Chemical Nitrite Oxidation in Acid Solutions as a Consequence of

Apr 30, 2005 - Chemical Nitrite Oxidation in Acid Solutions as a Consequence of Microbial Ammonium Oxidation. Kai M. Udert,*Tove A. Larsen, andWilli G...
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Environ. Sci. Technol. 2005, 39, 4066-4075

Chemical Nitrite Oxidation in Acid Solutions as a Consequence of Microbial Ammonium Oxidation KAI M. UDERT,* TOVE A. LARSEN, AND WILLI GUJER Swiss Federal Institute of Environmental Science and Technology (EAWAG) and Swiss Federal Institute of Technology (ETH), 8600 Du ¨ bendorf, Switzerland

In long-term experiments with membrane aerated biofilm reactors we observed complete nitrite oxidation in highly concentrated ammonium nitrite solutions with a contaminant pH decrease to values below 3. The maximum initial concentration for ammonium was 42 mM and for nitrite was 41 mM. We hypothesized that (1) acid-tolerant ammonium oxidizing bacteria were responsible for the pH decrease, and (2) chemical processes caused complete nitrite oxidation at low pH values. To test this hypothesis we set up a mechanistic computer model based on kinetic data from literature and we validated the model with additional experiments. The simulations fitted the measurements very well. Additionally, an experiment with the inhibitor allylthiourea showed that ammonium-oxidizing bacteria were active at pH values far below 5.5. Experiments in a sterile reactor confirmed the chemical nitrite oxidation to nitrate. Nitrogen balances revealed that 8 ( 4% of the initial nitrogen (ammonium, nitrite, and nitrate) were lost during the cycles. On the basis of measurements and simulations we concluded that volatilization was responsible for the significant nitrogen loss. We estimated that about half of the lost nitrogen volatilized as nitrous acid HNO2. The rest mainly volatilized as dinitrogen N2 and nitrous oxide N2O.

Introduction Microbial nitrification is an important process in wastewater treatment (1). Because of the acid production during nitrification, the pH value and the alkalinity are critical process parameters, especially for wastewater with high ammonium concentrations (2). When the pH reaches values around 5.5, nitrification is usually inhibited (3). In contrast to this observation, we found ammonium and nitrite oxidation at pH values as low as 2.6. We will show that besides biological ammonium oxidation, chemical processes (4, 5) caused nitrification in our reactors. Biological Oxidation of Ammonium and Nitrite. Lithoautotrophic bacteria, heterotrophic bacteria, and fungi oxidize ammonium to nitrite, but only lithoautotrophic bacteria are known to gain energy from ammonium oxidation and are therefore the most common ammonium oxidizers * Corresponding author current address: Massachusetts Institute of Technology, Ralph M. Parsons Laboratory, 77 Massachusetts Ave., Cambridge MA, 02139; tel: 617-253-3994; fax: 617-253-7475; e-mail: [email protected]. 4066

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(6, 7). Microbial ammonium oxidation proceeds in two main steps (6)

First step: NH4+ + 0.5 O2 f

NH2OH (hydroxylamine) + H+ (1)

Second step: NH2OH + O2 f NO2- + H+ + H2O (2) For every ammonium molecule oxidized two protons are released, leading to a pH decrease that will eventually stop the process. In wastewater treatment plants, ammonia oxidation is generally not observed below a pH value of 5.5 (3). This is in accordance with several studies showing that Nitrosomonas and Nitrosospira species, the most common ammonium oxidizing bacteria (AOB), were unable to oxidize ammonium in liquid medium with pH values below 5.5 (8). This pH limit is usually attributed to the limitation by ammonia, NH3, the true substrate of AOB. However, recent work has shown that the loss of AOB activity does not correlate with the ammonia concentration (2). Therefore, the pH limit must be caused by other inhibition or limitation effects, as there are inhibition by nitrous acid (9), limitation by carbon dioxide (10), or inhibition by NOx compounds. Several authors (11, 12) showed that Nitrosomonas eutropha are very sensitive to NO and NO2, compounds that can be produced by chemical processes at low pH values (see below). Additionally, AOB are subject to physiological limitations at low pH values (13). So far it is not clear which effect causes the strict cessation of AOB activity observed at pH 5.5. Nitrification at low pH has been described for acid environments such as acid soils (8) or acidified lakes (14). Several hypotheses for the activity of lithotrophic AOB in acid environments exist (8): acid-sensitive cells are surrounded by pH-neutral micro sites; or they keep the internal pH neutral by intracellular urea hydrolysis. Acid-tolerant cells can develop when the bacteria encounter frequent pH fluctuations or if they are immobilized. In general, high cell densities enhance the ability to nitrify at low pH values. Additionally, heterotrophic ammonium oxidizers may be active in acid soils if organic substrate is available (15), but their activity is assumably too low for a significant contribution to nitrification in acid soils (8). Similar to ammonium oxidizing microorganisms, most nitrite oxidizing microorganisms are lithoautotrophic bacteria. Additionally, some heterotrophic microorganisms, mainly bacteria, oxidize nitrite to nitrate (6). Nitrite oxidizing bacteria (NOB) mostly occur jointly with the nitrite producing AOB. There is some evidence that nitrous acid and not nitrite is the real substrate for NOB (16). Accordingly, the basic equation for the bacterial nitrite oxidation is

HNO2 + 0.5 O2 f NO3- + H+

(3)

NOB are widespread in the environment and in wastewater treatment, but they are very sensitive to high concentrations of nitrogen compounds (2). Incomplete nitrification often occurs in systems with high ammonium concentrations and low alkalinity (2, 17, 18). In these systems, nitrite oxidizers are inhibited by high concentrations of nitrous acid and hydroxylamine. Low pH values also affect the nitrite oxidation by NOB. Nitrous acid is the most important inhibitor (16). NO and NO2 are other possible inhibitors (11, 19) and carbon dioxide may limit NOB growth at low pH values (10). Physiological limitations such as protein damage or destabilization of the 10.1021/es048422m CCC: $30.25

 2005 American Chemical Society Published on Web 04/30/2005

cytoplasmic pH also occur at low environmental pH values (13). In contrast to the AOB, no pH limit for NOB exists in the literature. Several researchers stated that NOB, especially heterotrophic nitrite oxidizers, are responsible for nitrite oxidation in acid soils (8). Recently, Tarre and co-workers (20, 21) showed that complete nitrification is possible in liquid media with pH values below 5.5. They described complete nitrification in acid fluidized bed reactors (average pH values as low as 4.3 ( 0.1) and in an acid suspended biomass reactor (average pH value 3.8 ( 0.3) (21). Fluorescence in situ hybridization (FISH) analysis showed that the main AOB were Nitrosomonas oligotropha and Nitrosomonas mobilis, while the main NOB were Nitrospira species. Similar to our experiments, the initial biomass was taken from a wastewater treatment plant and not from acid soils. In contrast to our experiments, the nitrogen concentrations were low with ammonium nitrite concentrations below 0.7 mM and 0.02 mM, respectively. Accordingly, inhibition by nitrous acid could be excluded. Chemical Oxidation of Nitrite and Ammonium. Compared to ammonium, nitrate, or nitrite, nitrous acid is a very reactive compound. Nitrous acid undergoes a variety of chemical degradation processes, known from soil science and atmospheric chemistry. In soil science, the term chemodenitrification is used to summarize chemical nitrite degradation processes (4). Chemodenitrification is considered to be a major source for nitrogen volatilization. The different chemical nitrite degradation processes are poorly characterized. Ammonium nitrite solutions spontaneously decompose to N2 and H2O (5), a reaction enhanced by UV radiation (22). The chemical ammonium oxidation is usually much slower than the microbial ammonium oxidation, but at low pH values it contributes significantly to the loss of nitrogen as N2, as we will show in this paper.

Experimental Section Experimental Setup. Two identical membrane aeration bioreactors were run simultaneously in sequencing batch mode. The glass reactors contained membranes of silicone tubing fixed on a stainless steel rack (Figure 1). The experiment lasted for more than 45 months, of which the first 16 months were thoroughly observed. The refill interval was 6 days during the first half of the observation period and 7 days during the second half of the observation period. The fill volume was 1.6 L with a remaining headspace of 0.65 L. The silicon tubing was completely immersed in solution for the entire experiment. The reactors were closed to prevent gas exchange with the lab atmosphere. The input solution was partially nitrified diluted urine from an activated sludge sequencing batch reactor (SBR) (2). In the SBR nearly all biodegradable dissolved organic substrate was eliminated, and half of the initial ammonium was oxidized to nitrite. The nitrate concentration was generally less than 5% of the sum of ammonium, nitrite, and nitrate (Nsum). The measured Nsum (representing more than 90% of total nitrogen in the solution) was between 55 and 85 mM, with higher concentrations at the end of the observation period. The maximum initial concentration for ammonium was 42 mM and for nitrite was 41 mM. The concentration of dissolved phosphate was about 2.5 mM. The concentration of bicarbonate was generally below 0.8 mM, with typical values of 0.15 mM. The initial pH was always above 6 with maximum values of 7.2. The higher initial pH values were caused by denitrification in the SBR. The reactors were inoculated with biomass from the SBR (2). During the first month one reactor (reactor deep) was filled with nonnitrified urine (influent of the SBR) to grow a deep heterotrophic biofilm. The biofilm stayed thick during the whole experiment, but it partly detached from the silicone

FIGURE 1. Reactor thin with silicone tubing fixed on a steel rack. Day 295 of the experiment. tubing (sloughing). The other reactor (reactor thin) was filled with SBR effluent from the start. In this reactor the biofilm was very thin after one month, but its thickness increased during the experiment. The silicone tubing served for diffusive aeration and as substratum for the biofilm. Since the membrane aeration kept the bulk oxygen concentration always above zero, anoxic compartments in the biofilm could not develop (23). The diameter of the tubing was 4 mm, the wall thickness was 0.5 mm, and the length was 4.1 m. Before usage the surfaces of the membranes were roughened with sandpaper and covered with soil for several weeks to condition them for biofilm growth. The solutions were mixed with magnetic stirrers. An external water bath kept the temperature at 25 °C. Daylight had no significant influence on the processes in the solutions since the reactors were operated in a basement laboratory. Sampling, Conservation, and Preparation. Samples of 25 mL were taken with a glass pipet and filtered with glass fiber filters (Whatman, GF/F 0.7 µm). To prevent chemical nitrite oxidation, the pH value in the sample was elevated above 6 by adding 1 M NaOH. All samples were stored at 4 °C for less than 4 days before analysis. The dilution factor of the samples was up to one hundred. To minimize dilution errors, all dilutions including the base addition for the nitrite analysis were weighed on scales. Analytical Methods. All samples were analyzed for nitrite, nitrate, and ammonium. Colorimetric flow injection analysis FIA (Ismatec AG, Switzerland) was used to measure nitrite (sulfanilamide method) and ammonium (bromocresol purple method). Ion chromatography (Ion Chromatograph DX-300, Dionex IonPac AS12A; eluent: 2.7 mM Na2CO3, 0.3 mM NaHCO3) was used for nitrate analysis. Some samples were analyzed for phosphate, carbonate, hydroxylamine, and iron. Colorimetric flow injection analysis (Ismatec AG, Switzerland) was used for phosphate (molybdate method). The total carbonate concentration was determined with a TOC analyzer (Shimadzu TOC Analyzer 5000-A). Hydroxylamine was measured with the method by Frear and Burrell (24). To eliminate interference with nitrite, the samples VOL. 39, NO. 11, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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were treated with amidosulfonic acid (0.1 M amidosulfonic acid per sample). The detection limit was 10-5 M for total hydroxylamine. Dissolved iron was determined by ICP-OES (model Ciros, Spectro Analytical Instruments). Since urine is a highly concentrated and complex solution, we tested and proved all analytical methods with recovery tests (25). The pH values and oxygen concentrations were measured in situ with electrodes (combination pH electrode Mettler HA405-DXK-S8/225, WTW oxygen electrode CellOx 225) and were logged continually. In one cycle NO was continually measured in the exhaust air with a chemiluminescence detector (Nitrogen Oxides Analyzer model 8840, Monitor Labs Inc.). The measurement range was 1-10 ppm. In a further cycle N2O was monitored in the exhaust air with a gas filter correlation spectrometer (N2O-Monitor model 48, Thermo Environmental Instruments). Air Flow Control. The air flow through the silicone tubing was fixed with mass flow controllers to approximately 4.0 m3/d (Sierra Instruments and Bronkhorst Hi-Tec). Mass flows and atmospheric pressures were measured and later used to calculate the volumetric gas flows. Inhibition of Nitrifiers. We tested the activity of AOB by adding the inhibitor allylthiourea (ATU) to the reactor thin at a pH of 4.75. The ATU concentration in the reactor was 0.086 mM (10 mg ATU/L). After 2 h of inhibition we changed the reactor solution two times to prevent permanent detrimental effects on the bacteria. Experiments without Biological Activity. We validated a model for chemical nitrite oxidation, calibrated the gas transfer coefficients KL for HNO2 and NO, and determined the KL for CO2 in experiments without biomass, but with the same reactor configuration as above. The significance of the chemical nitrite oxidation processes was examined with a synthetic ammonium nitrite solution (35.7 mM NH4Cl and 35.7 mM NaNO2). Before the experiment, the reactor was sterilized with a 4% hypochlorite solution. Pure oxygen was used for aeration. The oxygen concentration was close to saturation during the whole experiment. The proton production by AOB was simulated with a continuous input of 0.1 M hydrochloric acid. The temperature was constant at 23 °C. To calibrate the KL for HNO2 and NO, we used filtered SBR effluent (0.7 µm glass fiber filters, Whatman GF/F) as reactor solution. Since the chemical oxidation of nitrite, NO production, and HNO2 volatilization are only significant at low pH values, the pH was initially adjusted with 1 M hydrochloric acid to a value of 4.2. Pure oxygen was used for aeration. The computer model described below was used to fit the gas transfer coefficients for HNO2 and NO. pH values in the solution and NO concentrations in the off-gas were used as fitting data (Figures 4 and 5). To test the carbon dioxide volatilization, the reactor was completely filled with a carbonate standard solution of 0.42 mM (5 gC/m3) and aerated with synthetic air (80% N2 and 20% O2). The duration of the experiment was 100 min. The pH was measured continually and four samples (30 mL each) were taken to follow the carbonate concentration. The KL values were calculated by fitting the pH values and the total carbonate concentrations in a simple computer model based on the two film theory.

FIGURE 2. Cycle starting on day 454 in reactor thin: pH (without symbols), total nitrite (2), nitrate (b), ammonium (9). the rest of the observation period, the reactors steadily produced a 1:1 ammonium nitrate solution with only brief disturbances caused by inhibitory effects of the influent. In this paper we describe the processes during complete nitrite oxidation. Figure 2 shows a typical cycle measured at the end of the observation period. In the beginning of the cycle, the pH decreased rapidly, reached a plateau at pH values around 4, and declined further to values of 3 and below. At the end of the cycle, nitrate was completely degraded, and the pH value was very low. During the observation period, the lowest pH values were reached in reactor deep with a minimum of 2.5, but in later measurements even lower pH values were found. Very low pH values slowed the process in the next cycle. Two buffer systems, bicarbonate/carbon dioxide (pK 6.35) (26) and nitrite/nitrous acid (pK 3.29) (27), controlled the form of the pH curve. In most cycles, however, the bicarbonate concentration was too low to act as substantial buffer. The nitrite concentration slightly increased at cycle start. After this initial increase the nitrite concentration declined following a sigmoid curve with the highest rate at pH values around 4. Except for the initial phase, nitrate concentration mirrored the trend of the nitrite concentration. Ammonium degradation was very slow compared to the nitrite decrease. In the experiment shown in Figure 2, the ammonium concentration decreased by about 12%. Inhibition of the ammonium oxidation seemed to occur at low pH values, but interpretation is hampered by the variance of the measurements (Figure 10). Nsum declined during the cycle. The average nitrogen loss, calculated from 10 separate cycles, was 8 ( 4%. The reactors were aerobic at all times, having minimum oxygen concentrations of 0.097 and 0.041 mM in reactor thin and reactor deep, respectively. These minima occurred only for short periods. The average oxygen concentrations were 0.184 ( 0.038 mM in reactor thin and 0.163 ( 0.059 mM in reactor deep. In Figure 3, the HNO2 concentration and the nitrite degradation rate in one cycle are compared. The maximum HNO2 concentration (1.4 mM) and nitrite degradation rate (8.4 mM/d) occur at a pH of 4.25. Since NOB are strongly inhibited by HNO2 at such high concentrations (28), the strong correlation of HNO2 concentration and total nitrite degradation suggests that the nitrite oxidation is not microbial.

Hypothesis Results from the Steady-State Phase In the early phase of the experiments, nitrification stopped at a pH of 5.8 and hardly any nitrate was produced. However, nitrate production slowly increased and the pH value fell below 5.8. After 5.3 and 8.9 months of operation in reactor deep and reactor thin, respectively, the pH decreased to very low values and nearly all nitrite was oxidized to nitrate. For 4068

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To our knowledge, the phenomena which we observed have not been described in detail before. To explain the processes we made the following hypothesis and tested it with a computer model and additional experiments: The oxidation of nitrite to nitrate is caused by chemical processes, which are triggered and mediated by microbial ammonium oxidation.

dinitrogen trioxide N2O3 with NH3 is the rate-determining step:

N2O3 + NH3 f N2 + NO2- + H+ + H2O

(4)

N2O3 is in equilibrium with NO and NO2

NO + NO2 T N2O3

FIGURE 3. HNO2 concentration (4) and nitrite degradation rate rnitrite (b) in reactor thin (cycle starting on day 454). The HNO2 concentration was calculated using the pK_nitrite given in Table 3. The nitrite degradation rate was interpolated with quadratic functions.

Model Four basic processes were examined to describe pH decrease and nitrite degradation: (1) microbial ammonium oxidation; (2) microbial nitrite oxidation; (3) chemical ammonium oxidation; and (4) chemical nitrite oxidation. Additionally, we modeled the gas exchange processes for nitrogen oxide NO, nitrous acid HNO2, and carbon dioxide CO2 and all necessary acid-base equilibria. The computer model was written in the simulation software Aquasim 2.0 (29). All kinetic constants for the chemical processes were taken from literature. The constants are given for 25 °C, if not stated otherwise. We assumed that all literature data were valid for zero ionic strength. Ion activity coefficients were calculated with the approach of Davies as stated in Stumm and Morgan (26). The gas transfer coefficients KL for the gas exchange were either measured or fitted. The parameters for the microbial processes were fitted with the experimental data. The process rates, kinetic constants, and equilibrium constants are given in Tables 1, 2, and 3. Microbial Processes. Since it was beyond the scope of this study to model the biofilm kinetics in detail, we used simple kinetic expressions common for reactors with suspended biomass (30). The stoichiometry and the kinetic expressions for the bacteria are given in Table 1. The following kinetic parameters were unknown and had to be fitted: • The activities AAOB and ANOB for AOB and NOB, respectively. These parameters include the biomass concentration, the growth rate and the biomass yield. The combination of these factors was possible due to the very low bacterial growth rate and the short cycles to be simulated. However, the fitted activities AAOB were slightly different in all simulated cycles because of the increase in biomass. • A purely empirical non-competitive proton inhibition factor Ki proton for AOB. As stated above, pH may cause a large number of different inhibition effects. Including them all would result in a grossly overparameterized model. • A switching function with the nitrous acid half saturation constant KHNO2 for NOB. We used KHNO2 to stop the NOB activity at low substrate concentrations, thereby preventing negative concentrations of nitrite. As we will discuss later, values for ANOB and KHNO2 could not be determined because the activity of the NOB was too low. Chemical Ammonium Oxidation. Ammonium nitrite solutions decompose to N2 and H2O (5). UV radiation, which enhances this reaction (22), could be neglected in the glass reactors. Harrison et al. (5) showed that the reaction of

(5)

In eq 4, we used the rate constant by Harrison et al. (5). The kinetic constant for the forward reaction in eq 5 has been given by Gra¨tzel et al. (31), and the equilibrium constant was recommended by Schwartz and White (32). Chemical Nitrite Oxidation. We surveyed the literature for chemical nitrite degradation processes. Due to the conditions in the reactors, we could exclude the reaction with transition metals such as Fe, Cu, and Mn in their reduced state (4, 33), the photolytic degradation of nitrite catalyzed by Fe(III) (34), and reactions with organic substances and amino compounds (4, 33). Nitrous acid also reacts with hydroxylamine (35, 36) to produce nitrous oxide, but this process could not explain the strong decrease of nitrite. We will discuss this process later with respect to the volatilization of nitrogenous compounds. Finally, we focused on the self-decomposition of nitrous acid under aerobic conditions. Nitrite self-decomposition has been widely discussed in soil science as the most important process of chemodenitrification (4, 33). Additionally, the reaction pathways have been intensively investigated in atmospheric chemistry (32, 37) and in connection with the removal of NO and NO2 from flue gas and the wet analysis of these gases (32). The complete reaction to nitrate involves three steps (38)

2 HNO2 T NO + NO2 + H2O

(6)

NO + 0.5 O2 f NO2

(7)

2 NO2 + H2O T HNO2 + NO3- + H+

(8)

The general mechanism for the chemical oxidation of nitrite to nitrate was postulated as early as 1925 by Reinders and Vles (39). The reactions in eqs 6-8 have been proven by several authors (32, 33, 38, 40). Equations 6-8 can be combined to give

HNO2 + 0.5 O2 f NO3- + H+

(9)

if no intermediates volatilize. The overall stoichiometry of the chemical nitrite oxidation is therefore essentially the same as that for the microbial nitrite oxidation (eq 3). N2O4 is often mentioned as an intermediate of the processes involved in the chemical nitrite oxidation to nitrate (32). We calculated the equilibrium of 2 NO2 T N2O4 with kinetic data from Gra¨tzel et al. (31) and Schwartz and White (32). The resulting N2O4 concentrations were very low, therefore we neglected this compound in further simulations. For the forward and backward reaction of eq 6 and for the forward reaction of eq 8 we used kinetic constants from Park and Lee (40). The constants were interpolated for 25 °C assuming a quadratic temperature dependence. Park and Lee (40) did not give any value for the backward reaction of eq 8, probably because this process is negligibly slow compared to the forward reaction. For the sake of completeness, we included this reaction in our model using the average value of three reaction constants listed in ref 32. The simulations showed that the nitrite oxidation is highly sensitive to the kinetic constant of eq 7. We applied the constant by Awad and Stanbury (41). An overview of other constants for the same process is given by Pires et al. (42). VOL. 39, NO. 11, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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TABLE 1. Processes and Process Rates in the Computer Modela process

process rate Acid-Base Equilibria k_H2PO4_back‚fA1‚[H2PO4-]‚fA1‚[H+] k_H2PO4_back‚10-pK_H2PO4‚[H3PO4]‚1000 k_HPO4_back‚fA2‚[HPO42-]‚fA1‚[H+] k_HPO4_back‚10-pK_HPO4‚fA1‚[H2PO4-]‚1000 k_NH3_back‚[NH3]‚fA1‚[H+] k_NH3_back‚10-pK_NH3‚fA1‚[NH4+]‚1000 k_nitrite_back‚fA1‚[NO2-]‚fA1‚[H+] k_nitrite_back‚10-pK_nitrite‚[HNO2]‚1000 k_HCO3_back‚fA1‚[HCO3-]‚fA1‚[H+] k_HCO3_back‚10-pK_HCO3‚[H2CO3]‚1000

H2PO4- + H+ f H3PO4 H3PO4 f H2PO4- + H+ HPO42- + H+ f H2PO4H2PO4- f HPO42- + H+ NH3 + H+ f NH4+ NH4+ f NH3 + H+ NO2- + H+ f HNO2 HNO2 f NO2- + H+ HCO3- + H+ f H2CO3 H2CO3 f HCO3- + H+ NO + NO2 + H2O f 2 HNO2 2 HNO2 f NO + NO2 + H2O HNO2 + NO3- + H+ f 2 NO2 + H2O 2 NO2 + H2O f HNO2 + NO3- + H+ N2O3 f NO + NO2 NO + NO2 f N2O3

Nitrogen Compounds Equilibria k_NO_back‚[NO]‚[NO2] k_NO_for‚[HNO2]2 k_NO3_back‚[HNO2]‚fA1‚[NO3-]‚fA1‚[H+] k_NO3_for‚[NO2]2 k_N2O3_for‚10pK_N2O3‚[N2O3]‚1000 k_N2O3_for‚[NO]‚[NO2]

N2O3 + NH3 f N2 + HNO2 + H2O 2 NO + O2 f 2 NO2

Chemical Nitrogen Conversion k_NH3_nitro‚[NH3]‚[N2O3] k_NO_ox‚[NO]2‚[O2] Gas Exchange H_CO2‚([CO2] -[CO2 sat])‚Qgas/V‚(1 - e-KL_CO2‚a‚V/H_CO2/Qgas) H_HNO2‚[HNO2]‚Qgas/V‚(1 - e-KL_HNO2‚a‚V/H_HNO2/Qgas) H_NO‚[NO]‚Qgas/V‚(1 - e-KL_NO‚a‚V/H_NO/Qgas)

CO2(aq) f CO2(g) HNO2(aq) f HNO2(g) NO(aq) f NO(g)

Microbial Processes Ki proton/(Ki proton + [H+])‚AAOB [HNO2]/(KHNO2+[HNO2])‚ANOB

NH4+ + 1.5 O2 f NO2- + 2 H+ + H2O HNO2 + 0.5 O2 f NO3- + H+

a All concentrations in [mM]. f , f : activity coefficients. V: volume of reactor. a: ratio of tubing surface area to reactor volume. Q A1 A2 gas: gas flow through tubing. CO2_sat: CO2 concentration in water at equilibrium with air.

TABLE 2. Kinetic Constants in the Computer Modela kinetic constant

value

unit

reference

Acid-Base Equilibria k_H2PO4_back 3.5‚1011 1/mM/d based on (43) p 194 k_HPO4_back 3.5‚1011 1/mM/d based on (43) p 194 k_NH3_back 3.5‚1011 1/mM/d based on (43) p 194 k_nitrite_back 3.5‚1011 1/mM/d based on (43) p 194 k_HCO3_back 3.5‚1011 1/mM/d based on (43) p 194

k_NO_back k_NO_for k_NO3_back k_NO3_for k_N2O3_for

N Compounds Equilibria 1.4‚1010 1/mM/d (40), interpolated 1.6‚103 1/mM/d (40), interpolated 7.3‚10-4 1/mM2/d (32) p 26, average 6.9‚109 1/mM/d (40), interpolated 9.5‚1010 1/mM/d (31), 20 °C

k_NH3_nitro k_NO_ox

Chemical Nitrogen Conversion 7.7‚107 1/mM/d (5) 1.8‚105 1/mM2/d (41)

KL_HNO2 KL_NO KL_CO2

Gas Exchange 1.4‚10-2 m/d 9.8‚10-2 m/d 0.5 m/d

AAOB ANOB Ki proton KHNO2

Microbial Processes 0.37 to 0.75 fitted for every cycle no fit possible 0.5 mM fitted mM no fit possible

a

fitted fitted experimentally determined

The constants are rounded to two significant digits.

Gas Exchange Processes. The gas exchange processes were described according to the two film theory. We did not model any reactions of the gaseous compounds in the biofilm or silicone membrane. We consider the errors to be insignificant compared to the variations in the thermodynamic data and measurements. Gas exchange processes were 4070

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modeled for nitrogen oxide NO, nitrous acid HNO2, and carbon dioxide CO2. For oxygen, the on-line measurements in the reactor were used as input data. We did not simulate the volatilization of NO2, N2O3, and NH3 because the concentrations were too low to cause a substantial nitrogen loss. N2 volatilization was not modeled because N2 was assumed to be inert in the investigated systems and to degas very fast. The CO2 volatilization had some importance for the carbonate buffer system. Its KL was determined experimentally. The KL for NO and HNO2 were used as main fitting parameters during calibration (see below). Acid-Base Equilibria. The acid-base equilibria were calculated by splitting the equilibria in a backward and forward reaction. The backward reactions are generally diffusion-controlled and have a rate constant of about 3.5. 1011 1/mM/d (43). Initial Concentrations. All initial concentrations were measured except for the bacteria, carbonate, and phosphate concentrations. Only occasional tests for carbonate and phosphate were done. For all cycles, the same initial carbonate and phosphate concentrations were used, unless specific measurements were available. Errors in the concentration of these buffers may be one reason for deviations of the pH values. Calibration. The four constants that had to be fitted were the KL for NO and HNO2, the activity of the AOB, AAOB, and the inhibition constant for the pH inhibition effects on AOB, Ki proton. The fit for the parameters of the NOB activity always resulted in values close to zero. We therefore postulated the NOB activity to be negligible and did not consider it for further calibration. Data from the experiment without biomass were used to fit the KL for NO and HNO2. The parameters were fitted to time-dependent NO off-gas concentrations and on pH values, respectively. Since both coefficients are interdependent, they

TABLE 3. Equilibrium Constants in the Computer Modela

a

parameter

description

value

unit

reference

H_CO2 H_HNO2 H_NO pK_H2PO4 pK_HPO4 pK_NH3 pK_nitrite pK_HCO3 pK_N2O3

Henry coefficient for CO2 Henry coefficient for HNO2 Henry coefficient for NO H3PO4 T H2PO4- + H+ H2PO4- T HPO42- + H+ NH4+ T NH3 + H+ HNO2 T NO2- + H+ H2CO3 T HCO3- + H+ NO + NO2 T N2O3

1.2 8.3‚10-4 21 2.15 7.21 9.24 3.29 6.35 -4.5

M(g)/M(aq) M(g)/M(aq) M(g)/M(aq)

(26) (27) p 16 (27) p 33 (44) (44) (45) (27) p 16 (26) (32) p 91

Values given for T ) 25 °C and I ) 0 M.

FIGURE 4. Calibration result for the fitting of the gas transfer coefficient KL for NO and HNO2 on pH (O) and NO (4) in an experiment without biological activity.

FIGURE 5. Measurements and simulated values for total nitrite (4), nitrate (O), and ammonium (0) in an experiment without biological activity. The high nitrite degradation rates are due to the high oxygen concentrations. were fitted iteratively. The KL for NO and HNO2 are the only fitting parameters for the abiotic processes. The fitting of these constants also compensates for the variability and possible errors of the literature data. The fit of nitrite, nitrate, and pH was good and the extent of NO degassing could be matched (Figures 4 and 5). Ki proton was fitted together with AAOB on the pH of a late cycle in reactor thin. The two parameters correlated with a correlation coefficient r of 0.87 in this fit (secant method), which was assumed to be low enough to separate both coefficients. The fits of pH, nitrite, nitrate, and ammonium were very good (Figures 6 and 7). Validation. The model was validated on pH, nitrite, nitrate, and ammonium in five single cycles. One cycle was measured in reactor thin and four were measured in reactor deep. An example is given in Figures 8-10. The fit of nitrite, nitrate, and pH was very good in all cycles, slight deviations were only found for ammonium.

FIGURE 6. Calibration result for the fitting of Ki pH (reactor thin, cycle starting on day 454).

proton

and AAOB on

FIGURE 7. Simulated total nitrite (4), nitrate (O), and ammonium (0) in the cycle used for calibration of Ki, proton and AAOB (reactor thin, cycle starting on day 454). Activity of NOB was not needed to simulate the oxidation of nitrite to nitrate, but the pH decrease could not be simulated without the activity of AOB. The chemical ammonium oxidation was too slow to explain the strong pH decrease at the beginning of the cycle.

Additional Experimental Validation of Main Results Microbial Ammonia Oxidation at Low pH Values. The specific AOB inhibitor ATU was used to determine whether microbial ammonium oxidation could be responsible for the pH decrease to values far below 5.5. ATU inhibits the enzyme ammonia monooxygenase (AMO) of lithoautotrophic AOB (46), which catalyzes the first step of the ammonium oxidation (eq 1). Some, but not all, heterotrophic AOB are also inhibited by ATU (7, 47, 48). The ATU experiment was conducted 43 months after reactor start, when the reactors were steadily producing ammonium nitrate. The pH usually reached values below 3. VOL. 39, NO. 11, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 8. Validation of pH on reactor deep (cycle starting on day 454).

FIGURE 11. pH decrease in reactor thin with (O) and without ATU dosing (b). Data from two consecutive cycles, 43 months after cycle start.

FIGURE 9. Validation of total nitrite (4) and nitrate (O) (reactor deep, cycle starting on day 454).

FIGURE 12. Simulated and measured concentrations for total nitrite (4), nitrate (O), and ammonium (0) in the experiment under sterile conditions.

FIGURE 10. Validation of ammonium (reactor deep, cycle starting on day 454). Simulations had shown that below pH 4.0 the chemical processes alone were able to decrease the pH. Therefore we chose a pH value of 4.75 for ATU addition. This pH value is far below the assumed limit for nitrification in wastewater systems (pH value 5.5) and it is still above the value where the chemical processes could cause a pH decrease. The addition of ATU caused an immediate stop of the pH decrease (Figure 11), proving that AOB are responsible for the pH decrease. Furthermore, the results suggest that the AOB are lithoautotrophic. However, the activity of some heterotrophic AOB cannot be excluded. Validation of Chemical Nitrite Oxidation under Sterile Conditions. The hypothesis of chemical nitrite oxidation was tested in a sterile membrane aerated reactor filled with a synthetic ammonium nitrite solution. To simulate the proton production by AOB, a 0.1 M HCl solution was added continually. The initial pump rate was 20 mL/d. To reflect 4072

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FIGURE 13. Measured and simulated pH values in the experiment under sterile conditions. the inhibition of AOB, the pump rate was reduced daily according to the AOB inhibition term given in Table 1. Acid addition was finally stopped at a pH of 3.5. The nitrogen and pH measurements in the sterile reactor exhibit the same course as the values in the biofilm reactors (Figures 2 and 12). The slightly faster nitrite degradation is due to higher oxygen concentrations. These data were used to validate the computer model which was enhanced with an input term describing the acid addition. Nitrite, nitrate, and pH measurements were simulated very well (Figure 12), whereas the simulated ammonia values were slightly higher than those measured (Figure 13). Possibly, the kinetic constant for the chemical ammonia degradation is too low. However, this small deviation has hardly any influence on the simulation of the total nitrogen conversion. A second experiment gave similar results (not shown).

Confirmation of the Hypothesis The above-described simulations and experiments confirmed that the oxidation of nitrite to nitrate is caused by chemical processes and that the chemical processes are triggered by acid-tolerant AOB. (1) The chemical nitrite oxidation to nitrate could be simulated with kinetic constants reported in the literature without introducing NOB. Experiments in a sterile system further confirmed the hypothesis. (2) The initial pH decrease could not be simulated with any chemical oxidation process found in the literature. The introduction of acidtolerant AOB was necessary. We finally confirmed the significance of acid-tolerant AOB with the addition of ATU. Although the chemical ammonium oxidation contributes substantially to the ammonium degradation (at least 27% in the experiment presented in Figure 10), the process is only important when NH3 and N2O3 occur concomitantly at high concentrations. Since NH3 is mainly found under alkaline conditions and N2O3 is found under acid conditions, the maximum rates are found in a narrow pH range around 4.4. The simulations showed that chemical nitrite oxidation causes a further pH decrease as soon as pH values of 4.0 were reached. Above this limit, microbial proton production was necessary to decrease the pH value. The model is highly reliable, because only two physical and two microbial parameters had to be fitted. The simulation of pH, nitrite, nitrate, and ammonium could be validated in five single cycles. The experiment under sterile conditions showed that the assumed rate of chemical ammonium oxidation might be slightly too low and the model could be improved by determining a new rate constant. However, the chemical ammonium oxidation is only of minor importance for the overall process. The acid-tolerant AOB might have adapted to the low pH values due to the pH fluctuations during the cycles, as previously reported by DeBoer et al. (49). Nitrogen Losses. According to measurements in five cycles of each reactor, the average loss of Nsum was 8 ( 4%. That means 16 ( 8% of the transformed nitrogen was lost by volatilization. The nitrogen loss was significantly higher in reactor thin with 11 ( 2% compared to reactor deep with 5 ( 2%, probably due to longer retention times in the deep biofilm, where volatile nitrogen compounds could be transformed to nitrate. The nitrogen loss in the experiment under sterile conditions was 10.4%, which is very close to the average value of reactor thin. This suggests that the nitrogen loss in the experiments with biomass is not caused by biological processes. Simulation Results on Nitrogen Volatilization. Six of the modeled nitrogen compounds were volatile: HNO2, NO, NO2, N2O3, NH3, and N2. The simulations showed that only three of the volatile nitrogen compoundssHNO2, NO, and N2s occurred in a concentration range that was high enough to contribute substantially to nitrogen volatilization. Generally, the simulations fitted the measured nitrogen loss very well. On the basis of the simulations of five cycles of each reactor, the average loss of Nsum was 8.6 ( 0.4%. This value is very close to the measured 8 ( 4%. The contribution of HNO2, N2, and NO to the nitrogen loss was 61 ( 8%, 37 ( 8%, and 2.0 ( 0.3%, respectively, suggesting that the nitrogen loss was mainly due to HNO2 and N2 volatilization. N2 was produced by chemical ammonium oxidation. Since the modeled rate of chemical ammonium oxidation might be slightly too low, the contribution of N2 to the total nitrogen volatilization could be slightly higher. NO and NO2 are also intermediates of microbial ammonium oxidation (11, 50) and NO can be produced by NOB during nitrite reduction (6). Those processes were not included in the model. However, these processes were assumed to be negligible because the bacterial activity was

FIGURE 14. pH (without symbol) and N2O (b) in reactor deep, cycle starting on day 510. low and the nitrogen losses in the biological and sterile experiments were in good agreement. N2O Measurements. One important volatile nitrogen compound had not been included in the model, although it was detected in the off-gas: nitrous oxide N2O (Figure 14). High N2O concentrations in the beginning had been caused by denitrification during storage of the influent. They decreased quickly by volatilization. Thereafter the N2O concentrations remained near the background level (0.3 ppm) until the pH value fell below 4. At these very low pH values, significant amounts of N2O were produced in the reactors. The lost N2O corresponded with 17% of the Nsum loss in this cycle. We surveyed the literature for chemical processes which might be responsible for this N2O production. All mechanisms found require hydroxylamine as reactant. Hydroxylamine may be released by AOB as intermediate of the ammonium oxidation. We assume that hydroxylamine was present in the reactors, but its concentration must have been very low, because we did not find any hydroxylamine at a concentration level of 10-5 M. Some of the processes found in the literature could be excluded because they require high pH values: disproportionation of hydroxylamine (51), autoxidation of hydroxylamine (52), and oxidation of hydroxylamine with NO (53). The reaction of hydroxylamine with nitric acid HNO3 could be excluded because it occurs only at very high acidity (54). Finally, we found two processes which we considered in more detail. (1) Reaction of hydroxylamine with nitrous acid (35, 36). Using kinetic data by Do¨ring and Gehlen (35), we estimated that at a pH value of 3, the total hydroxylamine concentration should have been in the range of 0.1 M. This concentration is 4 orders of magnitude higher than the detection level. Therefore, this reaction cannot explain the high N2O production. (2) Fe(III) catalysis of hydroxylamine oxidation (55, 56). We used kinetic data by Butler and Gordon (56) to estimate the necessary concentration of hydroxylamine. The concentration of dissolved iron was assumed to be 2.5 µM as measured in a later cycle. The concentration of the reacting Fe(III) species Fe(OH)2+ was calculated with the program PHREEQC (57). The calculation revealed that the necessary total hydroxylamine concentration must have been higher than 10-3 M, which is at least 2 orders of magnitude above the detection level. Thus, catalysis by dissolved Fe(III) cannot explain the high N2O concentrations. Nevertheless, Fe(III) catalysis may have been responsible for the N2O production. Instead of dissolved iron, solid iron from the steel rack may have reacted with hydroxylamine. At low pH values the protective passive oxide film on stainless steel surfaces is dissolved (26), so that the solid iron is in direct contact with the solution. Unfortunately, we did not find any reaction rates for such a process in the literature. VOL. 39, NO. 11, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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Besides the abiotic processes, N2O may also be produced by bacteria. Whereas denitrification by autotrophic and heterotrophic bacteria (58) is unlikely due to the high oxygen concentrations in the reactors, N2O occurs during the microbial oxidation of hydroxylamine to nitrite (50), especially when the hydroxylamine oxidation is inhibited. Therefore, AOB may have been a possible source for N2O under the inhibiting conditions at low pH values. The total nitrogen loss amounts to 10.4%, if the measured N2O degassing is added to the simulated nitrogen volatilization. This value is still in the range of the measured 8 ( 4%. Considering N2O as an additional degassing nitrogen compound, about 1/2 of the degassing nitrogen was HNO2, about 1 /3 was N2, and about 1/6 was N2O. NO degassing was below 2%. Nitrogen Emissions from Acid Soils. Simulations with our model might also help to better understand the processes that cause nitrogen volatilization in acid soils. Nitrogen emissions from acid soils are important sources of atmospheric NOx and N2O (59). In contrast to our experiments, ammonium and nitrite concentrations are generally lower in acid soils. Therefore, microbial nitrite and ammonium oxidation might be more important for the production of gaseous nitrogen compounds than in our reactors (59). However, it is also possible that the contribution of chemical nitrite and ammonium oxidation has been underestimated so far. In contrast to the general assumption for acid soils (4), we could show that chemical nitrite oxidation is not the main cause for nitrogen volatilization in acid nitrite solutions under aerobic conditions. Its products NO and NO2 contribute only marginally to the nitrogen volatilization. Interestingly, HNO2 has not been considered as a main degassing nitrogen oxide in acid forest soils before. One reason may be that HNO2 is rapidly transformed to NO and NO2 in the atmosphere. Additionally, HNO2 may be wrongly detected as NO2 in NOx analyzers (37). Besides the HNO2 volatilization, the production of N2 by chemical ammonium oxidation caused substantial nitrogen losses. Chemical ammonium oxidation has not been considered to be an important N2 source in acid soils so far. Possibly, this might have led to wrong interpretations of the microbial denitrification capacity.

Acknowledgments We thank Martin Biebow, Norbert Baumga¨rtel, Sabine Frommhold, Mariska Ronteltap, Steffen Zuleeg, and our laboratory staff Claudia Ba¨nninger, Karin Rottermann, Annemarie Mezzanotte, and Irene Brunner for their support and great commitment during the experimental work. We are also grateful to Beat Schwarzenbach, Lukas Emmenegger, and Peter Honegger from the Swiss Federal Laboratories for Materials Testing and Research (EMPA) for providing us with the analytical devices for the NOx and N2O measurements.

Nomenclature AAOB

activity of AOB (includes biomass concentration, yield, and growth rate)

AMO

ammonia monooxygenase

AOB

ammonium oxidizing bacteria

ATU

allylthiourea, inhibitor for AOB

KHNO2

nitrite half-saturation coefficient for NOB

Ki proton

proton inhibition coefficient for AOB

KL

gas transfer coefficient

NOB

nitrite oxidizing bacteria

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NOx

sum of nitrogen oxide NO and dinitrogen oxide NO2.

Nsum

sum of the dissolved nitrogen compounds total nitrite, nitrate, and ammonium

SBR

sequencing batch reactor (used for producing ammonium nitrite)

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Received for review October 7, 2004. Revised manuscript received March 21, 2005. Accepted March 23, 2005. ES048422M

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