DECOMPOSITION OF NITRITE UNDER VARIOUS pH AND AERATION CONDITIONS WASHINGTON BRAIDA and SAY KEE ONG∗ Iowa State University, Department of Civil and Construction Engineering, 490 Town Engineering Building, Ames, IA 50011, U.S.A. (∗ author for correspondence, e-mail:
[email protected])
(Received 27 February 1998; accepted 17 February 1999)
Abstract. Studies on the decomposition/oxidation of nitrite at different pH values and aeration flow rates were investigated using a bench-scale batch reactor. The conditions were pH 2.85, 3.50, 5.80, 7.0, and 11.6, and with or without aeration at airflow rates of 1.50, 2.25, and 3.25 L min−1 . Decomposition/oxidation of nitrite may be described by a pseudo-first order expression, and the rate constants for nitrite decomposition/oxidation ranged from 1.2 × 10−6 s−1 to 1.12 × 10−4 s−1 depending on the experimental conditions. The rate of decomposition/oxidation of nitrite was found to increase for low pH conditions and for high airflow rates. The experimental results showed that the dominant reaction in the decomposition/oxidation of nitrite in low pH solutions and in the presence of some aeration was most probably the decomposition of nitrous acid to NO and NO2 . Oxidation of nitrite to nitrate appeared to proceed in a smaller proportion in comparison to the breakdown of nitrous acid to nitrogen oxide compounds. Results from this study showed that emissions of nitrogen oxides from nitrite-containing solutions are possible if the solutions were agitated and the pHs of the solutions were less than 6. Keywords: aeration, decomposition, nitrite, nitrogen oxides, oxidation
1. Introduction Nitrite is an intermediate product in the nitrogen cycle and in ‘healthy’ biological systems, its presence is a relatively fleeting one. However, under certain conditions nitrite may reach toxic concentrations in aquatic ecosystems and aquaculture environments. The toxicity of nitrite in vertebrates is related to the nitrite ability to oxidize hemoglobin to methemoglobin, a form incapable of carrying oxygen (Freeman et al., 1983). Reported nitrite toxicity (96-hr LC50 ) to fathead minnows (Pimephales promelas) ranged from 1 to 70 mg L−1 depending on the fish weight and water hardness (Palachek and Tomasso, 1984). The effects of nitrite on human health were recognized in the preliminary version of the U.S. Safe Drinking Water Standards, in which nitrite-nitrogen was limited to 1 mg L−1 of the allowable 10 mg L−1 of NO− 3 -N for infants ingesting the water (U.S. EPA, 1973). Nitrite has been implicated as a possible cause of esophagus cancer (Manahan, 1984; Tannenbaun et al., 1978). In many metal finishing industries, nitrite is used as a corrosion inhibitor in the preparation of metal surfaces for painting. Rinsewater from metal surface preparWater, Air, and Soil Pollution 118: 13–26, 2000. © 2000 Kluwer Academic Publishers. Printed in the Netherlands.
14
W. BRAIDA AND S. K. ONG
ation processes could potenially contain high concentration of nitrite which are then discharged to wastewater treatment plants for treatment. Typical wastewater treatment of metal finishing wastewater includes a conventional flocculationsedimentation system which is usually not adequate for the rapid oxidation of nitrite to nitrate. Nitrous acid and nitrites have both oxidizing and reducing properties and may be strong oxidizing agents in acid solutions. The oxidizing properties of an aqueous waste stream containing nitrite was used by Muller and Kuhn (1970) to detoxify an aqueous solution of cyanide by mixing the two solutions at pH less than 5 in the presence of a catalyst. Oxidizing agents, such as chlorine gas, sodium hypochlorite or in certain cases hydrogen peroxide may be used for the detoxification of nitrite and cyanide-containing wastewater. Electrochemical and photochemical decomposition of nitrite have been the focus of some recent research. The electromechanical oxidation of nitrite using a rotating gold disk electrode was studied by Xing and Scherson (1988). They found that the decomposition of nitrite proceeded by a second-order disproportionation rate with the formation of nitrogen dioxide. Fischer and Warneck (1996) investigated the photodecomposition of nitrite and undissociated nitrous acid in aqueous solutions. The photochemical decomposition of nitrite at pH 6 and nitrous acid at pH 2 using benzene as scavenger for hydroxyl radicals yielded phenol and nitrate as final products. It is commonly accepted that under well-aerated conditions but in the absence of biotic and catalytic surfaces, nitrite rapidly oxidizes to nitrate. For this reason, it is rare to find nitrite concentrations greater than 0.1 mg L−1 in drinking water. Under gently aerated conditions and depending on the pH of the solution, oxidation of nitrite may proceed by the decomposition of nitrous acid to nitrogen oxides and nitrates. Although the reaction mechanism for the decomposition of nitrous acid may be fairly established, information regarding the relative importance of one reaction over the other at different pH values and agitation rates (e.g., aeration) is not available. The objectives of this study were to investigate the oxidation of nitrite under various pH conditions and aeration rates and to determine the dominant reaction for the oxidation of nitrite. Experiments were conducted using nitrite anions in sterile deionized water in the absence of any biotic or abiotic catalytic effects.
2. Materials and Methods Batch experiments to determine the decomposition/oxidation of nitrite at various pH conditions and aeration rates were conducted in one-liter beakers which were rinsed with sterile water and dried for two hr at 150 ◦ C. Five hundred milliliters of approximately 50 mg L−1 of nitrite solution were used in each experimental run. The solutions for the experiments were prepared by diluting 6.25 mL of stand-
DECOMPOSITION OF NITRITE UNDER VARIOUS pH AND AERATION CONDITIONS
15
ard 4000 mg L−1 sodium nitrite solution to 500 mL with deionized water filtered through a sterile 0.22 micron membrane. The solution was gently stirred using a magnetic stirrer. Air, supplied by a rubber diaphragm air pump, was bubbled into the solution through a fritted glass diffuser. Air flow rates used were 1.50, 2.25, and 3.25 L min−1 . Gas flow rates were measued using a Gilmont size 13 flowmeter (Gilmont Instruments, Barrington, IL). The oxidation-reduction potential (ORP) was measured with an ORP probe (Cole-Palmer Cat. Number 5994-21, Vernon Hilss, IL) and the dissolved oxygen of the solution was determined using YSI Model 58 dissolved oxygen meter (YSI Incorporated, Yellow Springs, OH). The pH of the solutions ranged from 2.85 to 11.6 and was measured with an Orion 710 pH meter (Cole-Palmer, Vernon Hills, IL). The pH of the solution was adjusted using concentrated sulfuric acid or 50% sodium hydroxide solution. The pH of the solution was monitored continuously throughout the experiments and was found to change on the average by ± 0.15. The temperature of the solution was kept constant at 20 ± 1 ◦ C. During each experimental run, aliquots of 2.5 mL were collected at various times and analyzed for nitrite, nitrate and sulfate using a Waters 501 ion chromatograph (IC) with a Waters 431 conductivity detector. Nitrite/nitrate determinations were performed according to Satndard Method No. 4110 C (Greenberg et al., 1992). Eluent flow rate was set at 1.2 L min−1 and sample volume used was 0.1 mL. Monitoring of the sulfate concentration (from the pH adjustment with sulfuric acid) throughout the experiments served as a check on the accuracy of the quantitative determinations of the IC and sampling consistency. Nitrogen mass balance experiments were conducted using a system as shown in Figure 1. The experiments were conducted to demonstrate that decomposition/oxidation of nitrite may result in the formation of nitrogen oxides. Off-gases from the 500 mL reaction flask were bubbled into a solution of 14 M ammonium hydroxide. Any nitrogen oxides stripped from the reaction flask would react with the ammonium hydroxide to form ammonium nitrite as given by the reaction below (Flagan and Seinfeld, 1988). NO(g) + NO2(g) + 2NH4 OH → 2NH4 NO2 + H2 O.
(1)
The initial and final concentrations of nitrite and nitrate in each flask were determined. Two experimental runs were conducted for different volumes of reacting and collecting solutions, and time of reaction. A blank run was conducted without adding any nitrite to the reaction vessel to check for the presence of nitrogen oxides in the air. After 2.5 hr of bubbling air through the system, no detectable amounts of nitrite or nitrate were measured at the collection vessel.
16
W. BRAIDA AND S. K. ONG
Figure 1. Reactor and collection vessel for nitrite decomposition and mass balance experiments.
3. Results and Discussion 3.1. E FFECTS
OF AERATION RATES AND DECOMPOSITION / OXIDATION
pH
ON NITRITE
The changes in the total concentrations of nitrite and nitrate are presented in Figures 2 and 3. Figure 2 presents the effects of pH on the change of nitrite and nitrate for an air flow rate of 1.5 L min−1 while Figure 3 presents the effects of aeration rates on nitrite decomposition/oxidation at pH 2.85. Figure 2 shows that decomposition/oxidation of nitrite was faster at lower pH values. For pH 2.85, approximately 40% of nitrite was removed within 180 min while at pH 11.6 removal of nitrite was less than 5%. Correspondingly, a higher concentration of nitrate was formed at lower pH conditions than at higher pH conditions. A commonly accepted explanation for the observed results is that nitrite was oxidized to nitrate as given by the following equation (reaction 2): − NO− 2 + 1/2 O2 → NO3 .
(2)
As shown in Table I, the dissolved oxygen concentrations were close to saturation throughout the experiments indicating that the oxygen partial pressure in the solution was maintained fairly constant throughout the experiment. Even though reaction 2 did not show any pH dependence, the measured ORP values for a flow rate of 1.5 L min−1 were 590 mV at pH 2.85, 265 mV at pH 7.0, and 89.6 mV at pH 11.6 (see Table I). The oxidizing conditions as measured by the ORP probe
DECOMPOSITION OF NITRITE UNDER VARIOUS pH AND AERATION CONDITIONS
17
Figure 2. Change in nitrite and nitrate concentrations with pH and at an aeration rate of 1.5 L min−1 − (open symbols show ratios of measured:initial NO− 2 concentrations, closed symbols show net NO3 concentrations).
were high at low pH values. The stronger oxidizing conditions at lower pH values would therefore provide conducive conditions for the oxidation of nitrite to nitrate as evident by the higher rate of oxidation at low pH values. Another possible explanation is the presence of nitrous acid in solution at different pH values. At pH 2.85, 67% of the total nitrite was in the form of nitrous acid while the remainder was in the form of nitrite anion. The pKa value for nitrous acid is 3.2 at 20 ◦ C. At pH 3.50, 31% of the total nitrite was in the form of nitrous acid while at pH values greater than 5.30 only 1% of the total nitrite was in the form of nitrous acid. Based on the relative amounts of nitrite and nitrous acid present, there appeared to be a correlation between the concentration of nitrous acid in the solution and the higher removal rates of nitrite at lower pH values. Nitrous acid is
18
W. BRAIDA AND S. K. ONG
Figure 3. Influence of aeration on nitrite removal at pH 2.85 (closed symbols show ratios of − measured:initial NO− 2 concentrations; open symbols show net NO3 concentrations).
DECOMPOSITION OF NITRITE UNDER VARIOUS pH AND AERATION CONDITIONS
19
TABLE I Estimated pseudo-first order rate constants for nitrite decomposition at 20◦ C and their 90% confidence limits pH
Dissolved oxygen (mg L−1 )
ORP (mV)
Gas
Aeration rate (L min−1
k × 105 (s−1 )
Correlation coefficient (r)
Aeration rate/vol. of solution (min−1 )
2.85 2.85 2.85 2.85 3.50 5.80 5.80 7.00 11.6
8.5 8.4 8.4 8.6 8.4 8.5 8.3 8.3 8.3
570 590 590 580
none air air air air air none air air
0 1.50 2.25 3.25 1.50 1.50 0 1.50 1.50
2.42 ± 0.08 5.12 ± 0.35 11.2 ± 0.55 10.5 ± 0.45 2.13 ± 0.18 0.53 ± 0.33 0.14 ± 0.11 0.26 ± 0.14 0.12 ± 0.09b
0.99 0.99 0.99 1.00 0.99 0.90 0.88 0.95 0.98
0 3.0 4.5 6.5 3.0 3.0 0 3.0 3.0
a a a
265 89.6
a Measurements were not made. b 80% confidence limit (only 3 experimental points).
unstable and decomposes by disproportionation to nitrogen oxides as given by the following (reaction 3): 3HNO2 ↔ H+ + NO− 3 + 2NO(g) + H2 O
(3)
In both reactions 2 and 3, nitrate is formed as a product. In reaction 2, the molecular ratio of nitrite to nitrate is 1:1. In reaction 3, the molecular ratio of nitrate formed to nitrous acid removed is 1:3. While the above two reactions may occur, the main issue was to determine the importance of each reaction under various solution conditions. For reaction 3 to proceed satisfactorily, nitrogen oxides must be removed from solution so that the equilibrium in reaction 3 would be driven towards the formation of more products. According to the data presented in Figure 3a, an increase in aeration flow rate from 0 to 2.25 L min−1 at pH 2.85 resulted in higher removal of nitrite (see Table I). For both conditions, the ORP values and dissolved oxygen concentrations were similar. As presented in Table I, an increase in the aeration rate from 2.25 to 3.25 L min−1 resulted in a slight increase in the rate of nitrite removal. This may indicate that the mass transfer of nitrogen oxides from the liquid phase to gas phase was limiting at the two airflow rates. As such, the data presented in Figures 3a and 3b showed some evidence that reaction 3 may be occurring. A visual inspection of the change in nitrite concentration over time as presented in Figures 2 and 3 showed that the decomposition of nitrite followed an exponential decay pattern which may be described by a pseudo-first order reaction.
20
W. BRAIDA AND S. K. ONG
Furthermore, a pseudo-first order reaction may be assumed since the dissolved oxygen concentration (i.e., oxygen partial pressure) and pH remained fairly constant throughout the experimental runs. Rate constants for a pseudo-first order reaction were determined using linear regression of the natural logarithm of the measured nitrite concentrations against time and are presented in Table I. For a fixed aeration rate of 1.5 L min−1 , the reaction rate constant at pH 2.85 was 5.12 × 10−5 s−1 which was approximately 40 times larger than the removal rate at pH 11.6. As presented in Table I, rate constants were lower for higher pH values but tapered off smoothly at values of 5.3 × 10−6 s−1 at pH 5.8, 2.57 × 10−6 s−1 at pH 7.0, and 1.2 × 10−6 s−1 at pH 11.6. A comparison between conditions with and without aeration showed that at pH 2.85 with an aeration rate of 2.25 L min−1 , the reaction rate constant was four times larger than without aeration. Figure 4 presents a three-dimensional plot of the change in pseudo-first order rate constants for the decomposition of nitrite as a function of pH and aeration rate. In general, the pseudo-first order rate constant decreased with increasing pH and increased with increasing aeration rate. As shown in Figure 4, the rate constant reached a fairly constant value for a given pH and aeration rate indicating that the mass transfer of the NOx from the solution to the atmosphere may control the overall reaction process. Using the computed rate constants, 90% removal of nitrite may be achieved in less than 6 hr at pH 2.85 with an aeration rate of 2.25 L min−1 . However, more than 24 hr would be needed to achieve the same percent removal without aeration. 3.2. N ITRITE - NITRATE
MASS BALANCE
To assess the reaction stoichiometry, the number of moles of nitrite removed and the number of moles of nitrate formed were computed and plotted as shown in Figure 5. Since the nitrate concentrations produced for the experiments conducted at pH 7.0 and 11.6 were too low to be analyzed accurately, the data from these experiments were not plotted in Figure 5. The data in Figure 5 were regressed linearly by forcing the regression line through the origin. The slope of the regression represents the nitrate to nitrite stoichiometry for the overall reaction. Results of the regression analysis are presented in Table II. As mentioned earlier, in reaction 2, one mole of nitrite will produced one mole of nitrate while for reaction 3, three moles of nitrite will be needed to form one mole of nitrate. Comparing the theoretical values with the computed values from the regression analysis, it can be seen that the computed values were not in accordance with either the stoichiometry of reactions 2 or 3. Except for the conditions at pH 2.85 with an aeration rate of 3.25 L min−1 , the slopes of the regression were between 0.366 and 0.603. Therefore, it can be assumed that decomposition/oxidation of nitrite in absence of biotic and catalytic reactions would proceed by a combination of reactions 2 and 3. If both reactions were assumed to proceed simultaneously, then the fraction of each reaction based on a linear combination of both mechanisms may be computed
DECOMPOSITION OF NITRITE UNDER VARIOUS pH AND AERATION CONDITIONS
21
Figure 4. Influence of pH and aeration rate on pseudo-first order rate constant (k) on nitrite decomposition.
TABLE II Estimation of fractions of different reactions using regression analysis for nitrate-nitrite stoichiometric balance pH
Aeration rate (L min−1 )
Slope of linear regression
Correlation coefficient (r)
Fraction of reaction 3
Fraction of reaction 2
2.85 2.85 2.85 2.85 3.50 5.80
0 1.50 2.25 3.25 1.50 1.50
0.603 0.390 0.435 0.196 0.366 0.473
0.98 0.93 0.99 0.97 0.98 0.98
0.59 0.91 0.84 >1.0 0.89 0.79
0.41 0.09 0.16 – 0.11 0.21
22
W. BRAIDA AND S. K. ONG
Figure 5. Ratio of change in nitrate concentrations (mmole L−1 ) to change in nitrite concentrations (mmole L−1 ).
DECOMPOSITION OF NITRITE UNDER VARIOUS pH AND AERATION CONDITIONS
23
as follows. Using the experimental run for pH 5.8 and an aeration rate of 1.5 L min−1 with a regression slope of 0.473 as an example, then, 0.473 = a(0.33) + b(1.0) 1.0 = a + b with a ≥ 0and b ≥ 0 where a is the fraction of reaction 3 and b is the fraction of reaction 2. Values of 0.33 and 1.0 are the stoichiometric coefficients for reactions 3 and 2, respectively. For pH 5.8, the fraction of nitrite oxidation attributed to reactions 2 and 3 were 0.21 and 0.79, respectively. Based on the calculations, the bulk of the oxidation reaction of nitrite was due to reaction 3 in which nitrous acid was decomposed to nitrogen oxides followed by volatilization as a result of aeration. Even when 99% of the nitrite was in the form of nitrite anion (for a pH solution of 5.8, see Table II), decomposition of nitrous acid appeared to account for 79% of the overall oxidation of nitrite. The larger stoichiometric ratio (0.603) for pH 2.85 without aeration was probably due to the larger role played by reaction 2. Under these experimental conditions, stripping of nitrogen oxide from the solution as in reaction 3 was minimized due to a lack of aeration. Even for this set of experimental conditions, the role played by reaction 2 was only 41% in the overall reaction. The low stoichiometric ratio (0.196) for pH 2.85 and an air flow rate of 3.25 L min−1 may be explained as follows. There are several proposed reaction mechanisms for reaction 3. One of the proposed mechanism is the initial formation of NO and NO2 as shown below (Abel and Schmid, 1928; Abel et al., 1928). 2HNO2 ↔ NO + NO2 + H2 O
(4a)
2NO2 ↔ N2 O4
(4b)
N2 O4 + 2H2 O → H3 O+ + NO− 3 + HNO2 .
(4c)
If the aeration rate is large enough, the nitrogen oxides produced in reaction 4a may be stripped out before reactions 4b and 4c may proceed to an appreciable extent. This would mean that less nitrate would be formed, and therefore, would indicate a low stoichiometric ratio. 3.3. N ITROGEN
MASS BALANCE
Results of the nitrogen mass balance experiments to confirm the formation of gaseous nitrogen oxides are presented in Table III. Nitrite was measured in the collection vessel demonstrating that nitrous acid was decomposed to gaseous nitrogen oxides in the reaction vessel during aeration. Table III shows the total mass of nitrogen at the start of the experiments was similar of the total mass of nitrogen measured at the end of the experiments. The mass balance differences at the start
24
W. BRAIDA AND S. K. ONG
TABLE III Estimation of nitrogen in mass balance experiments (temperature 20 ± 1 ◦ C)
pH in reactor Solution vol. in reactor, mL Solution vol. in collector, mL Aeration, L min−1 Reaction time, min. −1 NO− 2 in reactor, mg L (mg-N/L) −1 NO− 3 in reactor, mg L (mg-N/L) −1 NO− 2 in collector, mg L (mg-N/L) −1 NO− 3 in collector, mg L (mg-N/L) Aeration rate/Vol. of solution, min−1 Total N, mg
Experiment 1 Initial Final
Experiment 2 Initial Final
2.85 100 100 2.25 130 115.3 (35.1) 23.8 (5.4) 4.6 (1.4) 0.0 (0.0) 22.5
2.89 100 100 2.25 130 43.6 (13.3) 28.9 (6.5) 84.0 (25.5) 0.0 (0.0) 22.5
2.85 300 300 2.25 45 40.9 (12.4) 0.0 (0.0) 0.0 (0.0) 0.1 (0.02) 7.5
2.90 300 300 2.25 45 31.2 (9.5) 5.3 (1.2) 5.6 (1.7) 0.2 (0.05) 7.5
4.2
4.5
3.7
3.7
and at the end of the two experiments were 7.1 and 0.2%. The stripping effect produced by aeration played an important role in the decomposition of nitrite and the low concentrations of nitrate found in the reaction vessel for the two experiments supported the hypothesis that with high aeration rates, formation of nitrate was minimized but volatilization of nitrogen oxides was maximized. In addition, the aeration rate per unit volume played an important role in determining how much nitrate was formed. In experiment 1, very low amounts of nitrate were formed in comparison with the earlier kinetic experiments (Figure 2). A possible reason was that the aeration rate per unit volume for experiment 1 (22.5 min−1 ) was about 7.5 times larger than the aeration rate per unit volume for the kinetic studies (3 min−1 , see Table I) as presented in Figure 2. The release of nitrous oxide (N2 O) from wastewater treatment plans during biological denitrification has been the subject of recent research (Cepiel et al., 1995; Thorn and Sorensson, 1996; Spector, 1998). Researchers have found that 50 to 80% of the nitrogen reduced was in the form of N2 O, prior to its reduction to molecular nitrogen which leads to a higher estimate of nitrous oxide emissions than previously considered (Spector, 1998). Although the experiments of this research were conducted using an aqueous solution of nitrite in the absence of biotic and
DECOMPOSITION OF NITRITE UNDER VARIOUS pH AND AERATION CONDITIONS
25
catalytic surfaces, the implications of the decomposition of nitrite in wastewater through the formation of nitrogen oxides (NO and NO2 ) are evident. As a result of the disproportionation of nitrous acid, solutions containing nitrite from industrial production, wastewater containing nitrites as in metal finishing industries, or municipal wastewater treatment plants with nitrification, may be potential sources of nitrogen oxides (NOx ) depending on the pH of the solution and the extent of aeration or mechanical agitation.
4. Conclusion The experimental results showed that the dominant reaction in the oxidation of nitrite in solutions at low pH was most probably the decomposition of nitrous acid to NO and NO2 when aided with some agitation/aeration. For pH values greater than 6, the rate of oxidation of nitrite was dramatically reduced. The reason for the low reaction rate was that the concentration of nitrous acid available for the decomposition was low. For acidic conditions, the oxidation of nitrite to nitrite by oxygen occurred in a smaller proportion in comparison to nitrous acid decomposition. The decomposition of nitrite in aqueous solution for a given pH and constant dissolved oxygen concentration may be described by a pseudo-first order kinetic expression. Experimentally determined reaction rate constants ranged from 1.2 × 10−6 s−1 at pH 11.6 for an aeration rate of 1.5 L min−1 , to 1.12 × 10−4 s−1 at pH 2.85 for an aeration rate of 2.25 L min−1 . The rate of agitation by bubbling air affects the rate of reaction by stripping the nitrogen oxides produced from the decomposition of nitrous acid. The presence of nitrogen oxides removed by stripping was confirmed by nitrogen mass balance experiments whereby the volatilized nitrogen oxides were trapped using ammonia hydroxide solution. Nitrite is a common chemical used in many industrial processes and nitrite is also a degradation product of ammonia in municipal wastewater treatment plants. Results from this study showed that nitrogen oxides may be emitted if the pH of the solution is less than 6 and there is sufficient agitation.
References Abel, E., Schmid, H, and Babad, S.: 1928, Z. Phys. Chem. 136, 135. Abel, E. and Schmid, H.: 1928, Z. Phys. Chem. 134, 55. Flagan, R. C. and Seinfeld, J. H.: 1988, Fundamentals of Air Pollution Engineering, Prentice Hall, Inc., NJ, p. 517. Freeman, L., Beitinger, T. L. and Huey, D. W.: 1983, Comp. Biochem. Physiol. 75B, 27. Manahan, S. E.: 1984, Environmental Chemistry, 4th Ed., Brooks/Cole Publishing Company, Monterey, CA. Muller, J. and Kuhn, R.: 1970, U.S. Patent 3,502,576 assigned to Deutsche Gold und Silber Scherdeanstalt, Germany. Palachek, R. M. and Tomasso, J. R.: 1984, Bull. Environ. Contam. Toxicol. 32, 238.
26
W. BRAIDA AND S. K. ONG
Spector, M.: 1998, Water Environmental Research 70, 1096. Tannenbaun, S. R., Fett, D., Young, V. R., Land, P. D. and Bruce, W. R.: 1978, Science 200, 1487. Thorn, M. and Sorensson, F.: 1996, Water Research 30(6), 1543. U.S. EPA: 1973, Nitrogen Compounds in the Environment, U.S. EPA, SAB-73-001, Washington, D. C. Xing, X. and Scherson, D. A.: 1988, Analytical Chemistry, 60, 1468.
