical factor which can be obtained by standardization with known amounts of alkali sulfide. Attempts were made in the authors' laboratory to employ this method for the analysis of hydrogen sulfide and iron sulfide using solutions of commercial calcium hypochlorite (bleaching powder). The oxidation was, in many cases, closer t o 50 than t o 10056, and the empirical factor seemed t o vary with the strength of the hypochlorite solution as well as with the temperature and the amount of sulfur in the sample. Therefore, t o understand the reasons for this nonstoichiometric yield and t o learn how this method
chlorite solutions of pH between 8 and 11 and between 14 and 15. In the intermediate pH range the oxidation is incomplete, and colloidal sulfur is formed. The lowest yield, at 0' C. and pH 12, corresponds to about 50Yo oxidation to sulfate. As most favorable for analytical work a hypochlorite solution of pH 14.4 is suggested. The method can be used for determining sulfur whenever it is present as hydrogen sulfide, or can be converted into it. could be best utilized, a systematic study of the effect of pH and temperature was made. EXPERlM ENTA L
Preliminary tests had shown that potassium hypochlorite behaved in the same manner as calcium hypochlorite and was equally stable nith respect t o time. As commercial bleaching powder was suspected t o be of variable quality, the systematic work was based on the use of potassium hypochlorite solution.
In L small gas generator, similar to the one shown in Figure 1, about 3.2 grams of potassium permanganate were treated with 6 M hydrochloric acid and gently heated. The evolving chlorine was, by a flow of nitrogen, swept into 1 liter of 0.15 M potassium hydroxide solution, in which i t was absorbed to give a solution of equal (0.5) molarity of potassium hydroxide, potassium hypochlorite, and potassium chloride. This served as a stock solution and i t was stable when kept in a dark place a t room temperature. After 2 months its strength had decreased about 4.5%. For actual use the stock solution was diluted five times to give a solution which was 0.01 iM with respect to potassium hypochlorite or 0.02 N as to oxidizing power. T o obtain hypochlorite solutions of various pH values, dilute sulfuric acid or potassium hydroxide pellets were added. The p H of the hypochlorite solutions to which acid had been added was measured potentiometrically with a Fisher Titrimeter using a Beckman glass electrode standardized against a buffer solution of pH 7. Without acid added, the solution showed by this method a p H of about 13. Its theoretical p H could be calculated from the composition (0.01 M potassium hydroxide, 0.01 M potassium hypochlorite, and 0.01 M potassium chloride) to correspond to 12.03. As the glass electrode used was recommended only for work a t p H below 9, values at higher pH may be expected to be in error. Therefore, the pH of the original solution was taken as equal to the calculated value, 12.03, and the measured values were corrected downward by an amount which varied linearly from about one unit for the original solution to zero for p H 7. For the solutions to which potassium hydroxide had been added, the pH was calculated aseuming complete dissociation, and disregarding any change in activity coefficients as the concentration of potassium hydroxide increased. These values will, therefore, be rather formal, but give an approximate picture of the true pH. The small variation of pH with temperature was neglected throughout the entire composition range. During the absorption and oxidation of hydrogen sulfide a small amount of sulfuric acid is formed which will change the pH of the solution somewhat. This effect can be calculated and is negligible for most of the pH range. It becomes significant only below pH 9. A preparation of acid-soluble sulfide was made by mixing and grinding 1 part of iron sulfide of known sulfur content with about 350 parts of zinc dust (c.P.,Eimer and iimend, New York, 9. Y.). T o obtain uniformity the mixture was agitated and rotated in a jar mixer for several n.eeks. The mixture thus obtained contained about 0.1% sulfur. For each test an accurately 17-eighed amount (0.63 to 0.64 pram) of this sulfide mixture was used, corresponding to 0.63 to 0.64 mg. of sulfur. I n the apparatus shoxqn in Figure 1 the mixture was treated with 6 M hydrochloric acid. During this operation the reaction flask ' i ~ as immersed in cold water, the tempera-
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V O L U M E 24, NO. 2, F E B R U A R Y 1 9 5 2
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ture of which was gradually raised to boiling. When the sample had been completely dissolved, a vigorous stream of nitrogen was passed through the apparatus to assure complete transfer of hydrogen sulfide into the Volhard-type gas absorption flask. This contained 25 ml. of the 0.02 N hypochlorite solution to which acid or base had been added to give the proper pH. The absorption flask was immersed in a constant temperature bath and was, for runs a t 60" C., equipped with a reflux condenser. A full hour was allowed for the solution of the sample and absorption of the gas. To reduce the possibilities for photochemical reactions, direct sunlight was avoided and some yellow dye was dissolved in the constant temperature bath. Simultaneously, in a separate Erlenmeyer flask, a blank consisting of an equal amount of hypochlorite solution of equal basicity was kept in the constant temperature bath for the same length of time. When the absorption was completed, 1 to 2 grams of potassium iodide were added to both solutions, and the solutions were acidified with acetic acid. About 5 ml. of 5Oy0 acetic acid were used, except for the more basic solutions where 7 ml. of 50% acid were used for each gram of potassium hydroxide added. The blank as well as the solution in the absorption flask was then titrated with 0.01 N thiosulfate solution using 1 ml of 1% starch solution as indicator. The thiosulfate solution had been standardized against iodine, liberated with acid from a mixture of a weighed amount of potassium iodate and an excess of potassium iodide. The amount of hypochlorite consumed by the oxidation of hydrogen sulfide was then given by the difference between the titer for the blank and for the absorbing solution. The results were calculated as percentage of the theoretical oxidation of the hydrogen sulfide to sulfuric acid In Figure 2 are given the results obtained for four different temperatures and for pH values betxeen 8 and 15. At 0' C. some runs were made in which the pH had been adjusted with hydrochloric acid instead of sulfuric acid. The two sets of values are in good agreement. DISCUSSION
It follows from Figure 2 that approximately quantitative oxidation is obtained in the pH ranges from 8 t o 11and from 14 t o 15. In the intermediate range the oxidation is incomplete. The oxidation is least complete a t a pH of about 12-Le., close t o the pH of the hypochlorite solution with neither acid nor base added. The least oxidation occurs a t the lowest temperature, 0" C., where the minimum corresponds very closely t o 50% oxidationi.e., an average exchange of 4 electrons per atom of sulfur. .4t higher temperatures the oxidation t o sulfate is more complete,
but certain deviations from a smooth curve appear a t all temperatures in this intermediate p H range. For the deviations a t 40' t o 60" C. there are not enough data t o say whether they correspond t o reproducible maxima and minima, or are due t o factors beyond control, such as traces of catalyzing impurities, photochemical reactions, or the rate of hydrogen sulfide introduction. In any case the spread of the points is definitely larger than the titration error. A t 20" C. a narrow but reproducible peak appears at pH 12. At pH around 8, the formtion of undissociated hypochlorous acid becomes pronounced and some chlorine is swept out of t h e solution and is lost. This is the obvious reason for the apparent yield of oxidation higher than 1 0 0 ~ oa t high temperature and low pH. At very high pH, about 15, t h e yield of oxidation seems once more t o decrease. This is probably connected with the form% tion of potassium chlorate. The titer for the blanks with p H below 14.5 remained practically constant during the run, whereas it decreased significantly a t higher pH. The products of oxidation in the intermediate range were determined qualitatively. It was apparent that colloidal sulfur was formed. In some cases, corresponding t o the intermediate "maxima" a t 40" t o 60' C., sulfur could be observed during t h e initial part of the absorption, but it disappeared later on. In addition t o colloidal sulfur, sulfate was formed. This was shown for some solutions of pH around 12: After absorption of hydrogen sulfide, the solution was treated with mercury or zinc dust and acidified with hydrochloric acid t o remove hypochlorite and colloidal sulfur. After filtration, sulfate ion could be identified in the filtrate with barium chloride in the usual way. As other sulfur compounds are readily oxidized by hypochlorite, it seems t h a t Tvhen the oxidation is incomplete, the end product is a mixture of sulfur and sulfate. I t is not the purpose here t o t r y to explain the mechanism of this reaction or the reason for the incomplete oxidation. However, a few points should be mentioned: 1. A number of sulfur compounds are known with degree of oxidation between hydrogen sulfide and sulfuric acid, and it is likely that the oxidation goes through one or several of these compounds. From stereochemical considerations a possible intermediate compound would be sulfoxylic acid, H2S02. The formation of this compound would correspond to the exchange of 4 electrons Der atom of sulfur or 50% 0x1dation to sulfuric acid. As in no case less than 50% oxidation was observed, 100 intermediate oxidation of hydrogen sulfide to elementary sulfur (corre.90 sponding to 25% oxidation to sulfuric acid) is regarded as less probable. 0°C. .eo 2. Several of these intermediate 100 ' 70 compounds are knoivn either to dis90 proportionate and give elementary .BO sulfur or to react with each other to o'c. 8 0 give sulfur. I00
-
1009
100"
1008
90
2H2S02
80 100'
70 80
50
8
9
10
II
12
I3
14
15
PH-
f i g u r e 2.
Percentage of Oxidation z's. pH of Hypochlorite Solution for Four Different Temperatures
Crosses represent solutions i n which pH h a d been adjusted w i t h hydrochloric acid (0' C. only). Dotted lines indicate possible details in curves
H2SOz
ItIzSOa
+ So + HsO
+ HzSOa = 2 H + + Sod-- + So + H20
[According to the standard oxidationreduction potentials listed by Latimer ( 4 ) these reactions may proceed as indicated, even in a basic solution. Whether they actually occur in the present case remains to be proved.] 3. Hydrogen sulfide, and several of the intermediate sulfur compounds as well as hypochlorous acid, are weak acids with a t least one of their dissociation steps taking place in the pH range in question. I t is therefore possible that in one pH range the
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ANALYTICAL CHEMISTRY Table I.
Results for Various Standard Samples Sample Weight
Material
Grams
NBS standard sample of steel 15d (0.054’% Cu) NBS standard sample of steel 8g (0.019% C U I New Jersey Zinc Co. ZnS‘
FeS 5
+ Zn. 1: 350
Sulfur Found
Sulfur Found
Accepted Value
Mg.
%
%
2.0561 2.0668 3.4072
0.650 0.657 1.084
0.0316 0.033-0.034 0.0318 0.0318
2.1192 4.7387 0.00128 0.00260 0.00401 0.7029 0.7622 3.0091b
0.532 1.180 0.422 0.855 1.321 0.748 0.806 3.195
0.0251 0.025-0.026 0.0249 32.90 32.90 32.87 32.95 0.1063 0.106-0.107 0.1057 0.1062
Weighed on microbalance and mixed with 1 to 2 grams of zinc dust.
b 50 ml. of KOCl solution with 7 grams of
KOH.
reaction is primarily between undissociated acids, in another between the ions, and in intermediate ranges partly betn-een acids and partly between ions of varioue types. As stereochemical conditions probably are important for the choice of mechanism, the reaction path may be strongly affected by a slight change in pH. 4. Colloidal sulfur is more rapidly oxidized by hypochlorite in a strongly basic than in a neutral solution. The over-all yield of oxidation as shown in Figure 2 may be a result of a combination of the various factors mentioned above. A tentative explanation for the incomplete oxidation would then be initial formation of sulfoxylic acid or its ion, Tvhich may react further according t o the scheme outlined under point 2. The peak a t 20’ C. for pH 12 still represents an open question.
solution has t o be acidified before titration, a buffer may require unreasonable amounts of acid for acidification. The best results were obtained in the very basic range, and as standard basicity an addition of 3.5 grams (26pellets) of potassium hydroxide per 25 ml. of solution was used. This corresponds t o a 2.5 N potassium hydroxide solution, and a pH of about 14.4. Various sulfides of known sulfur content and two Bureau of Standards steel samples were tested. The results of these analyses are listed in Table I. It is apparent from the data given in Table I t h a t under the right conditions this is a sensitive and accurate method. Its limitation for steel analysis lies in the fact that sulfur in steel may be present as sulfides, which are not soluble in acids. For example, steel 15d, for which slightly too lox values were found, contained as much as 0.054% of copper. Steel 8g, on the other hand, contained only 0.019% copper and the analysis showed a perfect agreement with Bureau of Standards values. SUMMARY
The oyidation of hydrogen sulfide with potassium hypochlorite was studied as a function of temperature and pH of the hypochlorite solution. Nearly quantitative oxidation t o sulfate was obtained in the pH ranges from 8 t o 11 and from 14 t o 15. At intermediate pH values the oxidation was incomplete, and a t 0’ C. and pH around 12 the least oxidation was found, corresponding t o the average exchange of 4 electronsper atom of sulfur. The products of oxidation were, in this range, a mixture of sulfate and rolloidal sulfur. The most favorable solution for the analysis of microgram quantities of hydrogen sulfide was found t o be hypochlorite in a 2.5 X potassium hydroxide solution, corresponding to a pH of about 14.4. ACKNOWLEDGMENT
ANALYTlCAL APPLICATIONS
The incomplete Oxidation and poor reproducibility which were experienced when a solution of commercial bleaching powder was used as oxidizing agent may now readily be understood. Such a solution has a pH near 12. For analytical work the ranges between pH 9 and 10 and between 14 and 15 seem t o give the best possibilities. Attempts to work out a useful method in the lower basicity range were not very successful. It is difficult t o adjust the pH of the solution t o be just right. Some sulfuric acid is formed by the reaction and some hydrochloric acid may come over from the evolution flask and cause a large change in the pH. It might be possible t o add an appropriate buffer t o maintain correct basicity, but since the
The authors are indebted t o S . H. Xachtrieb for valuable suggestions and criticism. LITERATURE CITED
(1) Kitchener, J. A., Bockris, J. O’hl., and Liberman, A,, Discussions Faradau SOC.,4, 57 (1948). (2) Kolthoff, I. M., and Sandell, E. B., “Textbook of Quantitative Inorganic Analysis,” p. 538, New York, Macmillan Co., 1947. (3) Ibid., p. 634. (4) Latimer. W. M.. “Oxidation States of the Elements and Their Potentials in Aqueous Solutions,” pp. 64-74, New York, Pren-
tice-Hall, 1938. ( 5 ) Pepkowitz, L.. ANAL.CHEM.,20, 968 (1948). RECENEDJ u n e 15. 1951.
Spectrographic Determination of Beryllium in Urine and Air R. G. SMITH, Bureau of Industrial Hygiene, Detroit Department of Health, A. J. BOYLE, Wayne University,
W. G . FREDRICK, Bureau of Industrial Hygiene, Detroit Department of Health,
AND
BENNIE ZAK,
Wayne University, Detroit, Mich.
T
HE recognition of beryllium as an extremely toxic metal when inhaled in minute quantities has stimulated analytical research directed a t the accurate estimation of submicro amounts of this metal in air and urine. Various spectrographic procedures have been described in recent years (1,3, 4 , 67, all of which employ the direct current arc. A fluorometric method using morin, which achieves the desired sensitivity by using large quantities of urine, has also been described ( 5 ) .
The present study was undertaken in an effort t o develop a somewhat more rapid method, of acceptable accuracy and sufficient sensitivity t o permit use of the relatively small samples of urine and air that might be collected routinely in beryllium-using industries. It was further desired to utilize an alternating current spark technique for solution analysis, on which research has been in progress for the past 2 years. This technique, which utilizes the -4pplied Research Laboratories rotating electrode assembly,
