ANALYTICAL CHEMISTRY
'924
With seventy-five aliquots containing 0.25 to 4 mg. of dihydrostreptomycin a regression line was established using this procedure. The equation of the line was Y = 104.011847 X - 0.4236 where X is the difference in absorbance and Y the micrograms of dihydrostreptomycin in 1 ml. of the h a 1 solution. The correlation coefficient was +0.9998 and the standard error of estimation 10.97.
Y,
,
Comparison of Methods for Determination of Dihydrostreptomycin
Table I. Sample
Potency, Micrograms/hlg. Bioassay Ultraviolet
54550 53827 54281 55278 50207 50323 50419 50381 55626 5038 5040 5072 50317 50318 50327 55245 55477 51051 51825 50830 50256 5020 55936 50513
702 701 736 690 696 742 676 775 739 766 807 826 798 792 793 736 73 1 733 737 834 696 791 790 697
~ l t ~ ~ Ls ~ . i ~ l ~ t of Bioassay
733 685 697 751 700 741 668 760 746 772 798 793 811 770 741 752 755 712 696 764 742 770 756 710 Mean
WAVELENGTH
Figure 1. Spectrograms of Streptomycin Sulfate and Dihydrostreptomycin Heated in 0.25 N HzSO4 for 2 hours in boiling water Dihydrostreptomycin. 30 micrograms/ml. 0.06 NHkKh Dihydrostreptomycin. 90 micrograms/ml. 0.06 NHYSO~ Streptomycin sulfate. 90 microgramslml. 0.06 NHgSO4
104.4 97.7 94.7 108.8 100.6 99.9 98.8 98.1 100.9 100.9 98.9 96.1 101.6 97.2 93.4 102.2 103.3 97.1 94.4 91.6 106.6 97.3 95.7 101.9 99.5
tivrty of the method is less A reading at 235 mp is useful as N check on the identity of the compound. In Table I results by the ultraviolet assay are compared with those by the bioassay on twenty-four commercial samples of dihydrostreptomycin. There is no significant difference between the potencies obtained by the two methods, as the t value is 1.050. The fourth column of the table lists the ultraviolet determinations as per cent of the bioassay. The bioassay figures mere determined by Miss K. Fitzpatrick. LITERATURE CITED
Although one reading must be made a t 265 mg, there is a choice of wave lengths for the second reading; 380 mp was chosen a8 being well within the horizontal part of the curve and yet not too close to the upper limit of the ultraviolet range. Readings may be made a t 235 and 265 mp for the quantitative determination of dihydrostreptomycin, but the difference in absorbance is not as great as with the two wave lengths chosen and the sensi-
(1) Boxer, G. E., Jelinek, V. C., and Leghorn, P. M.,J . B i d . Chem., 169, 153 (1947). (2) Eisenman, W., and Bricker, C. E., ANAL.CHEM.,21, 1507 (1949). (3) Peck. R. L., Hoffhine, C. E., and Folkers, K., J . $m. Chem. soc.. 68, 1390 (1946). (4) Schenk, J. R., and Spielman, hf. A , , Ihid., 67, 2277 (1945). (5) Titus, E., and Fried, J., J . B i d . Chem., 174, 57 (1948).
RECEIVED September 15, 1950
Quantitative Test for Nornicotine LOUIS FEINSTEIN AND EDWARD T . MCCABE Bureau of Entomology and Plant Quarantine, U. S. Department of Agriculture, Beltscille, Md.
NEW quantitative test has been found and developed for the Nicotine fails to give any similar color even when taken in amounts of 80 times that of nornicotine. Anabasine develops a color only 80% as intense as nornicotine Fvhen taken in an amount 80 times that of the nornicot rnr Synthetic and natural nornicotines react quantitatively like cach other. I t 15 as observed that nornicotine produced a very intense violet color when added to an acetone solution of l,3-diketohydrindene if diipopropyl ketone was also present (1). The violet color, howwer, failed to appear with the use of a new batch of diisopropyl ltrtonc~. The old bottle of ketone had a cork stopper and an experirnc,nt was set up to extract a new cork with the neiv ketone. This cork extract caused the violet color to appear with nornicotine and the acetone solution of 1,3-diketohydrindene. Other experiments showed that tannic acid, gallic acid, or p-hydrouytmizoic acid could be used to make the new ketone reactive. The present method employs as reagents, acetone, diisopropyl bctone, p-hydroxybenzoic acid, and 1,3-diketohydrindene. Us-
A Nicotiana alkaloid, nornicotine.
40 42 44 46 40 50 52 54 56 58 60 62 64 66 6 8 7C FILTER NUMBER
Figure 1. Determination of Nornicotine
V O L U M E 23, N O . 6, J U N E 1 9 5 1
925
ing p-hydroxybenzoic acid and omitting the diisopropyl ketone allowed the violet color due to nornicotine to appear, but color development continupd even over 20 hours. Further experiments showed that diisopropyl ketone (new) added to the reagents acted as a reaction stabilizer and the violet color 1 hour after starting the reaction was stable for over a 1-hour period.
shake to mix, and place the flasks in the dark for 60 minutes. Determine color density of the solution with the colorimeter, using a No. 54 filter, and plot the readings against the concentration.
ANALYTICAL PROCEDURE
ACKNOWLEDGiMENT
Reagents. Acetone, A.C.S. Diisopropyl ketone, Eastman's P3244. p-Hydroxybenzoic acid, Eastman's 1520, 2% by weight/ volume in diisopropyl ketone. l,%Diketohydrindene, Eastman's P3565, 0.3% by weight/volume in diisopropyl ketone. Nornicotine. Preparation of Standard Curve. Dissolve 0.0400 gram of nornicotine in 500 ml. of acetone. Place 2, 3, 4, and 5 ml. of the resulting solution in glass-stoppered flasks and add acetone to bring to 5 ml. where necessary. Add to each flask 15 ml. of diisopropyl ketone followed by 2 ml. each of the p-hydroxybenzoic acid and 1,3-diketohydrindenc reagent solutions. Stopper the flasks,
The authors wish to thank -4bner Eisner, Eastern Regional Research Laboratory, Bureau of Agricultural and Industrial Chemistry, Agricultural Research Administration, United States Department of Agriculture, Phihdelphia, Pa., for the synthetic nornicotine.
Figure 1 shows that natural and synthetic nornicotine react alike. The same quantity of nicotine and anabasine gives very little more color than the reagent blank.
LITERATURE CITED
( I ) Feinsteiii, L., and McCabe, E. T., Science, 112, 534 (1950). RECEIVED July 29, 1950.
Contribution from the United States Department of Agriculture, Agricultural Research Administration, Bureau of Entomology a n d Plant Quarantine, Beltsville, Md. Report of a s t u d y under the Research a n d Alarketing Act of 1946.
Determination of Formal Oxidation Potentials of Ferric-Ferrous and Dichromate-Chromic Systems Selection of Substituted 1,lO-Phenanthroline Indicators for Determination of Iron ut Low Acidity G. FREDERICK SMITH, University of Illinois, Urbana, I l l .
'HE determination of iron in hydrochloric acid solution by dichromate oxidation, according to existing procedures, is accompanied by irksome inhibitions. Diphenylamine-type indicators, which are ordinarily employed, leave much room for improvement. Because the oxidation of iron by dichromate gives a green solution due to the chromic ion, 1,lO-phenanthroline-type indicators would be preferable. The formal potentials involved have not been previously studied systematicallr, particularly in the lower range. The problem is complicated by the low oxidation potential of the dichromate-chi omic ion system, particularly at low acid concentrations. The present work has for Its objectlve determination of the requisite formal oxidation potentials and selection of suitable 1,lOphenanthroline-type ferrous complex ions to serve as indicators in the determination of iron by dichromate oxidation. The procedures described involve the use of 0.1 to 1.5 F hydrochloric acid. Previous Investigations. The study of polysubstituted I ,lophenanthrolines for use as ouidation-reduction indicators in the form of their ferrous complex ions together with the determination of various pertinent physical constants has been made by Brandt and Smith (1). The determination of iron by dichromate oxidation in 2 F sulfuric or hydrochloric acid employing the 5,6dimethyl-l,l0-phenanthroline ferrous ion as indicator has been described by Smith and Brandt ( 3 ) . Phenanthroline-type ferrous complex ions for use as oxidation-reduction indicators are now known which have color transition potentials covering the range 0.84 to 1.27 volts. The lower potentials, 0.84 to 1.10 volts, are represented by the various methyl substituted types ( I ) , and the higher values, 1.10 to 1.28volts, by the chloro-, bromo-, and nitrosubstituted types described by Smith and Richter ( 5 ) . -4pplications in volumetric microdeterminations of iron, arsenic, calcium, and the oxalate ion have been described by Smith and Fritz ( 4 ) and by Salomon, Gabrio, and Smith ( 2 ) . Ferroin-type indicators have not previously been known and available for use in the titration of ferrous iron by dichromate in 0.1 to 0.25 F hydrochloric acid. The dichromate titration of iron using a series of familiar indicators has been studied by Stockdale (6).
DETERMINATION OF OXIDATION POTENTIALS
Solutions of known or determined concentration of ferric chloride and potassium dichromate of 0.1 .V strength are required. In addition, hydrochloric acid solutions of graded formalities of 0.1 to 4.0 with small increments of increase in strength between 0.1 and 1.5 F were to be prepared. Solutions of the indicators, 4,7-dimethyl- and 3,4,7,&tetramethyl-l,IO-phenanthroline ferrous complex ions, which were 0.01 Ar are prepared by reaction of weighed portions of the organic base with hydrated ferrous sulfate, addition of water to promote solution arid complex formation, and dilution to calculated final volume. The ferric chloride solutions were prepared as follows: One-tenth molecular weight of ferric chloric hexahydrate was dissolved in 8.245 ml. of reagent hydrochloric acid (specific gravity 1.19, 37.5% hydrochloric acid) and diluted wit'h ivater with stirring. The solution thus obtained was transferred to a 1000ml. graduated flask, diluted to the mark, and thoroughly mixed. The procedure was repeated with appropriate increase in added hydrochloric acid t,o prepare approximately 0.1 N ferric iron in 0.25, 0.50, 0.75, 1.0, 1.5, 2.0, 3.0, and 4.0 F hydrochloric- acid in addition to the 0.1 F hydrochloric acid solution described. Solutions of the same hydrochloric acid formality w r v prcpared to serve for volume dilutions. Solutions of potassium dichromate in 0.1 to 4.0 F hydrochloric acid were prepared as follows: Samples of pure potassium dichromate weighing 4.9035 grams (0.1 equivalent weight) were accurately weighed and transferred to 1000-ml. beakers, dissolved in water, and diluted to 800 to 900 ml. with water. With constant stirring the proper volume of reagent hydrochloric acid was then added (8.245 ml. for 0.1 F acidity, 82.45 ml. for 1 F acidity, etc.). Finally, the beaker contents were transferred quantitatively to a 1000-ml. graduated flask, diluted to the mark, and thoroughly mixed. Solutions of potassium dichromate above 1 F in hydrochloric acid form some free chlorine upon storage and should be made up just before use. The presence of a very low chlorine content will become evident through the attainment of high results for the oxidation potential of the chromate-chromic ion couple. For this reason the solutions are prepared precisely as described to eliminate the presence of free chlorine, and solutions more than a few hours old are discarded. The potential and potentiometric titration studies were evaluated using an assembly of a Leeds & Northrup student potentiometer and decade resistance box, Weston standard cell, and lamp and scale galvanometer. The electrode system was a saturated calomel reference electrode and a platinum mire indicator elec-
