retary of Transportation:Washington, DC, 1976; TAD-443.1. (50) U S . Environmental Protection Agency “Denver Air Pollution Study”, 1976, EPA-600/9-77-001,EPA-600/9-76-007a. (51) U S . Environmental Protection Agency “Federal Air Quality Control Regions”, 1972, AP 102. (52) U S . Environmental Protection Agency “Proceduresfor Quantifying Relations Between Photochemical Oxidants and Precursors’’, 1978, EPA 450/2-77-002b. (53) U S . Environmental Protection Agency “Air Quality Criteria for Ozone and Other Photochemical Oxidants”;EPA-600/8-78-004, 1978;Vol. I, Section 4. (54) Scully, R. D. Enuiron. Sci. Technol. 1980,14,1271. (55) Trijonis, J. C., Jr. “Oxidant and Precursor Trends in the Metropolitan Los Angeles Region: An Update”; Conference on Air Quality Trends in the South Coast Basin, California Institute of Technology, Pasadena, CA, Feb 21-22,1980. (56) Trijonis, J. C.; Peng, T. K.; McRae, G. J.; Lee, Lester, EQL Memorandum No. 16; California Institute of Technology, Pasa-
dena, CA, 1976. (57) Dimitriades, B. Enuiron. Sci. Technol. 1977,11,80-7
Received f o r reuiero January 9,1980, Accepted March, 20,1981
Supplementary Material Available: Calculated vs. Experimental ozone values for Pitts et al. smog-chamberstudy on Los Angeles basin surrogate compositions (4pages) will appear following these pages in the microfilm edition of this volume of the journal. Photocopies of the supplementarymaterial from this paper on microfiche (105 X 148 mm, 24X reduction, negatives) may be obtained from Business Operations, Books and Journals Division, American Chemical Society, 1155 16th St., N.W., Washington, DC 20036. Full bibliographic citation (journal, title of article, author)and prepayment, check or money order for $4.00 for photocopy ($5.50 foreign) or $4.00 for microfiche ($5.00foreign), are required.
Kinetics and Mechanism of Ethanolamine Chlorination J. M. Antelo, F. Arce, F. Barbadillo, J. Casado,” and A. Varela Departamento de Quimica fisica, lnstituto de lnvestigaciones Quimicas del C.S.I.C., Facultad de Quimica, Universidad de Santiago de Compostela, Spain
The kinetics and the mechanism of ethanolamine chlorination in aqueous solution a t high p H were studied by using sodium hypochlorite in an alkaline medium as chlorinating agent. The results obtained show the reaction to take place in two stages with very different velocities, N-chloroethanolamine being formed in the first stage and undergoing decomposition in the second. The rate of formation of N-chloroethanolamine is proportional to the concentration of hypochlorite and amine and inversely proportional to that of NaOH. The rate of decomposition of N-chloroethanolamine is proportional to its concentration and to that of NaOH and is independent of ionic strength when different concentrations of NaN03 or NaC104 are used. The influence of temperature upon the rate of decomposition of N-chloroethanolamine was studied, and the activation parameters were determined ( E , = 71 f 3 kJ mol-1). A reaction mechanism compatible with the experimental results is proposed.
Introduction In the last few years the number of synthetic chemical compounds has increased enormously, with a corresponding increase in their release into the environment. The importance now given to preserving the quality of the environment has consequently prompted the desire for more knowledge about compounds resulting upon the introduction of chemical agents into natural systems. One aspect of this problem is the investigation of the stability of the compounds which arise during chlorination of water supplies of effluent and which end up in rivers, lakes, and the sea (1). In the work described in this paper, we have studied the process of formation and decomposition of N-chloroethanolamine (ClNHCH2CH20H) in aqueous solution consequent upon the mixture of ethanolamine NH2CH2CH20H and sodium hypochlorite under various experimental conditions. Ethanolamine was chosen because its occurrence in laboratory and industrial processes makes it a likely polluting agent in effluent. I t is therefore of some interest to know its behavior with respect to the chlorinating agents employed in water treatment. The oxidation of amines by halogens has been studied by several researchers, and in particular Kovacic et al. ( 2 )have published an extensive survey of the chemistry of N-chlo912
Environmental Science & Technology
ramines. Recently, kinetic studies of the oxidation of alcohol amines by chloramine-?’ ( 3 , 4 ) have thrown light upon the mechanism of formation of N-chloroethanolamines. These reactions are analogous to those occurring in hypochlorite solution but much slower owing to a series of equilibria involving species active in chloramine-?‘ solution ( 5 ) .
Experimental Technique The reaction was followed by measuring absorbance at 292.5 and 252 nm. The spectrophotometric study of reactants and products, and of those reactions with half-lives t 1/2 longer than 1 min, was carried out by using a PYE-UNICAM SP 1700 spectrophotometer equipped with a PYE-UNICAM AR-25 plotter and a cell holder held a t a temperature constant to f O . l OC by means of water flow from a HETO-CB-7 thermostat. Very slow reactions with half-lives longer than 1 h were studied iodometrically by mixing a constant volume (1mL) of the reaction system with a fixed volume (20 mL) of acidic 0.1 M potassium iodide solution and determining the iodine released either by titration against sodium thiosulfate solution of known concentration or by measuring the absorbance of triiodide a t 353 nm. Fast reactions ( t l / z < 1min) were followed with a Nortech SF-3A stopped-flow spectrophotometer coupled to a Tektronix T-912 oscilloscope. The output displayed on the oscilloscope screen was recorded photographically. The mixing cell and a part of the flow system were surrounded by a bath ensuring a temperature constant to f O . l “C. The reaction was started by mixing such volumes of thermostated alkaline solutions of hypochlorite and amine to achieve the desired conditions of concentration and basicity. The hypochlorite solution was prepared by bubbling chlorine through 1 M NaOH solution to which, once saturated, sufficient NaOH was added to hold the pH a t -11 and so inhibit the decomposition of the hypochlorite. For each experimental run fresh hypochlorite solution was prepared and determined by iodometric titration before use. The pH of solutions and reaction mixtures was measured with a Beckman Model 4500 pH meter equipped with a Beckman Futura 39501 combined electrode. Solutions of reactants were prepared from Merck p.a. products.
0013-936X/81/0915-0912$01.25/0
@ 1981 American Chemical Society
Spectrophotometric Studies. An initial spectrophotometric study of reactants and products of reaction showed 292.5 nm to be the optimum wavelength for following the reaction, only hypochlorite presenting absorption a t this wavelength (6).The Lambert-Beer law for solutions weaker than 6 X 10-3 M was found to obtain under the experimental conditions employed. A series of preliminary experiments designed to establish the most suitable experimental conditions agreed with published findings (7) in demonstrating the existence of two distinct reaction processes with very different rates. Figure 1 shows a typical oscilloscope curve of voltage against time obtained by the stopped-flow technique at 292.5 nm. The voltage falls rapidly to V , within 0.5 s, a t least 90 min being required for the subsequent drop from V , to the final value Vf. (In the FRO-1 module of the SF-3A spectrometer used in the experimental setup, the voltage read off the oscilloscope is proportional to the absorbance of the solution present in the reaction cell.) Figure 2 shows the results of experiments employing conventional spectrophotometry scanning between 220 and 350 nm every 70 s. As can be seen, by the time the first curve ( n = 1) is recorded a minute after initiation of the reaction, 292.5-nm absorption has practically disappeared, while at the
same time a strong peak has appeared at 252 nm. Since this wavelength does not correspond to any initial reactant or final product, the peak must be produced by some intermediary whose concentration, according to Figure 2, decreases with time. The literature (2) indicates that the intermediary should be an N-chloramine, and this is supported by the results of iodometric titration, which show that the amount of active chlorine present in the system is greater than that assigned by spectrophotometric determination a t 292.5 nm to the hypochlorite. Accordingly, we propose the following reaction scheme:
+
N H ~ C H Z C H ~ O HHClO
2ClNHCHzCHzOH + HzO kz
ClNHCHzCHzOH +products
(fast)
(slow)
The experimental results show the difference between the values of k l (investigated by stopped-flow spectrophotometry at 292.5 nm) and hz (investigated by conventional spectrophotometry at 252 nm or by iodometry) to be sufficiently great for the use of these different methods to be permissible. According to published findings (3, 7) the products of reaction in excess oxidant are given by 3HC10
+ NHzCH2CHzOH
+
2HCOOH
+ 3HC1+ NH3
In excess amine the initial product of oxidation is formaldehyde, which may subsequently undergo further oxidation to form formic acid: HClO
+ NHzCHzCHzOH
+
2HCOH
+ HCl + NH3
The presence of formic acid in the reaction mixture was shown by the chromotropic acid test ( 8 ) ,and the presence of formaldehyde by the formation of a white precipitate of melting point 187-189 "C on addition of dimedone solution (7).
1
I
I
Figure 1. Typical oscilloscope trace at = 292.5 nm, [CIO-] = 3 X M, [NaOH] = 0.4 M, [EA] = 0.06 M, T = 24 O C . Trace a: 0.50 s/division Trace b: 0.05 sldivision.
Results Fast Reaction. Formation of N-Chloroethanolamine. The isolation method was used to determine the order of the reaction with respect to hypochlorite, the initial concentration of which was thus kept low with respect to the other reactants. The spectroscopic studies establish the following relationship between absorbance (Ab) and concentrations at 292.5 nm:
where 1 is the length of the mixing cell, € 1 and € 2 are the molar absorptivities of the hypochlorite ion and N-chloroethanolamine, respectively, and [C10-]0 and [C10-lt are the concentrations of hypochlorite initially and at time t . This relationship enables the voltage-time data of the oscilloscope curves to be interpreted. If Ab, is the absorbance of the solution when all hypochlorite has disappeared and no appreciable decomposition of N-chloroethanolamine has yet taken place, then Ab, = EZ~[CIO-]O =E~~[C~NHCH~CH~OH]~
so that
230 250 270 290 310 h(nm) Figure 2. Variation of absorbance with time. [CIO-] = 1.1 X [NaOH] = 0.4 M, [EA] = 0.1 M; (---) hypochlorite solution.
M,
where c is a constant and V, the voltage registered on the oscilloscope screen, is directly proportional to the absorbance, differing from it only by a fixed factor depending on the Volume 15, Number 8, August 1981
913
FRO-1 module of th'e stopped-flow system and on the oscilloscope. The value of V , may be found graphically as the voltage of the point at which the oscilloscope curve flattens out, and the values of V , and V t obtained yield graphs (Figure 3) whose linearity shows the reaction to be of order 1 with respect to hypochlorite. A least-squares fit of the experimental results gives the experimental first-order rate constants kexptlshown in Table I. keXptlcan be seen to depend linearly on the concentration of amine; i.e., the reaction is first-order with respect to ethanolamine. The influence of the concentration of NaOH was investigated in a series of experiments in which the initial concentration of NaOH was varied while that of the other reactants was kept constant. Table I1 shows the results obtained: the first-order rate constant is found to depend linearly upon the reciprocal of [NaOH]. Taken all together, the experimental behavior described above indicates that the kinetics of the reaction comply with the equation: u = k [HClO][NH2CH2CH20H]/[OH-]
In
( V t - Vm)
Slow Reaction. Decomposition of N-Chloroethanolamine. The kinetics of the decomposition of N-chloroethanolamine were studied by measuring the absorbance a t 252 nm of reaction systems with different initial concentrations of hypochlorite. Table 111shows the initial rates for different values of [ClO-1. The linear relation between these quantities shows the reaction to be first-order with respect to N-chloroethanolamine, the decomposition of 1 mol of C10- entailing the formation of 1 mol of ClNR. The experimental absorbance-time curves agree with this interpretation, and the first-order rate constants derived from them are also shown in Table 111,where they can be seen to be independent of the initial concentration of hypochlorite, and therefore of that of N-chloroethanolamine. Varying the initial concentration of amine has no effect upon the value of the experimental rate constant in the range studied (Table IV). This result is only to be expected, since the initial concentration of N-chloroethanolamine, whose decomposition is the process here in question, depends solely upon that of hypochlorite if hypochlorite is kept short. The experimental values thus remain constant even for concentrations incompatible with the isolation method. The results of a series of experiments in which only the initial concentration of NaOH was varied (Table V) reveal a linear relationship between the rate constant and [NaOH];i.e., the decomposition of N-chloroethanolamine is first-order with respect to NaOH. Table II. Influence of NaOH Concentration upon the Rate of Formation of N-Chloroethanolamine ko = kexpt~[OH-l/ [NaOH]/M
kexpds-'
(mol L-Is-1)
0.20 0.15 0.09 0.07 0.05 0.03
13 18 28 36 44 91
2.6 2.7 2.5 2.5 2.2 2.7
ko = 2.5 f 0.2mol L - ' s - ~ a
[CIO-] = 2.8 X
[EA] = 0.1 M, T = 23.2'C
Table 111. Rate Constants of the Decomposition of NChloroethanolamine a I
I
I
I
0,2 0,3 0.4 t (sec 1 Figure 3. First-order plots at 24 O C for the disappearance of CIO- in alkaline medium. [CIO-] = 3 X loT3 M, [NaOH] = 0.4M, [EA] = 0.02, 0.03,0.04,and 0.05M for A, 8,C, and D,respectively.
0,l
Table 1. Influence of Amine Concentration upon the Rate of Formation of N-Chloroethanolamine a (EAIIM
*exptl/S-'
0.02 0.03 0.04 0.05 0.06
2.09 3.14 4.30 5.50 6.40
k 2 = (kex I/[EAI)/ (L 5-1)
105 105 108 110 107
k2 = 107 f 2 L mol-'^-^ a
914
[ClO-] = 3 X
M, [NaOH] =
0.4M, T = 24.5 'C
Environmental Science & Technology
5.3 4.6 4.5 4.2 4.0 3.5 3.0 2.5 2.2 2.0 1.8 1.4 1.1
10 9.5
5.7 5.7
8.9
5.4
8.2
5.2
5.1 5.0
6.1 7.1
11.0 11.0 12.2 11.0 12.2 12.4 11.4
3.4 3.1 1.8
5.4 6.4 5.0 T = 15.4 O C , kexptl = (5.8f 0.6)X T = 25 O C , kexptl = (11.7f 0.6)X
[EA] = 0.1 M, [NaOH] = 0.4 M.
min-'
lop3 min-l
Table IV. Influence of Amine Concentration upon the Rate of Decomposition of N-Chloroethanolamlne IEAIIM
1O2keXp,l*/min-'
PH
102kexptlb/min-l
PH
~~~k,,~tl/m~n-'
104 kexptl/min-l
0.30
1.3
12.55
29.79
10.91
1.oo
0.20
1.1
12.38
22.17
10.61
0.68
0.14
1.4
1 1.83
6.57
10.56
0.66
0.10
1.2
11.55
5.07
10.37
0.40
0.060
1.5
0.71
0.050
0.53
0.030
0.56
0.010
0.70
0.0050
0.67
0.0025
0.62
T = 25
OC,
kexptl = (1.3 f 0.2) X
T = 14.7 O C , kexptl = (0.63 a
Table VI. Influence of pH upon the Rate of Decomposition of N-Chloroethanolaminea
M, [EA] = 0.1 M, [H3B03] = 0.1 M, [CINa] = 0.1 a [CIO-] = 0.897 X M, [NaOH] = variable, T = 25 'C.
Table VII. Influence of the Concentration of NaN03 and NaCIO4 upon the Rate of Decomposition of NChloroethanolamine a
min-I
f 0.07) X
INaN0311M
M, [NaOH] = 0.4 M. T = 25 OC. [ClO-] = 3.2 X M, [NaOH] = 0.44 M. T = 14.7 'C.
0.5
[CIO-] = 1.7 X
Table V. Influence of NaOH Concentration upon the Rate of Decomposition of N-Chloroethanolaminea [ NaOH I/M
104kexp~l/mln-1
104kexp~l/min-1
PH
min-'
12.88
50.80
12.80
49.2
1.o
12.79
61.5
1.5
12.75
61.5
2.0
12.75
62.8
[ NaC104 I/M
PH
k = ((kexpti/[NaOHI) X I O 2 ] / (L mol-' min-1)
0.5
12.85
53.5
1.o
12.80
53.5
104kexpt~/min-'
0.034
3.7
1.1
0.084
7.8
0.9
1.5
12.82
58.5
16
1.2
2.0
12.83
50.9
0.135 0.21 1
18
0.9
0.263
31
1.2
0.337
33
1.o
0.388
41
1.1
0.464
48
1.o
0.515
50
1.o
0.591
60
1.o
0.641
68
1.2
0.692
77
1.1
0.717
79
1.1
0.742
72
1.o
0.768
79
1.o
k = (1.1 f 0.1) X 10-*L mol-' min-1 IClO-1 = 3.3 X
a
M, [EA] = 0.1 M, [NaOH] = 0.1 M. T = 25 OC.
[CIO-] = 0.897 X
Table VIII. Influence of Temperature upon the Rate of Decomposition of N-Chloroethanolaminea T/K
102kexp~l/mln-1
288.2
0.48
T/K
10' kexptl/min-
307.7
3.7
293.2
0.99
312.1
5.7
298.0
1.6
317.7
7.6
303.0
2.3
AM = 69 + 3 kJ mol-'
E. = 71 f 3 kJ mol-',
AS* = -83 f 8 J mol-' K - l a
[CIO-] = 3.2 X
M, [NaOH] = 0.4 M, [EA] = 0.1 M.
M. [EA] = 0.1 M. T = 14.7 'C.
CIOIn Table VI are displayed the results of a series of experiments using buffer solutions to bring out the influence of p H on the reaction. The results listed in Table VI1 for different concentrations of NaNO:%and NaC104 show the rate constant for the decomposition of N-chloroethanolamine to be independent of background ionic strength. Finally, a series of experiments carried out a t different temperatures (Table VIII) complied with Arrhenius' equation and with activated complex theory and was used to calculate the activation parameters for the decomposition of N-chloroethanolamine. The foregoing experimental behavior indicates that the rate equation for the reaction must be of the form: u = h [CINHCH#H*OH] [OH-]
Discussion
Fast Reaction. Formation of N-Chloroethanolamine. On the basis of published work on N-chlorination ( 2 , 7 , 9 ) ,the following mechanism for the formation of N-chloroethanolamine can be put forward:
'
+ H20
h
+HNHZCH~CH~+ OH OHNH2CH2CH20H
HClO
+ OH-
KI + NHzCH2CHzOH + H20
+ OH- +NHzCH2CH20- + H20 K2
N H z C H ~ C H ~ O+HHClO
2CINHCH2CH20H+ H 2 0
from which it f'ollows that u r = -d[HClO]/dt = d[CINHCHsCH2OH]/dt = hi [HClO][ N H ~ C H Y C H ~ O H ]
The concentrations of HClO and NH2CH2CH20H can be found in terms of the equilibrium constants Kh, K I , and K 2 [HCIO] =
K,[OH-I
[NH&H2CH20H] = [Am] 1
+ K,[OH-] + K I K 2 1 0 H T
Volume 15, Number 8, August 1981
915
The rate equation can therefore be expressed as
NH2CH2CH20H. Under these conditions of pH, the rate equation reduces to: (ur)max
where [ C ~ Tis] the total concentration of active chlorine, [Am] the total concentration of amine, and K a the ionization constant of hypochlorous acid. The values of Kh, K1 and K2 can easily be derived from known results for the constants of the following ionization equilibria: for the ionization of hypochlorous acid, pK, = 7.54 (10);for that of the protonated amine, PKbl = 9.45, where Kbl = K l K , ( 1 1 ) ; and for that of the alcohol group of the free amine, pKb2 = 16-17, where Kb2 = K2K, (12,13) (in these relations K , is the ionic product of water). Accordingly, the rate equation may be simplified in various ways depending on the p H of the medium. In alkaline solution with p H 10.5-13.5, we may write u, = kl[C1~][Arn]Kh/[oH-]= hl[c1~][Am][H+]/K, =
K.3 Slow Reaction. Decomposition of N-Chloroethanolamine. On the basis of the kinetic results obtained and the published information concerning similar processes ( 1 , 2 , 7 ) , the following mechanism can be put forward for the decomposition of N-chloroethanolamine: ClNHCH2CH20H
K3 + OH- F+ ClNHCH2CH20- + H20 hz
ClNHCH2CH20- -+ C1HN=CH2
+ HN=CH2 + CH2O
fast
+ H20 +NH3 + CH2O
Since ionization equilibria are usually very fast, the concentration of ClNHCH2CH20- can be related to the constant K3:
+
[ClNHCH2CH20-] = [C1~2]K3[0H-]/(l K3[OH-])
k exptl [CITI This equation adequately describes all of the experimental facts under the conditions employed in our work ([NaOH] = 0.02-0.4 M), including the increase in kexptlwith decreasing concentration of NaOH. This inverse dependency on [NaOH] limits our investigation of N-chloroethanolamine formation to p H >12, since the mixing time of the stopped-flow apparatus used is -2 ms, and processes whose half-lives are less than -1.5 ms cannot be studied accurately ( 1 4 ) . From the equilibrium constants the variation of log [HClO], log [ N H ~ C H ~ C H Z O H and ] , log ([HClO][NH2CH2CH20H]) against pH for a system in which [C~T]= 1 X M and [Am] = 1 X 10-1 M can be calculated (Figure 4). Log ([HClO] [NH2CH2CH20H]) can be seen to reach a maximum between pH 8 and 9. The reaction rate should also therefore be fastest in this p H range, since the reaction mechanism that we have proposed means that the rate is proportional to the product of the concentrations of the uncharged species HClO and
= k 1[ClT][Am] KIKw -- (kexpt1)max[C1~]
where [ C l ~ zis] the total concentration of N-chloroethanolamine. The rate of extinction of N-chloroethanolamine will be given by U,
= -d[ClNHCH2CHzOH]/dt = k2[ClNHCH2CH20-] =
~ z [ C ~ T Z ] K ~ [ ~+HK3[OH-]) -]/(~ The value of K3 is not known, but it seems reasonable to suppose that it is very close to that of K P .In the experimental conditions used ([OH-] < 0.1 M) we must therefore have K3[OH-] < 1,and the rate equation becomes ur
= k2K3[C1T21[OH-]
which adequately describes all of the experimental findings under the conditions employed.
Conclusions The foregoing considerations indicate that the variation in N-chloroethanolamine concentration will be reflected in two terms, one corresponding to its formation and the other to its decomposition:
I
Kh
K1PH-I K1K2[0H-J2
"'= kl[CIT1[Aml [OH-] 4-Kh 1 + K1[OH-]
+
-3
-5
-7
-9
I
I
I
1
I
3
5
1
9
I
PH
Figure 4. Variation of [HCIG] and [NH2CH2CH20H]with pH as predicted by the respective dissociation equilibria.
916
Environmental Science & Technology
This expression, which can be simplified for certain values of [OH-], reveals a complex relationship between the reaction rate and the pH of the system: increasing the basicity of the medium discourages the formation and favors the decomposition of N-chloroethanolamine. In the range of concentrations studied, however, the difference between the values of the rate constants k l and h 2 is so great that the overall oxidation process depends fundamentally upon ha, which is the slow process and therefore constitutes the limiting stage. From the experimental results of Table I, the rate constant for N-chloroethanolamine formation can be calculated as hl = 12.7 X 107 s-l a t 24.5 "C. On the basis of this result, the maximum experimental rate, occurring a t pH 8-9, can be = 1.5 X lo5 s-l for 0.1 M roughly estimated as (hexptl)max ethanolamine solution. The rate of formation of N-chloroethanolamine should fall off on either side of this p H range. The results of Table VI, on the other hand, show that for the decomposition of N-chloroethanolamine the rate de-
(Table VIII). In the range studied a 10 “C rise in temperature roughly doubles the reaction rate.
Table IX. Rates of Decomposition of NChloroethanolamine in Neutral Media kexp~~/mln-’
PH
x 2.9 x 3.4 x 3.5
7.91 7.67 6.88
10-5 10-5
creases with pH. For the range p H 10.3-12.5 we find the linear relation
- 13.26
We acknowledge the helpful suggestions made by the reviewers.
10-6
[CIO-] = 0.897X 10-3M, [EA] = 0.1 M, [NaCIO4] = 1 M, T = 25 O C .
log heXptl= 0.86 X p H
Acknowledgment
( r x y= 0.998)
Extrapolating these results suggests that the rate must be very slow in the range p H 6-8. In order to calculate the rate constant for naturally occurring conditions, experiments were carried out a t pH 6.88,7.67, and 7.91 and 25 O C . The decomposition of N-chloroethanolamine under these conditions is indeed very slow. For 20 days we each day determined pH and active chlorine (by iodometry) and traced the absorption spectrum between 190 and 300 nm. The results obtained are shown in Table IX, and they show that a discrepancy exists between experimental values and those obtained by extrapolation of the data in Table VI, underlying the caution with which such extrapolation should be used in this kind of studies. Finally, the marked influence of temperature upon the rate of decomposition of N-chloroethanolamine is to be noted
Literature Cited ( 1 ) Stanbro, W. D.; Smith, W. D. Environ. Sei. Technol. 1979, 13,
446. (2) Kovacic, P.; Lowery, M. K.; Field, K. W. Chem. Rev. 1970, 70, 639. (3) Chandra, M.; Bansal, 0. P. J. Inorg. Nucl. Chem. 1978, 40, 1185. (4) Chandra, M.; Lal, S.; Bansal, 0. P. J . Indian Chem. SOC.1977, 54,1040. (5) Jennings, V. J. CRC Crit. Rev. Anal. Chem. 1974,3,407. (6) Hussain, A.; Trudell, P.; Repta, A. J. J . Pharm. Sei. 1970, 59, 1168. (7) Dennis, W. H.; Hull, L. A.; Rosenblatt, D. H. J. Org. Chem. 1967, 32, 3783. (8) Feigl, F. “Spot Test in Organic Analysis”, 6th ed.; Elsevier: Amsterdam, 1960; p 349. (9) Weil, I.; Morris, J. C. J. Am. Chem. SOC.1949, 71, 1664. (10) Morris, J. C. J . Phys. Chem. 1966, 70, 3798. (11) Cachaza, J. M.; HerrBez,M. A.; Pedrares, M. D. An. Q u h 1972, 68,1341. (12) Masure, F.; Schaal, R. Bull. SOC.Chim. Fr. 1956, 1138, 1141, 1143. (13) Jacquinot-Vermese, C. C. R. Hebd. Seances Acad. Sei. 1962,254, 3679. (14) Caldin, E. F.; Crooks, J. F.; Queen, A. J. Phys. E. 1973,6,930. Received for review April 11,1980.Accepted April 3,1981
Volatile Liquid Hydrocarbon Characterization of Underwater Hydrocarbon Vents and Formation Waters from Offshore Production Operations Theodor C. Sauer, Jr. t Department of Oceanography, Texas A & M University, College Station, Texas 77843
Underwater hydrocarbon vent and formation water samples, two discharges from offshore production operations in the Gulf of Mexico, were compositionally characterized for volatile liquid hydrocarbons (VLHs). Hydrocarbons in surface samples of an underwater vent were not detected with carbon numbers greater than 10 ( n - C l o ) . Alkanes were the major components of all of the VLHs in vent samples with less than 10% being aromatic hydrocarbons. Hydrocarbons in samples of a formation water discharge were considerably more extensive and complex than vented hydrocarbons. Total VLH concentrations were -20 mg/L, 80%of which were aromatic hydrocarbons (mostly benzene, toluene, and xylenes), close to the percentage found in coastal waters of the Gulf of Mexico. Considerable amounts of C3 and Cq alkylbenzenes (100-400 pg/L per component) were evident. Estimations of the amount of VLH discharged into the outer continental shelf of Louisiana and upper Texas from these two discharges were made for underwater hydrocarbon venting, 400 X 106-1200 X lo6 g/yr, and for formation waters, 750 X 106-1100 X 106 dyr. Present address: Exxon Production Research Co.,P.O. Box 2189, Houston. TX 77001. 0013-936X/81/0915-0917$01.25/0
@
Introduction
Hydrocarbon venting and formation water discharges from offshore production operations are two major sources of gaseous hydrocarbons (Cl-Cj aliphatics) and volatile liquid hydrocarbons (aliphatics, C6-Cl4; aromatics-benzene, naphthalene, and alkylbenzenes; cycloalkanes alkylcyclopentanes, and alkylcyclohexanes) in coastal waters ( I 1, particularly in the Gulf of Mexico, where almost two-thirds of the world’s offshore production platforms and one-eighth of the offshore production occurred during 1976 (2). Compositional characterization of the gaseous hydrocarbon fraction of these discharges has been sufficiently studied ( I , 3 , 4 ) . However, little information is available on the full characterization of the volatile liquid hydrocarbon (VLH) fraction of these offshore discharges. Only the major components of the VLH fraction (benzene and toluene) have been identified by some authors ( 3 , 4 ) . From an environmental standpoint, this fraction of hydrocarbons is the most important because of its higher immediate toxicity to organisms than the other fractions of hydrocarbons (5-7). This paper presents the results of the determination of VLHs in underwater vented and formation water discharges sampled from production platforms in the Gulf of Mexico.
1981 American Chemical Society
Volume 15, Number 8, August
1981
917
