652
INDUSTRIAL AND ENGINEERING CHEMISTRY
Each titration figure in Table I was obtained from two or more titrations differing by 0.02 ml. or less. The tabulated results show that outside and inside indicators give identical values with a pure sulfate. Apparently these slightly low results tend t o be compensated by titrimetric standardization of the barium chloride solution. This is to be expected, with unknowns of equivalent sulfate-ion concentration, as the effect of probable slight incompleteness of barium sulfate precipitation in the course of a rapid titration is canceled out. For highest accuracy, the best procedure appears t o be standardization of the barium chloride by titration, coupled with dilution of the original sample to permit use of volumes and sulfate concentrations closely similar t o those used in standardization. The results with ammonium sulfate, and other observations, do not entirely convince the authors that addition of alcohol is essential in titrations of this order. Reliable figures for inside indicator titration of liver extract fractions were not obtained because of the indeterminate end points previously mentioned. The tolerance for phosphate ion unfortunately could not be
VOL. 10, NO. 11
raised by the outside indicator method appreciably above the 60 p. p. m. limit determined by Sheen and Kahler (2). Xumerous buffers were tried, b u t the shifting or indeterminate nature of the end points, caused by changes in p H and different buffer systems or concentrations in t h e presence of phosphates, indicated the impracticability of these procedures.
Summary A ne\v and economical tetrahydroxyquinone indicator solution is described which, used externally, greatly facilitates the direct titration of small amounts of sulfate ion in certain unashed samples containing organic constituents. The interference of phosphate ion is again pointed out.
Literature Cited (1) Schroeder, W. C., IND. ESG.CHEX,Anal. Ed., 5 , 403 (1933). (2) Sheen, R. T., and Kahler, H. L., I h i d . , 8 , 127 (1936). (3) I b i d . , 10, 206 (1938). (4) Winsor, H. T.V., Ibid., 9, 453 (1937). RECEIVED June 20, 1938.
Determination of Small Amounts of Potassium A Simpler and More Rapid Variation of the Sodium Cobaltinitrite Method D. S. BROWN, R. R. ROBIXSON, AND G. &I. BROWNING West Virginia Agricultural Experiment Station, Morgantown, W. Va.
T
HE procedure herein described for the determination of potassium by precipitation with sodium cobaltinitrite is simpler and more rapid than any of the many similar procedures which have been proposed since the original work of Adie and Wood (1). The potassium is precipitated in a relatively short time at room temperature. A centrifuge is used t o separate the precipitate, as in the procedures of Kramer and Tisdall (7) and others (5, 8, 9 ) , and obviates the filtration employed elsewhere (1-4, 11, 13). The precipitate is washed only once, whereas two or more washings are required in other procedures ( 1 , 2 , S,6-9,11, I S ) . The use of ceric sulfate (4) instead of potassium permanganate for the determination of the nitrites in the precipitate is advantageous. The end point with ceric sulfate is very sharp. When potassium permanganate is used, often a precipitate of hydrated manganese dioxide is formed, which necessitates the addition of a n excess of sodium oxalate to effect its solution, after which the end point is reached by an additional titration with permanganate. This difficulty is, of course, not encountered when ceric sulfate is used. The procedure is satisfactory for the determination of potassium in amounts ranging from 0.2 t o 1.0 mg., the error, in general, not exceeding 2 per cent. It is especially applicable t o the determination of potassium in plant material and in soil extracts. Reagents Precipitating Reagent. Mix together 46.2 grams of sodium cobaltinitrite, 18.9 grams of sodium acetate, 120.0 ml. of distilled water, and 18.0 ml. of glacial acetic acid. Prepare this solution 48 hours before using. Keep stoppered and in a cold, dark place. Before using, centrifuge to remove any precipitate. Ethyl Alcohol. 95 and 70 per cent by volume. Ceric Sulfate. Dissolve about 9 grams of anhydrous ceric sulfate in 500 ml. of distilled water to which have been added 30 ml. of concentrated sulfuric acid. Make up to 1liter. This solution, which is approximately 0.02 N , may be standardized with sodium oxalate.
Ferrous Ammonium Sulfate. Dissolve 8 grams of FeSO4(KH4)&304.6H,O in 500 ml. of distilled water to which have been added 10 ml. of concentrated sulfuric acid and make up to 1 liter. Sulfuric Acid. Concentrated sulfuric acid diluted 1 to 1. Indicator. 0.025 M o-phenanthroline ferrous complex.
Procedure To 1.5 ml. of 95 per cent ethyl alcohol in a 15-ml. centrifuge tube add a 5-ml. aliquot of the potassium solution. Mix thoroughly. Add dropwise, with continuous shaking, 2.0 ml. of the precipitating reagent. Allow to stand for 1 hour at a temperature of from 20 to 25' C. Centrifuge for about 10 minutes at about 2000 r. p. m., so that the precipitate is firmly packed in the bottom of the tube. Pour off the supernatant liquid and allow the tube to'drain for about 5 minutes. Wash the precipitate with 5 ml. of 70 per cent alcohol, breaking up the bulk of the precipitate by forcing the wash solution in a fine stream from a pipet. Centrifuge for 5 minutes and drain as before. Dry the precipitate for 0.5 hour at 80" to 85' C. to remove all the alcohol. Add 5 ml. of the ceric sulfate reagent and 1 ml. of 1 to 1 sulfuric acid. Heat in a water bath at 90" t o 100" C. until all the precipitate is oxidized, as indicated by its disappearance (usually within about 5 minutes). Maintain an excess of ceric sulfate throughout the reaction (5 ml. of 0.02 N ceric sulfate are sufficient for precipitates containing no more than 0.5 mg. of potassium). Cool to room temperature and titrate the excess ceric sulfate with ferrous ammonium sulfate, using one drop of o-phenanthroline ferrous complex as indicator. The end point is very sharp, the color of the solution changing from pale blue to red. CALCELATION. Milligrams of K = ml. of Ce(SO4)l used in oxidation of the precipitate X normality of Ce(SO& X 6.52. O
Discussion The precipitating reagent is similar in composition to that used by Adie and Wood ( I ) , but is more easily prepared. It involves one solution only, whereas t h a t used in some other procedures (1, 2, 3, 5-8, 11) involves two solutions. Experience has shown that, after its preparation, i t is best to let the reagent stand for 2 days before using. The results will be high if i t is used too soon after being prepared. There is no deterioration of the reagent within 2 or 3 weeks if
XOVEMBER 15, 1938
AKALYTICAL EDITION
it is kept in a cold, dark p1ace-e. g., a refrigerator. It is advisable, however, to make daily determinations with known potassium solutions-e. g., solutions containing 0.2 and 0.8 mg. of potassium per 5-ml. aliquot as a precautionary check on the reagent, especially if it is more than a week old. The precipitation can be carried out a t room temperature. Piper (11) noted that the temperature during precipitation influences the recovery of potassium, but concluded that, in his procedure, the recovery was satisfactory a t room temperature. I n other procedures (6, IS), on the other hand, temperatures from 0" to 6" C. are recommended. Average results showing the effect of temperature on the recovery of potassium for this procedure are reported in Table I. The recovery tends to be high a t low temperatures and low a t high temperatures. It is apparent, however, that the variations are significant only in the recovery of the 0.2 mg. of potassium. Therefore, it is recommended that when determining small amounts of potassium (0.2 and 0.3 mp.) the temperature be maintained a t about 20" C. For larger amounts of potassium, recovery will be eatisfactory over the ordinary range of fluctuation in room temperature. TABLEI. EFFECTOF TEMPERATERE ON
THE
RECOVERYOF
POTASSIUM FROM SOLETIONS O F POTASSIUM SULFATE Actual Amount
of K Mg.
0.200 0,200 0,200 0.200 0.200 0.600 0 600 0.600 0.600 0.600
om i ,000 1
a
Temperature
Amount Recovereda
0 18 20 25 30 0 18 20 25 30 0 18 20 25 30
0.216 0.206 0.203 0.192 0.118 0.615 0.602 0.607 0.601 0.597 1 024 1,000 0.993 0.996 1.000
c.
1,000 1.000 1.000 Arerage of four determinations.
Per Cent Recovered
Mg.
108 103 102 96 89 103 100 101 100 100 102 100 99 100 100
The variation in the recovery a t different temperatures results from a difference in the solubility of the precipitate and from the effect of temperature upon the composition of the precipitate. The proportion of potassium to sodium in the potassium-sodium cobaltinitrite precipitate increases with temperature (11). As the proportion of potassium to sodium increases, the factor used in the calculation must be increased to compensate for the change in composition. The factor 6.52 used in this procedure is satisfactory for room temperatures. A smaller factor would offset the high results a t lower temperatures, and conversely, a larger factor, the low results a t higher temperatures. The precipitation is more convenient at room temperature, of course, although the sensitivity of the reaction is greater a t lower temperatures (12). One washing of the precipitate with 5 ml. of 70 per cent alcohol is adequate. Table I1 shows average results in the recovery of potassium from solutions of potassium sulfate when the precipitates were washed with from one to four successive 5-ml. portions of 70 per cent alcohol. The differences are within the limits of experimental error conceded to this procedure and are not significant. There is a considerable saving of time by washing only once. K h e n an unknown solution is so dilute that a 5-ml. aliquot does not contain a t least 0.2 mg. of potassium, it is necessary to concentrate the solution by the evaporation of a larger aliquot. Satisfactory results cannot be obtained with dilute solutions when precipitation is attempted in aliquots larger than 5 ml., even though the amounts of alcohol and reagent are increased in the proportions used for 5 ml. Fair recovery may be secured with 5-ml. aliquots containing more than 1.0 mg., but the most consistent results are obtained with aliquots
653
containing between 0.2 and 1.0 mg. of potassium. The error, on the average, does not exceed 2 per cent, particularly with the larger amounts of potassium. As might be anticipated, the exact recovery of small amounts (0.2 and 0.3 mg.) is more difficult] resulting occasionally in a greater error which, however, does not exceed 3 per cent. T A B L E 11. EFFECT O F x;UI\lBER OF WASHINGS UPON RECOVERY O F POTASSIUM FROM SOLUTIONS OF POTASSIUM SULFATE (Washed with 5-ml. portions of 70 per cent alcohol) Actual Amount Number of Amount Per Cent of K Was hi ngs Recovered" Recovered
M g. 0.200 0,200
Me. 1
0.200
0.200 0.600 0.600 0.600 0.600 Q
0.202 0.201 0.197 0.201 0.605 0.602 0.602 0.612
101 101 99 101 101 100 100 102
.4verage of three determinations.
Kramer (6) and others (2, I O , 11) have demonstrated that the sodium, calcium, magnesium, barium, strontium, zinc, iron, sulfate, chloride] nitrate, and phosphate ions do not interfere in the volumetric determination of potassium with sodium cobaltinitrite. Similar results have been obtained with this procedure. Ammonia is the only substance which will seriously interfere. In soil extracts ammonia is eliminated from the sample by evaporation and ignition. If there is any danger of contamination from ammonia fumes in the laboratory, the centrifuge tubes should be stoppered during the precipitation. It should be noted, with respect to sodium, that the ratio of sodium to potassium in the reaction mixture influences the composition of the precipitate (14). According to Schueler and Thomas ( I S ) , an excess of sodium over potassium gives the best results. Within limits, a t least, the reaction of the unknown is unimportant, the recovery of potassium being equally satisfactory from solutions ranging in pH from 1.5 t o 12. TABLE 111. EFFECTOF COBALT ON CERICSCLFATE REQUIRED. FOR SUBSEQUENT TITRATION Trial NO.
1 2 3 4
.I\?.
0.021 N Ce(SOa)z Cobalt Cobalt present absent IMZ. MI. 3.72 4.14 3.68 4.07 3.72 4.15 3.66 4.11 3.69 4.12 Ratio 3.69 : 4.12 = 10.8 : 12
The factor 6.52 used in the calculation of the potassium in the precipitate is empirical. It was notkd by Drushel (3) and later by others (8,6, 7, 8, 11) that :n the titration of the nitrites in the precipitate the cobalt is reduced, accounting for one equivalent of the nitrites, the other eleven being accounted for by the oxidizing agent (ceric sulfate in this procedure). If the precipitate has the composition KzNaCo(KO& as determined by Adie and Wood ( I ) , the stoichiometric factor is 7.10 when the action of the cobalt is considered. However, when the cobalt is not considered and all the twelve nitrite equivalents are accounted for by the ceric sulfate, the stoichiometric factor is 6.52, which is also the factor used in this procedure. Since this indicates that the cobalt might not act in an oxidizing capacity under the conditions of this procedure, analyses were made in which the nitrites in two sets of precipitates containing 0.5 mg. of potassium were titrated, one in the presence of cobalt as usual, and the other in the absence of cobalt which was removed by filtration after being precipitated with sodium hydroxide. The data are shown in Table 111.
INDUSTRI4L AND ENGINEERIXC CHEMISTRY
654
The average number of milliliters of 0.021 N ceric sulfate necessary to oxidize the nitrites in the presence of cobalt was practically eleven-twelfths (10.8/12)of that necessary when cobalt was removed-i. e., the cobalt accounts for the oxidation of one of the twelve nitrite equivalents in the precipitate. This confirms the results of Drushel ( 3 ) and others ( 2 ) 6, '7, 8, 11) and establishes the factor 6.52 as empirical for a precipitate of the composition KZSaCo(X0&; its identity with the stoichiometric factor for a precipitate of this composition is merely coincidental. An explanation of the relation may be found in the fact that a variability in the composition of the precipitate between KNa&o(KOz)6 and K 2 ~ a C o ( S o 2 )as6 a result of variations in the conditions of precipitation has been noted by several investigators (8-12, 14). It is possible that the precipitate in this procedure approaches the relative composition Kl.~aSal.16Co(Tu'Oz)6,which is the approxiinate formula for which the factor 6.52 is stoichiometric in the presence of cobalt. No analyses have been made of the composition of the precipitate, and the above formula is merely suggested as a possible means of explaining a seemingly empirical relation. Until the actual composition of the precipitate formed in this procedure is determined, the factor 6.52 must be considered as empirical. The procedure is of definite practical importance, especially when applied to soil extracts, and while the nature of the precipitate and the character of the factor are of interest, the practical application of the procedure should not be neglected because of any obscurity as to the theoretical considerations latent in the method.
VOL. 10, NO. 11
Acknowledgment The authors thank W. H. Pierre, formerly head of the Department of Agronomy and Genetics a t West Virginia University, now head of the Department of Agronomy a t Iowa State College, G. G. Pohlman, head of the Department of Agronomy and Genetics, and R. S. Marsh, head of the Department of Horticulture, a t West Virginia University, for their suggestions and criticisms of the work. They also acknov-ledge the assistance of the American Potash Institute in financing the project of which some of this work is a part.
Literature Cited (1) hdie, R. H., and Wood, T. B., J . Chent. SOC.,77, 1076 (1900). (2) B o w e r , L. T., J. IND.ESG. CHEU.,1, 791 (1909). (3) Drushel, W.A., Am. J . Sei. (Ser. 4),24, 433 (190i). (4) Harris, H. C., Soil Sci., 40, 301 (1935). ( 5 ) Hubbard, R. S., J . Riol. Chem., 100, 557 (1933). (6) Kramer, B., Ibid., 41, 263 (1920). ( 7 ) Kramer, B., and Tisdall, F. F., Ibid., 46, 339 (1921). (8) Lewis, A. H., and Marmoy, F. B., J . SOC.Chem. I n d . , 52, 177T (1933). (9) Lohse, H. TV.,
ISD. ESG.CHEJI.,Anal. Ed., 7, 272
(1935).
(10) Morgulis, S.,and Perley, A., J . Biol. Chern., 77, 647 (1928). (11) Piper, C. S., J . SOC.Chem. I n d . , 53, 392 T (1934). (12) Robinson, R. J., and Putnam, G. L., ISD.ENG.CHEX.,Anal. Ed., 8, 211 (1936). (13) Schueler, J. F., and Thomas, R. P., Ibid., 5, 163 (1933). (14) Tan Rysselberge, P. J., Ibid., 3, 3 (1931).
RECEIVED M a y 9, 1938. Published with t h e approval of the Director of t h e K e s t Virginia Agricultural Experiment Station as Scientific Paper No. 206.
Microscopic Identification of Some Important Substituted Naphthalenesulfonic Acids WILLET F. WHITMORE
T
Ai%D ARTHUR
I. GEBHART, Polytechnic Institute of Brooklyn, Brooklyn,N. Y.
HE commercial importance of the naphthalenesulfonic
acids has resulted in numerous methods for the identification of these difficultly characterized compounds. I n most cases their salt-forming properties have been utilized in preparing metallic or arylamine salts (1, 2, 3, 6-8, 10, I d ) . Chambers and Scherer (4) used the base benzylisothiourea for characterization of a- and 0-naphthalenesulfonic acids and the 1,5-,1,6-, 2,6-,and 2,7-disulfonic acids. Hann and Keenan (11), using microscopic methods with this reagent, reported the optical data on the derivatives of these same acids. The limitation of the above methods is their lack of applicability to large groups of the acids. By a combination of such procedures-metallic salt formation, benzylisothiourea salts, and free acids-Garner (9) outlined a procedure for the microscopic identification of twenty substituted naphthalenesulfonic acids. His method requires the preparation and examination of a number of derivatives of each acid. The method which is here reported is based upon the fact that benzoylation of a number of naphthylamine-, naphthol-, and aminonaphtholsulfonic acids yields characteristic, readily isolated test forms. Photomicrographs of the derivatives and the free acids or their sodium salts are included (all of the same magnification, approximately 70). The latter two are generally poorly described in the literature and are in many cases character-
istic and serve as additional proof of identity. Optical data are given for the derivatives.
General Procedure PURIFICATION OF SAMPLES.A11 acids insoluble in mater are dissolved in strong sodium carbonate solution and treated with activated carbon (Darco). The free acid precipitated with hydrochloric acid is filtered by suction, washed with a little cold water, reprecipitated similarly a second or third time, and dried at 40" t o 50" C. In what follows, unless otherwise noted, this is the purification procedure used. PREPARATION OF BENZOYL DERIYATIVES.For monosubstituted acids-naphthylamine- or naphtholsulfonic a c i d s 4 . 2 gram of the acid or its sodium salt is dissolved in 10 ml. of spproximately normal sodium carbonate solution in a 125-ml. glass-stoppered Erlenmeyer flask and 0.2 ml. of benzoyl chloride (reagent quality) added. With disubstituted acids such as the aminonaphthol type the quantities of sodium carbonate and benzoyl chloride are doubled; otherwise the procedure is the same. Contrary to usual procedure, sodium carbonate appears to serve better than sodium hydroxide for the benzoylation, the derivative frequently precipitating more easily when the carbonate is used. KO trouble is experienced from gaseous carbon dioxide, as it is absorbed in the excess of carbonate used. The flask is stoppered and vigorously shaken until all odor of benzoyl chloride has disappeared. This may take as long as 5 minutes in some cases and a precipitate may or may not form, depending on the sulfonic acid.
