ANALYTICA L EDITION
402
Vol. 6, No. 6
shellac, and further study may reveal varying amounts as due t o season, climatic conditions, and the like as they affect
fraction and would account for the poor bleaching qualities. These experiments illustrate the need for a suitable bleaching the insect. test (1) to be used in the importation of shellac for the express purpose of manufacturing white grades. No complete TABLE 11. SOLUBILITY OF SHELLAC IN BOILING ACETONE standardization of white grades can be set up until extensive INSOLINIOLstudies of the bleaching quality of raw lac have been made. UBLE UBLE SAMPLE RESIN SAMPLE RESIN A more complete study is in progress at the bureau, where % % these methods will be used in conjunction with a wide variety T.N. 6.22 U.S.S.A.V.S.O. 3.46 T.N. 7 . 7 0 Lemon 1.99 of solvents. Preliminary results show that the varied fracU.S.S.A.T.N. (poor bleaching) 7.44 Lemon 0.28 tions have different chemical constants (Kennxahlen). U.S.S.A.T.N. 4 . 3 2 Lemon 0.14
TABLE 111. COMPARISON OF SOLUBILITY IN BOILINGACETONE WITH BLBACHING QUALITY OF LAC SAMPLE
D
INSOLUBLE EXCESS BLEACH Rasrnr LIQUORRBQUIRND % % 16.76 11.62
A B E
10.00 9.32 6.91 4.82
C
F
64 31 22 19 9 0
Of even greater interest are the results shown in Table 111, where the amount of resin insoluble in hot acetone has been determined for the different types of lac which were visually separated from a shellac showing poor bleaching qualities, These samples were described in a previous paper (4). A direct relation appears to exist between the amount of insoluble resin and the amount Of bleach liquor required to give a definite color of cream in the bleached product. The un-
extracted resin from hot acetonewas in mostcases much darker than that which dissolved in this solvent. Any “ k e d ” coloring matter would probably be found in this insoluble
LITERATURE CITED (1) Barry, T. H., “Natural Varnish Resins,” E. Benn, Ltd., London, 1932. (2) Burnett, H., Thesis, The Polytechnic Institute of Brooklyn (June, 1930). (3) Dieterichs, K., and Stock, E., “Analyse der Harse, Balsame, und Gummiharze,” Julius Springer, Berlin, 1930. (4) Gardner, W. H., IND.ENG.CHEM.,25, 550 (1933). (5) Gardner, W. H., Whitmore, E. F., and Harris, H. J., Zbid., 25, 696 (1933). (6) U.8. Shellac Importers’ ABSOC.and Am. Bleached Shellac Mfg. Assoc., “Official Methods of Analysis, Revision,” New York, 1930. (7) Wolff, H., “Die natnrlichen Harse,” Wissenschaftliche Verlegsgesellschaft, m. b. H., Stuttgart, 1928. RECEIV~D February 20, 1934. Presented before the Division of Paint and Varnish Chemistry at the 85th Meeting of the American Chemical Society, Washington, D. C., March 26 to 31, 1933. I n part based upon a theais by Harry J. Harris, submitted in partial fulfilment for the degree of maater of science in chemistry, The Polytechnic Institute of Brooklyn, June, 1933 Contribution N ~ 22, . Department of Chemistry, and No. 10, Shellac Research Bureau, U. s. Shellac Importers’ Association.
Determination of Sulfide Sulfur in Alkaline Solutions Containing Other Sulfur Compounds E. L. BALDESCHWIELER, Standard Oil Development Co., Elizabeth, N. J.
T
HE various methods for determining sulfide sulfur found in the literature are generally based on either alkalimetric and iodometric titrations alone or in conjunction with precipitation by the addition of salts of
some of the heavy metals. Sutton (16)in analyzing a mixture of alkali sulfides, sulfites, thiosulfates, and sulfates, recommends a method worked out by Richardson and Akroyd (11) in which the sulfide sulfur is determined by titration with ammoniacal zinc solution. Another method mentioned by Sutton, that of Sander ( l a ) ,involves iodine titration to obtain the total thiosulfate and sulfide sulfur while the thiosulfate sulfur is determined by precipitation with mercuric chloride and titrating the liberated acidity. Griffin (6) determines sodium sulfide in black liquor by titration with an ammoniacal zinc solution using nickel ammonium sulfate as outside indicator. For white liquor Griffin recommends Moe’s method (8) whereby the total sulfide, sulfite, and thiosulfate radicals are determined by iodine titration and the sulfide is precipitated with alkaline zinc chloride. Scott (14) determines sodium sulfide by first treating the solution with barium chloride to remove any sulfite resent, then distilling the sample with ammonium chloride w%ch will remove the sulfide sulfur in the form of ammonium sulfide. The distillate is collected in ammoniacal cadmium chloride solution, forming cadmium sulfide which is titrated with iodine according to the procedure given for determining sulfur in steel by the evolution method. Other methods for determining sodium sulfide are those of Budnikoff and Krause (3) who oxidize the sodium sulfide with ferric sulfate and titrate the resulting ferrous sulfate with potassium permanganate. Sulfites and thiosulfates will also be oxidized in this procedure. Muller (9) determines the sulfur present as sulfide and thiosulfate by titration with iodine solutions, and the thiosulfate alone by acidifying with acetic acid
and driving off the hydrogen sulfide by evacuating with a pump. Jarvinen (7) precipitates the sulfide with copper sulfate and determines the excess copper colorimetrically. Pauli (10) determines the sulfide sulfur in a mixture by titrating the solution with standard iodine solution before and after precipitation of the sulfide with mercuric chloride. Schulek (13) boils a solution containing sulfide, polysulfide, and thiosulfate sulfur with boric acid and potassium cyanide whereby the polysulfide sulfur is converted to potassium sulfocyanide, the sulfide sulfur to hydrogen sulfide, and the thiosulfate remains unchanged. Cantoni (4)precipitates the sulfide with an excess of potassium arsenite, acidifies the solution, and titrates back the unused potassium arsenite with iodine. Atkin and Hugonin (1) treat with lime water and titrate the sulfide sulfur with zinc sulfate, determining the end point by spot tests on lead acetate paper. Horst (6) adds the sodium sulfide from a buret to Fehling’s copper sulfate solution taken as standard, using sodium sulfide as outside indicator. Billheimer and Reid (2) analyze a mixture of sodium hydroxide, sodium sulfide, and mercaptide by determining the total alkalinity to methyl orange, the total. reactive sulfur with iodine, and the hydrogen sulfide formed in the decomposition of the mercaptide (by repeatedly boiling the alkaline solution) with a second iodine titration.
PROCEDURE In the analysis of spent soda solutions used in petroleum refining, it is often necessary to obtain the sodium sulfide content with a fair degree of accuracy, particularly if the treating reactions are being studied. These solutions are strongly alkaline and contain besides sulfates, sulfites, thiosulfates, and sulfides, mercaptides of the type MSR and sulfonates of the type MSO,R, where M is the alkali
November 15,1934
I N D UST R I A L AN D E N G I N E E R I N G C H E M IST R Y
metal used and R an organic radical. The presence of these compounds, particularly the mercaptides, the properties of which are similar to those of the sulfides, greatly complicates the analysis. Thus cadmium, mercury, zinc, and copper salts all form insoluble mercaptides with either or both lower and higher mercaptans. Again, the lower mercaptans are volatile and will be driven off as the ammonium salt together with ammonium sulfide when ammonium chloride is added to the solution. Mercaptides will also react with iodine with the added complication that iodine may combine with the organic group R if the latter happens to be an unsaturated aliphatic. A procedure has been developed whereby the sulfide sulfur can be determined in an alkaline solution containing sulfates, sulfites, thiosulfates, mercaptides, and sulfonates. The method, while admittedly tedious, has been found to be accurate and has the great advantage that it gives the sulfide content directly. The procedure depends upon the fact that the lead salts of the above sulfo-acids are all soluble in ammonium acetate, the only exception being lead sulfide. This property of lead sulfide was observed after various experiments with the mercaptides of the heavy metals. The determination is accordingly carried out as follows: To a known amount of the unknown, such as spent soda solution, add 10 per cent lead acetate solution until no further precipitation takes place. The precipitate will contain, in addition to the lead salts of all the sulfo-acids present, lead hydroxide and the lead salts such as the carbonate and the silicate of some of the impurities generally present in commercial caustic soda. Filter through filter paper and wash with hot water. Unfold the filter over the beaker in which the precipitation had been made and wash the precipitate into the beaker with a stream of hot water. Make the volume up t o about 100 cc. with hot water, add 15 grams of solid ammonium acetate and acetic acid until distinctly acid to litmus paper. Boil gently for 15 minutes and filter, washing with hot water. The insoluble portion will consist of lead sulfide together with some lead silicate if sodium silicate was present. For refinery control work filter the ammonium acetate insoluble through a Gooch crucible dried at 105’ C. and weigh as lead sulfide, the presence of silica giving somewhat higher results. For more accurate work, the sulfur content of the lead sulfide precipitate must be determined. For this purpose, filter the ammonium acetate solution through filter paper, wash with hot water, and transfer the precipitate back into the beaker. Oxidize the lead sulfide to lead sulfate by adding bromine or potassium chlorate and nitric acid. Boil off the excess halogen, dilute with hot water, and cautiously neutralize with small portions of solid sodium carbonate, finally adding about 5 grams in excess. Boil for about 10 minutes and let settle. Decant through a filter into a beaker. Add about 25 cc. of 10 per cent sodium carbonate to the insoluble in the first beaker, boil again and filter, washing thoroughly with hot water. The insoluble on the filter will consist of lead carbonate, while the sulfate radical will pass into the filtrate as sodium sulfate. Carefully neutralize with hydrochloric acid, adding a fair excess, evaporate to dryness t o decompose nitrates and dehydrate the silica, take up with water and a little hydrochloric acid, and determine the sulfur as usual by recipitation with barium chloride and weighing as barium sulyatate. The method has been tested against known mixtures made up as follows: A standard solution of sodium sulfide was prepared by dissolving about 5 grams of the chemically pure salt in water and making up to 1 liter. The total sulfur was determined in aliquot samples by oxidation with bromine and precipitation with barium chloride. The sulfide sulfur was then determined by the method just described. According to the writer’s experience, low results are sometimes obtained when the total sulfur in a fairly concentrated sodium sulfide solution is determined by adding an excess of bromine to the sglution. This is probably due t o loss of hydrogen sulfide. For this reason, it is preferable to carry out the pxidation by first adding 4 to 5 grams of sodium hydroxide, then lust sufficient
403
bromine to still keep the solution alkaline. In this case the oxidizing agent is sodium hypobromite according t o the equation: 4NaBrO
+ NazS
= NaaSOi
+ 4NaBr
The alkaline solution need only be warmed for one hour for completing the reaction before acidifying with hydrochloric acid and proceeding with the barium chloride precipitation. Several solutions were made up by using known volumes of the above standard sodium sulfide solutions and adding about equal amounts of the various types of sulfo-salts. The resulting mixtures were then analyzed for their sulfide content, the results being given in Table I. TABLEI. SODIUM SULFIDECONTENT OF SYNTHETIC MIXTUREB SOLUTION I X P U R I T IA ~ SD D ~ D
SODIUMSULFIDECONTENT Added Found Average Error
c./1.
.... .... .... ....
UJl.
Q./i
%
.... .... ....
3.5045“ 3.5028 None 3.5012’ Check determination 3.4878b 3.4827 None Check determination 3.4777 b .... Sodium ethyl mercaptide0 1.7414 1.7322d .... -0.5 Check determination 1,7414 1.7556d +0.8 1.7414 1.7472 Check determination +0.3 1.7414 1.7459 Check determination -I-0.3 Sodium ethyl mercaptide and sodium thiosulfate 1.1648 1.1671 +0.2 D Sodium benzene sulfonate, sodium butyl mercaptide, sodium thiosulfate. and sodium sulfite’ 0.3483 0.3461 -0.2 0.3483 0.3504 C Check determination +0.2 Sodium sulfite, sodium benE eene sulfonate and sodium sulfate sodium ethyl merea tidi and sodium thio3.4827 3.4945 50.3 &ate .... -0.15 3.4827 3.4778 E Check determination a Calculated from total sulfur obtained by oxidation with eodium hypobromite. b Calculated from a aulfide sulfur determinatiqn. C Sodium mercaptide solutions prepared by addin 1 or 2 CC. of the pure mercaptans to 100 cc. of a 10 per cent solution of socfhm hydroxide. d Data obtained on 10-cc. samples; other determinations made on 50-eo. aliquot portions. A A A A B B B B C
....
.... .... .... .... . . . a
.... .... e . . .
The results obtained on solution A by the bromine-sodium hypobromite method are higher than those obtained on the same solution by the sulfide sulfur method. This is explained by the fact that sodium sulfide always contains other sulfosalts, particularly sodium thiosulfate, which are oxidized to sodium sulfate by the hypobromite solution. Hence the sodium hypobromite method gives the total sulfur in the sodium sulfide solution which should be equal to (if the salt is pure) or higher than the sulfide sulfur obtained by the method just described. The actual sodium sulfide content of solution A, 3.4827 grams per liter, has therefore been used as a basis for caIculating the theoretical amount of sodium sulfide present in the other solutions. The results show that the presence of other sulfur compounds do not interfere with sodium sulfide determinations made according to the procedure just described and t h a t a n accuracy of 1 0 . 3 per cent can be claimed for themethod.
LITERATURN CITED (1) Atkin and Hugonin, Boll. uficia2e in& peEZi. mat. concianti, 3, 298-9 (1925). (2) Billheimer, E. C., and Reid, E. E., J. Am. Chem. SOC.,52, 4338 (1930). (3) Budnikoff,P., and Krause, K., Z. anorg. Chem., 122, 171 (1922). (4) Cantoni, O.,Giorn chim. ind. applicala, 8, 316-8 (1926). (5) Griffin, R., “Technical Methods of Analysis,” p. 393, 2nd ed., 1927. (6) Horst, F. W., Ledertech. Rundschau, 21, 180-1 (1929). (7) Jarvinen, K. K., Z. anal. Chem., 63, 369-92 (1934). (8) Moe, Carl, Paper, 17, 30 (1916). (9) Muller, J. A., Bull. soc. chim., 19, 8-9 (1916). (10) Pauli, K., Z. anal. Chem., 68, 268-99 (1926). (11) Richardson and Akroyd, J . SOC.Chem. Ind., 15, 171 (1896). (12) Sander, A., Chem.-Ztg., 39, 945 (1915). (13) Schulek, E., 2. anal. Chem., 65, 352-8 (1925). (14) Scott, W. W., “Standard Methods of Chemical Analysis,” D. 1762. 4th ed.. 1925. (15) Suiton, F., “Systematic Handbook for Volumetria Analysis,’’ p. 329, 11th ed., 1924. RE~CE~IVE~D July 7, 1934.
