Kinetics and Mechanism of the Decomposition of N-Chloroalanine in Aqueous Solution William D. Stanbro” and Wendy D. Smith’ The Johns Hopkins University, Applied Physics Laboratory, Johns Hopkins Road, Laurel, Md. 20810
The kinetics and mechanism of the decomposition of N chloroalanine (a rapidly formed chlorination product of the common amino acid, alanine) have been investigated. The decomposition products have been shown to be acetaldehyde, ammonia, carbon dioxide, and chloride ion or, depending on pH, pyruvic acid, ammonia, and chloride ion. The kinetics were studied by UV-visible spectrophotometry and by the DPD-FAS titrimetric technique. The reaction is first order in N-chloroalanine and independent of alanine concentration over three orders of magnitude of concentration. There is a complex dependence on p H (however, there is no p H dependence over the range in most natural water, p H 5 to 9). The half-life of the reaction in the 5-9 p H range is 46 min at 25 “C. The rate constant shows a marked temperature dependence a t all p H values, changing by a factor of more than three for a 10 “ C temperature change. The p H dependence of the first-order rate constant and the molar excitation coefficient have been used to infer the ionic state of the reactants and the mechanism. Our society’s increased concern over the quality of the environment requires an understanding of the fate of compounds generated by the addition of chemicals to natural systems. One area of considerable interest is the stability of chlorine compounds produced when chlorine is added to sewage or cooling water to control undesirable microbiota and the product is then discharged into rivers, lakes, or estuaries ( I ) . It is desirable to be able to predict the lifetimes of these potentially toxic materials under various environmental conditions. In this report, we examine, for a range of concentrations, pHs, and temperatures, the kinetics and mechanism of the decomposition of N-chloroalanine,one of the products of chlorination. This compound was selected because of the high frequency of occurrence of alanine among the amino acids found in natural waters (2, 3 ) and the rapid rate of chlorination of alanine relative to other naturally occurring nitrogen compounds ( 4 ) . The interaction of chlorine with amino acids has been studied by several investigators. An extensive review of halamine chemistry was published by Kovacic et al. ( 5 ) .The decomposition of N-chloro-a-amino acids was first discussed by Langheld (6) who noted that a-amino acids react with hypochlorous acid salts in the same manner as they do with amines to form monochlorinated or dichlorinated derivatives. Subsequently, the chloroamino acids can decompose, and this decomposition is affected by the nature of the a-amino group. If the group is primary, then aldehydes or ketones, ammonia, carbonic acid, and sodium chloride are formed. If the group is secondary, the same products are obtained except that the corresponding amine replaces the ammonia product. Langheld formulated the following mechanism (Scheme I) for the decomposition, assuming an intermediate imine formation. These products were confirmed in studies by Dakin et al. ( 7 ) , Dakin ( 8 ) ,and Wright (9). In their study of the synthesis of indoleacetic acid by glutamic acid conversion, Fox and Bullock (10) formulated a mechanism for the decomposition of N-chloro-tu-amino acids
’
Present address, SCM Corporation, P.O. Box 389, ,Jacksonville, Fla. 332201. 446
Environmental Science & Technology
Scheme I
Scheme I I (I!
0 il
(II!
0
H
R-C-COO-
//
”;:,.-
/Ill)
C
H+
H ‘0
R-Ck
I
I
HN
‘cl’
r
R-CENH
t
CO? + HCI
\ i :
H 2 0 , H+
R-C-H
IlVI
R-C-CO
R-C-COOH
Scheme Ill
that expands the original one suggested by Langheld. Their mechanism is presented in Scheme 11. If a sodium salt of an N-chloro-tu-amino acid (I) is acidified, it yields a free acid. The acid can form a six-membered ring (11) that can decompose after a shift in electrons to an intermediate imine (111) and finally to an aldehyde (IV). On the other hand, if a proton is removed from the initial N-chloro-tu-amino acid salt, a carbanion (V) is produced that immediately stabilizes itself by losing the chloride ion to form the imino intermediate (VI). The final product in this case is an cy-keto acid (VII). The amount of an (\-ketoacid formed is related to the basicity of the reaction mixture. In their bactericidal studies of chlorine, Ingols et al. ( 1 1 ) found that pyruvic acid as well as acetaldehyde are formed upon the decomposition of N-chloroalanine a t p H 8. Dakin ( 8 ) ,Wright (9),and Pereira et al. (12) have investigated the decomposition products of N,N- dichloro-tu-amino acids. Their results indicate rapid formation of carbon dioxide, chloride ion, and the corresponding nitrile compound (Scheme 111). The UV absorption spectroscopy of N-chloroamino acids was examined by Metcalf (13).Morris ( 4 ) used UV absorption to study the formation of these species following methods used in earlier studies of chloroamines by Morris ( 4 ) , Weil and Morris (14, 1 5 ) , Czech et al. (16),and Kleinberg e t al. (17). Nikol’skii et al. (18)used this method to examine the stability of hypochlorous acid and hypochlorite solutions. Kantouch and Abdel-Fattah (19), Stankovic (20), and Stankovic and Vasatko (21) have investigated the rates of chloroamino acid decomposition. However, they have neither
0013-936X/79/0913-0446$01 .OO/O
@ 1979 American Chemical Society
differentiated between the decomposition of N-chloro-aamino acids and N,N-dichloro-a-amino acids, nor made an investigation of the kinetic rate law. Experimenta 1 The kinetics of N-chloroalanine decomposition were examined in two different concentration ranges. The first was a high concentration range (2.9 X lop2to 7.1 X M alanine) that allowed the use of UV spectroscopy to monitor the concentration of N-chloroalanine, The second was a low concentration range (7.1 X M alanine) that is more representative of natural systems. The low-range studies used the FAS-DPD titrimetric method. Determination of N-Chloroalanine Decomposition Rate. High Concentration Range. The decomposition rate of N-chloroalaniiie was determined by preparing a reaction mixture of alanine and chlorine and observing the decrease in its absorbance a t 250 nm on a Varian Techtron UV-visible spectrophotometer (Model 635) a t a slit width of 0.2 nm. Reaction mixture cuncentration, pH, and temperature were altered to obtain decomposition rates for various conditions. Solutions were prepared with 0.40 m NaCl to control ionic strength. The water used was first distilled and then passed through a Barns tead research-type demineralizer. Alanine (Eastman Chemical Co.) was prepared to yield a final reaction to 7.1 X 10-* mixture concentration in the range 2.9 X M. The chlorine solution was made by bubbling chlorine gas (Matheson Ultra-High Purity) through a 0.01 N NaOH solution and determining its concentration by iodometric titration (22).The final chlorine concentration in the reaction mixture was chosen to ensure a t least a 15:l molar ratio of alanine to chlorine in order to prevent the formation of N,N-dichloroalanine. The pH of the reaction mixture was corrected using NaOH or HCl. The temperature was controlled to hO.1 “C by a Masterline Forma Scientific 2095 bath and circulator that was connected to the spectrophotometer. The temperature in the spectrophotometer sample compartment was monitored with a Varian Techtron temperature readout. When thle alanine and chlorine solutions were combined in the reaction mixture, the time was noted and the solution was placed in a 1.0-cm quartz cell in the spectrophotometer. Absorbance readings were recorded on a stripchart recorder. Lou: Concentration Range. The low concentration range reaction mixture was prepared in the same manner as the high concentration case except that 7.1 X 10-5 M alanine and proportionately less chlorine were used. All measurements were done in 0.4Cl m NaCl as they were in the high concentration range and the molar ratio of alanine to chlorine was 6:l or greater. Temperature was controlled by immersing the flask containing the reaction mixture in a constant temperature bath. Periodically, aliquots were removed for analysis using the FAS-DPD titrimetric technique (22). The aliquot was analyzed for both “free residual” and “combined residual”. At no time was any free residual observed. Determination of N-Chloroalanine Extinction Coefficients. The solutions were prepared the same way as for the high concentration decomposition rate determination. After the reactants were mixed, the UV absorbance was read on the spectrophotometer a t 250 nm a t the same time that an iodometric titration (22) was performed to determine the N chloroalanine concentration. Reaction Product Analysis. The formation of pyruvic acid and acetaldehyde during the decomposition reaction was confirmed by observation of their UV absorbance in a quartz cell having IO-cm path length, following dechlorination with sodium sulfite. The formation of pyruvic acid was not observed below a p H of approximately 7. Aliquots of reaction mixture were also analyzed for acetaldehyde and acetonitrile
0 50
0
(I)
030
m
a
i
0 20
0
00006
00002
00010
00014
00018
CON C E NT RAT I 0 N (molesA i t e r ) Figure 1. KChloroalanine absorbance vs. concentration at pH 6.53 in phosphate buffer
-- 5001 I
- c
1
I
I
I
I
I
I
I
1
I
1 Lu
8
o 0i
I
I
I
I
I
I
I
I
1
2
3
4
5
6
7
8
I 9
10
PH
Figure 2. N-Chloroalaninemolar extinction coefficient vs. pH
by gas chromatography using a Chromasorb 105 column on a Tracor 560 gas chromatograph equipped with a flame ionization detector. Some acetonitrile was occasionally observed immediately after the reaction was started, but it did not increase with time. Apparently it was due to the transitory formation of N,N-dichloroalanine a t the time the chlorine solution was added and before complete mixing was achieved. Yields of acetaldehyde and ammonia were consistent with the amount of chlorine used from the time the initial N-chloroalanine was measured (-1 min after mixing) to the time the reaction was stopped by dechlorination with sodium sulfite. Ammonia was determined by nesslerization following distillation of a dechlorinated aliquot of reaction mixture (22),or by use of the phenate method (23) on a neutralized, dechlorinated sample. Results and Discussion The concentration of N-chloroalanine was monitored by measuring the absorbance a t 250 nm. N-Chloroalanine follows Beer’s law a t constant pH as does monochloramine ( 2 7). This is illustrated in Figure 1. However, as shown in Figure 2 , the molar extinction coefficient ( e ) shows a strong pH dependence. The molar extinction coefficients reported here are in agreement with Metcalf‘s measurements a t neutral pHs ( 2 3 ) . Figure 3 shows a representative change in the logarithm of N-chloroalanine absorption as a function of time. The decomposition of N-chloroalanine appears to be first order in N-chloroalanine under the conditions studied. Figure 4 shows a similar plot for the change in N-chloroalanine as determined by the DPD technique in the low concentration range. Tables I and I1 give the first-order rate constant at 25 “C as a function of pH, initial N-chloroalanine concentration, and alanine Concentration. The rate constant was calculated with a weighted least-squares computer routine, W S T A T . The Volume 13, Number 4, April 1979
447
10
P
w 0
2a 2 2
I
-
z
1
0.9 0.8 0.7 0.5 0.6
t-
0
10
20
30
0 40
50
60
70
TIME (minutes) Figure 3. Log of Kchloroalanine absorbance vs. time at 25 "C f 0.1,
0
1 10
1
1 20
1
30
1 60
1
50
40
J
70
TIME (minutes) Figure 4. Log of concentration of a-Kchloroalanine vs. time at 25 O C , low concentration range, pH 8.02
pH 8.62
Table 1. Observed First-Order Rate Constant for the Decomposition of N-Chloroalanine at 25 "C, High Concentration Range [alanine], moilL
7.1 X 7.1 X 7.1 X 7.1 X 7.1 X 7.1 X 7.1 X 7.1 X 7.1 X 7.1 X 7.1 X 7.1 X 7.1 X 5.7 X 5.7X 5.7X 5.7X 5.7X 4.3X 2.9 X 2.9X 2.9 X 2.9X 2.9X 2.9X a
lo-' lo-' lo-'
lo-' IO-' IO-'
lo-' lo-' lo-' lo-'
lo-' lo-' lo-'
lo-' lo-' lo-' lo-*
lo-' lo-'
lo-' IO-'
PH
8.62 8.35 8.15 6.51 5.59 4.72 3.99 3.50 3.39 3.20 2.76 2.65 1.25 8.21 7.31 3.20 2.99 2.71 3.00 8.21 7.65 3.00 2.74 2.72 2.41
initial [ N-chloroalanine], mol/L
3.74x 9.61x 3.74x 1.69x 1.14X 7.01X 1.31x 5.20 x 3.23 x 3.27x 3.81 x 3.50x 1.94x 3.38x 3.64x 4.00x 5.91 x 6.00x 3.64 x 3.38x 3.38x 3.44x 5.24x 2.38x 6.29x
10-3 10-4 10-4 10-3
loW3 10-3 10-4 10-4 10-4 10-4 10-4 10-3 10-4 10-4 10-4 10-4 10-4 10-4 10-4 10-4 10-4 10-4 10-4 10-4
kobsd f 10, a min-'
0.017f 0.00025 0.017f 0.00054 0.016f 0.00040 0.015f 0.00023 0.015f 0.00034 0.0078 f 0.00032 0.013 f 0.00044 0.017f 0.0014 0.017 f 0.0013 0.020f 0.00089 0.030f 0.0042 0.039f 0.0045 0.098f 0.0027 0.015f 0.00053 0.015f 0.0017 0.017 f 0.0021 0.017 f 0.0014 0.026f 0.00091 0.036 40.0047 0.016f 0.00052 0.016f 0.00056 0.024f 0.0029 0.029f 0.0017 0.041 f 0.0098 0.044 f 0.0029
inltlal
Environmental Science & Technology
[ N-chloroalanine],
[alanine], moilL
pn
7.1x 10-5 7.1 X 7.1x 10-5 7.1 X loW5 7.1 X lop5 7.1 X lop5 7.1 X lop5 7.1 X lop5 7.1 X loT5 7.1 X lop5 7.1 X 7.1 X 7.1 X lop5
9.00 8.02 7.35 6.78 6.40 6.38 5.90 5.80 5.46 4.36 3.00 2.99 2.39
a
mol/L
6.31X 7.44 x 6.20X 7.61 X 6.12X 5.75 x 5.86X 5.07X 4.79x 2.53 X 1.69 X 1.46 X 1.23x
10-6
loe6 10-6
10-6
10-5
kobsd f la, a
m1n-1
0.016f 0.0004 0.018f 0.0007 0.014f 0.001 0.014 f 0.0009 0.018f 0.0005 0.018f 0.0007 0.016f 0.0002 0.019f 0.001 0.017f 0.002 0.0072zk 0.002 0.044f 0.007 0.032f 0.002 0.002f 0.001
10 = one standard deviation.
Table 111. Observed First-Order Rate Constant as a Function of pH and Temperature, High Concentration Range
10 = one standard deviation
weighting factor was taken to be proportional to the square of the absorbance at the high concentration range and to the square of the volume of titrant a t the low concentration range (24).The results indicate a complex dependence of the rate constant upon pH, but none upon initial N-chloroalanine concentration or alanine Concentration. Of particular importance is the similarity of the results in the two concentration ranges. This allows the use of the more accurate and convenient high concentration techniques t o study the mechanism. Tables I11 and IV show the temperature dependence of the observed first-order rate constant a t different pHs. Again the low and high concentration data are consistent. Tables I through IV show that in most natural waters, which are in the 5-9 p H range, there will be little pH dependence of the rate 446
Table II. Observed First-Order Rate Constant for the Decomposition of N-Chloroalanine at 25 "C, Low Concentration Range
a
temp, OC
kobsd f 1 b, a
PH
2.99 3.12 3.90 4.39 8.04 8.19 8.35 8.51 2.99 3.11 3.81 4.98 8.30 8.35
34.7 34.7 34.7 34.7 34.7 34.7 34.7 34.7 19.2 19.2 19.2 19.2 19.2 19.2
0.059f 0.0026 0.067 f 0.0030 0.049f 0.0062 0.022f 0.0013 0.065f 0.0026 0.062f 0.0026 0.059 f 0.0018 0.064f 0.00032 0.022f 0.0023 0.013 f 0.0012 0.0062f 0.00045 0.0039 f 0.00029 0.0052f 0.00042 0.0053f 0.00034
m1n-1
1 u = one standard deviation.
constant. However, there is a very marked dependence on temperature at this range. The rate constant more than triples with a 10 "C rise in temperature.
Table IV. Observed First-Order-Rate Constant as a Function of pH at 34.7 "C, Low Concentration Range PH
kob.d f 10, a min'l
k&.d f 10, a rnin-1
PH
8.95 0.050 f 0.003 7.00 0.060 f 0.003 6.05 0.056 f 0.003 a 1a = one standard deviation.
AE = -26 600 f 1200 cal-mol-l-K-1 and In A = 40.8 f 1.9 where AE is the activation energy, A is the frequency factor in min-', and In indicates the natural logarithm. The pH dependence of the observed molar extinction coefficient and the observed high concentration range firstorder rate constant are shown in Figures 2 and 5, respectively. In the discussion below, we will suggest a mechanism consistent with the above patterns. We will then use this model mechanism to obtain an expression for the observed firstorder rate constants as a function of pH. N-Chloroalanine, like alanine, can exist in a number of ionic forms (VIII to XI), due to the presence of both an amino group and a carboxyl group. However, the relative acidities of these functional groups, as compared to alanine, are perturbed by the presence of the chlorine atom on the amino group. CH3
CH3
0 //O
H-N-C-C
I
H -N - C - C ,
I
ci
H-N-C-C
I ; lo 0
(XI
I OH
H CH3 0
CI
4
I
H
Cl
H CH3
n 0.10 a '-
-
OE
COMPUTED
G L 0.08
0.023 f 0.004 0.036 f 0.002
4.40 3.99
The average observed rate constants in the 5-9 pH range at the temperatures examined were used to obtain a leastsquares fit to the Arrhenius equation. The results were:
I
W-
01
I 4 AI
0
H-N-C-C
I
CI
OH
1x11
01 0
I 1
I
I
,
2
3
4
I
1
I
I
1
6
7
8
9
10
o 5 PH
Flgure 5. NChloroalanine observed first-order rate constant vs. pH at 25 "C, high concentration range
since they can be interconverted without loss or addition of a hydrogen ion. At even lower pHs, XI would be the dominant species. We suggest that since the drop in the rate constant takes place a t a pH of around 5 and the drop in the extinction coefficient due to protonation of the amino group takes place below a pH of 4, species IX greatly predominates over X. This assumes that the change in ionic form is responsible for the variation in reactivity. The fact that the equilibrium between IX and X should lie in the direction of IX is not surprising in light of the effect of chlorine substitution on the basicity of other amines (15). Also, the protonation of the carboxyl group occurs at a higher pH than in the case of alanine because of the smaller electron-withdrawing capability of the chlorine atom compared with the positively charged amino group of normal alanine. This is confirmed by the relative values of K , a t 25 "C for a-chloropropionic acid, P-chloropropionic acid, and ala1.04 X and 4.49 x 10-3, renine, which are 1.47 X spectively (25). We suggest the following mechanisms (Scheme IV) for the decomposition of species VIII and IX to form "3, CO2, C1-, and acetaldehyde. The imine is then hydrolyzed to ammonia and acetaldehyde (22) (Scheme V). The reaction should be viewed as a spontaneous decarboxylation followed by rapid hydrolysis of the imine. The decarboxylation of VI11 is probably by an SEI mechanism (26) where the carbanion is stabilized by the chlorine atom through incipient chloride ion formation. Species IX decomposes by a cyclic mechanism analogous to the decarboxylation of a P-keto acid in which the chlorine fulfills the role of the a-carbonyl oxygen (26,27).This Scheme I V
Examining the pH dependence of the molar extinction coefficient, we see a pattern similar to the one found by Weil and Morris (15) for N-chlorodimethylamine and N-chlorodiethylamine. They attributed the drop in the molar extinction coefficient with decreasing pH to the protonation of the amines. The molar extinction coefficient data do not indicate other changes that could be assigned to the protonation of the carboxyl group, and we may conclude that this has no impact. The observed first-order rate constant undergoes two changes as one moves from neutral toward lower pH. First, at a pH above the transition point in the extinction coefficient, there is a marked decrease in the rate constant. Next, at the same point as the extinction coefficient transition, there is a sharp and continued increase in the rate constant. This pattern is seen at all temperature and concentration ranges in the data in Tables I through IV. Species VI11 is certainly the predominant form at high pH. Addition of a proton to VI11 will form either IX or X, which must exist (at equilibrium) in a constant ratio to each other
H-N-C-c
--+
I H-N' H
Scheme V
0 OH? H20
H-C-CH3
II
N-H
H-C-CH3
I
0 N-H
-I OH
I
H-C-CH3
H-C
tNH3 'CH3
"2
Volume 13, Number 4, April 1979
449
mechanism, which is the one proposed by Fox and Bullock (IO),is only applicable over a short pH range due to the absence of an appreciable concentration of the neutral molecule a t most pHs. If we take VI11 as a standard for comparison, we can explain the slower relative rate of I X by the requirement to form the cyclic intermediate. This is difficult because of the weakness of the chlorine-hydrogen bond. As one goes to still lower pH, one sees a rapid rise in the first-order rate constant. The pH of the start of the rise is correlated with the pH of the start of the drop in the molar extinction coefficient. Interpreting this as the formation of XI, it would appear that X I is less stable than VI11 or IX. However, the continued rise in the rate constant cannot be explained solely in terms of the formation of XI. It appears that there is also a mechanism involving acid catalysis in this pH range. The increased instability of XI vs. VI11 and I X may be related to the large electron-withdrawing capability of both the chlorine atom and the positively charged amino group. The acid-catalyzed mechanism may involve protonation of the chlorine atom to allow the neutral HCl molecule to be the leaving group. On the basis of the above discussion, we propose that the observed first-order rate constant can be written in terms of the fundamental processes as: kobsd
=
kl[VIII]
+ k2[IX] + k3[XI] + k.ia~+[XI] (1) [VIII] + [IX] + [XI]
where kobsd is the observed first-order rate constant, the brackets indicate units of moles per liter, and aH+ is the hydrogen ion activity. In addition, we have assumed that ks[X]
