ANALYTICAL EDITION
September 15, 1942
(6) Kohlrausch, F., 2. phyaik. Chem., 44, 233 (1903). (7) Mellor, J. W., “ComprehensiveTreatise on Inorganic and Theoretical Chemistry”, Vol. XI, p. 268, London, Longmans, Green and Co., 1931. ( 8 ) Meschezerski, J., 2.anal. Chem., 21, 399 (1882). (9) Miller, F. T.V., “Laboratory Manual of Qualitative Analysis”, p. 121, New York, D. Appleton-Century Co., 1930. 110) Noll, W., Z . anorg. allgem. Chem., 199, 193 (1931).
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(11) Noyes, A. A., and Swift, E. H., “Qualitative Analysis”, 10th ed., p. 283, New York, Macmillan Co., 1942. (12) Reichard, C., Chem.-Ztg., 27, 877 (1903). (13) Ricci, J. E., and Davis, T . W., J.Am. Chem. SOC.,62, 407 (1940). (14) Seidell, A., “Solubilities of Inorganic and Metal Organic Compounds”, p. 296, New York, D. Van Nostrand Co., 1940. (15) Stewart, P. B., and Kobe, K. A., ISD. Exo. CHEM., ANAL.ED., 14, 298 (1942).
Volumetric Determination of Iron and Titanium WALLACE nl. RICNABB AND HERMAN SKOLVIK Department of Chemistry and Chemical Engineering, University of Pennsylvania, Philadelphia, Penna.
A
METHOD for the determination of iron and titanium has been described b y Lundell and Knowles (6, 7 ) . The iron is determined b y the stannous chloride reduction method or precipitated as the sulfide from a n ammoniacal tartrate solution and both iron and titanium are determined by means of the Jones reductor, the titanium being found b y difference. Kakazono (8) recommends the use of liquid zinc amalgam for the reduction of an aliquot portion of the solution and titration with standard permanganate to give the total iron and titanium. A second portion of the solution is reduced, and added to the first solution, and the combined solutions are titrated for excess iron necessarily present in using this method. According to Gooch and Newton (5) the use of bismuth trioxide to oxidize only the titanium, after the titanium and iron have been reduced by heating with zinc and acid, gives excellent results for the iron. The method requires the filtration of the bismuth trioxide preceding the titration of the ferrous iron. Thornton and Roseman (12) reduced titanium and iron by passing the solution through a Jones reductor (a zinc-amalgam column), using air as the differential oxidizing agent for estimating iron in the presence of titanium. They suggest that “it may be well to introduce the proviso that the procedure is most apt to succeed when the iron is equal to, or preponderates over, the titanium.” Axt and LeRoy (1) state that they were able to air-oxidize reduced titanium within 3 to 8 minutes. They Jvorked with 0.04485 and 0.116 gram of titanium but gave no direct results of their experiments other than a curve. According to this method the titanium !vas reduced Tvith liquid zinc amalgam and finally air or oxygen was bubbled through the reduced titanium solution by means of an apparatus so constructed that a porous plate divided the air or oxygen into a multitude of small bubbles. Brandt (3) used titanium trichloride as the reducing agent for the iron, the excess titanium trichloride being destroyed with copper sulfate. In this procedure, too, it is necessary to carry out a filtration before the titration of the ferrous iron. The object of this investigation \vas to find a more rapid, accurate, and convenient method for the determination of iron and titanium. Reduction of the iron and titanium and subsequent titration with permanganate would give the combined amounts of each. Air-oxidation of the titanium alone, following the reduction of both the iron and titanium, would make i t possible t o titrate the iron and determine the titanium b y difference. This required a study of the air-oxidation of titanous and ferrous solutions.
Reagents Titanium dioxide, purity 98.62 per cent as determined by the cupferron method (10). (Spectrographic analysis gave no positive test for other reducible materials such as chromium and vanadium but did show presence of some magnesium which does not interfere.) The following materials were c . P. reagent grade: ferric ammonium sulfate, 100 grams per liter of 0.5 M sulfuric acid; ammonium sulfate; sulfuric acid, 36 N ; mercuric chloride, saturated solution in distilled water; zinc, 30-mesh. Zimmermann-Reinhardt solution: 70 grams of manganous sulfate, 125 ml. of concentrated sulfuric acid, and 125 ml. of 85 per cent phosphoric acid diluted to a liter.
Potassium pcrmanyanatc, approximately 0.1 and 0.05 A’ solutions prepared in the usual way and standardized against sodium oxalate according to the Sational Bureau of Standards method
(4)* Air-Oxidation of Titanium A standard solution of titanium sulfate was prepared by dissolving titanium dioxide in concentrated sulfuric acid and ammonium sulfate. The solution was adjusted to an acidity of 3 to 4 N with respect to sulfuric acid, and analyzed by reducing the titanium, pouring through a zinc reductor, and receiving into a ferric alum solution ( I O ) . The reduction was repeated bvith solutions containing from 0.0203 to 0.2512 gram of titanium dioxide, the reductor was removed from the suction flask, and a rubber stopper containing a glass tube was fitted into the mouth of the flask with the end of the tube dipping into the solution. Air was bubbled through the solution by means of water suction for 25 to 30 minutes. The unoxidized titanium was determined by titration with potassium permanganate. The results are given in Table I. A more rapid method for the oxidatioii oi Ti1rrto Ti1” was investigated. Fifty milliliters of a saturated solution of mercuric chloride were added to a reduced titanium solution. After standing for one-half hour, with occasional shaking, the solution showed incomplete oxidatiori of titanium. A solution of titanium sulfate in 4 S sulfuric ;tcid ivvits reduced, by running through a reductor, and received into a 4 N sulfuric acid solution. To this were added 50 ml. of a saturated solution of mercuric chloride, and air was bubbled through the solution until the violet color disappeared and then for 5 minutes longer in an attempt to ensure complete oxidation of the titanium. The unoxidized titanium \vas determined by permanganate titration.
T.AHI.F: I. AIR-OXIDATION O F TITANOUS SULFATE SOLUTIOK TiOi Taken Gram 0.2512 0,1252 0.0626 0.0203 ‘1
h
Air
.Lf in, 400
40b 30 26
Ti02 Oxidized Gram 0 . 1991 0,0946 0.0526
o ozno
Solution had a pronounced violet color. Solution had a faint violet color.
TABLE
11. AIR-OXIDATION O F TI,rAh-ous SUI,FATESOLTJTIOS IN THE PRESENCE OF XERCURIC CHLORIDE
Ti02 Taken Gram 0,3425 0,2612 0 . 1557
0.1252 0.0626 0.0203
HgCln
Air
‘MI.
‘lli71.
50
15
50
16 10 10
eo
00 50 50
Ti02 Oxidized Gram 0,3423 0,2512 0. 1657 0 . 1250 0.0625 0.0203
Number of Determinations
Individual values: a0.1657, 0.1557, 0 . 1 5 5 5 , 0.1387, 0.1656, 0.1956, 0 . 1 5 3 b 0.0624, 0.0627.
4 3 7a :1 2b
2
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INDUSTRIAL AND ENGINEERING CHEMISTRY
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was titrated with standard permanganate, and from this the TABLE111. AIR-OXIDATIONOF FERROUS SULFATESOLUTION number of milliequivalents of permanganate equal to the iron and titanium was obtained. The reduction wm repeated, and the Number of reduced iron and titanium were poured throu h the reductor, and FezOa Taken Air Fen08 Found Determinations received into a solution of 4 N sulfuric acif. To this solution Gram Min. Gram were added 50 ml. of a saturated solution of mercuric chloride. 4 15 0.6219 0.6231 Air was bubbled through the solution until the violet color dis20 0.6224 3 0.6231 appeared, then for 5 minutes longer. The ferrous iron was ti15 4 0,2963 0.2961 20 40 0,2963 0.2963 trated with standard permanganate and calculated as ferric oxide. 20 0.1732 2 0,1732 The difference between the number of milliequivalents in this 25 2 0.1732 0.1732 and the first titration represents the number of milliequivalents 15 2 0.1226 0.1226 0.1226 2 20 0,1226 of titanium dioxide. The results obtained are shown in Table V. 0.0635 15 3b 0.0835 The procedure was repeated for the determination of iron in the 0.0754 20 3 0.0754 presence of titanium by reducing and bubblin air through the Individual values: a 0.2963, 0.2961, 0.2964, 0.2964. solution in the absence of mercuric chloride (Ta%leVI). b 0.0832, 0.0635, 0.0637.
TABLEIV. AIR-OXIDATIONOF FERROUS SULFATE SOLUTIOX IS THE PRESENCE OF MERCURIC CHLORIDE Fez03 Taken HgClz Air FerOs Found Gram M1. Min. Gram 0,3424 50 120 0.3418 0,3243 50 30 0.3243 0.1835 50 10 0.1635 50 20 0.1836 0,1835 50 30 0.1836 0.1635 50 10 0.1697 0.1698 50 20 0.1696 0.1696 50 30 0.1697 0.1698 0.1696 50 40 0.1696 50 30 0.0971 0.0971 50 30 0.0666 0.0666 50 30 0.0463 0.0463 50 30 0.0173 0.0174 Individual values: 0. 0.1635 0.1835 0.1636 0.1637. b 0.0970: 0.0970: 0.0972: C 0.0463, 0.0464, 0.0462, 0.0464.
Number of Determinations
2 4 3 2 4a 3 2 3 2 3b
4 4c 4
The results are given in Table 11. Thus. although air alone or mercuric chloride alone was not sufficient for the oxidation of trivalent titanium, the action of the two combined was very effective and resulted, in a short time, in a catalytic oxidation of the titanium. In Tables I1 to VI the final values are the averages of the results of the number of determinations indicated in the last column. The accuracy of the method may be judged further by examination of the individual values given in Tables 11, 111,IV, and V as typical.
Air-Oxidation of Ferrous Iron It has been known for some time that ferrous salts are only slowly oxidized by bubbling air through their solutions. Baskerville and Stevenson (2)observed no oxidation of ferrous to ferric iron upon passing air through neutral or acid solutions of ferrous sulfate for 3 hours. The results were shown to be influenced only negligibly by dust in the air, glass wool, pumice stone, and various salts. They also observed no oxidation on reducing a sulfuric acid solution of ferric sulfate with zinc, filtering off the zinc, and passing air throu h the solution for 3 hours. More recently (9) Pound concluded tgat the oxidation of ferrous salts was a minimum in the presence of a strong nonoxidizing acid such as sulfuric or hydrochloric. When the concentration exceeded 0.2 N for the hydrochloric and 6.0 N for the sulfuric acid, the tendency for the oxidation of ferrous iron by air became appreciable. A solution of ferric alum was prepared in 4 N sulfuric acid, analyzed for iron ( I I ) , and calculated as iron oxide. Aliquot portions of this solution were poured through the reductor in the usual way, and air was bubbled through the solutions for 15 to 25 minutes. The results given in Table I11 show some oxidation of the ferrous sulfate when large quantities of iron were present. I n another series of experiments a saturated solution of mercuric chloride was added, after the reduction of iron, and air bubbled through the solutions for 10 to 120 minutes. The ferrous iron was titrated with standard permanganate and calculated tis iron oxide. Results are shown in Table IV.
Air-Oxidation of Iron and Titanium Together Accurately measured volumes of standard solutions of ferric sulfate and titanium sulfate were transferred to a beaker. The solution was adjusted t o approximately 4 N in sulfuric acid, reduced with zinc, poured through the reductor with moderate suction, and received into a ferric alum solution. The ferrous iron
TABLE V. AIR-OXIDATIONOF TITAXOUS SULFATE SOLTJTIOX IS PRESENCE OF FERROUS SULFATE AND MERCURIC CHLORIDE
THE
Number of Ti02 Taken Fen03 Found TiOz Found Determinations Gram Gram Gram 0.2782 0,7795 2 0.2776 0.1391 2 0.7798 0.1385 0.2782 0.3904 0.2760 5 0.1391 0.3903 0.1392 35 0,3902 0.0927 0,0929 2 0,3902 0.0685 0.0668 2 0.0343 0.3898 0.0345 2 0.3899 2 0.0171 0.0173 0.1948 0.1391 0.1390 4 0.0975 0.1388 0.1391 25 0.1392 0.1391 0.0326 3 7ralues: a Fen08 found: 0.3905, 0.3903, 0.3901 Ti02 found: 0.1390, 0.1392, 0.1394. b FezOa found: 0.0977, 0.0973 TiOz found: 0.1369, 0.1387.
1?el08 Taken Gram
0,7600 0.7800 0.3900 0.3900 0,3900 0.3900
0,3900 0.3900 0.1950 0.0975 0.0325 Individual
TABLEVI. AIR-OXIDATIONOF TITANOUS SULF.4TE PRESENCE OF FERROUS SULFATE FerOs Taken
TiOz Taken Air Gram Gram Min. 0.2512 40" 0.3962 0.2369 0.2512 40 a 0.2512 40" 0,0637 0.1256 25: 0.2369 0.0628 15 0.2369 a Solution had a pronounced violet color. b Solution became colorless within 13 minutes. C Solution became colorless within 2 minutes.
I S THE
Ti02 Unoxidized Gram 0.1672 0.0786 0.1954 0.0428 0.0061
Discussion The simplicity of construction of the zinc reductor prepared by sealing a glass tube on the end of a 2.5-cm. funnel and stocking the funnel with 30-mesh zinc makes it available at very little cost to any laboratory. Air alone is not sufficient for the complete oxidation of titanium by this method. However, when the reduced titanium solution is aerated in the presence of a large amount of mercuric chloride, catalytic oxidation takes place in a short time. The results obtained by this method are not in agreement with those of Thornton and Roseman ( I d ) , and a part of their work was repeated using a Jones reductor (zinc-amalgam column). Although their results were reproducible, it was found that a considerably larger volume of solution remained to be aerated and no definite time could be set for complete oxidation of the titanium. According to the results and statement given in the summary of their paper, oxidation of the titanium is more likely to succeed if the amount of iron is equal to or greater than the titanium present. A definite procedure is given in this paper whereby the time factor for the quantitative osidation of titanium is controlled irrespective of the amounts of iron present. It might appear at first that it is not necessary to add so much as 50 ml. of a saturated solution of mercuric chloride. It is true that smaller amounts of mercuric chloride can be used for the oxidation of small quantities of titanium, but such quantities cannot be assumed in R sample of unknown composition. A thorough investigation resulted in the use of 50-ml. of mer-
ANALYTICAL EDITION
September 15, 1942
curic chloride in order to give the best results for any amount of titanium up t o about 0.6 gram of titanium dioxide. Since such large amounts of titanium would not be used in a single analysis, these values are omitted from the tables. The volumetric estimation of iron and titanium generally used in the titanium pigment industry involves reduction of solution (zinc-amalgam reductor), titration of the combined iron and titanium with standardized potassium permanganate, and separate titration of titanium with standardized ferric ammonium sulfate. The method described in this paper has the advantage of requiring the use of only one standardized solution.
Acknowledgments The authors gratefully acknowledge a grant from the Faculty Research Committee of t h e University of Pennsylvania and wish to thank R. D. Summers of the DeDartnient of for making the Physics Of the University Of spectrographic analyses.
713 Literature Cited
k i t , M., and LeRoy, M., Ing. chim., 24, 28 (1940). Baskerville, C.,and Stevenson, R., J . Am. Chem. SOC.,33, 1104 (1911).
Brandt, L., Chem.-Zto., 42, 433,450 (1918). Fowler, R. M.,and Bright, H. A., J . Research Natl. Bur. Standards, 15, 493 (1935). Gooch, F. A, and Newton, H. O., Am. J . Sci., 23, 365 (1907). Lundell, G. E. F., and Knowles, H. B., IND.ENQ.CHEM.,16, 723 (1924). Lundell, G. E. F., and Knowles, H. B., J . Am. Chem. SOC.,45, 2620 (1923). Nakazono, T., Sci. Repts. Tohoko Imp. Univ., 16,733 (1927). Pound, J. R., J. Phys. Chem., 43,955 (1939). Skolnik, H., and McNabb, W. M., IND. Eso. CHEM., B x . 4 ~ED., . 12, 672 (1940). Skolnik, H., and McNabb. W.M.. Netal Fin.. 39.241 (1941). Thornton, W.M., Jr., and Roseman, R., J . Am. Chem.'Soc,;57, 619 (1935). before the Division of Analytical and Micro Chemistry at t h e 102nd Meeting of the BYERICAN C H E m c A L SOCIETY, Atlantic City, N.J.
PREsENTED
Analysis of Tin-Base Bearing Metal Permanganate and Iodometric Methods for Antimony and Copper JOSEPH R. ANDREWS AND ALEXANDER J. BENDER Chemical Laboratory, United States Navy Yard, Portsmouth, Va.
UhIEROUS combinations of standard methods can be used to obtain accurate and reliable values for antimony and copper in tin-base bearing metal. However, since these methods have been found t o be too time-consuming for routine work, various methods have been investigated in this laboratory.
Experimental When copper was separated from 10 to 20 times as much tin plus antimony by the nitric acid-hydrolysis method (8,IO), some copper invariably remained with the tin and antimony acids. The quantity of copper retained b y different precipitates varied so much that a constant correction could not be used, and each precipitate had to be treated to recover the occluded copper. Fusion of the precipitate with sodium carbonate and sulfur ( 5 ) ,leaching with sodium sulfide and alkaline tartrate solution, and also volatilization of the tin and antimony b y heating with ammonium iodide at 500" to 560" C. were found to give quantitative recoveries of the copper. Separation of copper from a dilute sulfuric acid solution as the thiocyanate (11) required much less time than any of the above procedures, while separation of the copper as the sulfide from a n alkaline tartrate solution (7) was the fastest method found in the literature. Attempts to determine copper in tin-base bearing metal b y the iodometric method (1-4, 10) led to low and variable results. The voluminous precipitate of tin and antimony apparently interfered, perhaps b y retaining some copper from the solution and b y adsorbing some of the iodine formed in the reduction of the cupric salt b y potassium iodide. Kolthoff and Sandell (9) discuss the iodometric determination of arsenic in a tartrate solution of Paris green. They add that it is possible to determine the copper iodometrically in the same solution after adding sulfuric acid t o give the proper acidity. The proper acidity is not stated. Since antimony can be determined by permanganate titration of a n acid
solution containing tartaric acid (6), the iodometric titration of the copper in the resulting solution should permit the rapid determination of antimony and copper in tin-base bearing metal.
A series of tests was carried out along these lines. When the solution was acidified with acetic acid and the usual amount of potassium iodide added, 3 grams per 80 ml. of solution, the titration was not satisfactory because of drifting end points and incomplete reduction of cupric salts. The concentration of cupric ion evidently was too low, owing to the complex tartrate of copper. In another series of tests the added potassium iodide was varied up to 10 grams per 80 ml. of solution. Increasing the potassium iodide concentration favored the reduction of cupric ions and permitted quantitative titrations when 8 to 10 grams were added to each 80 ml. of solution, but the reaction was slow and the drifting end points were troublesome. Since satisfactory titrations were not obtained on weakly acidified tartrate solutions, another series of tests was carried out in which the pH was lowered by the addition of various amounts of sulfuric acid. At a pH of approximately 2.65 and a potassium iodide concentration of 6 grams per 80 ml. of solution, the reduction of cupric ion was both quantitative and relatively rapid. Antimony mas not reduced even when twice the quantity of sulfuric acid was added, although the additional acid accelerated the air oxidation of the potassium iodide and caused an iodine blank equivalent to 0.15 t o 0.20 ml. of 0.1 N thiosulfate. When still more acid was added the reduction of antimony became apparent. Evidently, the tartrate complex of antimony requires a relatively high acid concentration for reaction with potassium iodide. Because of the large Concentration of potassium iodide required, Griffin's procedure for antimony (6) was modified so as t o reduce the volume of solution and thereby reduce the quantity of potassium iodide required for the reduction of cupric ion. Several difficulties were encountered in the determination of antimony by this modified method, and the four main difficulties are explained below. 1. When the sample was digested in a too rapidly heated sulfuric acid-potassium sulfate mixture, some sulfur sometimes collected high up on the sides of the flask. Because of this, it has been found advisable to heat the acid mixture moderately while solution of the sample is taking place. 2. Some difficulty due to fading of the end point was experienced in the permanganate titration, owing to the presence of
