ANALYTICAL CHEMISTRY
144 teristics of the 103-0 plates used were measured by rotating a logarithinic step sector (ratio 1 to 1.5) a t Sirks’s focus (6) before the slit. An iron arc sector pattern was photographed on one of each new box of plates used and the H and D curve ( 2 )thereby obtained. Limitations imposed by the procedure, as herein outlined, are that the compound used should be (preferably) of a crystalline nature, and it should possess a suitably low volatility, such that oven or other drying will not volatilize the compound along with the alcohol carrier. Doubtless other modifications of the technique might be made toward investigating liquids and volatile compounds. The technique should prove of value in the study of nitrogenous wastes, etc., where a semiquantitative value would suffice.
Measurements of density .with the visual-type densitometer used indicate that a precision of *2.0% nitrogen may be obtained. This precision could conceivably be improved by the use of a photoelectric microphotometer wherein human error is eliminated or greatly reduced. LITERATURE CITED
(1) Gerlach, Walther, “Foundations and Nethods of Chemical Analysis by the Emission Spectrum,” Chap. 5 , London, Hilger, Ltd.. 1929. (2) Hurter, F., and Driffield,V. C., J . Soc. Chem. Ind.. 9, 455 (1890). (3) Sachtrieb, N. H., Ph.D. dissertation, University of Chicago, 1941. (4) Sawyer, R. A., “Experimental Spectroscopy,” pp. 151-4, New York, Prentice-Hall. 1946. (5) Sirks. J. L., - 4 S t ~ ~ n o ? ? and ? Y dstrophysics, 13,763 (1S94). RECEIVED September 2 1 , 1950.
Alphazurine G and Some N,N’-Substituted Benzidines as Redox Indicators *
RALPH N. ADAMS’
AND
ELLWOOD M. HAMMAKER2, Rutgers University, New Brunswick, .V.J .
The purpose of this study was to investigate the usefulness of a typical triarylmethane dyestuff as a redox indicator. The dye chosen for this purpose was Alphazurine G, C.I. No. 712, which had already proved to be an excellent and reliable indicator. The transition potentials of Alphazurine G in various concentrations of sulfuric, hydrochloric, and perchloric acids are reported for the titration of ferrous iron and ferrocyanide with cerate. Alphazurine G
T
H E triarylmethane dyes were introduced as reversible oxidation-reduction indicators in 1931 by Knop ( 7 ) . Several investigators have applied a few of these indicators to specific titrations (4,5,8, 9, 17). Details as to the chemical constitut’ion, trade names, C,olour Index number, etc., of the dyes are given by Knop ( 7 ) . In contrast to the widely acclaimed o-phenanthroline or ferroin series of indicators also introduced in 1931 by Walden, Hammett, and Chapman (16) and exhaustively studied by Smith and coworkers ( 2 , Is), the triarylmethane dyes have received onlv secondary interest in volumetric analysis. I n view of their striking color changes, excellent reversibility, and extensive applicability to cerate oxidimetry, this is hardly justified. The limited use of these indicat’ors is due in part to B lack of experimental evidence concerning their precise elect,rochemical and indicator behal-ior. Beyond the original approximate determination of the transition potentials of the dyes of Knop ( 7 ) and the re-evaluation of the transition potentials of Patent Blue V by Yoe and Boyd (18) and Setopalin C by hliller and Van Slyke (11), very little has been published concerning the electrochemical behavior of these indicators. KO study has been made of the variation of transition potential with type and concentration of acid or in any titration other than the iron-permanganate reaction. The purpose of this study was to investigate the redox indicator properties of a typical triarylmethane dyestuff with respect to the t,ransition pot,ential and its variation with type and concentration of acid, and the application of the dye to specific 1 Present address, Chemistry Department, Princeton University, Princeton, N. J. 2 Present address, Research Center, Johnson and Johnson, Yew Brunswirk. N. J.
has also been found to give reliable results as a risual indicator for the titration of ferrous iron. Furthermore, two new oxidation indicators have been proposed, N,N’-tetramethylbenzidine and N,.Y’-tetramethylbenzidine-3-sulfonic acid. The former was found to be stable at temperatures up to 100” C. This information and these indicators may prove valuable in filling in the gaps in the existing knowledge concerning redox indicators. titrations. The dye chosen for this purpose was hlphazurine G, C.I. S o . 712. I t was introduced by Whitmoyer ( 1 7 ) for the titration of ferrous iron, ferrocyanide, and hydroquinone with cerate. The indicator properties of the triarylmethane dyes are similar, and though Eriogrren and Erioglaucine are perhaps the t x o best known dyes of the series, Alphazurine G was chosen as representative of the group because it had already proved t o be an excellent and reliable indicator and should possibly receive more attention. REAGENTS
Alphazurine G, commercial sample, National Aniline Co., 74% active dye, was used without further purification. Hexanitratoammonium cerate, c.P., G. F. Smith Chemical Go. Ferrous sulfate, FeS04.7H20,c.P.,Merck. Stannous chloride, SnCI2.2H20,c.P.,hlerck. Potassium ferrocyanide, Kpe(CN)~.3H20,c.P., Srerck. Sulfuric, hydrochloric, and perchloric acids of desired strength were prepared by proper dilution of t,he C.P. concentrated acids. Cerium(1V) solutions in tmheabove acids were prepared by dissolution of ceric hydroxide in the warm acid, using hesanit’ratoammonium cerate as a rimary standard, according to the method of Smith and Fly &SI. APPARATUS
All potentiometric tit.rations and related e.m.f. measurements were made with a Leeds & Northrup Type K1 potentiometer and enclosed lamp and scale galvanometer. The calomel electrode was of the ground-glass cap salt-bridge type manufactured by Leeds & Northrup. Magnetic stirring and a titrating stand with fluorescent.lamp were provided for all titrations. EXPERIMENTAL
The transition potentials were determined by a modification of Knop’s original procedure ( 7 ) . Port,ions of approximately 0.1 N ferrous solutions in varying concentrations of sulfuric, hydro-
745
V O L U M E 2 3 , NO. 5, M A Y 1 9 5 1 Table I. Transition Potentials of Alphazurine G in Sulfuric, Hydrochloric, and Perchloric Acids Sulfuric Acid Molarity Et (v.), usualu Et ( v . ) , Z-R Molarity Et (v.), usual Et (v.j, Z-R Molarity Et (v.), usual Et (v.), Z-R
1.0 1.03 0.92
Hydrochloric Acid 1.0 1.02 0.96 Perchloric Acid
2.0 1.02
4.0
0.93
0.93
2.0 1.02 0 95
4.0 1 01 0.93
2.0 1.06 1.02 1.02 Ferrocyanide Titrations 1.0 1.07
1 . 0 molar Sulfuric acid 1 . 0 molar Hydrochloric acid Ferrous solutions titrated directly. Zimmerliiann-Reinhardt procedure.
4.0 1.07 1.03
1.02
6.0 1.06 1.03
0 . 9 9 v. 0 . 9 8 v.
chloric, and perchloric acids (prepared in all cases b y dissolving the required weight of ferrous sulfate heptahydrate in the proper acid of desired molarity) were titrated under potentiometric control with 0.1 N cerate prepared in the same acid and concentration, uaing 2 drops of 0.01 Alphazurine G as visual indicator. Potential readings were taken usually a t 1-ml. intervals, except in the region of the mid-point of the titration where several readings were taken to interpolate for the formal potential of the imic-Ferrous system. Within 0.1 to 0.2 nil. of the known esuivalence Doint, 0.01 N. and finally 0.002 A', cerate was added dropwise t,o raise the potential slowly. The transition potential was recorded a t the indicator color change as compared to a color standard containing the same quantities of acid, ferrous solution, and indicator and previously titrated to within 0.5 nil. of the end point. I n all cases the titration was carried beyond the color change and then reversed by addition of 0.002 ;V ferrous solution. This was repeated several times, narrowing down the transition interval as much as possible with each reversal. .It no time did a variation of transition potential occur Kith repeated reversal of the color. In some cases the reverse color change (addition of ferrous) was more distinct than the forir-ard and was of grester value in evaluating the transition interval. The saturated calomel electrode was used as the reference electrode. llamer ( 6 ) has reviewed the e.m.f. values and liquidjunction potent,ials encountered with the various calomel electrodes a i d gives 0.246 volt as an appropriate value for the saturated calomel rvhen working with solutions of strong acids. This value ~ 1 - used a ~ throughout as a standard e.ni.f. for the reference electrode and ail e.m.f. values were referred to the hydrogen electrode by adding the above value to the observed e.m.f., making no ot,her corrections for liquid-junction potent,ials. However, in the measureinen+ in perchloric acid media, a saturated sodium perchlorate d t bridge was interposed between t.he calomel electrode and the titration mixture, because perchloratoccrate rapidly oxidizes chloride resulting from diffusion from the calomel salt bridge. The re,sulting unknown liquid-junction potential was evaluated by removiiig the perchlorate salt bridge midway through the titration and pkicing the reference electrode directly in the titrat,ion mixture just long enough to make a measurement. This was done again a t the end oi the titration and the two values were averaged to represent the mean correction. The magnitude of this corrrction varied from io to 90 niv. in the various perchloric acid concentrations studied. In order t,o test the effect of mercuric chloride and other compounds on the transition potential, ferrous solutions, previous to titration, were treated as described by Rieman, Seuss, and Naiman (12) for the determination of iron by the ZimmermanIieinhardt procedure. The comparison standard was treated in exactly the same manner and the titration made as previously described. The silky, white precipitate of mercurous chloride did not interfere in any n-aylvith the evaluation of the color change. DISCUSSION
The formal potentials of the ferric-ferrous system were checked in all titrations to serve as references. In 1.0, 2.0, and 4.0 M aulfuric acid the values were 0.68, 0.68, and 0.69 volt, respectively. I n 1.0, 2.0, 4.0, 6.0, and 8.0 M hydrochloric acid the values were 0.69, 0.68, 0.66, 0.61, and 0.56 volt. In similar concent#rationsof perchloric acid the values were 0.74, 0.74, 0.77,
0.i8, and 0.87 volt. These rcsults are in excellent agreement with similar data obtained by Smith and Richter (14). The transition potentials of Blphazurine G under various conditions are summarized in Table I. All values are in volts referred to the hydrogen electrode as zero, and all values listed are the mean of three or more determinations. In concentrations greater than 4 ;If in sulfuric or hydrochloric acid media, the intense yellow color of the reduced form of Alphazurine G masked the transition interval to a point where it could not lie determined with any accuracy. The transition potentials w r c also determined for the titration of ferrocyanide wit'h cerate in 1 Jf sulfuric and hydrochloric acids. .in attempt was made to determine the transition potential by differential potentiometric titration of the dye and a reference redox system in the manner used by Smit,h and Richter (14>, for o-phenanthroline and the substituted phenanthroliiies. The stannic-stannous couple &-as chosen as the reference redox system because its fairly lo^ formal potential could be expected to aid in obtaining a sharp differential Oxidation. Dichromate in 1 Jf hydrochloric acid was used as the oxidant and :t ~ v c l l defined differential oxidation was obtained. However. the titration was not reproducible, the shape of the upper half o i the titration curve and the volume increment between the osiilation of stannous ion and the dye being dependeiit upon rate of stirring, manner and rate of addition of oxidant, etc. In viea- of later studies, this is recognized as due to a lack of stability of the oxidized form of Alphazurine G. Alphazurine G was tested as an indicator in the visu:il titrations of 0.1, 0.01, and 0.001 N ferrous solutions. Portions of these solutions were titrated ( I ) using 2 drops of approximately 0.007 J!f water solution of Alphazurine u, and (21 potenTi0metrically. 811 titrations were made in 1 JI sulfuric avid. I n the case of 0.1 S and 0.01 2%- solutions, thc correction \\-:i* tir,gligible, amounting to only 0.03 nil. of 0.01 .\- cerate for the latter. For 0.001 S solutions the indicator correction vas 0.27 nil. of 0.001 -V cerate. The triphenylnwthane dyes :tre subject t o destructive side oxidation by the action of excess uf strong oxidizing agents, which may result in variable indicator corrections if the oxidant is added in a variable or too rapid nianner during the titration. steady, dropwise addition of the oxidant. is recommended for titrations using thew iiitlicators. Brennecke (3) h:is given a thorough review of this frequilntlj- overlooked problem. TETRARIETHYLBEhZIDINE A S D
TETRAMETHYLBENZIDINE-3-SULFONIC ACID 111 connection with \vork on the mcchariisni of the oxidation of rilphazurine G it \viis dc~sir;al)let o prcpare J,.\-'-tetraniethylbenzidine and study its indicator properties. The compound was prepared according to the method of UIIniann and Dieterle ( 1 6 ) and the purified product melted at 190" C. The 3-sulfonic acid was also prepared hy the action of fuming sulfuric acid on the amine according to the directions in Beilitein (1). These compounds were made up as 0.27, solutions i n 1 3f hydrochloric acid for use as indicators. The amine is very soluble in dilute hydrochloric acid to give a clear, colorless solution. The ammonium salt of the sulfonic acid is dissolved in hydrochloric acid by the addition of a drop of concentrated sulfuric acid, giving a clear, pale yellow solution. The tetraethylbenzidine \\-as also prepared, but was found to be unsatisfactory as a redos indicator.
The transition potentials of these compounds i n molar sulfuric and perchloric acids were obtained by the same procedure used for hlphazurine G. These values for the tetraniethylbenzidine are 0.86 volt in 1 :21sulfuric and 0.90 volt in 1 JI perchloric acid. For the 3-sulfonic acid derivative, the transition potentials are 0.88 in 1*If sulfuric and 0.91 volt in 1 M perchloric acid. These compounds were then tested for their application to visual titrations, comparing visual and potentiometric titrations as before. For both indicators the corrections are negligible with 0.1 and 0.01 S solutions, while with 0.001 N solutions the corrections amount to +0.54 and f0.60 ml. of 0.001 A' cerate
ANALYTICAL CHEMISTRY
746 for tetramethylbenzidine and tetramethylbenzidine-3-sulfonic acid, respectively. The use of these compounds as redox indicators has not been previously reported. This is due in all probability to the original observation by Michler and Pattinson (10) and Knop (7), that tetramethylbenzidine is only very slowly reversible in its color change-Le., it does not immediately reverse its color when I drop of reducing agent such as ferrous ion is added to the oxidized form. However, it was found to be rapidly reversible when an excess of reducing agent is added, and with a countertitration the end point can be readily determined. The sulfonic acid derivative, on the other hand, is truly reversible. One drop of even 0.001 N ferrous or cerate solutions causes an immediate color shift in 1 M sulfuric acid. The color change of both of these indicators is from colorless to a deep yellow. A faint pink develops within 0.5 to 1ml. of the end point and serves as a “warning color” preceding the end point. In accord with their color change, they are not serviceable for titrations of ferrous iron in hydrochloric acid, because of the yellow color of ferric ion in this acid, which obscures the indicator color. They are likewise unsatisfactory for titrations of ferrocyanide, owing to the color of ferricyanide ion formed in the oxidation. Tetramethylbenzidine is more reliable than the sulfonic acid derivative in titrations of dilute solutions. The sulfonic acid was found superior to the tetramethylbenzidine in perchloric acid media by virtue of its more rapid color change. As with the triphenylmethane dyes, the titration must be carried out slowly within 1 ml. of the end point, as the indicator response is slugglsh in this region. This is more pronounced in very dilute solutions and a t least 30 seconds should elapse between drops in titrations of 0.001 N solutions. However, in contrast t o the triphenylmethane dyes, there is absolutely no fading of the end-point color. Of special interest is the fact that the oxidized form of tetramethylbenzidine is stable at high temperatures. Two drops of the tetramethylbenzidine indicator solution and I drop of 0.1 Ai cerate in a volume of 100 ml. of 1 M sulfuric acid gave a deep yellow solution which showed no diminution in color after 10 minutes a t the boiling point. This suggests that this indicator may be serviceable for titrations a t elevated temperatures. SUMMARY
The transition potentials of Alphazurine G have been studied in varying concentrations of sulfuric, hydrochloric, and perchloric
acids for the titration of ferrous iron and ferrocyanide. The effect of mercuric chloride and other compounds present in the Zimmermann-Rcinhardt iron titration on the transition potential of hlphazurine G has been determined. Alphazurine G has been tested as a visual indicator for the titration of ferrous iron and found to give reliable r e d t s when solutions from 0.1 to 0.001 S are used. 90indicator corrections are involved with 0.1 S solutions. Y,N’-tetramethylbenzidine and ~V,S’-tetraniethylbenzidine-3sulfonic acid are proposed as two new oxidation-reduction indicators. These indicators have transition potentials in the 0.9-volt range and are satisfactory for titrations of ferrous iron in sulfuric and perchloric acid media, using solutions from 0.1 to 0.001 N . Tetramethylbenzidine map be used as a redox indicator a t temperatures up to 100”C. LITERATURE CITED
(1) Beilstein, F., “Handbuch der organischen Chemie,” Vols. 12, 13, Berlin, Julius Springer, 1929, 1930. . 21, 1313 (1949). (2) Brandt, W. W., and Smith, G. F., ~ A LCHEhc., (3) Brennecke, E., and Bottger, W.C., “Newer Methods of Volu-
metric Analysis,” tr. by R. E. Oesper, New York, D. Van Sostrand Co., 1938. (4) Furman, N. H., and Evans, 0. AI., J . A m . Chem. SOC.,51, 1128 (1929).
( 5 ) Furman, K.H., and Wallace. J. H., Jr., Ibid.. 52, 2347 (1930). (6) Hamer, JT. J., Trans. Electrockem. Soc., 72, 45 (1937). (7) Knop, J., 2. anal. Chem., 85, 253 (1931). (8) Knop, J., and Kubelkova, O., 2. anal. Chern., 77, 125 (1929); 85,401 (1931). (9) Kolthoff, I. M., and Sandell, E. B., IXD.ENG.CHEM.,ANAL. ED., 2, 140 (1930). and Pattinson, S.,Ber., 14, 2161 (1881); 17, 118 (10) Michler, W., (1884). (11) Miller, B. F., and Van Slyke, D. D., J . Biol. Chern., 114, 583 (1936). (12) Rieman, W.,111, Neuss, J. D., and Naiman, B., “Quantitative Analysis,” New York, McGraw-Hill Book Co., 1981. (13) Smith, G. F., and Fly, W.H., ANAL.CHEM.,21, 1233 (1949). ENG.CHEM.,AN.%L. ED.. (14) Smith, G. F., and Richter, F. P., IND. 16, 580 (1944); “Phenanthroline and Substituted Phenan-
throline Indicators,” Columbus, Ohio, G. Frederick Smith Chemical Co., 1944. (15) Ullmann, F., and Dieterle, P.. Ber.. 37, 23 (1904). (16) Walden, G. H., Jr., Hammett, L. P., and Chapman, R. P., J . Am. Chem. SOC.,53, 3908 (1931). (17) Whitmoyer, R. B., IND.ENG.CHEM., AXAL.ED.,6 , 268 (1934). (18) Poe, J. H., and Boyd, G. R.. Jr., Ibid., 11,492 (1939).
RECEIVED January 27,
1951.
Modified Colorimetric Assay of Pteroylglutamic Acid RUTH ABBOTT KASELIS’, WLADIMIR LEIBMANN, WILLIAM SEAMAN, J. P. SICKELS’, E. I. STEARNS, AND J. T. WOODS American Cyanamid C o . , Calco Chemical Division, Bound Brook, X. J .
I
N ADDITION to biological methods, various chemical and physical methods have been reported in the literature for the determination of pteroylglutamic acid, N- [4-([(2-amino-4hydroxy-6-pteridyl)methyl]amino ) benzoyl]glutamic acid, either in purified preparations or in the presence of other substances. These include fluorometric (1, 14), polarographic ( 3 , 6, I I ) , and two colorimetric methods; one colorimetric method involves reduction with titanous ehloride ( 7 ) , and the other involves reduction with zinc in the presence of gelatin (8). In the colorimetric methods the amine formed by reduction is diazotized and Present address, 148 West Franklin St.. Bound Brook, N. J. Present address, Department of Chemistry, University of Miami, Coral Gables, Fla. 2
coupled to the Brstton and Marshall ( 8 ) reagent, l-naphthylethylenediamine, to produce a color which is measured spectrophotometrically. The color obtained by diazotizing p-aminobenzoic acid and coupling to the Bratton and Marshall reagent is used as the standard. The zinc reduction method is most generally used and is included in the “Pharmacopeia of the United States of America” (IS). The U.S.P. method has been of great value for the determination of pteroylglutamic acid in body fluids, natural products, capsules, elixirs, and other pharmaceutical preparations, but it is subject to several errors which limit its usefulness as a precise method of assay for the isolated material. These errors are caused by the use of zinc and acid for reduction, of paminoben-
