Environ. Sci. Technol. 1904, 18, 348-353
LaVoie, R. J.; Tulley-Freiler, L.; Bedenko, V.; Hoffman, D. Mutat. Res. 1983, 116, 91-102. LaVoie, R.J.; Tulley-Freiler, L.; Bedenko, V.; Hoffman, D. Mutat. Res. 1981, 91, 167-176. Pool, B. L.;Lin, P. Z. Food Chem. Toxicol. 1982, 77, 383-391. Garrett, W. N.;Coppock, R. H. In "Agricultural Residue Management: A Focus on Rice Straw"; Residue manage-
ment Task force, University of California: Davis, CA, 1981; pp 1-85. (35) Rappaport, S.M.; Wang, Y. Y.; Wei, E. T. Environ. Sci. Technol. 1980, 14, 1505-1509.
Received for review June 3, 1983. Revised manuscript received October 3,1983. Accepted October 28, 1983. This project was funded by the California Rice Research Board.
Kinetics of Monochloramine Oxidation of N,N-Diethyl-p -phenylenediamine Howard E. Moore, Maria J. Garmendla,+ and William J. Cooper* Department of Physical Sciences and Drinking Water Research Center, Florida International University, Miami, Florida 33199
The reaction kinetics of monochloramine (NH2C1)and N,N-diethyl-p-phenylenediamine(DPD) were studied. The method of initial rates was used to study the rate of formation of a colored intermediate product and the rate of fading of that product to colorless end products. Equilibrium constants were determined for protonated and nonprotonated DPD and the intermediate products. The reaction is first order with respect to both NH2Cland DPD and showed a complex order with respect to pH. This study confirmed that NH2Cl oxidizes DPD to the colored intermediate a t a rate of 5.6 and 6.0% of the NH2Cl concentration, in the first minute, a t 25 OC and 5 X lom6M (3.5 mg as C12/L)and 1 X lo4 M (7.0 mg as Cl,/L) NH2C1, respectively. Chlorine is used extensively in water treatment primarily for disinfection. Chlorine reacts with ammonia to give a complex series of reactions, the initial product of the reaction being monochloramine (1-3). Monochloramine, a weaker disinfectant than chlorine (4-6), is used as an alternative disinfectant to minimize the formation of trihalomethanes during drinking water treatment (7-11). It is the predominant chlorine species in wastewater treated with chlorine. Analytical determinations for chlorine species in water should be able to differentiate between free available chlorine (HOCl/OCl-) and combined chlorine (most commonly NHzCl and NHCl2) (12). One of the most commonly used colorimetric procedures for chlorine determinations is the DPD test, which uses N,N-diethyl-p-phenylenediamine (DPD) as the indicator. It has been shown operationally that DPD reacts directly with both free chlorine and monochloramine (13-17). The reaction between DPD and free chlorine is very rapid while the reaction with monochloramine is slower. The DPD test procedure shows a positive interference in the free chlorine measurement when monochloramine is present, the extent of the interference being proportional to the monochloramine concentration. Potentiometric studies have shown that p-phenylenediamine and its alkylated derivatives undergo acid-base reactions (18). They also undergo oxidation reactions to yield colored intermediate oxidation products (19-23) Figure 1 details the sequential oxidation reactions and acid-base equilibria for N,N-diethyl-p-phenylenediamine. The intermediate products (Wurster's dyes) are semiquinoid cationic radicals (20,23,24). In the case of aromatic diamines, the quinonoid structures (VII, VIII, and IX) are relatively unstable forms (It?), while the semiquinoids (IV, V, and VI) are surprisingly long-lived because of the increased resonance stabilization. Present address: La Salle High School, Miami, FL 33133. 348
Environ. Sci. Technol., Vol. 18, No. 5, 1984
Alkylphenylenediamine cation radicals exhibit distinct absorption maxima in the red region and less distinct maxima at longer wavelengths. Each alkyl group substituted on the amino group results in a bathochromic shift of the absorption maxima (18). The colored oxidation product of DPD is unstable (15). This instability leads to fading of the solution produced in the colorimetric procedure and can result in an anomalous measurement. The generalized reaction can be written, as shown a t the top of Figure 1. A, would be the unreacted N,N-diethyl-pphenylenediamine, Az the colored intermediates, and A3 the final colorless reaction product. kl,obsdis the overall rate constant for the appearance of the colored intermediates, while k2,0bsdis the overall rate constant for the fading reaction. The results of potentiometric titrations performed on p-phenylenediamine and derivatives show that, when normal potentials are plotted against pH, a straight line is obtained with a slope of 0.06 V/pH unit (20). This has been taken as an indication that the oxidized and reduced forms differ not only by one electron but also by a hydrogen atom. The slope of such a line is constant in a pH range that is characteristic of each compound. Outside such pH ranges the potentials drift. This drifting is caused by a change in the dissociation state of the diamine being occurs titrated and for N,N-diethyl-p-phenylenediamine at pH below 3 and above 8 (20). When potentiometric titrations are performed rather rapidly or extended past an initial end point, the color of the intermediate radical disappears (20),apparently due to oxidation of the intermediate radical, to a substance that does not absorb in the visible region. This study reports the kinetics of the oxidation of N,N-diethyl-p-phenylenediaminewith monochloramine in aqueous solution. The reaction order was studied as a function of DPD and NH2Cl concentrations and pH by using the method of initial reaction rates. Studies of the equilibria between the DPD species in solution were conducted as a function of pH by the spectrophotometric methods described by Miller and Wilkins (25). The implication of these reactions with respect to the analytical determination of chlorine species in water is discussed. Reagents and Instrumentation
Sodium hypochlorite and monochloramine solutions were prepared in chlorine demand free water and stored in demand free glassware as described elsewhere (15). Buffers of varying pH were prepared by adding 0.2 M K2HP04to 0.2 M KH2P04to chlorine demand free water until the desired pH was obtained. N,N-Diethyl-pphenylenediamine (DPD) solutions are extremely air sensitive (15). The solutions prepared with chlorine demand free water turned pink after standing a few minutes.
00 13-936X/84/0918-0348$0 1.50/0
0 1984 American Chemical Society
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,
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1
'+
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I
, \ H t i
ti
H
11
/3
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H
/N\
N
H
1
I,
WAVELWTH tnml
Flgure 2. Absorption spectra of 9.6 X M N,N-diethyl-pphenylenediamine solution. (1) pH 1.9; (2) pH 7.0; (3) pH 12.0.
Table I. Absorbance of 9.6X M DPD at pH 2.8-3.8 at 240 nm absorb- [ DPDH' ] X [ DPDH,Z+3 X pH ancea 105,M 105,M Figure 1. Oxidation and acid-base equilibrium of N ,N-diethyl-pphenylenediamine.
Therefore, solutions of the desired concentrations were prepared daily by dissolving the appropriate amount of DPD oxalate (Eastman Chemical 7102) in water degassed with Nz.Solutions were kept in stoppered brown bottles flushed with N2. All reactions were run at 25 f 0.2 OC and at an ionic strength of 0.28 f 0.02.
N,N-Diethyl-p-phenylenediamineDissociation Constants
The dissociation constants (K, and K z , in Figure 1)of N,N-diethyl-p-p-phenylenediamine (DPD) were determined from the variation of absorption spectra with pH (26-28). The effect of pH on the absorbance of a 9.6 X M DPD solution is shown in Figure 2. To determine the molar absorptivity of the DPDH', five 9.6 X M DPD solutions were buffered at pH 3.0-5.8, and the absorbance at 240 nm was obtained. It was determined that the absorbance at 240 nm and pH 4 provided the best estimate of the molar absorptivity of DPDH+. At this pH the contributions to the 240-nm peak of the DPD and DPDHz2+were assumed to be negligible. The validity of this assumption can be shown from the values for molar absorptivities and equilibrium constants obtained below. At pH 4 the [DPD]/[DPDH+] would be approximately 3.5 X and the contribution to the total molar absorptivity at pH 4 would be less than 1M-l cm-'. If all of the absorption at pH 1.9 were assumed to be due to DPDH+, then the maximum contribution to the total molar absorptivity at pH 4 of [DPDH+] would be 90 M-' cm-l. This is approximately the same value as the error obtained; thus, the assumption appears valid. Five DPD solutions were prepared at pH 4, and from the slope of absorbance vs. DPDH+ concentration a molar absorptivity of 21 258 f 84 M-l cm-l was calculated. To determine the molar absorptivity of the DPD, a series of
2.8 3.0 3.3 3.8 a
0.846 1.065 1.460 1.787
3.98 5.01 6.77 8.41
5.63 4.59 2.83
K,,M 1.12 x 10-3
1.09 x 10-3 1.19 x 10-3 1.12 x 10-3 1.19 Average of triplicates, u < 0.001 absorbance unit
4.8 X M DPD solutions were prepared, and the absorbance was measured at 240 nm. At pH 12 the only DPD species in solution would be the nonprotonated form, and from the slope of absorbance vs. DPD concentration the molar absorptivity was determined as 18 335 f 153 M-l cm-l. To determine the value of K1, 9.6 X 10" M solutions of DPD were prepared at pH 2.8, 3.0,3.3, and 3.8. The absorbance of each solution was measured at 240 nm by using the proper phosphate buffer as a blank (Table I). DPDH+ concentrations were determined from the measured absorbances, and the molar absorptivity of DPD in acidic media was calculated in this study. The DPDHZ+ concentrations were calculated by difference between DPDbd and DPDH+. From the calculated DPDH+ and DPDHz2+ concentrations, K , = (1.13 f 0.04) X M. At higher pH values, DPD undergoes a second dissociation process as shown in Figure 1, for which the equilibrium constant is K2. The equilibrium constant corresponding to this process was determined by using the method described by Miller and Wilkins (25). Between pH 6.8 and pH 10.9 the DPD spectrum is assumed to be due to the contributions of DPD both in its monoprotonated form, DPDH+, and the nonprotonated form, DPD. At any specific pH value EA[A] 4- EB[BI = Eobsd([A] P I ) (1) where EA, E g , and Eobd represent the molar absorptivities of DPDH+, DPD, and the mixture of both at 240 nm. Nine 9.6 X M DPD solutions were prepared from pH 6.8 to pH 10.9. The absorbance of each solution was measured at 240 nm (Table 11). From the measured abwas calculated for each solution. By sorbance, the ICobsd use of EA and Eg, molar absorptivities calculated for DPDH+ and DPD, the dissociation constant K 2 was calEnviron. Sci. Technol., Vol. 18, No. 5, 1984
349
Table 11. Absorbance of 9.6 X pH 6.8-10.9 a t 240 nm pH
absorbancen
Eobsd
6.8 7.0 7.4 7.9 8.0 8.3 8.7 9.4 10.9
2.035 2.030 2.013 1.981 1.966 1.935 1.856 1.807 1.762
2 1 200 2 1 150 20 950 20 640 20480 20 160 1 9 330 1 8 840 1 8 350
Average of triplicates, u
M DPD at K,, M 4.26 3.84 4.70 3.40 3.63 3.01 3.86 1.90 2.50
x
x x
x 10-9 x
x x x x
10-9 10-9 10-9 10-9 < 0.001 absorbance unit.
d
!
I
I
I
i
4.6
5.0
5.4
5.8
-Log INITIAL DPD CONCENTRATION
Flgure 4. Effect of initial DPD concentration on initial rate of reaction as a function of pH. The values given for each symbol are pH, slope, and r , respectively: (A) 5.4, 1.0058, 0.9997; (0)5.8, 0.9283, 0.9957; ( 0 )6.2, 1.0177, 0.9993; (0)7.0, 1.0347, 0.9969.
7.41
43
4.7
-LOG INITIAL NHICl
5.1
5.5
CONCENTRATION
Flgure 3. Effect of initial monochloramine concentration on initial rate of reaction as a function of pH. The values given for each symbol are pH, slope, and r , respectively: A 4.5, 1.049, 0.9993; (0)5.4, 1.011, 0.9999; (0)5.8, 0.986, 0.9985; (0)6.2, 1.006, 0.9999; (m) 7.2, 1.048, 0.9992.
culated. The value obtained for K 2 was (3.50 f 0.87) X lo* M. The errors in Kl and K2 are one standard deviation errors for the data shown in Tables I and 11. These values of K1 and K 2 are consistent with the values cited for potentials a t which drifting occurs during potentiometric titrations (20). Determination of the Order of the Initial Reaction The order of the initial reaction was determined by varying, in turn, the monochloramine concentration, the DPD concentration, and the pH, while holding the other two reactants constant. Kinetic runs were repeated from 1-4 times at each pH and each DPD or monochloramine concentration. For duplicate runs the initial .reaction rate varied by less that 2% in all cases. The concentration of monochloramine was varied from 4.51 X lo4 to 1.15 X M, while the DPD concentration was held constant at equal to or greater than a 10-fold excess. Figure 3 shows a plot of the negative log of the initial NH&l concentration vs. the negative log of initial rate as a function of pH. The order of the reaction with respect to monochloramine was 1.02 f 0.03. 350
Environ. Sci. Technol., Vol. 18, No. 5, 1984
'1
71 40
I
50
I
60 Log
,
70
J
80
H30+ COnCenlrallOn
Figure 5. Effect of pH on initial rate of reaction of DPD and monoChloramine.
The concentration of DPD was varied from 3.00 X lo4 to 1.20 X M, while the monochloramine concentration was held constant at equal to or greater than a 10-fold excess. Figure 4 shows a plot of the negative log of the initial DPD concentration vs. the negative log of initial rate as a function of pH. The order of the reaction with respect to DPD was 0.996 f 0.047. The effect of pH was studied by varying the pH from 4.0 to 8.5. The concentration of DPD was varied from 2.40 X to 5.50 X M with the concentration of monochloramine in 10-fold excess. The results of these experiments are shown in Figure 5. The shape of this curve strongly indicates that at least one of the reactants is involved in acid-base equilibrium (25). The overall observed rate constant, kl,obsd, at a constant pH follows the expression [-d([DPDHZ2'] + [DPDH'] + [DPD])]/dt = k,,obsd([DPDH22+]+ [DPDH'] + [DPD])([NH2C1]) (2) According to Wilkins (29),from the equilibrium constant
'61
"'I
I
70
TIME IMinulesI
1
I
Figure 7. Growth and decay of the Oxidation product of DPD by monochloramine at 515 nm: (-) observed; (0)calculated. [DPD] = M, and pH 5.5. 6.8 X M, [NH2CI] = 6.8 X
[DPDHZ2+]=
[DPDH+][H+] (3)
K1
(4)
It can then be demonstrated that the order of reaction with respect to H+ concentration follows the expression k3[H+I2+ k4K,[H+] + k5K1K2
(5) [H+I2 Kl[H+] K1K2 The rate constants kat k4, and k5 (see Figure 1) were calculated by solving sets of simultaneous equations in the form of eq 5 by using K1 and K2values obtained above and kl,oMdetermined as a function of pH. The results of these calculations are k3 = 1.14, k4 = 10.4, and k5 = 0.68 M-' s-l. The data are summarized in Figure 6. The oxidation product of DPD is probably a cationic radical of the type known as Wurster's salts. The proposed mechanism involves the loss of one electron and hydrogen atom from any of the three acid-base-related species of DPD to form a cation radical (Figure 1). In moderately acidic solutions the free radical that exists is primarily V. The stability of this substance is reported to be greatest between pH 4.5 and pH 6.0 (20) and is produced directly from I, the predominant diamine at pH 7. The increased resonance stabilization of the product could explain why k4, the rate constant of I1 to V, is approximately 10 times larger than k, or k5, the rate constants of I and I11 to IV and VI, respectively. When the radical was prepared in moderately acidic solutions at pH 5.0 with no excess of monochloramine and kept under N2 atmosphere, the product was stable for approximately 3 days. After 3 days the intensity of the color diminished. The solution turned violet, but the absorption maxima at 515 and 554 nm were still present. kl,obsd
=
+
+
Fading Studies The initial reaction between DPD and monochloramine leads to the formation of the red products. This reaction is then followed by subsequent steps, resulting in further oxidation or dimerization. The products of these subsequent steps depend upon the conditions under which the reaction is carried out. In the presence of excess monochloramine, the red color decreases with time until the
solution becomes colorless. An example of the growth and decay of the color is shown in Figure 7. The progress of the fading process was followed by recording the change in the absorbance with time at 515 nm. Three separate sets of solutions were prepared. In the first experiment, 5.78 X lov5M DPD solutions were mixed with 5.78 X M monochloramine solutions. Kinetic runs were made at pH 4.5,5.2,and 5.5. In a second experiment M monoM DPD was mixed with 6.8 X 6.8 X chloramine at pH 5.5, 5.8, 5.8, 6.5, 7.0, and 7.4. A third experiment consisted of 2.64 X M DPD solutions and 2.63 X loT4M monochloramine at pH 8.3. When absorbance was measured as a function of time, the time necessary to reach the maximum absorbance, tmax,obsd, was determined. Determination of Equilibrium and Rate Constants for the Fading Reaction The equilibrium constants for the products of the first oxidation and the rate constants for the fading of the red color were evaluated by considering the reaction as a series of pseudo-first-order consecutive reactions represented by the equations at the top of Figure 1,where the change in concentration with time of A2 is (25) dIA21/dt = kl,obsd[All - k2,0bsd[A21 (6) [A,] and [A2] represent the concentrations of the initial reactant and intermediate product, respectively. It can be shown that (7)
where the subscripts 0 and t refer to concentrations initially and at some specified later time. Equation 7 assumes [A,], = 0. If kl,omis smaller than k2,0m,the concentration of Az will pass through a maximum, which is given by tm, =
1 IZ2,0bsd-kl,obsd
k2,0bsd
In kl,obsd
Obtaining the t, from the absorbance vs. time traces in the spectrophotometer and using eq 8, the value of k2,0bsd was obtained at various hydrogen ion concentrations. These data appear in Figure 8. The variation oft, with pH indicates that the mechanism for the formation of the quinoid compound involves three acid-base-related species of the free radical and that the product is formed by the loss of a hydrogen atom from the original semiquinoid-like compound. Environ. Sci. Technol., Vol. 18, No. 5, 1984 351
Table 111. Summary of Monochloramine Interference in the Measurement of Free Chlorine Using Methods Employing N,N-Diethyl-p-phenylenediamine interference, %/min temp, “C ref 2.7 3.6 3.7
0 3
40
50
IO
60
80
4.5 4.5 6.0 8.0
,
90
PH
Flgure 8. Effect of pH on the fading rate constant: (-) calculated; (0) experimental.
6.0
The expression for the pH dependency of kz,obsdwas obtained in the same way as kl,obsd,the observed rate constant for the first step of the oxidation process. Therefore, k2,0bad as a function of pH will be given by the expression k2,0bad
=
+ k8K3K4 [H+I2+ K3[H+]+ K3K4
k6[H+I2+ k,K3[H+]
(9)
where K3 and K4 are acid-base equilibrium constants and k6,k,, and k8 are the rate constants for the oxidation of IV, V, and VI to VII, VIII, and IX, respectively. Using for k2,0bsdthe values obtained from the experimentally measured t,, values, applying eq 9, and solving a set of simultaneous equations, the values of KB,K4, k,, k,, and k8 were calculated. The results obtained are K 3 = 1.03 X M, K4 = 1.58 X M, k6 = 0.102, k, = 0.8, and k8 = 0.152 M-l s-l. By use of all of the rate constants and equilibrium constants obtained above for the formation and fading of the colored oxidation product of DPD, the calculated absorbance vs. time curve shown in Figure 7 was obtained. The difference between the calculated values and observed decay curve in Figure 7 may be due to the dimerization of the free radical species, thus increasing the apparent rate of fading. Dimerization Reaction
Dimerization of semiquinoid amino radicals has been observed (30, 31). Since dimerization is a bimolecular reaction, it is concentration dependent. In general, the dimerization of semiquinoid radicals is much less frequent in ionization states where equivalent resonance forms prevail (19). The formation of the dimer 4,4’-bis(dimethy1amino)azobenzene by oxidation of concentrated solutions of N,N’-dimethyl-p-phenylenediaminewith monochloramine has been reported (32). If oxidation of the DPD by monochloramine was carried out by allowing M DPD buffered at pH 7.0 to react with lo4 M monochloramine and the mixed solutions were kept stoppered and under N2 atmosphere, a black solid formed after 1 day. This solid product was dissolved in ether and was purified by recrystallization from ether solutions. The residue was analyzed by mass spectrometry. The mass spectrum obtained is consistent with the dimer of N,N-diethyl-p-phenylenediamine, with m / z = 324 for the molecular ion. Asymmetric molecular cleavage at the diazo nitrogen-aromatic carbon bond results in a peak at m / z = 177 and also m / z = 149. Monochloramine Interference in Free Chlorine Measurement
The practical significance of the present study is to provide the data necessary to quantitatively determine the 352
5.6
Envlron. Scl. Technol., Vol. 18, No. 5, 1984
18
25 unspecified 20 20 35 unspecified Monochloramine a t 5 x 25 Monochloramine at 1 x 25
33 37 35 34 15 16 36 M this study M this study
positive interference that monochloramine can have on a free available chlorine measurement with test procedures employing N,N-diethyl-p-phenylenediamine.This monochloramine interference has been reported by several investigators and is summarized in Table 111. On the basis of the rate constants determined in this study the interference of monochloramine can be calculated at any given pH. All of the present experiments were conducted at 25 “C. The pH recommended for the colorimetric DPD procedure is 6.2-6.5 (12); the kl,obsdat pH 6.2 of 10.38 M-l s-l was chosen. The concentration of the DPD reagent used in the colorimetric procedure is approximately 1 X lom4M (12). The concentration of the colored product (A2) was determined at t = 60 s and expressed as the percent per minute of the starting monochloramine concentration. Two monochloramine concentrations were chosen, 5 X 10” and 1 X lo4 M, 3.5 and 7.0 mg as ClZ/L,respectively. From the rate equations for the second-order reactions, it is calculated that the DPD interference is 5.6 and 6.0% of the monochloramine after 60 s. These values agree well with observations reported by others (Table 111). The variability in the percent interference shown in Table I11 is undoubtly due to variations in monochloramine initial concentrations, DPD concentrations, and pH as shown by this study. Acknowledgments
We thank Arthur Herriott and David H. Rosenblatt for helpful discussions. Registry No. I, 88868-44-4; 11, 65614-33-7; 111, 93-05-0; IV, 88868-45-5;V, 50470-97-8; VI, 88868-46-6; VII, 83848-64-0;VIII, 29812-35-9; IX, 88868-47-7; monochloramine, 10599-90-3.
Literature Cited (1) Morris, J. C. In “Principle and Applications of Water Chemistry”; Faust, S. D.; Hunter, J. K., Eds.; Wiley: New York, 1967; pp 23-53. (2) Wei, I. W.; Morris, J. C. In Chemistry of Water Supply”; Rubin, A. J., Ed.; Ann Arbor Science Publishers, Inc.: Ann Arbor, MI, 1974; pp 297-332. (3) Saunier, B. M.; Selleck, R. E. J.-Am. Water Works Assoc. 1979, 71, 164-172. (4) Scarpino, P. V.; Cronier, S.; Zink, M. L.; Brigano, F. A. 0. Proceedines of the 5th Water Qualitv Technoloav Conference, ALerican Water Works Assoc“iation, Den;&, CO, 1977, paper 2B-3. (5) Hoff, J. C.; Geldreich, E. E. J.-Am. Water Works Assoc. 1981, 73, 40-44. (6) Committee Report J.-Am. Water Works Assoc. 1982, 74, 376-380. ( 7 ) Brodtman, N. V.; Russo, P. J. J.-Am. Water Works Assoc. 1979, 71, 40-42.
Environ. Sei. Technol. 1984, 18, 353-358
Norman, T. S.; Harns, L. L.; Loayenga, R. W. J.-Am. Water Works Assoc. 1980, 72, 176-180. Mitcham, R. P.; Shelley, M. W.; Wheadon, C. M. J.-Am. Water Works Assoc. 1983, 75, 196-199. Brodtman, N. V., Jr.; Kroffskey, W. E.; DeMarco, J. In “Water Chlorination: Environmental Impact and Health Effect”; ed. Jolley, R. L.; Brungs, W. A.; Cumming, R. B., Eds.; Ann Arbor Science Publishers, Inc.: Ann Arbor, MI, 1980; Vol. 111, pp 777-788. Hubbs, S. A,; Amundsen, D.; Olthius, P. J.-Am. Water Works Assoc. 1981, 73, 97-101. “Standard Methods for the Examination of Water and Wastewater”, 15th ed.; American Public Health Association: Washington, DC, 1980; pp 277-301. Guter, K. J.; Cooper, W. J.; Sorber, C. A. J.-Am. Water Works Assoc. 1974, 66, 38-43. Cooper, W. J.; Sorber, C. A.; Meier, E. P. J.-Am. Water Works Assoc. 1975, 67, 34-39. Cooper, W. J.; Roscher, N. M.; Slifker, R. A. J.-Am. Water Works Assoc. 1982, 74, 362-368. Johnson, J. D. In “Water Chlorination-Environmental Impact and Health Effects”; Jolley, R. L., Ed.; Ann Arbor Science Publishers, Inc.: Ann Arbor, MI, 1978; Vol. I, pp 37-63. Strupler, N. Proc. A W W A Water Qual. Technol. Conf. 1978, p2A-6. Michaelis, L. Chem. Rev. 1935,16, 243-286. Michaelis, L. Ann. N.Y. Acad. Sci. 1940, 40, 39-71. Michaelis, L. J . Am. Chem. SOC. 1931,53, 2953-2961. Michaelis, L. J . Biol. Chem. 1932, 96, 703.
Elama, B. J . Biol. Chem. 1939, 100, 149. Michaelis, L.; Schubert, M. P.; Granick, S. J . Am. Chem. SOC.1939, 61, 1981-1992. Weitz, E.; Fischer, K. Ber. Dtsch. Chem. Ges. 1926,59,432. Miller, F.; Wilkins, R. G. J. Am. Chem. SOC.1970,92,2981. Brown, H. C.; McDaniel, D. N. J . Am. Chem. SOC.1955, 77, 3752. Rehm, C.; Bodin, J. I.; Connors, K. A.; Miguahi, T. Anal. Chem. 1959, 77,483. Brode, W. R. J . Am. Chem. SOC.1924,46, 581. Wilkins, R. G. “The Study of Kinetics and Mechanisms of Reactions of Transition Metal Complexes”; Allyn and Bacon: Boston, MA, 1974; p 46. Michaelis, L.; Fetcher, E. S., Jr. J . Am. Chem. SOC.1937, 59, 1246. Michaelis, L. J . Biol. Chem. 1938, 123, 527. Jaffari, G. A.; Nunn, A. J. J. Chem. SOC.C 1971,93,823-826. Palin, A. T. J.-Am. Water Works Assoc. 1957,49,873-880. Snead, M. C.; Olivieri, V. P.; Dennis, W. H. In “Chemistry in Water Reuse”; Cooper, W. J., Ed.; Ann Arbor Science Publisher, Inc.; Ann Arbor, MI, 1981; Vol. 1, 401-427. Fiquet, J. M. Tech. Sci. Munic. 1982, 77, 243-249. Nicolson, N. . Analyst (London) 1965, 90, 187-198. Johnson, J. D.; Overby, R. Anal. Chem. 1969,41,1744-1747.
Received for review June 6,1983. Accepted November 4,1983. This work was partially supported by the Physical Sciences Department and the Drinking Water Research Center, Florida International University.
Polycyclic Aromatic Hydrocarbons in the Clam Tridacna maxima from the Great Barrier Reef, Australia J. David Smith,?John Bagg,**$and Brian M. Bycroft’ Marine Chemistry Laboratory, School of Chemistry, University of Melbourne, Parkville, Victoria 3052, Australia, and Department of Industrial Science, University of Melbourne, Parkville, Victoria 3052, Australia
The concentrations of eight polycyclic aromatic hydrocarbons (PAH), anthracene, pyrene, chrysene, benzo[klfluoranthene, benzo[a]pyrene, benzo[ghi]perylene, fluoranthene, and perylene, were measured in clams, Tridacna maxima, collected from sites on the Great Barrier Reef ranging in latitude from 14’31’ S to 23’33’ S. At most locations the concentrations of PAH were not significantly above the limit of detection, e.g., pyrene C 0.07 pg/kg wet weight, benzo[a]pryene C 0.01 pg/kg, and chrysene < 0.07 pg/kg. These levels of PAH appear to be the lowest reported for clams anywhere in the world, indicating the pristine nature of the Great Barrier Reef a t the present time. Concentrations significantly above detection levels were found at only two sites, Lizard Island First Beach (anthracene, 3.2 pg/kg; pyrene, 1.4 hg/kg) and Heron Island Harbour (pyrene, 1.2 pg/kg; benzo[a]pyrene, 0.02 pg/kg). Both sites are frequently visited by power boats which are the most likely source of hydrocarbon contamination. These low levels of contamination would not have been demonstrated by the measurement of only the most commonly studied PAH, benzo[a]pyrene. Simultaneous determination of several PAH was necessary to show clearly that some localized pollution had occurred. Introduction The Great Barrier Reef is a chain of individual reefs stretching 2000 km approximately north-south from laMarine Chemistry Laboratory, School of Chemistry. of Industrial Science.
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titude 8 O S to 2 4 O S. It lies on the eastern continental shelf of Australia facing the Coral Sea. With the exception of some small resorts, and the research stations on Lizard and Heron Islands, the reef has no permanent habitation. The coastal region adjacent to the Great Barrier Reef has a population of only 300000 and is not highly industrialized. Consequently, inputs of polycyclic aromatic hydrocarbons (PAH), generated by combustion of fuels, into the atmosphere or into rivers with waste discharges are small and would have to be transported for distances greater than 20 km in order to reach the reef. This paper describes an investigation that shows the Great Barrier Reef to be almost unaffected by coastal discharges with the most likely sources of contamination arising from human activity on the reef itself. The Great Barrier Reef is being studied with increasing intensity because of its importance as the largest coral reef system in the world and the possibility that oil reserves under the reef may be exploited. Petroleum exploration permits issued by the Queensland State Government cover more than 60% of the Great Barrier Reef region. These permits have been suspended following an inquiry by State and Federal Royal Commissions (19741, but the possibility of future exploration still exists, and such activity could lead to some contamination of the reef. The determination of background hydrocarbon levels was regarded as essential and urgent. Petroleum is a complex and variable mixture containing up to 50% aromatic hydrocarbons. Aromatic hydrocarbons have more undesirable effects on biota than do the ali-
0 1984 American Chemical Society
Environ. Sci. Technol., Vol. 18, No. 5, 1984 353
