Environ. Scl. Technol. 1984. 18. 97-100
Engelhardt, G.; Wallnofer, P. R. Appl. Environ. Microbiol. 1978,35,243-246. Engelhardt, G.; Wallnofer, P. R.; Hutzinger, 0. Bull. Environ. Contamin. Toxicol. 1975,13,342-347. Keyser, P.; Pujare, G.; Eaton, R. W.; Ribbons, D. W. EHP, Environ. Health Perspect. 1976,18,159-166. Kurane, R.; Susuhi, T.; Tahahara, Y. Agric. Biol. Chem. i977,4i,i031-103a. Kurane, R.; Susuhi, T.; Y. Tahahara, Y. Agric. Biol. Chem. 1977,41,2119-2123. Taylor, B. F.; Curry, R. W.; Corcoran, E. F. Appl. Environ. Microbiol. 1981,42,590-595. Benckiser, G.; Ottow, J. C. G. Appl. Environ. Microbiol. i982,44,m-582. Gledhill, E.; Kaley, G.; Adams, J.; Hicks, 0.;Michael, R.; Saeger, V. W.; LeBlanc, G. A. Environ. Sci. Technol. 1980, 14,301-305. Jacobs, L. W.; Phillips, J. H.; Zabik, M. J. Final Report to Michigan Department of Natural Resources, 1981. Strachan, S. D.; Nelson, D. W.; Sommers, L. E. J . Environ. Qual. 1983,12,69-74.
not previously been observed.
Acknowledgments We thank Brian Musselman of the MSU Mass Spectroscopy Center for his assistance in obtaining the mass spectra and Sue Frazier for her technical assistance. Registry No. Dimethyl phthalate, 131-11-3; diethyl phthalate, 84-66-2; di-n-butyl phthalate, 84-74-2; butyl benzyl phthalate, 85-68-7; di-n-octyl phthalate, 117-84-0; bis(2-ethylhexyl)phthalate, 117-81-7.
Literature Cited Graham, P. R. EHP, Environ. Health Perspect. 1973,3, 3-12. “Synthetic Organic Chemicals. U.S.Production and Sales, 1979”;US.International Trade Commission: Washington, DC, 1979. Atlas, E.; Giam, C. S. Science (Washington,D.C.) 1981,211, 163-165. Chem. Eng. News 1983,61(9),36-37. Mathur, S.P. J . Environ. Bull. 1974,3, 189-197. Johnson, G.; Lulves, W. J . Fish. Res. Board Can. 1975,32, 333-339. Saeger, V.W.; Tucher, E. S. Appl. Environ. Microbiol. 1976, 31,29-34.
Received for review December 20,1982. Accepted July 13,1983. This work was supported in part by a Biomedical Research Support Grant and an EPA-Batelle Grant. Contribution of the Michigan Agricultural Experiment Station, East Lansing, MI 48824. AES J . Ser. No. 10688.
Kinetics and Mechanism of the Decomposition of N-Chloro-2-( methylamino)ethanol in Aqueous Solution Juan M. Antelo, Florenclo Arce, Julio Casado,” Ramen Castro, Maria E. Sanchez, and Angel Varela Departamento de Quimica Fisica, Facultad de Qdmica, and Departamento de Investigaciones Qdmicas, C.S. IC., Universidad de Santiago de Composteia, Santiago de Composteia, Spain
A kinetic study of the system 2-(methy1amino)ethanol plus hypochlorite has been carried out under various conditions. The results obtained show that the formation of N-chloro-2-(methylamino)ethanoltakes place very rapidly, in under a minute, and that its stability varies in a complex fashion according to the pH of the medium. Over the range pH 9.5-13 the rate constant depends linearly upon the reciprocal of the concentration of protons, but at lower pH the results deviate from linearity, and in the range pH 6-8 the half-life is greater than 4.8 days. Comparison of the results obtained for various N-chloro alcoholamines shows their relative stabilities to follow the order C1N(H)CH2CH20H > C1N(CH3)CH2CH20H > ClN(CH&H,)CH&H20H > ClN(CH2CH20H)CH2CH20H.
Introduction The presence of nitrogenated organic compounds in waters subjected to chlorination treatment results in organic N-chloramines being formed in these systems (1). The ever-increasing use of alcoholamines in laboratories and industry (2) is particularly relevant to this process. N-Chloramines are apparently not biologically inert, and some research workers have even warned of mutagenic activity whose medium- and long-term effects are difficult to estimate (3). It is accordingly of interest to study the stability of N-chloramines, since this will influence any evaluation of the hazard they present (4, 5). This paper reports the results of a study on the stability and mechanism of decompsotion of N-chloro-2-(methylamino)ethanol (NClMAE) obtained by chloration of 20013-936X/84/0918-0097$01.50/0
(methy1amino)ethanol(MAE) in aqueous solution. These results are compared with those obtained for other Nchloramines (6-8) so as to arrive at an ordering by stability which might be borne in mind when an alcoholamine is selected for industrial use.
Experimental Section According to the available published findings (6,9,10) for this kind of system the reaction between MAE and hypochlorite may be written as HN(CH,)CH&H20H + HClO C1N(CH3)CH2CH20H+ H 2 0 (fast) ClN(CH3)CH2CH20H products (slow)
-
-
In the work described here NClMAE was generated by mixing known volumes of solutions of amine and sodium hypochlorite together with a suitable quantity of NaOH solution or buffer to maintain the reaction medium at the desired working pH. The decomposition of NClMAE was studied by recording the absorbance at 266 nm. Slow reactions, of half-life longer than 5 h, were also followed iodometrically (9). The preparation of solutions and experimental method have been described in an earlier article (6).
Spectrophotometric Study A preliminary spectrophotometric study of the reagents, products, and reaction mixture showed the formation and decomposition of NClMAE to take place at very different rates. Figure 1 shows the results obtained in one such experiment confirming the presence of both stages in the
0 1984 American Chemical Society
Environ. Sci. Technol., Vol. 18, No. 2, 1984 97
1
Table I. Influence of the Concentration of NClMAE upon the Rate of Reactiona
Ab
[cio-1x
103, M kexp, min-I Ab, 3.71 0.446 1.223 2.96 0.441 0.980 2.47 0.438 0.822 1.73 0.436 0.576 1.24 0.441 0.415 0.99 0.440 0.331 0.49 0.446 0.164 k e x p = 0.441 i 0.004 min-' = 331.3 f; 0.4 M-'cm-I
0.4-
a [MAE] = 0.05 M, [Cl'] = 0.5 M, and [NaOH] = p H 12.83.T = 25.1'C. and ill = 0.6 M.
0.2-
I
240
280
X(nm)
Figure 1. Spectrophotometric curves of the reagents and the reaction mixture. Dotted line: [CIO-= 2.13 X lo3 M. Unbroken line: [CIO-] = 2.13 X lom3M, [MAE] = 23.3 X lo3 M; (1)at 3 min, (2)at 23 min, (3)at 58 mln, (4)at 1 1 8 min, and (5)at 240 mln.
reaction. The dotted curve in this figure represents the spectrophotometric profile of a solution of hypochlorite and the unbroken curves those of the reaction mixture, C10- + MAE, at various times. On mixing the amine and hypochlorite solutions there occurs a fast process (tl,? < 30 s) during which the 292.5-nm hypochlorite absorption band rapidly vanishes and a 266-nm band due to the NClMAE formed appears. By the time the first spectrophotometric profile of the reaction mixture has been recorded all hypochlorite has already disappeared, so that the formation of NClMAE can only be studied by means of special techniques suitable for fast reactions. The presence of active chlorine after the C10- absorption band has disappeared can be demonstrated by using an iodometric method. The second process, the decomposition of NClMAE, is much slower and can be followed by using conventional spectrophotometry to measure the absorption due to the NClMAE at 266 nm. Reaction Products According to the results obtained by Dennis et al. (10) the reaction of alcoholamines with HClO in aqueous solution involves fragmentative oxidation in which the C-N bond is broken. In the present work all the kinetic experiments were carried out by using an [amine]/[ClO-] ratio greater than 10, so that the expected products of reaction were formaldehyde and a primary amine. The presence of formaldehyde was tested for using chromotropic acid (11)and its concentration determined via the precipitate ( p 183-184 "C) obtained on adding dimedone (12). The yield of formaldehyde was found to be 0.65 mol/mol of hypochlorite (65%). As for the primary amine, two products are possible depending on which of the C-N bonds is broken. Gas chromatography showed that both possibilities (H2NCH3and H2NCH2CH20H)are produced simultaneously and also provided evidence that the formaldehyde produced itself reacts with the amines, which would explain the less than 100% yield measured gravimetrically. Results In all the kinetic experiments carried out the 266-nm absorbance time data were found to fit a first-order integrated rate equation. The initial absorbance and rate 98
Environ. Sci. Technol., Vol. 18,No. 2, 1984
0.1 M ;
constant were estimated by the ordinary least-squares method. The coefficients of regression are greater than 0.99 at all pH values and greater than 0.999 at pH >9.5. To check the reproducibility of the results, some experiments were carried out in triplicate, and the rate constants obtained varied by less than 3 % for pH near 12. Whenever possible, reactions were followed until 75% complete, but this was not always possible due to the long half-life of some of the reactions. For the kinetic analysis of the data the value A , = 0 was used both for the relatively fast experiments (tlIz< 2 days) in which it was experimentally confirmed and for the slower ones in which experimental difficulties prevented its demonstration. The influence of the concentration of NClMAE was studied in a series of experiments in which the intitial concentration of C10- used to generate it was varied while all other experimental conditions were kept constant. The results, listed in Table I, show that the initial concentration of NClMAE does not affect the first-order rate constants. The values obtained for the initial absorbance show that the Lambert-Beer law is obeyed by solutions of NClMAE and allow its coefficient of extinction to be determined. It is found that at pH 12.83, ,A, = 266 nm, emax = 331.3 f 0.4 M-l cm-l, and the ordinate at the origin of the Lambert-Beer equation is statistically not different from zero, so that all C10- is incorporated in NClMAE at this PH. To study the influence of the concentration of protons on the rate of reaction, experiments were carried out over a wide range of pH. Table I1 shows the results obtained in the range pH 1.5-13, no attempt having been made to protect the reaction mixture from light radiation. It should be pointed out that at pH 38.5 h), so that only 20-50% of the reaction was studied, the kinetic data are more widely scattered (with coefficients of regression below 0.999, though greater than 0.99), and the reproducibility of the results is also poorer than that achieved at higher pH. Figure 2 shows the existence of a complex relationship between log k,, and pH. Above pH 9.5, however, the relationship is linear, and the following empirical equation has been obtained with a regression coefficient of 0.998: log k,, = (0.98 f 0.02)pH - (12.9 f 0.2) In another series of experiments light was excluded by shielding the reaction flask with aluminum foil. Periodic samples were taken whose absorbance at 266 nm was measured or whose concentration of NClMAE was determined icdometrically. The results obtained in this series of experiments appear to fit the first-order rate equation rather better, with coefficients of regression between 0.997 and 0.9998. The difference between the rate constants
Table 11. Influence of pHa .
.
r k e x p , min-' PH 0.9999 12.76 0.405 0.396 0.9999 12.74 0.199 0.9998 12.47 0.193 0.9999 12.47 0.9999 12.22 0.118 0.0742 0.9998 12.02 0.0597 0.9998 11.97 0.0308 0.9998 11.67 11.45 0.0183 0.998 11.03 0.0079 0,999 0.00319 10.74 0,999 10.13 0.999 0.00128 9.70 0.999 0.00053 0.999 0.00031 9.48 9.32 0.998 0.00023 9.20 0.998 0.00020 0.997 0.00020 9.16 0.00020 9.10 0.999 8.31 0.997 0.00010 0.000071 7.49 0.997 5.79 0.0000091 0.996 0.0000086 5.40 0.999 4.48 0.998 0.0000077 0.000011 3.96 0.992 0.000013 3.49 0.996 0.00005,approximate 2.49 2.49 0.00012,approximate In the Absence of Light Radiation 9.00 0.00042 0.999 8.52 0.00046 0.9999 8.53 0.00043 0.9998 8.00 0.00034 0.9997 0.000080 0.9997 7.27 6.97 0.000022 0.998 6.50 0.0000066 0.997 6.45 0.0000049 0.998 3.54 0.0000068 0.998 3.26 0.0000098 0.999
[ClO-] = 1.98 lo-' M, [MAE] = 2.4 M, and [NaCl] = 0.5M. Buffer concentration: borate, 0.2 M; acetate, 0.3 M. T = 25 OC.
Table I11 summarizes the results obtained in various series of experiments in which the first-order rate constant was found- to be independent of the concentration of amine, the ionic strength, the concentration of bromide, and the presence of small quantities of divalent metal ions (Cu2+,Ni2+,and Hg2+). For each series the value of k,, given is the mean after correction to the pH shown. Al! the experiments were carried out at pH >lo, where log kexp depends linearly on pH. Finally, the influence of temperature on the reaction rate was studied. A 10 deg rise in temperature was found approximately to treble the value of the rate constant. The results obtained allow the thermodynamic activation parameters at pH 12.13 to be calculated as E, = 80.2 f 0.6 kJ mol-l, AH* = 77.3 f 0.6 kJ mol-l, and AS* = -40.4 J mol-' K-l.
Reaction Mechanism The solutions of NClMAE used in the kinetic studies were prepared by mixing excess amine with hypochlorite, so that the equilibria present in the reagent solutions and the process of formation of NClMaE must be borne in mind. Published findings support the following mechanism for the formation of NClMAE (6, 13): K
+ H 2 02HClO + OHH2N+(CH3)CH2CH20H + C10-
2HN(CH3)CH2CH20H+ H 2 0
OH-
+
HN(CHJCH2CH2OH
OH-
& HN(CH3)CH2CH20-+ H 2 0
H2N+(CHJCH2CH20H
+ HClO
k4 +
ClN+H(CH&H2CH20H
+ HClO
HN(CH3)CH&H20H
+ H2O
ks -*
ClN(CHJCH2CH20H
+ H2O
HN(CH3)CH2CH,O- + HClO -* ClN(CHJCHZCH20-
+ H2O
k6
ClN+H(CH3)CH,CH20H OH-
2C1N(CH3)CH2CH20H+ H 2 0
ClN(CH,)CH2CH20H
-1-
+
OH-
+
& C1N(CH3)CH2CH20-+ H 2 0
Any of the forms of NClMAE may lead to the formation of the find products, so that three parallel decomposition processes must be considered:
-3-
-
C1N+H(CH3)CH2CH20H C1N(CH3)CH2CH20H
-5-
1
C1N(CH3)CH2CH20-
I
2
6
'0
pH
k9
k10
kll
products
products
products
Flgue 2. Influence of the pH of the medlum upon the rate of reaction. (a) NChloroethanolamlne(6); (b) Nchloro-2-(methylarnino)ethanol (this paper): (c) Nchloro-2-(ethylamlno)ethanol (8);(d) N-chlorodiethanolamine (7).
The overall rate equation is therefore of the form d [products] u= = k9[C1NfH(CH3)CH2CH20H] + dt k,,[ClN(CH3)CH&H20H] + kil[ClN(CH3)CH2CH20-]
found in the presence and in the absence of light seems to show that when the reaction is very slow (pH > 1 >> K,[OH-]. The linear relationship between log ke, and pH found experimentally at pH >9.5 may therefore be explained if in addition it is assumed that under these conditions k, and klo are insignificant beside kll, for then calculation of the concentration of C1N(CH3)CH2CH20-shows that the rate equation may be written as
0.4 M-l cm-l. (c) Over the range of maximum environmental interest, pH 6-8, the formation of NClMAE takes place in much less than a minute, and the rate constant for its decomposition is less than 1 X mi&, with tlIz longer than 4.8 days. (d) A t room temperature the rate constant for the decomposition of NClMAE is roughly trebled by a 10 deg rise. (e) Comparison of various Nchloro alcoholamines of the form C1N(R)CH2CH20H shows their stability to depend on R following the order H > CH3 > CHzCH3 > CH2CH20H.
where [Cl,] is the total concentration of N-chloro alcoholamine. Granted these assumptions, then the mechanism by which the final products are arrived at at pH >9.5 may be written, (17-20) as ClN(CH,)CH,CH20C1- H+ + CH2=NCH,CH20-
Literature Cited Scully, F. E.; Bempong, M. A. “Water Chlorination”; Ann Arbor Science Publishers: Ann Arbor, MI, 1980; Vol. 111, p 203. Mullins, R. M. Kirk-Othmer Encycl. Chem. Technol., 3rd Ed. 1978,1,944-960. Bempong, M. A.; Scully, F. E. “Water Chlorination”; Ann Arbor Science Publishers: Ann Arbor, MI, 1980; Vol. 111, p 817. Payne, J. F.; Martins, I.; Fayon, D.; Rahintula, A. “Water Chlorination”; Ann Arbor Science Publishers: Ann Arbor, MI, 1980; Vol. 111, p 845. Bull, R. J. Environ. Sci. Technol. 1982, 16, 554A. Antelo, J. M.; Arce, F.; Barbadillo, F.; Casado, J.; Varela, A. Enuiron. Sci. Technol. 1981, 15, 912. Antelo, J. M.; Arce, F.; Casado, J.; Varela, A. An. Quim.
-
+
+ H2NCH2CH20C1- + CH2O + CH,N=CH2 ClN(CH3)CH&H20CH,N=CHZ + H2O CH2O + CH3NH2
CH2=NCH&H20-
+ H2O +
+
CHzO
-
However, since the contributions of klo and k9 increase as pH falls, it is quite possible that at pH CH3> CH2CH3> CH2CH20H. At pH 10.5, for example, the values of k found are, in this same order, 0.005, 0.27, 2.19, and 4.17 min-’. Conclusions
The information obtained from our study of the reaction system formed by 2-(methy1amino)ethanol and hypochlorite can be summarized in the following points: (a) Spectrophotometric data reveal the Occurrence of two consecutive processes with very different reaction rates: the rapid formation of NClMAE is followed by its slow decomposition. (b) Spectrophotometric analysis of the reaction mixture shows that NClMAE exhibits maximum absorption at 266 nm. The coefficient of extinction at this wavelength calculated from kinetic data is ern= = 331.3 f 100 Environ. Sci. Technol., Vol. 18, No. 2, 1984
1982, 78, 63. Antelo, J. M.; Arce, F.; Casado, J.; Castro, R.; Shchez, M. E.; Vqela, A. XIX Reuni6n Bienal Real Sociedad Espafiola de Quimica, Santander, Spain, 1982. Hull, L. A.; Giordano, W. P.; Rosenblatt, D. H.; Davis, G. T.;Man, C. K.; Milliken, S. B. J. Phys. Chem. 1969,73,2147. Dennis, W. H.; Hull, L. A.; Rosenblatt, D. H. J. Org. Chem. 1967, 32, 3783. Feigl, F. “Spot Test in Organic Analysis”, 6th ed.; Elsevier: Amsterdam, 1960; p 349. Walker, J. F. “Formaldehyde”;Robert E. Krieger Publishing Co: New York, 1975; pp 494-497. Weil, I.; Morris, J. C. J.Am. Chem. SOC.1949, 71, 1664. K7 0.316/Kw. Calculated in our laboratory. Weil, I.; Morris, J. C. J. Am. Chem. SOC.1949, 71, 3123. Maswe, F.; Schaal, R. Bull. SOC.Chim.Fr. 1956,1138,1141, 1143. Douheret, G.; Pariaud, J. C. J. Chim. Phys. Phys.-Chim. Biol. 1962, 59, 1021. Kaminski, J. J.; Bodor, N.; Higuchi, T. J. Pharm. Sci. 1976, 65, 553. Patai, S. “The Chemistry of the Amino Group”; Intencience: London, 1968; Chapter 6. Stanbro, W. D.; Smith, W. D. Enuiron. Sci. Technol. 1979, 13, 446. Ogata, Y.; Kimuro, M.; Kando, Y. Bull. Chem. SOC.Jpn. 1981, 54, 2057.
Received for review March 3, 1983. Accepted August 1, 1983.
