dimethyiformamide (DMF) in 50ml. graduated centrifuge tubes. The tubes are heated in a boiling water bath for 30 minutes. T h e volume is brought t o 15 ml. with distilled water. The entire contents of the fat samples are extracted with 12 ml. of toluene and the toluene extract is discarded. For color development, 10 ml. of the aqueous DMF extract is added to 1 ml. each of 1.5% phenylhydrazine hydrochloride and 12N hydrochloric acid. This mixtiire is heated 25 minutes at 70" C., cooled, and extracted with 5 ml. of toluene. The absorbance of the toluene layer is determined with the spectrophotometer from 400 to 460 mp using pure toluene as a blank. Aqueous standards of nitrofuran containing 5, 10, and 25 pg. per 10 ml. are treated in exactly the same manner as described for f a t and muscle. A standard curve is constructed in the usual manner and
the control-corrected absorbance is used to calculate the nitrofuran content of the tissue. The difference curve must show a maximum of 440 mp t o prove the presence of the nitrofuran. Tissue Recovery Experiments. The accuracy and precision of this method were demonstrated by carrying out tissue recovery experiments. I n these experiments, known amounts of nitrofurazone or furazolidone were added to the tissue prior to extraction. Typical results of these experiments are presented in Table I. With fat and muscle the background absorbance of the control tissue was low enough (equivalent t o less than 1 p.p.m.) so that chromatographic separation of the phenylhydrazone derivative was not required.
LITERATURE CITED
(1) Belloff, G . B., Buzard, J. A., Roberts,
H. D. B., Poultry Sci. 37,223 (1958). (2) Bender, R. C., Paul, H. E., J. Biol. Chem. 191, 217 (1951). (3) Buzard, J. A., Ells, V. R., Paul, M. F., J. Assoc. O&. Agr. Chemists 39, 512 (1956). (4) Buiard, J. A., Vrablic, D. M., Paul, M. F., Antibiotics & Chemotherapy 6 , 702 (1956). (5) Maftin, J. E., Mich. State Univ. Vetennarian 19, 95-101, 119 (winter lQ5Q).
(6) Paul, M. F., Paul, H. E., Bender, R. C., Kopko, F., Harrington, C. M., Ells, V. R., Buzard, J. A., Antibiotics & Chemotherapy, in press.
(7) Proc. 2nd Natl. Symposium on Kitrofurans in Agriculture, Univ. of Georgia, Athens, Ga., Mar. 27-28, 1958.
RECEIVED for review February 23, 1960. Accepted June 10, 1960.
Photometric and Visual Titration of Certain Alkaloids in Glacial Acetic Acid Using Malachite Green as Indicator SAMUEL M. TUTHILL, ORLAND W. KOLLING', and KARL H. ROBERTS* Mallinckrodf Chemical Works, St. louis, Mo.
,Certain naturally occurring opium alkaloids, such as codeine, cryptopine, morphine, thebaine, narcotine, and papaverine, may b e assayed by titration with perchloric acid in glacial acetic acid solutions with malachite green as the indicator. The change of color at the end point is more distinct and more nearly coincidental with the potentiometric end point than is that of the often used crystal violet. The end point with malachite green may b e determined visually or spectrophotometrically. The visual procedure is more rapid and can be used for routine work, if desired. The method is particularly advantageous when applied to the weakly basic alkaloids such as papaverine and narcotine, since they cannot b e titrated satisfactorily in aqueous solution.
0
THE MANY METHODS utilizing titration in nonaqueous media which have been developed in recent years, the titration of weak bases with a solution of perchloric acid in glacial acetic acid is probably the most familiar. In such titrations, crystal violet is the indicator most commonly used for the colorimetric determination of the end 1 Present address, Department of Chemiqtry, Southwestern College, lVhbeld, F
Kan
Present address, !vlaGe:in c'o , L'leveland 15, Ohio. 1678
('hvmlc:t'
ANALYTICAL CHEMISTRY
point. The major difficulty in its use is the variety of color changes which it undergoes during a titration ( 5 ) . Seaman and Allen ( 7 ) concluded that a preconceived color change generally cannot be used in the titration of different substances; instead they stated that the correct color change for a given determination must he sciected by obserhlng the color of crystal violet a t the potentiometric end point. The photometric titration of weak acids in aqueous solutions with and without added indicators has been studied by Goddu and Hum(: ( 2 ) , and Reilley and Schwizer (6)have shoivn that certain bases, which have no absorbance in eithcr the acidic or basic form, may be titrated photomctrically in glacial acetic acid by adding a wenkcr base, which does absorb. During a study of the absorption spectra 0: crystal violet and malachite green in glacial acetic acid in the presence of varjjng quantities of perchloric acid, it was observed that as the mole ratio of perchioric acid to crystal violet increased, the atxorption mnsimum of crystal violet a t 5% nw decreased, and an nh.cor:)tion mnsini:im XppPared a t 6?0 m s . Tllc, 1sttc.r mzsimuin decreased in inti,r!sity, .in(! :t tliird z~xirr.cm~:Lt 440 i7.u ; 1 1 ( . 1 c : i d i:l intcnric tr
These observations are similar to those reported by Conant and Werner (1). The corresponding family of absorption spectra for malachite green shoxed that an absorption niaximum'&t 622 mp decreased as the ratio of perchloric acid to the indicator increased and that a maximum a t 450 inp increased at the same time. The sharp changes in absorption of these dyes with an increase in the quantity of added acid suggested their use as indicators for the photometric titration of organic bases in glacial acetic acid. Since 3Iallinckrodt Chcmical Works is engagrd in the commercial extraction of naturally occurring opium alkaloids, the possibility of assaying certain of these alkaloids by photometric titration, using the above-mentioned dyes as indicators was investigated. This paper describcs the development of a procedure for such titrations. Basic impurities, including other alkaloids, interfere in the procedure, since they are titrat,ed Kith the alkaloid being assayed.
RECOMMENDED PROCEDURE
90 W 0
z
$ 70 5 $50 t L
W
n
-
"t IO
Photometric Titration. Transfer about 0.3 gram of the dried alkaloid t o a 150-ml. beaker. Dissolve in 80 ml. of glacial acetic acid, with efficient stirring t o avoid a gummy, di5cultly soluble residue. Add from a buret 15.00 ml. of 0.05N perchloric acid solution and 1.0 ml. of malachite green indicator solution, stir well, and place the beaker in the cell compartment of the spectrophotometer. Set the spectrophotometer wave length at 620 m p , adjust the dark current, and set the instrument to read the desired a b s o r b ance. Continue the titration with the perchloric acid solution. Stir thoroughly after each addition, and read the absorbance, taking care to readjust the dark current. Correct the absorbance readings for dilution (3)and plot these data us. the milliliters of perchloric acid. Determine graphically (3) the end point of the titration and the volume of perchloric acid corresponding to it. Apply to the volume so obtained a temperature correction of 0.11% per degree difference between the temperature prevalent during the titration and the temperature during the standardnation of the perchloric acid (7). Visual Indicator Titration. Transfer about 0.3 gram of the dried alkaloid t o a 250-ml. Erlenmeyer flask and dissolve in 100 ml. of glacial acetic acid. Add 1.0 a l . of the malachite green indicator solution and titrate with the perchloric acid solution until the color changes from blue t o bluegreen. Carefully continue the titration dropwise until the blue-green color changes to green. Correct the volume used for the indicator blank and apply the temperature correction given above. Determine the indicator blank by titrating 100 ml. of glacial acetic acid containing 0.10 gram of dried sodium perchlorate to the green end point.
450
500
550 600650x)o
WAVE LENGTH, M p
Figure 1. Variation of absorption spectrum of malachite green with increasing acidity 5 X 10 ' M Malachite green Mole ratio, acid to Curve indicator 1 0 2 10 3 15 4 20 5 30
ances during the spectrophotometric titrations. This instrument was modified to permit the use of a 15Ckml. beaker as a titration vessel in a manner similar to that described bv Goddu and Hume (3). A Warren SDectracord was used to record a u t o m h a l l y the absorption spectra of malachite green and crystal violet in glacial acetic acid solutions containing varying amounts of perchloric acid. Standard perchloric acid solution, 0.05N perchloric acid in glacial acetic acid, was prepared by adding 8.6 ml. of 70% perchloric acid to lo00 ml. of glacial acetic acid in a %liter volumetric flask. The solution was mixed well, 50 ml. of acetic anhydride was added, and the mixture was then diluted to volume with glacial acetic acid. After standing 24 hours, the acid was s t a n d a r d i d potentiometrically against primary standard potassium acid phthalate as recommended by Seaman and Allen (7). Indicator solutions were 1 x 10-aM and were prepared by dissolving the required weight of the compound in 100 ml. of glacial acetic acid. All alkaloids used were obtained from the Narcotics Department, Mallinckrodt Chemical Works. The codeine, morphine, narcotine, and thebaine samples were from finished lots, the papaverine sample was taken from a semipure process intermediate, and the cryptopine sample was one which had been prepared a t an earlier date. The thebaine used in the experiments was dried to constant weight in a vacuum desiccator containing sulfuric acid. The other alkaloids were dried to constant weight a t l O 5 O C. Malachite green hydrochloride was obtained from Xational Aniline Division, Allied Chemical and Dye Corp. Crystal violet (gentian violet U.S.P.) was obtained from Fisher Scientific Co.
EXPERIMENTAL
To determine the effect of perchloric acid concentration on the absorption spectra of malachite green and crystal violet in glacial acetic acid, two series of solutions were prepared by adding varying amounts of 1 X 10-3M perchloric acid in glacial acetic acid to suitable aliquots of each indicator solution and diluting to volume. The indicator concentration of solutions in both series was 5 X 10-M. Absorption spectra for solutions containing malachite green and for those containing crystal violet are shown in Figures 1 and 2, respectively. For the titration of the alkaloids, stock solutions of each alkaloid were prepared by dissolving an accurately weighed quantity of the alkaloid in glacial acetic acid, quantitatively transferring the solution to a 2Wmi. volumetric flask, and diluting to volume with more acetic acid. The weight used was such that 25-r-1. aliquots contained about 0.3 gram. The assay value of each alkaloid w a s determined by p i p e t ting 25-mi. aliquots into a 150-mi.
8 z
90
s 70 82 50
" I L IO
450
500
550 600650
00 W A M LENGTH, Mp
Figure 2. Variation of absorption spectrum of crystal violet with increasing acidity 5 X 10-M cryrtol vbld Mole ratlo, add to b
e
indkata 0 10
1
2
12
3 4 5 6
15 30 60
beaker and titrating potentiometrically with a solution of 0.05N perchloric acid in glacial acetic acid. Additional aliquots of each solution were also titrated according to the photometric and visual titration procedures described above. To approximate the ionic strength prevailing in solutions of the alkaloids which have been titrated to the end point, the indicator blank for the visual titration was determined in a solution containing 0.10 gram of sodium perchlorate in 100 ml. of glacial acetic acid. The blank was about 0.15 ml. of 0.05N perchloric acid in glacial acetic acid. When the pipet was used to deliver aliquots of the glacial acetic acid solution of the alkaloids, it was observed that the manner of drainsge differed from that encountered with aqueous solutions. Calibration of the pipet with glacial acetic acid showed that it delivered a volume 0.03 ml. greater than when water was used. This correction was applied in calculating the results of the titrations. With a buret, however, the difference between the delivered volumes of water and glacial acetic acid was negligible. DISCUSSION
The effect of added perchloric acid on the absorption spectrum of a dilute solution of malachite green in glacial acetic acid is shown in Figure 1. The principal absorption maximum for this dye occurs at 622 mp and decreases sharply with added acid. The correspondicg set of curves for crystal violet is shown in Figure 2. The latter dye exhibits a more complex spectrum. VOL
32.
NO. 12, NOVEMBER 1960
1679
Initially there is an absorption maximum at 588 my which decreases in intensity as the acid-base ratio increases. I n addition, the absorption at 630 mp increases from its initial value until a distinct masimum occurs and then decreases sharply as the acid-base FZ&O increases further. The requirements involved in the selection of indicators for photometric titrations differ from those pertaining to the visual titration of acids and bases, although in both cases the indicator is a significantly weaker acid or base than the substance being titrated. For the visual titration, the indicator should be chosen so that it changes cblor as nearly at the equivalence point as possible. The color change then takes place during the conversion of 10 to 90% of the indicator from either the acidic or basic form to the conjugate species. On the other hand, an indicator selected for a photometric titration should, ideally, begin to change only when the equivalence point is reached. I n this latter case, the photometric titration provides automatic elimination of the indicator blank, which is often encountered in visual titrationa. In their paper on the photometric titration of weak acids, Goddu and Hume ( 2 ) discuss the titration of mistures of weak acids. I n the potentiometric titration of such mixtures, the location of the inflection point corresponding to the first equivalence point is difficult when the ratio of the ionization constants is about 100, and the ratio must be greater than several thousand to obtain satisfactory results. These authors demonstrated experimentally that photometric titrations of acids are possible when the ratio of the constants is only 20. Such a titration leads to an ~~
Table 1. Assay Values of Certain Alkaloids Determined by Potentiometric, Photometric, and Visual Titrations
Potentiometric 100.1 100.3 10.1
.4lkaloid Codeine Cryptopine Morphine
Narcotine
Papaverine
99.9
100.0 100.0 100.0 100.1 100.0
Photometric 100 1 100 1 100.1
... ... 99.7 09.7
100.1 100.1 100.2
99.8 99.8 99.9 99.8 99.9 ...
100.0
100.2 ...
...
100.5 100.5
100.4 100.4 100.5
100.2 100 2
99.9
...
1680 *
99.9
100.1
99.9 99.8 99.9 99.9 100 5
Thebaine
Visual 100.1
99.9 99.9
99.8 99 8
ANALYTICAL CHEMISTRY
99 6
100.2
10.3
100.4 99 7 X.7
18.0
17.0
19.0
20.0
Y U A P I I T E slEEN INDICA101
MIU
0.20
mafr
INDICATOI
I
I
I
I
18.0
19.0
20.0
21.0
ML. OF PERCHLORIC ACID
Figure 3.
Photometric titration curves
error in the first end point of 2 to 3%. However, if the indicator has a further color change beyond the first break in the photometric titration curve, the desired acid still may be accurately determined by adding a known amount of the indicator and determining the total acidity from the second break in the curve. The quantity of the desired acid may then be determined by difference. I n principle, the same procedure may be applied to the titration of weak bases. Curves A and B of Figure 3 were obtained during the photometric titration of aliquots of a solution of morphine with malachite green and crystal violet, respectively, 3,s the indicators. Absorbances were read a t 620 mp in both cases. Thc results based on the first break of curve B were almost 1% low, whereas the results based on correcting the total basicity value given by the second break for the known number of equivalents of indicator added agree within 3 parts per thousand with the potentiometric results and the results obtained with malachite green. The low results given by the first break in the crystal violet curves are similar to those obtained by Goddu and Hume in the photometric titration of mixtures of weak acids. Such results with crystal violet probably occur because the first basicity constant of this indicator in glacial acid is not sufficiently different from that of morphine to permit the use of the first break in the photometric titration curve. llalachite green, on the other hand, gives only one break, as shown in curve A , in the photometric titration curve. The results based on this break agree well with those obtained from a potentiometric rnd point. Therefore, since a blank correction is not required, malachite green posscsscs a marked advantage over crystal viokt as a photomctric indicator, and, therefore, is specified in the Ficcornmc~ndrd Procedure.
The results summarized in Table I show the applicability of the procedure using malachite green as the indicator. All of the results for each alkaloid are for the same sample. The photometric titration procedure cannot be used for cryptopine because a white suspension forms in the vicinity of the end point. The slightly high result obtained by titration for papaverine is attributed to traces of alkali remaining in this semipure sample. The presence of a base of low equivalent weight would cause high results, since the total basicity of the sample is titrated. The visual titration, using malachite green a6 the indicator, is considered to give an end point superior to that obtained with crystal violet. Even so, the visual determination of the end point is subjective rather than objective. For this reason, although the visual method is more rapid, the photometric method provides a sounder approach to the problem of determining the end point, and, therefore, is preferred. The alkaloids included in the present study are relatively weak bases. For example, codeine, morphine, and thebaine have values of p& (4) of 6.05, 6.13, and 6.05, respectively, while those of narcotine and papaverine are 7.82 and 8.10. Codeine, morphine, and thebaine commonly are titrated in aqueous medium, using methyl red as the indicator, or potentiometrically with a pH meter. However, these bases are about as weak as can be titrated in aqueous solution to yield reasonably satisfactory results. Thus the nonaqueous titration, while not essential, is a t lrast on a theoretical basis, an improvemcnt ovcr the aqueous methods. The wcaker bases, narcotiric and papaverine, cannot be titrated satisfactorily in aqueous solution and, therefore, for the assay of these alkaloids, the proposed method of nonaqueous titration, utilizing malachite green as the in-
dicator, possesses the advantages enumerated earlier.
addition, they thank Lester McKenzie for making available the alkaloids which wcre used.
ACKNOWLEDGMENT
LITERA1URE CITED
The authors are indebted to Thomas C. Boersig, Joseph R. Simmler, and (1) Conant, J. B., Werner, T. H., J. Am. Chem. SOC. 52, 4436 (1930). Howard W. Ziegler of the Department of chemical Control, ~ ~ l l i ~ ( 2~) Goddu, k ~ R. ~ F., d ~Hume, D. N., ANAL. CHEM.26, 1679 (1954). (3) &d., p. 1740. Chemical Works, for assisting with the (4) Lange, N. A,, "Handbook of Chemisexperimental work of this paper. I n
try," 9th ed., pp. 1203-4, Handbook Publisbm, 1nc.t Sandusky, Ohio, 19%. (5) Nadeau, G. F., Branchen, L. E., J . Am. Chem. SOC.57. 1363 (1935). (6) Fkilley, C. N., Schweizer, B., ANAL. CHEM.26, 1124 (1954). (7) Seaman, W., Allen, E., [bid., 23, 592 (1951). RECEIVED for review April 22, 1960. Accepted August 15, 1960. Presented in part, Division of Anakytical Chemistry, 130th Meeting, ACS, Atlantic City, N. J., September 19, 1956.
Primary Standards for Titration af Anionic Detergents Using Quaternary Ammonium Halides BENJAMIN VELDHUIS General Chemical Division, Allied Chemical Corp., Morrisfown, N. 1. The alkaline salts of polyhalogenated benzenesulfonic acids have been found suitable for direct standardization of quaternary ammonium halides. These stable sulfonates can be prepared easily.
D
with quaternary ammonium halides is the most widely applied method for assaying anionic surfactants ( 1 , 3 ) . I n spite of gcneral use, present reagcnts for primary standardization of the quaternaries are some\\ hat iriconvenient. Sodium alkarylsulfonate ( I ) is of inexact chemical composition and purity; the corresponding sulfohic acid (3) is hygroscopic. Anotlicr rccommended standard ( b ) , a 0.25% aqueous solution of Aerosol OT (American Cyanamid Co.), the clioctyl ester of sodium sulfosuccinic acid, was found to undergo slow deterioration shortly after preparation. The alkaline salts of polyhalogenated benzenesulfonic acids appear free of these objections. They are relatively easy to prepare and purify, and appear indefinitely stable both as solids and in solution. I n addition, they are pure compounds of known chemical composition. One of these materials, potassium 2, 4, 5-trichlorobenzenesulfonate, is being made commercially available (Baker and Adamson Division, Allied Chemical Corp.). IRECT TITEATION
EXPERIMENTAL
Preparation of Reagents. Exactly 363 grams of lJ2,4-trichlorobenzene, practical grade, were placed in a threenecked laboratory reaction flask equipped with dropping funnel, mcchanical agitator. thcrmometer, and piovision for extcrnal cooling arid Iicxting Then 168 grmis of stabilizrri
liquid sulfur trioxide, Sulfan (Baker and Adamson Division, Allied Chemical Corp.), were added dropwise over 18 minutes. The temperature rose spontaneously to 105" C. and was held there by slight cooling. After addition of the trioxide, the dropping funnel was replaced by an azeotrope trap; 500 ml. of water were added and the mixture was refluxed with stirring for 2 hours to distill 23 ml. of unreacted starting material. The warm solution was then filtered to give 1337 grams of acid solution. A 668-gram portion of the acid solution was neutralized with 17% aqueous potassium hydroxide, cooled to room temperature, filtered, and washed with cold water to give the crude potassium sulfonate. The 403 grams of wet cake --ere redissolved in 3.5 liters of distilled water a t 95" C.; the solution was filtered hot to remove suspended solids, cooled, and refiltered. The product was airdried to constant weight, 157 grams. The other sulfonates w r e prepared similarly from the corresponding halogenated benzenes. Analytical and
Table
I.
Analytical Data for Standard Sulfonates
Hydration, >foles
S o . Benzenesulfonate 1 4-Bromo
3
2
2,4,5-Trichloro 2,4,5-Trichloroc
4
2-Bromn-5-
.5
6
chloro 2,5-Dibromo 2,4,5-Trihromo
titration data are given in Table I. The degree of hydration shown for sulfonates 1,2, 5, and 6 are in agreement with the literature values (4); salts 3 and 4 are new, although the sulfonic acids are known (4). None of these sulfonates has been prepared previously using liquid sulfur trioxide. The solubilities of salts 2, 4, and 6 in water at room temperature are approximately 0.9, 3.0, and 0.8 grams per 100 grams of solution, respectively. At 95" C., the corresponding figures are a p proximately 6, 12, and 7 grams per 100 grams of solution. The theoretical molecular weights in Table I were arrived a t by determining the water content and assuming the sulfonates to be pure as indicated by halogen and sulfur analysis. Originally chloride determinations were used to determine the Hyamine 1622 concentration. Since this method was not completely satisfactory, salt 2 became the usual standard. The water was determined by the Karl Fischer titration; the halogen by
Water Sone None 15
Molecular Wt. Theo. Founda 275 33Ob 300
310
Halogen Found
300 312 326
Tho. 29.0 35.4 34.3 26.4 43 0
1 0
328
1 0 1 0
372
374
451
149
53 1
Sulfur
Theo.
Found
28.9
11.6
34.6
10.3
11 3 10.7 0.7
42 6 .5Z 1
8 6 7 1
8 1 6.8
35.8 26 6
10.7
By titration with rtmdardized Hyamine 1622 (diisohutyl phenoxy cttiosy ethyl clirwttiyl bt~iizylanirnonirim chloride, monohydrate). End point w r y iiidistirict. Sotliiim d t used; all others were potassium ssirs.
VOL 32, NO. 12, NOVEMBER 1960
1681
