were placed in five individual atmospheres containing different known concentrations of ammonia. The manifold of the sampling system was air purged between each ammonia sample. Figure 4 is a chart record of absorbance readings for the five ammonia samples alternating with base line readings for the air purges. The ammonia concentration of each sample is proportional to the absorbance value indicated by the lower tip of the heavy pen marking on the chart. I n Figure 4, the sets of readings on the left ( A , B ) represent duplicate analyses of the five atmos. pheres of ammonia. The sequence of reading is the same for both groups. The degree of reproducibility of analyses is shown by comparison of the two readings for any of the five samples. The three sets of readings on the right (C, D,E ) are for the same atmospheres as the 2 sets on the left, but the sequence of individual readings was changed by manipulation of the connections at the terminal strip of the solenoid valve unit (Figure 3). The repetitive readings of the same atmosphere are in good agreement, irrespective of the sequence in which the samples are taken. T o evaluate the effects of drastic changes in ammonia concentration on the accuracy of the measurement, a 25foot length of epoxy-lined tubing was introduced into an ammonia-air atmosphere and the mercury pump energized. The recorded absorbance of the first sample drawn into the cell represented 70% of the true ammonia
A
B
C
D
E
corded concentrations were 92 t o 95% of the actual concentrations. The automatic sampling system described has operated satisfactorily for many months. Since the sample volume is only 0.25% of the atmosphere of a carton of oranges, it has been possible t o establish ammonia concentration gradients resulting from localized generation of ammonia within a carton of fruit. A complete record of application of this measuring system t o a biological problem is published elsewhere (1). LITERATURE CITED
(1) Eckert, J. W., Kolbesen, M. J., Garber, M. J., Phytopathology 53, 140 (1963). (2) Gunther, F. A., Barkley, J. H.,
Figure 4. Chart record showing effect of changes in sampling sequence
Kolbezen, M. J., Blinn, R. C., Staggs, E. A., ANAL.CHEM.28, 1985 (1956). (3) Gunther, F. A,, Blinn, R. C., Barkley, J. H.. Kolbezen. M. J.. Staggs. E. A.. J. Agr. Food Ch&. 7,489.(15g9). (4) Gunther, F. A,, Blinn, R. C., Kolbezen, M. J., Conkin, R. A., Wilson, C. W., Zbid., 7, 496 (1959). (5) Gunther, F. A,, Blinn, R. C., Kplbezen, M. J., Wilson, C. W., Conkin, R. A.,ANAL.CHEM.30, 1089 (1958). (6) Roistacher, C. K.,Eaks, I. L., Klotz, L. J.. Plant Disease Revtr. 39. 202 '
concentration in the atmosphere or 93% if corrected for the volume of the tubing. However, in our studies on the generation and distribution of ammonia gas in citrus fruit containers, the maximum length of tubing used was 4 feet and the concentration at any given sampling position changed slowly over a period of hours. Under these conditions, therecorded ammonia concentrations were about 90% of the true value for the first cycle, and in experiments where sampling was continued at 2hour intervals for many days, the re-
(1955 j. (7) Roistacher, C. N., Klotz, L. J.,
Kolbezen, M. J., Staggs, E. A., Zbid.,
42, 1112 (1958). (8) Tomkins, R. G., Trout, S. A., J . Pomol. Hort. Sci. 9, 257 (1931). RECEIVEDfor review July 25, 1963. Accepted November 26, 1963. Work supported in part by grants from Sunkist Growers, Inc., Pure Gold, Inc., and
Western Fruit Growers' Sales Company.
Spectrophotometric Determination of Ethyl-, Diethyl-, and Triethylamine in Aqueous Solution GEORGE DAHLGREN Department o f Chemistry, University o f Alaska, College, Alaska
b Ethylamine, diethylamine, and triethylamine can be determined in aqueous solution in the parts per billion concentration range. Several different procedures are described for the determination of amines in the presence or absence of ammonia and amides. In general, the procedure involves the reaction of the amine with hypochlorite to form the chloro derivative, followed by the destruction of the excess hypochlorite b y nitrite (the chloramine i s unaffected) and the subsequent color development of the sample b y the reaction of the chloramine with starch-potassium iodide (SKI) reagent. The absorbance of the resulting blue solution is read at 540 mp. At a pH of 5.2 (acetate buffer) amines, amides, and
596
0
ANALYTICAL CHEMISTRY
ammonia are detected; at a pH of 6.6 (phosphate buffer), amines and some amides are detected, while ammonia in low concentration is not; a t a pH of 8.1 (bicarbonate buffer), amines are detected, but ammonia and most of the amides tested are not. The presence of bromide ion in the pH 8.1 procedure can be used in some cases to render amides nondetectable while increasing the sensitivity for ethylamine, and not affecting the sensitivity for diethylamine. Triethylamine i s not detected in the presence of excess bromide ion. The procedure at pH 8.1 has been used to follow the kinetics of the hydrolysis of N-ethyl- and N,N-diethylmaleamic acid, where the products of the reaction include ethylamine and diethylamine.
A
have been determined by reaction with Xessler's reagent (il), color-developing reagents (i6), nitrite ion (0), other organic reagents (2, 6, IS), by polarography (19), and by many other methods found in the literature. The sodium phenate- sodium hgpochlorite method for ammonia (1) was used on ethylamine with poor results. The usual blue indophenol color developed but only at very high concentrations of the amine. The method appears to be less sensitive for amines than ammonia by a factor of one thousand. XesslerJsreagent also gave a weak positive color test but only a t high amine concentrations. The tanninsilver nitrate test for ammonia (12) was tried on ethylamine using tannin-silver nitrate and gallic acid-silver nitrate. hm-m
Both methods yielded positive results, but the results were far from quantitative since the amine probably acts only as a catalyst by forming the silveramine complex in the reduction of the silver ion to free silver. Water blanks yielded the same result on standing. Kone of the procedures investigated appear to offer the convenience and sensitivity of the method described by Zitomer and Lamber , (18) for the spectrophotometric dete:mination of ammonia as trichloramine. The methods described here for the determination of amines in the presence of ammonia and amides differ from the trichloramine procedure in p H and in the starchiodide reagent. The trichloramine method can be used on amines with good results a t a pH of 5.2 if ammonia and amides are nct present in the sample. At p H 6.6 ammonia in low concentration is not detected but some amides do react. Ai, p H 8.1 ammonia and most of the amides are not detected. The general scheme involves the reaction of the ,imines with hypochlorite to chloro derivatives, destruction of the excess hypochlorite by nitrite ion (the chloro derivatiws are not attacked under these conditions), and oxidation of iodine by the chloro derivatives in the presence of starch. EXPERIMENTAL
Reagents. No attempt was made to use ammonia-free water in any of the reagents or procedures. Reagent grade chemicals were used throughout unless otherwise stated. AMINES. Stock solutions of ethylamine, diethylamine, and triethylamine in water were prepared from Eastman White Label amine!;. Amine samples 1 to 40 p M were prepared from the stock solutions. NITRITE. Sodium nitrite, 0,5y0 (18). pH 5.2 BUFFER. ,3odium acetate trihydrate (136 grams) and 21.2 ml. of glacial acetic acid per liter of solution
arsenic trioxide containing excess sodium hydrogen carbonate and SKI solution (1 ml. S K I per 25 ml. arsenic trioxide solution). STANDARDIZATION OF AMINES. The stock solutions (approximately 0.336) were standardized by conductometric titration with hydrochloric acid. HYPOCHLORITE AND IODINE CALIBRATION CURVES. To 25-ml. samples of hypochlorite or iodine were added 1.00 ml. of glacial acetic acid and 1.00 ml. of SKI solution. After 10 minutes, the absorbances were read at 570 mp on a Beckman DU Spectrophotometer. AMINECALIBRATION CURVE,pH 5.2. To 25-ml. samples of amine were added 1.00 ml. of p H 5.2 buffer and 1.00 ml. of hypochlorite stock solution. After 4 minutes for ethylamine, and after 10 minutes for diethylamine (triethylamine cannot be analyzed by this procedure), 1.00 ml. of nitrite solution was added. One minute later 1.00 ml. of S K I solution was added, and the absorbances a t 540 mp were read after 15 minutes development. p H 6.6. Same as above replacing the 5.2 buffer solution with the 6.6 buffer solution. pH 8.1. To 25-ml. samples of amine were added 0.5 to 1.0 gram of solid sodium hydrogen carbonate and 1.00 ml. of hypochlorite stock solution. After 1 minute for ethylamine and diethylamine, and after 1 minute per pmole triethylamine per liter, mas added 1.00 ml. of nitrite solution, After 1 minute, 1.00 ml. of SKI solution was added and the absorbances were read a t 540 mp after 3 minutes. RESULTS AND DISCUSSION
The molar absorptivities obtained in this study are given in Table I. The values for hypochlorite and iodine are given a t 570 me instead of 540 mp since it was found that molar absorptivities a t 570 mM for solutions below p H 4 and the values a t 540 mp for solutions above pH 4 were identical. The SKI solution, when protected (18). from the light and stored a t IO" C. pH 6.6 BUFFER. Potassium dihywhen not in use, remained colorless for drogen phosphate (25 grams) per 100 several weeks and showed no signs of ml. of solution. To this solution was microbe attack. It has the advantage added concentrated sodium hydroxide over the cadmium iodide-linear starch to a pH of 6.05. Buffer solution (1 ml.) described by Lambert (10) in that it in 25 ml. of sample produces a solution can be used a t higher pH values, of p H 6.60 f 0.02 throughout the addition of all reagents. However, substitution of the ordinary STARCH-POTASSIUIII IODIDE (SKI). soluble starch for the linear-starch Fisher soluble starch (7.5 grams) in 1 fraction decreases the molar absorptiviliter of boiling water. After 15minutes, ties by as much as 40%, depending upon 10 grams of potacsium iodide were the wavelength used. added slowly and the boiling was conWhen titrating a standard solution of tinued an additiona five minutes. arsenic trioxide with 0.001M hypoHYPOCHLORITE. Xonite, 2.5 ml. per chlorite, an end point of less than 0.05 100 ml. of solution (18). Samples 1 t o 40 p X were prepared from the stock ml. excess hypochlorite in approxisolution. mately 50 ml. of total solution was Procedure. For comparison pureasily detected. The results obtained poses, calibration curves were run with this standardization procedure with iodine and hypochlorite solutions. were identical to those obtained using STAKDARDIZATIOK OF HYPOCHLORITE. the usual hypochlorite procedures (8). The stock solution hypochlorite was used to titrate a standard solution of The low results described elsewhere clf
Table 1. Molar Absorptivities at 540 rnp. T = 25 2 " C.
Compound Ethylamine
Procedure pH 5.2 6.6 6 . 6 (Br-) 8.1
8 . 1 (Br-)
Diethylamine 5 . 2 6.6 6.6 (Br-) 8.1 8 . 1 (Br-)
Triethylamine 8 . 1 Ammonia 5.2 Hypochlorite 2 . 7 Iodine
3.2
Molar
absorptivity 31,000 30,800 30,900 26,400 44,600 20,800 22,900 20,500 22 200 22;500 14,000 30,500 20,500 (570 mp) 19,600 (570 mp) I
(7') for a similar but acid procedure were not observed. I n the pH 5.2 procedure, the waterreagent blanks were light blue due to the residual ammonia in the distilled water used. In the pH 6.6 and 8.1 procedures, the blanks are colorless. The times of reaction are not critical. In general, however, primary amines react faster with hypochlorite than secondary amines, which in turn react faster than tertiary amines. Some sample fading was observed in the pH 8.1 procedure after 5 minutes of development. Interferences. Ammonia and amides interfere with amine analysis a t p H 5.2. Triethylamine does not react a t a n appreciable rate a t this p H and therefore it is not conveniently determined. Acetamide and pyridine do not interfere in the pH 6.6 procedure. Ammonia interferes in this procedure in concentration greater than about 70 p M . Once again a t pH 6.6, the reaction of triethylamine is so very slow that this procedure offers no particular advantage over the pH 8.1 procedure. Jlaleamic acids and aniline and their N-ethyl and N,N-diethyl derivatives give varying degree3 of interference. The pH 8.1 procedure is relatively free from interference by ammonia, acetamide, N-substituted maleamic acids, and ,V-substituted anilines. Maleamic acid and aniline give slight positive interferences. I n all cases the temperature of the samples was the same as that used in constructing the calibration curves. Higher temperatures cause a decrease in absorbance as observed in the ammonia determination (18). Effect of Bromide Ion. Throughout the following discussion, chloramine and bromamine are used to designate the chlorine and bromine derivatives of the amines and do not VOL. 36, NO. 3, MARCH 1964
597
necessarily imply t h e number of halogen atoms per molecule. T h e number of chlorine and bromine atoms per molecule of amine will be discussed in a later section. The presence of bromide ion in the hypochlorite had no effect on the molar absorptivity for hypochlorite. The presence of excess bromide ion in the p H 5.2 procedure causes a decrease in the absorbance. This decrease is due to the destruction by nitrite ion of the bromamines formed by the reaction of the amine with hypobromite. The bromamine obtained from ethylamine is reduced rapidly by nitrite ion, while that obtained from diethylamine reacts more slowly. Triethylamine does not react with hypobromite under these conditions. Similar bromide ion interference was observed in the analysis of ammonia by this procedure (18). Bromide ion in the p H 6.6 procedure causes a slight decrease in the absorbance of ethylamine and diethylamine when analyzed by this procedure. The molar absorptivities obtained with excess bromide ion present are recorded in Table I. Triethylamine is not detected by this procedure with excess bromide ion present due to the reduction of the bromamine by nitrite ion. The presence of bromide ion in the p H 8.1 procedure causes an increase in the absorbance for ethylamine, no change in the absorbance for diethylamine, and a decrease in the absorbance for triethylamine. At sufficiently high concentrations of bromide ion (equal to the hypochlorite ion concentration), triethylamine is not detected. I n the presence of bromide ion, the blanks obtained in the p H 8.1 procedure are light blue. This color is possiblv due to the presence of residual ammonia in the distilled water. Blanks prepared from ammonia-free water show only a very slight trace of color probably due to the ammonia in the reagents. This color is not a stoichiometric representation of ammonia since the addition of ammonia to the water causes an increase in the absorbance of the sample only up to a point ( l O - 5 - V ) and then the absorbance remains constant a t a relatively low value (0.07 at 540 mp). Ethylamine reacts very rapidly in the presence of excess bromide in the p H 8.1 procedure to yield a molar absorptivity 1.5 tinies that obtained without bromide present. The molar absorptivity of diethylamine under the same conditions is the same in the presence or absence of bromide ion. Triethylamine is not detected in the presence of excess bromide due to the reduction of the bromamine by nitrite ion. The molar absorptivities for ethylamine and diethylamine at 540 mw, the absorbance maximum, obtained with excess bromide present, are recorded in Table I. 598
ANALYTICAL CHEMISTRY
L 0.55
1
I
0.5 0
w
u z 4
m
0.30
1
0
Hypochlorite Concentration
1.53 x 1 6 4 M
Ethylamine Concentrotion
1.28 x I6'M
8 12 16 20 M O L A R I T Y OF BROMIDE I O N x 10's
4
24
Figure 1. Determination of bromide lion concentration effect on ethylamine analysis at pH 8.1
Acetamide, aniline, and N-substituted maleamic acids interfere with the procedure at p H 8.1 when bromide ion is present. Maleamic acid, pyridine, and N-substituted anilines do not interfere. In the case of ethylamine, the maximum effect of bromide ion was found when the concentrations of bromide and hypochlorite are about equal. A plot of the absorbance at 540 m p us. the bromide ion concentration for a solution 12.8 p M in ethylamine is given in Figure 1. The plateau is found a t approximately 1.5 X l O - 4 X bromide ion. The concentration of hypochlorite in the solution is 1.53 X 10-4V. Stoichiometry. Zitomer and Lambert (18) report a value of 5.94 for the ratio of the absorbance for ammonia to that for hypochlorite on a weight basis. This value leads to a molar absorptivity ratio of 1.42 compared to the theoretical 1.5 obtained from their equations for the formation of iodine from ammonia. 3H+ + NHa
+ 3-0C1+
+ 3H20 312 + 6C1- + Nz KCla
2sc13
+ 61- +
d comparison of the molar absorptivities for ammonia and hypochlorite obtained in this work (30,500 and 20,500, respectively) give a ratio of 1.49 in excellent agreement with the theoretical. Audrieth and coworkers (4. 16) have used the reaction of amines with hypochlorite as a means of preparation of substituted hydrazines. Their proposed reactions involve chloramine8 as intermediates. The following reaction accounts for the molar absorptivity of ethylamine observed in this work:
2RNClz
+ 613Iz
+ R z N ~ H+~4C1-
which suggests that hydrazine itself is an intermediate in the reaction of ammonia with hypochlorite via: 4H+
+ 2sc13 + 101- 4
NzHd
+ 212 +
512
+ NzH4 + 6C1-
41-
+ NO + 4H+
When added, these reactions give the equation presented by Zitomer and Lambert (18). ,411 alternate scheme based on ionic + species of the type RX;C13 can also be imagined. N'eil and AIorris (17) postulate this species as an intermediate in the reaction of tertiary amines with hypochlorite. However, there appears to be no advantage in proposing this ionic mode of reaction except in the case of triethylamine where the species Et&('% is considered. Support for the triethylchlorammonium ion can be found by comparison of the reaction of trieth>.lamine with hydrogen peroxide to form the amine oxide (14). HOWever, even in the case of triethylamine it is possible to write the reaction of triethylamine and hypochlorite to yield diethylchloramine and acetaldehyde as suggested by Ellis and Soper (6). For simplicity, the following reactions are offered to account for the observed molar absorptivities: Ethylamine EtNH2 2-0C1
+
--f
EtNClz
+ 2-OH
Diethylamine EtzNH
+ -OCl+ Et&Cl+
EkNCl
-OH
+ H20 + 21- + + EhKH + C1- + -OH 12
Triethylamine
+
EtsN HOC1 Et&Cl+ 21-
+
+EtaNClf + -OH +I2 + Et3X + C1-
The excew hypochlorite is destroyed by reaction with nitrite ion: -0C1
+ NOz-
+
i C1Nos-
and the color development reactions may be represented by the reaction:
I-
+ Ip + SKI + 13---EIKI complex (blue)
The ratio of molar absorptivities for ethylamine, diethyls mine, and triethylamine predicted from the above reaction scheme is 1,5:1.0:1.0, and the observed ratio is 31,000:22,500: 14,000 or 1.4:1.0:0.62. The low value for triethylamine and for ethylamine at p H 8 is probably due to incomplete reaction with hypochlorite because of some prior equilibrium. The very high molar absorptivity for ethylamine at p H 8 in the presence of bromide ion may be accounted for by postulating the formation of the ionic species EtNBr3+ whose molar absorptivity should be 1.5 times that for EtNCI?. Table I gives a value of 44,600:31.000 or 1.44 for this ratio, which is considered good agreement. At pH 6.6, apparently some equilibrium prevents the formaticn of the ethyltribromammonium ion. The fact that diethylamine does not form a similar
complex ion with bromine a t p H 8 may be due to the steric effect offered by the two ethyl groups and two large bromine atoms in the region of the nitrogen. Ammonia is not readily detected in the procedures a t p H 6.6 and 8.1 due to the destruction of the chloramines by hypochlorite and not due to their reduction by nitrite ion. Evidence for this view is found when the hypochlorite reaction time is decreased to 15 seconds or less and the nitrite reaction is held a t 1 minute. The samples determined under these conditions showed strong positive tests, indicating the presence of some chloramine, while blanks under the same conditions were colorless. Ammonia is probably oxidized to nitrogen at the longer hypochlorite reaction times. The procedures described here are easily carried out and the results are reproducible. I n several determinations on each amine with different technicians preparing their own reagents and carrying out their own procedures, the molar absorptivities were reproducible to within 1% and an average error on amine analysis of less than 2y0 was obtained. The method has been used very successfully in the study of the kinetics of the hydrolysis of N-substituted maleamic acids (3) in which ethylamine and diethylamine were determined in the presence of their respective maleic acid amides. ACKNOWLEDGMENT
The author gratefully acknowledges the technical assistance of Nancy L. Simmerman and the valuable comments of Fred Zitomer.
LITERATURE CITED
(1) Barnes, H., “Apparatus and Methods of Oceanography,” p. 140, Interscience,
New York, 1959. (2) Clark, S. J., Morgan, D. J., Mikrochim. Acta, 1956, 966. (3) Dahlgren, G., Simmerman, N. L. (unpublished data). See Dahlgren, G. and Simmerman, N. L., Science 140, 485 (1963). (4) Diamond, L. H., Audrieth, L. F., J . Am. Chem. Soc. 77, 3131 (1955). (5) Ellis, A. J., Soper, F. G., J . Chem. Soc., 1954, 1750. (6) Hawkins, W., Smith, D. M., Mitchell, J., Jr., J . Am. Chem. SOC.66, 1662 (1944).’ (7) Kolthoff. Kolthoff, I. M.. M., Belcher. Belcher, R.. R., ‘[Volii“VoluAnalysis, ’ Vol. 111,” ’p, metric Andysis, p. 58i, 581, Interscience, New York, 1957. (8)Kolthoff. (8) Kolthoff, I. M.. M., Sandell. Sandell, E. B., “Textbook of ’Quantitative Quantitative Analvsis.” Analysis,” p. n. 597, Macmillan, Macmdlan, New York, 1952. 1k2.’ (9) Lacy, W., Shemwell, R., Huckaby, J., Proc. Louisiana Acad. Sci. 18, 94 (1955). (10) Lambert, J. L., ANAL. CHEM. 23, 1247 (1951). (11) LiebhafRky, H. A., Bronk, L. B., Zbid., 20, 588 (1948). (12) Makris. K. G.. 2. Anal. Chem. 84. . 241 ( m i j . (13) Mitchell, J., Jr., Hawkins, W., Smith, D. M..J . Am. Chem. SOC. 66, 782 (1944).’ (14) Noller, C. R., “Chemistry of Organic Compounds,” p. 233, Saunders, Philadelphia, 1951. (151 Rowe. R. A.. Audrieth. L. F.. J . Am. Chem. SOC.78, 563 (1956). (16) Van Hoogstmten, C. W., Rec. Trav. Chim. 51, 414 (1932). (17) Weil, I., Morris, J. C., J . Am. Chem. SOC.71. 1664 (1949). (18) Zitomer, F:, Lambert, J. L., ANAL. CHEM.34, 1738 (1962). (19) Zuman, P., Nature 165, 485 (1950). \
,
- I
’
RECEVIEDfor review June 6, 1963. Accepted November 7, 1963. Work supported by the American Cancer Society through grant P-272.
S pect rophot o metric Dete rminatio n of 8-Hydroxypenillic Acid DAVID A. HULL, JEAN M. SHEPP, WILLIAM J. WEAVER, ROBERT F. RESCHKE, and JOSEPH BOMSTEIN Brisfol Laborafories, Clivision of Brisfol-Myers Co., Syracuser N. Y.
b A method has been developed for determining 8-hydroxypenillic acid, based on isolating the silver salt, followed b y reacting with vanillin to produce a visible chromophore. PreEffects of cision is better than 3%. operating variables and interferences are discussed. Detailed procedures are given for determinations in fermentation media.
R
(1, 4 ) demonstrate that 8-hydroxypenillic acid (HPA) [The name 8-hydroxypenillic acid, for the enolic form of 3,3-diECENT REPORTS
-
-
methyl 8 - oxo - 4 - thia 1,7 - diazabicyclo - (3.3.0) - octane - 2,6dicarboxylic acid, has been used in conformity with Johnson and Hardcastle ( 4 ) ] can be isolated from Penicillium chrysogenum fermentations in which 6-aminopenicillanic acid (ilPA) is produced. In these reports, and in a paper appearing simultaneously (d), the authors show that APB reacts with COz in water to form HPA in high yields. Thus, substantial degradation of .4PA may occur in the presence of dissolved C 0 2 in aqueous solutions and fermcntation media. For this reason a reliable method is required for determining
HPA in the presence of APA and impurities likely to be encountered. Ballio et al. ( I ) reported that HPA takes up six equivalents of iodine per mole, and that this reaction can be used for quantitative analysis. An iodine titration method is subject to serious limitations, however, because of the possible presence of other iodineabsorbing species, for example, the penicilloic acids of XPA and Penicillin G. A search of the literature revealed no other applicable methods. The technique we have developed depends on isolating HPA as the silver salt, followed by reacting with vanillin VOL. 36, NO. 3, MARCH 1964
599
