mixture twice with 15-ml. portions of 1-butanol, and follow with three extractions with 15-ml. portions cf carbon tetrachloride, discarding the organic solvent layer. The last portion of carbon tetrachloride need not be separated. Transfer the aqueous phase to the original Erlenmeyer flask, dilute with 55 ml. of water (use portions of the water to complete the transfer), and titrate with 0.05X potassium dichromate solution, using diphenylaminesulfonic acid indicator. Carry out a blank determination in the same manner, omitting the sample.
yo sulfoxide = (blank - titer) X A; dichromate X M . R . sulfoxide X 100 2000 X grams of sample RESULTS
The results of the determination of 2-hydroxydiethyl sulfoxide and some other sulfoxides ( 3 ) by Barnard’s method for saturated sulfoxides, and by the modification of this method, are shown in Table I. LITERATURE CITED
(1) Barnard, D., Hargrave, K. R., Anal. Chim.Acta 5 , 476 (1951).
(2) Ibid., p. 536. (3) Groves, Kermit, Legault, R. R., unpublished manuscript. (4) Treadwell. F. P.. Hall. W.T.. “Analytical Chemistry,” kol. 11, p. 699, Wley, Yew York, 1947. RECEIVED for review February 5, 1957. Accepted May 27, 1957. Portion of a paper delivered before Analytical Chemistry Section, ACS, Seattle, Wash., June 11, 1956. Investigation supported in part by funds provided for biological and medical research by the State of Washington Initiative Measure No. 171. Scientific Paper No. 1575, Washington Agricultural Experiment Stations. Work conducted under Project No. 1229.
Determination of Sulfur in Titanium MAURICE CODELL, GEORGE NORWITZ, and CHARLES CLEMENCY Pitman-Dunn laboratories, Frankford Arsenal, Philadelphia, Pa. Sulfur in titanium and titanium alloys can b e determined by dissolving the sample in a mixture of hydrochloric and hydrofluoric acids. The evolved hydrogen sulfide i s absorbed in ammoniacal cadmium chloride, which is then acidified and titrated with standard potassium iodate solution.
investigations have recently shown that sulfur may become an important alloying element for titanium because it produces grain refinement and increases the tensile and yield strength of titanium (4). It is not known in what form sulfur is present in titanium, although the titanium-sulfur system has been studied (f2,16). I t is certain that the sulfur is not present as elemental sulfur. The following titanium sulfides are known to exist: Ti& Tis. Ti&, Ti&, and TIS, (16). ETALLURGICAL
POSSIBLE METHODS
The first approach to the problem was to attempt to develop a barium sulfate precipitation method, because such a method is usually considered to be the referee method. Experiments were conducted with a 2.5-gram sample of titanium. However, it was not possible to determine sulfur in titanium gravimetrically by barium sulfate precipitation for the following reasons. Titanium inhibited the precipitation of barium sulfate. No precipitate a t all could be obtained if less than 0.1% sulfur were present. The authors believe that the solvent action of the titanium is probably due to the formation of the complex titanium sulfate ion, [Ti(S04)3]-- (3.4). Titanium ap1496
ANALYTICAL CHEMISTRY
parently resembles chromate, which interferes by forming a complex chromium sulfate ion (22, 37). The experiments on the precipitation of the barium sulfate were carried out using the minimum amount of hydrochloric acid (20 ml. of acid per 200-ml. volume) that could hold 2.5 grams of titanium in solution. The solutions were allowed to stand overnight a t room temperature. They could not be heated without causing hydrolysis of titanium. It is probable that the amount of hydrochloric acid used also had a solvent action on the barium sulfate. Recoveries were no better with solutions containing a smaller amount of acid and fluoride, tartrate, or citrate to complex the titanium. The addition of hydrogen peroxide seemed to help the precipitation of barium sulfate slightly. There were no satisfactory separations of the titanium from the sulfur that would permit the subsequent application of the barium sulfate method. The Meineke method by which steels (.2?5>29) and copper-base alloys (11) are dissolved in a 5% hydrochloric acid solution containing cupric ammonium chloride and the precipitated copper sulfide filtered off, is not applicable to titanium because the cupric ion in dilute hydrochloric acid exerts a passivating action on titanium and prevents its solution (7). The separation of 2.5 grams of titanium from sulfate by the use of cupferron did not seem feasible. Precipitation of titanium by hydrolysis -for instance, by fuming with perchloric acid-has, according to experience in this laboratory, proved valueless when applied to the separation of large amounts of titanium from any element. An ammonium hydroxide or
sodium hydroxide precipitation might be satisfactory for the separation of only small amounts of titanium. Ullrich (36) used an ammoniacal separation in the determination of sulfur in titaniferous coals containing about 4% titanium. The best separation of titanium from sulfate is no doubt the ion exchange method (28, $1). However, it too did not seem too feasible for the separation of 2.5 grams of titanium. The titanium metal could not be dissolved in a satisfactory oxidizing medium for the gravimetric determination of sulfur. Konoxidizing solvents could not, of course, be used because sulfur would be lost as hydrogen sulfide. The following oxidizing media n 111 not dissolve titanium: nitric acid, bromine, a mixture of nitric acid and bromine, a mixture of nitric and hydrochloric acids, and a mixture of hydrochloric acid and hydrogen peroxide. Fuming nitric acid attacks titanium but the reaction may be explosive. The only oxidizing medium that could be used was 3 mixture of hydrofluoric and nitric acids, or a mixture of hydrofluoric acid and hydrogen peroxide. Both the hydrofluoric acid and nitric acid 1%ould be objectionable because the fluoride and nitrate could contaminate the barium sulfate precipitate. Decomposition of a 2.5-gram sample by fusion with sodium peroxide or mixture of sodium carbonate and potassium nitrate did not seem practical. Because the barium sulfate precipitation method showed no promise, efforts were turned toward other procedures. Luke (23) has proposed a method for the determination of sulfur in many alloys whereby the sulfur is converted to sulfate, which is then reduced t o
hydrogen sulfide with hydriodic and hypophosphorus acids. The method was not applicable to titanium because it could not be brought into solution with the hydrochloric-nitric acid misture used by Luke. Sulfur has been determined in copper (3) and in steels and ferroalloys (6, 29) by heating in hydrogen and determining the hydrogen sulfide produced. Another method that has been used for steels and ferroalloys (25, 29) is to heat 0.1 to 1 gram of the finely divided sample in a tube through which is passed hydrogen that has been saturated with hydrogen chloride gas. N o experiments along these lines were conducted with t'itanium. d commonly used method for the determination of sulfur in metals is the ignition method, whereby the sample. is burned in oxygen and the sulfur converted to sulfur dioxide, rvhich is determined by titration with iodate. The method has been applied to t'he determination of sulfur in steels ( 2 , I S , 29), copper-base alloys ( 2 , I S ) , nickel-base alloys (f4),and ores ( I O , SI). -2 method somewhat similar in principle is the procedure for sulfur in copper, whereby the sample is heated wit'h 20 cc, of oxygen a t l l O O o C. for 30 minutes, and the sulfur dioxide determined by the mass spectrometer (17). Experiments were conducted on the application of the ignition method to titanium-sulfur alloys prepared by melting weighed amounts of ferrous sulfide rvit'h tit,anium sponge metal. Experi-
nients in this laboratory showed that no sulfur was lost in preparing these alloys; thus, they constitute a satisfactory synthetic standard. A similar technique in preparing synthetic standards has been used for titanium-carbon (5) and titanium-oxygen samples ($6). The recovery of sulfur in applying the ignition method to the titanium-sulfur alloys mas poor and the results were erratic. S o improvement in recovery was attained by the use of lead, copper, or tin as fluxes. The failure to obtain complete recovery of the sulfur on burning the titanium alone was not surprising because titanium often burns superficially, a coating of titanium dioxide protecting the inner part of the metal. The poor recovery with lead as a flux was unexpected because the flux has been used successfully in the determination of hydrogen (8) and carbon (9) in titanium by the ignition method. The failure to obtain complete recoveries of sulfur when the fluxes were used could not be due to incomplete ignition of the titanium, as the samples burned completely and homogeneous melts 11ere obtained. During the ignition with fluxes the temperature rose to about 1600° C., as judged by a n optical pyrometer. The heat was so intense i t was necessary to place the sample in Vycor tubing on clay supports in the reaction tube to prevent cracking. The high teniperature of ignition is desirable because a high temperature favors the formation of sulfur dioxide rather than sulfur trioxide (13, 2f ). Porous materials
A '
Figure 1,
Apparatus for determination of sulfur
A . Helium tank B . Tpgon tubing C. Reaction flask (500 ml ) D . Ground-glass joint
E. F. G. H. I. J.
Dropping funnel Stopcock Hydrochloric acid condensation flask (200 ml. I Rubber stopper Test tube (50 ml.) Bunsen burner
such as glass wool absorb sulfur dioxide and cause low results ( I S , 26, 3f); therefore, they were not used. The cause of the low results probably involved a reaction between the sulfur dioxide and the titanium dioxide. A search of the literature showed that such a reaction does, in fact, occur. Neumann (26), in studying catalytic processes, found that titanium dioxide a t red heat partially catalyzes the oxidation of sulfur dioxide to sulfur trioxide. The sulfur trioxide is then absorbed by the titanium dioxide to form titanium sulfate (26). In view of this, the ignition method cannot be recommended for the determination of sulfur in titanium. Another frequently used method for the determination of sulfur in metals is the hydrogen sulfide evolution method, in which the sample is dissolved in a nonoxidizing acid and the hydrogen sulfide evolved is determined. The method has been applied to the determination of sulfur in steels ( 2 , 26, 29), copper-base alloys (W), antimony ( 2 ) , cadmium (2f),aluminum ( I ) , manganese (fg), cobalt (19), tin (33), and chromium (20). The hydrogen sulfide evolution method gave practically 100% recovery with the synthetic titaniumsulfur alloys previously mentioned.
HYDROGEN SULFIDE
EVOLUTION METHOD
Apparatus. Figure 1 shows t h e apparatus for applying this method t o titanium and titanium alloys. It consists of a reaction flask with a ground-glass head which contains a dropping funnel a n d a n inlet t u b e for a carrier gas. Following t h e reaction flask is a 200-ml. Erlenmeyer flask containing a little water t o absorb a n y hydrogen chloride gas t h a t might distill over. Two absorption tubes containing ammoniacal cadmium chloride solution follon-. The second absorption tube is attached as a precaution. No precipitate of cadmium sulfide was ever obtained in the second tube in running samples containing as much as 10 mg. of sulfur. The authors prefer helium (Grade A) as a carrier gas because, unlike hydrogen, it does not contain sulfur compounds and, therefore, does not need to be purified. Also, there is no danger in using helium near a n open flame. Reagents. Potassium Iodate Solution, 0.03s. Dissolve 1.0701 grams of potassium iodate in about 800 ml. of water; add 1 gram of sodium hydroxide and 10 grams of potassium iodide. Dilute t o 1 liter in a volumetric flask. Ammoniacal Cadmium Chloride Solution. Dissolve 20 grams of cadmium chloride dihydrate in 1750 ml. of water and add 500 ml. of ammonium hydroxide. Starch Solution. hIake a paste of 5 grams of starch and 15 ml. of water and pour it into 500 nil. of boiling water. VOL. 29, NO. 10, OCTOBER 1957
1497
Cool and add 25 grams of potassium iodide. Procedure. Assemble t h e apparat u s as shown in Figure 1. Weigh t h e sample into t h e reaction flask. Use a &gram sample for 0 t o 0.2% sulfur, a 2-gram sample for 0.2 t o 0.5% sulfur, and a 1-gram sample for 0.5 t o 1% sulfur. Add 25 ml. of water t o t h e hydrochloric acid condensation flask, a n d 15 ml. of water and 15 ml. of ammoniacal cadmium chloride solution to the test tubes. Add a mixture of 80 ml. of hydrochloric acid, 5 ml. of hydrofluoric acid, and 10 ml. of water to the dropping funnel while the stopcock is turned off. Adjust the flow rate of helium to a few bubbles per minute. Open the stopcock and allow the acid mixture to flow into the reaction flask. Turn off the stopcock while a little of the acid mixture is still in the dropping funnel. When the reaction has slackened, increase the flow of helium and heat the reaction flask with a Bunsen burner. When the sample has completely dissolved, allow the helium t o flush the system for a few minutes. Wash the contents of the test tubes into a tall-form, 300-ml. beaker with water and dilute to about 200 ml. Add 50 ml. of hydrochloric acid (1 to 1) to the test tubes and then pour the acid into the beaker. Add 5 ml. of starch solution and titrate with standard potassium iodate solution (0.03-V) to a dark blue color. Calculate the per cent sulfur as follolvs: yo sulfur = A B XC 0.016 loo where
A B
=
C
=
ml. of potassium iodate solution
= normality of potassium iodate
solution grams of sample RESULTS
The results obtained for sulfur in three synthetic titanium-sulfur alloys prepared as described above are shown in Table I. The method shows good accuracy and precision. DISCUSSION
For dissolving the sample, a mixture of hydrochloric and hydrofluoric acids was necessary to ensure solution of titanium sulfides. The presence of the titanous ion probably helps in the dissolution of sulfides, as sulfides seem to be more soluble in acid in the presence of reducing agents (SO). A mixture of hydrochloric and fluoboric acids also gave acceptable results but required more time to dissolve the sample. The use of hydrochloric acid alone is not recommended because prolonged boiling is required to dissolve the sample. The use of dilute sulfuric acid or a n acid mixture containing sulfuric acid, as has been recommended for the determination of sulfur
1498
ANALYTICAL CHEMISTRY
in chrome steels (29), is not recommended because the titanous ion partially reduced the sulfate ion to sulfide. All the samples encountered dissolved in the hydrochloric-hydrofluoric acid mixture without any trace of a residue. If a residue were encountered in the analysis of unknown alloys, it might be necessary to filter it off, fuse it with sodium peroxide and precipitate the residual sulfur as barium sulfate. The amount of titanium that could remain in the residue would not interfere with the barium sulfate precipitation. As recommended by the National Bureau of Standards (25) and the American Society for Testing Materials (g), the theoretical factor was used in titrating the cadmium sulfide with potassium iodate. Interferences. Selenium would interfere by forming hydrogen selenide which would cause t h e precipitation of cadmium selenide. Arsenic would interfere by partially distilling as arsine and precipitating as arsenious sulfide. Selenium and arsenic a r e ordinarily not found in titanium or titanium alloys. It is not known whether large amounts of carbon in t h e titanium would interfere by reacting with t h e sulfur t o form dimethyl sulfide, as sometimes happens with high carbon steels (25, 29). It is sometimes stated that copper, tin, and molybdenum can interfere by precipitating as sulfides in the reaction flask, The copper or tin can cause no error in the method described because copper and tin sulfides are completely soluble at the acidity used (24). Molybdenum ordinarily forms a rather insoluble sulfide. However, when titanous ion is present the molybdenum is reduced to molybdenum(V) and does
Table I. Determination of Sulfur by Hydrogen Sulfide Evolution Method
Sulfur, % Present 0.05
0.20
1.00
__
Found 0,049 0,051 0,052 0.053 0.047
Bv. 0.050 0.196 0.204 0.208 0.212 0.196 Av. 0.203 0.95 0.99 0.98 0.98 0.95 0.96 0.96 Av. 0.97
not give a precipitate v-ith hydrogen sulfide (27). For the determination of very small amounts of sulfur in titanium i t would probably be desirable to collect the hydrogen sulfide in ammonium hydroxide, add lead citrate solution, and measure the intensity of the amber color produced by the lead sulfide (23). Because no satisfactory standards containing very small amounts of sulfur mere available, this possible modification of the method was not investigated. ACKNOWLEDGMENT
The authors wish to thank 0. W. Simmons who prepared the titanium-sulfur alloys used in this work. LITERATURE CITED
’
(1) .-lluminum Co. of America, New Kensington, Pa., “Chemical Analvsis of Aluminum.” I D.. 124 (1956). Am. Soc. TestingflMaterials, Philadelphia, Pa., ASTM Methods for Chemical Analysis of Metals,” pp. 92, 129, 207,287,288 (1950). (3) Bassett, W. H., Bed\\-orth, H. A , , Trans. Am. Inst. Mining M e t . Engrs. 73, 784 (1926). (4) Rerger, L. W.,Williams, D. N., Jaffee, R. I., Metal ProgT. 70, No. 3,224 (1956). Cadoff, I., Nielson, J. P., J . Metals 5 , 248 (1953). Clarke, B. L , Wooten, L. A . , Pottenger, C. H., IND. EXG. CHEM., Ax.4~.ED. 7, 242 (1935). Cobb, J. R., Uhlig, H. H., J . Electrochem. SOC.99, 13 (1952). Codell, M., Nornitz, G., ANAL. CHEM.28, 106 (1956). Codell, M., Norn-itz, G., Schneider, E. F., Anal. Chzm. Acta 15, 218 (lF).ifi\ \
-
-
-
-
I
Coller, 11. E., Leininger, R. K., =INAL. CHEM.27, 949 (1955). Cooney, J. O., IAD.EXG. CHEM., ANAL.ED.4,33 (1932). Dean, R. S., Silkes, B., U . 8.Bur. Mines, Inform. Czrc. 7381, p. 10, Washington, D. C., 1946. Division of Analytical Chemistry, ACS, Round-Table Discussion, ANAL.CHERI.24, 202 (1952). (14) Dusing, W., Winckelmann, K., 2. anal. Chem. 113,419 (1938). Handbuch der an(15) “Gmelins organischen Chemie,” System Nummer 41, p. 339, Verlag Chemie, GMBH, TTeinheim/Bergstrasse, 1951. Haughton, J. L., Prince, -4.,“Constitutional Diagrams of Alloys: A Bibliography,]’ p. 218, Institute of Metals, London, 1956. Hickam, W. M., ASAL. CHEM.24, 362 (1952). Holler, A. C., Klinkenberg, R., Ibid., 23, 1696 (1951). Holthaus, C., Arch. Eisenhtkttenw. 10,511 (1937). Horak. O., Z . anal. Chem. 140, 255 (19.52’1 \ - - - - I .
Isbell, H. G., IND.ESG. CHEW, -4NAL. ED. 4,284 (1932). Kolthoff, I. M . , Sandell, E. B., “Textbook of Quantitative Inorganic Analysis,” p. 338, Macmillan, New York, 1943.
(23) Luke, C. L., ANAL. CHEW21, 1369 (1949). (24) Lundell, G. E. F., Hoffman, J. I.,
“Outlines of Methods of Chemical Analysis,” p. 52, Wiley, New York. 1938. (25) Lundei, G. E. F., Hoffman, J. I., Bright, H. A,, “Chemical Analysis of Iron and Steel,” p. 238, Wiley, New York, 1931. (26) . , Neumann. B.. 2. Elektrochem. 35.
42 (1929). ’ (27) Norwitz, G., Codell, hl., ANAL. CHEM.25, 1438 (1953).
(28) Xorwitz, G., Codell, M., Anal. Chim. Acta 11,233 (1954). (29) Pigott, E. C., “Ferrous Analysis,
Modern Practice and Theory,”
p. 451, Chapman & Hall, London, 1953.
(30) Prescott, A. B., Johnson, 0. C.,
“Qualitative Chemical Analysis,”
p. - 503, Van Nostrand, New York, ^^^
lYj3.
w.
(31) Rice-Jones, c . , ~ A L CHEM. . 25, 1383 (1953). (32) Samuelson, O., “Ion Exchangers in Analvtical Chemistrv.” D. 142. Wile:, New York, 1933.
(33) Scott, F. W., Chemist Analyst 31, 52, (1942). (34) Sidgwick, K . W., “Chemical Ele-
ments and Their Compounds,” Vol. I, p. 644, Oxford University Press, London, 1950. (35) Ullrich, F., 2. anal. Chem. 117, 10 (1939). ~ ~ 22, . 297 (36) Walter, D. I., A 4 ~CHEM. (1950). (371 Willard. H. H.. Schneidewind,. R.., Trani. Electrochem. Sac. 46, 333 (1929). \
I
Se miqua ntita tive Esti mati o n of Dithionite THOMAS P. WHALEY’ and JOSEPH A. GYANl Research laboratory, Ethyl Corp. , Detroit, Mich.
b Ammoniacal Naphthol Yellow S solution i s a specific indicator for detecting small quantities of dithionite, the color changing rapidly from yellow to red. Although the color change is nearly permanent for small amounts of dithionite, an apparent reversal of this color change occurs when larger amounts of dithionite are present. The time required for the second color change to occur is roughly proportional to the amount of dithionite and can be used for estimating it semiquantitatively.
T
HE KEED for a rapid method for estimating sodium dithionite in the presence of sodium methoxide motivated a rather exhaustive search for potential analytical procedures. A specific qualitative test (1) for dithionite in the absence of cesium, mercury, potassium, rubidium, and tin ( 2 ) has been described. A small amount of dithionite changes the color of a dilute ammoniacal solution of Naphthol Yellow S (sodium salt of 2,4-dinitroI-naphthol-7-sulfonic acid) from yellow t o red. Inasmuch as other common anions do not effect this color change, the reaction can be used to detect the presence of trace amounts of dithionite. The color change is transitory-i.e., the red coloration slowly fades and eventually changes to the original yellow. The reversal of the color change is extremely slow when the dithionite concentration is low, but it is relatively rapid when dithionite concentration is high. Subsequent experimentation
1 Present address, Ethyl Corp., Baton Rouge, La.
showed that the rate a t which the color reversal from red to yellow took place was roughly proportional to the amount of dithionite present. This led to a study of the concentration-time relationship and subsequent adaptation to a semiquantitative method for estimating dithionite content. The test was unaffected by the presence of sulfite, hydroxide, alkoxide, chloride, sulfide, sulfur dioxide, or other sodium sulfoxy compounds, so long as all of the ammonia was not neutralized. PROCEDURE
-4solution of 0.8 gram of Naphthol Yellow S (Kational Aniline Division, Allied Chemical & Dye Corp.) and 10
0 1 1 1 1 1 1 1 1 1 1 1
IO0
50
0
Na2S204, WT. Figure 1. Color reversal after reduction of Naphthol Yellow S b y sodium dithionite
ml. of concentrated ammonium hydroxide in 1000 ml. of distilled water was prepared as the indicator solution. Synthetic mixtures of anhydrous sodium sulfite and anhydrous sodium dithionite (commercial grade, approximately 90% pure) were prepared over the entire 0, 20, 30, concentration range-Le., 40, , . . 100% commercial-grade sodium dithionite, diluted with sodium sulfite. One-gram samples of the dithionitesulfite mixtures were dissolved in 25-ml. portions of distilled water, and 25 mi. of the ammoniacal Naphthol Yellow S test solution was added to each. The color change from yellow to red occurred immediately and the time required for the color reversal to take place was determined by a stop watch. The color reversal was gradual, going through several color graduations from red to yellow; consequently, the time recorded for the complete reversal was somewhat arbitrary, depending on the judgment of the observer. Kevertheless, the times were reproducible for a given observer and they were used successfully for estimating semiquantitatively the concentration of sodium dithionite in unknown samples. Unknowns were analyzed for sodium dithionite content (based on commercial purity) by noting the time required for color reversal with 1-gram samples and reading the concentration directly from the prepared curve (Figure 1). $11 operations were conducted a t room temperature. LITERATURE CITED
( 1 ) Jelley, E. E., Analyst 55, 34-5 (1930). (2) Welcher, F.,, J., “Organic Analytical Reagents, pp. 18-20, Van Nostrand, Xew York, 1948.
RECEIVEDfor review October 24, 195G. Accepted May 15, 1957. VOL. 29, NO. 10, OCTOBER 1957
1499
