Reactions of Sulfur Dioxide with Ammonia: Dependence on Oxygen

Oct 8, 1996 - The resulting hydrogen sulfate and sulfate ions in the aqueous layer react ... for ammonia and sulfur dioxide in the reactor (Bai et al...
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Ind. Eng. Chem. Res. 1996, 35, 3362-3368

Reactions of Sulfur Dioxide with Ammonia: Dependence on Oxygen and Nitric Oxide Koichi Hirota,*,† Jyrki Ma1 kela1 ,‡ and Okihiro Tokunaga† Department of Radiation Research for Environment and Resources, Japan Atomic Energy Research Institute, 1233 Watanuki, Takasaki, Gunma 370-12, Japan

The influence of oxygen and nitric oxide on the reactions of sulfur dioxide with ammonia were studied in a simulated flue gas in the range of 0-20% oxygen and 0-300 ppm nitric oxide at temperatures in the range of 40-60 °C. A Fourier transform infrared spectrometer (FT-IR) analyzed the reaction products deposited on the reactor surface and revealed that ammonium sulfate was the main product of the reactions, with sulfamic acid and ammonium sulfamate as the minor products. The results showed that oxygen and nitric oxide enhanced the oxidation reactions of sulfur dioxide to form ammonium sulfate. The yield of the minor products markedly increased in the presence of nitrogen dioxide. The size and number concentration of product aerosols increased at lower temperature. The fraction of sulfur dioxide which formed aerosols increased with sulfur dioxide removal. Introduction To save our environment from the emission of harmful pollutants such as SO2, NOx, etc., treatment is applied to flue gases in Japan. High removals of the pollutants are achieved by methods that meet strict regulations established by the Japanese government and the local governments. The wet lime scrubber method and the selective catalytic reduction method can treat SO2 and NOx, respectively. However, the wet lime scrubber method needs the treatment of waste water, and the selective catalytic reduction method needs expensive catalyst. Combination of the two methods is necessary to treat the toxic pollutants. Investigation of a new method is expected for the simple and simultaneous treatment process of the pollutants. The electron beam irradiation process for flue gas purification has been proposed as an efficient method (Tokunaga et al., 1993). Sufficient removal performance for NOx and SO2 in the process was demonstrated for flue gases from the coal-fired power generator (Namba et al., 1995), the municipal waste incinerator (Osada et al., 1995; Hirota et al., 1995), and the sintering machine (Kawamura et al., 1980). The principal of the removal reactions of NOx and SO2 in the process is quite simple: active species, such as hydroxyl radicals produced by electron beam irradiation of humid flue gas, oxidize sulfur dioxides and nitrogen oxides in the flue gas to yield sulfuric and nitric acids, respectively. Then, the acids immediately react with ammonia added to the flue gas to give powdery ammonium sulfates and nitrates, which can readily be collected by electrostatic precipitators. These ammonium salts can then be used as fertilizers. In this process, some sulfur dioxide can be removed through direct reactions with ammonia in the absence of electron beam irradiation. The reactions are expected to save energy in the process because some sulfur dioxide is treated without the irradiation which needs electricity. Hirota et al. (1993) have done laboratoryscale studies on the reactions of sulfur dioxide with ammonia and proposed a feasible reaction mechanism, * To whom correspondence should be addressed. † Japan Atomic Energy Research Institute. Fax 81-273-469688. ‡ Current address: Department of Physics, P.O. Box 9, University of Helsinki, Helsinki, Fin-00014, Finland.

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where the reactions occur in a water layer formed on a reactor surface by water condensation out of the gas phase. In their study, the reactions were initiated by the dissolution of gaseous sulfur dioxide and ammonia into the water layer. The dissolution of sulfur dioxide forms hydrogen sulfite and sulfite ions in the layer. The hydrogen sulfite and sulfite ions may heterogeneously be oxidized to hydrogen sulfate and sulfate ions, respectively, by gaseous oxygen. The resulting hydrogen sulfate and sulfate ions in the aqueous layer react with ammonium ions to form ammonium sulfates. Thus, oxygen is expected to play an oxidation role in the reactions of sulfur dioxide with ammonia. Moreover, our preliminary experiments showed that oxygen depressed the formation of sulfamic acid (HOSO2NH2). This is desirable because HOSO2NH2 cannot be used as fertilizers. It is therefore essential to clarify the role of oxygen in the reactions. The present study characterizes the effect not only of oxygen but also of nitric oxide on the reactions. Nitrogen dioxide influence on the reaction products is also studied. Experimental Section Apparatus. The experimental setup used for the studies on NH3-SO2 reactions is schematically shown in Figure 1. The setup consists of a gas-supply system, a water-supply system, a reaction system, and analysis devices. In the gas-supply system, pressurized gas cylinders of 1% SO2, 3% NH3, and 3000 ppm NO diluted with nitrogen and pure oxygen, carbon dioxide, and nitrogen were used for the preparation of simulated coal-fired flue gas. The flow rates of these gases were controlled separately using flow meters. All tubes of the gases except for NH3 were connected to a reactor through a mixing vessel filled with glass-made cylindrical rings of radius 6 mm and length 25 mm. The reactor and the mixing vessel were placed in an oven (Eyela: WFO-1000ND). An ammonia gas tube was directly attached to the reactor. Reaction temperature was monitored with a thermocouple detector placed at the center of the reactor. Water vapor was charged to the simulated flue gas with nitrogen gas bubbled through water in a water bath (Eyela: Digital Uni Ace UA-100) placed prior to the mixing vessel. The temperature of the water bath was controlled precisely for the desired © 1996 American Chemical Society

Ind. Eng. Chem. Res., Vol. 35, No. 10, 1996 3363

Figure 1. Schematic flow chart of experiments.

Figure 2. Design of a reactor and a mixing vessel.

water concentration in nitrogen. A tube that carried nitrogen from the water bath to the oven was heated to 100 °C to avoid water condensation. The progress of the reactions was continuously observed by the measurement of the concentration of unreacted sulfur dioxide in the simulated flue gas with an infrared absorption type sulfur dioxide detector (Shimazdu: IRA 107). An ammonia scrubber installed upstream from the detector consisted of a filter and a glass fiber unit. The glass fiber was soaked with concentrated phosphoric acid for absorbing unreacted ammonia. This ensured the correct measurement of sulfur dioxide concentration with the detector. Gas tubes from the reactor to the ammonia scrubber and from the reactor to an aerosol measurement system were also heated to 100 °C to avoid further reactions. The size and the number concentration of aerosols formed by the NH3-SO2 reaction were also measured downstream from the reactor with an active scattering aerosol spectrometer (Model ASASP-300X (HB)) and a microlaser particle spectrometer mLPS (Particle Measuring System Inc.) when the sulfur dioxide concentration was kept constant. The total flow rate of the gas mixture was 10 L/min, and the reaction time in the reactor was 10 s. Reactor and Mixing Vessel. It is of importance to show a detailed design of the mixing vessel and the reactor since the mixing condition for ammonia and sulfur dioxide in the simulated flue gas has an influence on the reactions, depending on the relation of the inlet positions for ammonia and sulfur dioxide in the reactor (Bai et al., 1992). Figure 2 shows the reactor and the mixing vessel used for the present study. They are made of Pyrex glass, 80 mm in diameter and 330 mm in length. The reactor has a 1/2 in. inlet port for the simulated flue gas and a 1/4 in. port for ammonia gas.

These two ports are placed on a concentric circle. The reactor can be disassembled into two parts for the cleaning of its inner wall after experiments. The two parts are connected by a stainless steel clamp with a rubber o-ring. The mixing vessel has a 1/2 in. inlet port for the mixing gas of oxygen and nitrogen, and three 1/4 in. ports for SO2, NO, and CO2. These gases are sufficiently mixed with the cylindrical rings. Procedure. The temperature of the oven was first set to approximately 60 °C at which temperature the reactions never occur under our experimental conditions. The flow-controlled gases from the cylinders, except for ammonia, were introduced into the reactor through the mixing vessel. Ammonia was fed into the reactor when the gas temperature and the sulfur dioxide concentration in the reactor were stabilized. It took about 30 min to get the constant concentration of sulfur dioxide after the addition of ammonia. After the constant concentration, the oven temperature was lowered at 5 °C intervals to 45 °C and then at 2-3 °C intervals to 40 °C. Once the reactions occurred after the change in the oven temperature, the concentration of unreacted sulfur dioxide measured with the sulfur dioxide detector fell slowly and then came to a constant concentration. The constant concentration was obtained about an hour in the temperature range of 60-45 °C and less than 2 h in the temperature range of 45-40 °C after the temperature was changed to the new level. Uniformly-deposited reaction products were observed on the inside wall of the reactor during the reactions. A small amount of product was partially attached as a white film just on the ammonia gas outlet inside the reactor. The extent of the reaction of sulfur dioxide with ammonia was determined from the sulfur dioxide removed. Analysis of Reaction Products. White-colored reaction products were analyzed using an FTIR spectrometer (JEOL: JIR-100D). Samples were collected on the reactor surface for 2 h at a temperature of 45 °C and were scratched with a spatula from the surface. Then, a small amount of the product was mixed with KBr and pressed to make a thin film for the FTIR analysis. The spectra obtained were calculated from an average of 20 scans at a 1 cm-1 resolution. Results Sulfur Dioxide Removal Dependence on Oxygen. The influence of oxygen on the sulfur dioxide

3364 Ind. Eng. Chem. Res., Vol. 35, No. 10, 1996

Figure 3. SO2 removal efficiency as a function of reaction temperature for a variety of oxygen concentrations in the range of 0-20%. Inlet gas: 600 ppm SO2, 1200 ppm NH3, 225 ppm NO, 0-20% O2, 12.5% CO2, 5% H2O, and N2 (balance).

Figure 5. SO2 removal efficiency as a function of reaction temperature for a variety of nitric oxide concentrations in the range of 0-300 ppm. Inlet gas: 600 ppm SO2, 1200 ppm NH3, 0-300 ppm NO, 10% O2, 12.5% CO2, 5% H2O, and N2 (balance).

Figure 4. SO2 removal efficiency dependence on oxygen concentration at various reaction temperatures in the range of 40-50 °C. Inlet gas: 600 ppm SO2, 1200 ppm NH3, 225 ppm NO, 0-20% O2, 12.5% CO2, 5% H2O, and N2 (balance).

Figure 6. SO2 removal efficiency dependence on nitric oxide concentration at various reaction temperatures in the range of 4050 °C. Inlet gas: 600 ppm SO2, 1200 ppm NH3, 0-300 ppm NO, 10% O2, 12.5% CO2, 5% H2O, and N2 (balance).

removal through the NH3-SO2 reaction was studied in the range of 0-20% oxygen concentration in the gas mixture of 600 ppm SO2, 1200 ppm NH3, 225 ppm NO, 12.5% carbon dioxide, 10% water, and nitrogen (balance). Observed sulfur dioxide removal is shown in Figure 3 as a function of reaction temperature at various oxygen concentrations. The removal was observed below a certain temperature. For example, the removal was obtained at temperatures below approximately 52 °C at an oxygen concentration of 3%. The onset of removal was shifted to higher temperature with oxygen concentration. The reactions did not occur at temperatures above 56 °C for the oxygen concentrations investigated. For the interpretation of oxygen influence, the removals were replotted against oxygen concentration at temperatures of 40, 45, and 50 °C in Figure 4, where the values of the removal at 40 °C were estimated by extrapolation in Figure 3. Sulfur dioxide removal increased and leveled off up to 20% oxygen concentration. Even in the absence of oxygen, the reactions occurred at temperatures below 50 °C. Sulfur Dioxide Removal Dependence on Nitric Oxide. Experiments on the effect of nitric oxide in the range of 0-300 ppm were carried out in a gas mixture similar in composition to those of the mixture in the oxygen experiments. Sulfur dioxide removal is shown in Figure 5 as a function of reaction temperature for a variety of initial concentrations of nitric oxide. Nitric

oxide also moved the onset of removal to higher temperature. For example, the onset temperature in the absence of nitric oxide was shifted from 50 to 58 °C by the addition of 300 ppm nitric oxide. On the other hand, as much as 20% of oxygen was required to shift the temperature from 50 to 56 °C, as shown in Figure 3, indicating that sulfur dioxide removal was much more influenced by nitric oxide than by oxygen. Figure 6 shows sulfur dioxide removal dependence on nitric oxide, which was given by the same procedure as for the results of oxygen experiments. In contrast to the obtained curve shown in Figure 4, a continuous increase in sulfur dioxide removal was observed up to 300 ppm nitric oxide concentration. Reaction Products. An FTIR spectrum of the reaction products obtained under the condition of 5% oxygen and 225 ppm nitric oxide in the simulated gas is shown in Figure 7. Three strong absorption peaks (I-a-I-c) in the range of 600-1400 cm-1 correspond to ammonium sulfate ((NH4)2SO4). Ammonium sulfate is the main product of the NH3-SO2 reaction under this condition. This indicates that the oxidation of sulfur dioxide occurred during the reactions since sulfur has six valences in ammonium sulfate. Two weak absorption peaks (II-a and -b) are also observed in the range of 1200-1350 cm-1 on the spectrum. These correspond to standard sulfamic acid (HOSO2NH2) and ammonium sulfamate (NH4SO3NH2). Sulfamic acid exhibits par-

Ind. Eng. Chem. Res., Vol. 35, No. 10, 1996 3365

Figure 7. FT-IR spectrum of reaction products collected at reaction temperature of 45 °C. Inlet gas: 600 ppm SO2, 1200 ppm NH3, 225 ppm NO, 5% O2, 12.5% CO2, 5% H2O, and N2 (balance).

Figure 8. Relative abundance ratio of (NH4)2SO4 to HONH2SO2 from FT-IR spectra of reaction products formed at various oxygen concentrations. Inlet gas: 600 ppm SO2, 1200 ppm NH3, 225 ppm NO, 0-20% O2, 12.5% CO2, 5% H2O, and N2 (balance).

ticularly strong absorption peaks in the range of 12501350 cm-1. Two peaks (II-c and -d) around 800 and 600 cm-1 belong to the absorption of ammonium sulfamate. A weak peak (III) at 970 cm-1 is assigned to SO32- by comparison with the spectrum of standard sodium sulfite. In addition to the identification of reaction products, FTIR spectra also gave us quantitative information. The depth of an absorption peak is proportional to the amount of corresponding components. Accordingly, the depth ratio of I-b to II-b on the spectrum shown in Figure 7 approximately corresponds to the relative abundance ratio of ammonium sulfate to sulfamic acid in the products of the NH3-SO2 reaction. Figure 8 shows the relative abundance ratio of ammonium sulfate to sulfamic acid calculated from the peak depth of I-b and II-b on spectra obtained at various oxygen concentrations in the gas mixture. The ratio gradually increased from 2.6 in the absence of oxygen and reached about 8.0 for 15% oxygen concentration. The results indicate that the amount of ammonium sulfate increased at higher oxygen concentration. Reaction products were also analyzed by FTIR for a variety of initial concentrations of nitric oxide in the range of 0-300 ppm in the gas mixture containing 600 ppm SO2, 1200 ppm NH3, 10% oxygen, 12.5% carbon dioxide, 5% water, and nitrogen (balance). The resulting FTIR spectra showed that the absorption attributed to sulfamic acids faded away with an increase in nitric oxide concentration in the simulated flue gas. For example, the relative abundance ratio of ammonium sulfate to sulfamic acid was 3.2 in the absence of nitric oxide and 6.7 at 225 ppm of nitric oxide concentration. Nitrogen dioxide at a concentration of about 20 ppm is the minor component of nitrogen oxides in coal-fired

Figure 9. FT-IR spectrum of reaction products collected at reaction temperature of 45 °C in the presence of nitrogen dioxide. Inlet gas: 600 ppm SO2, 1200 ppm NH3, 215 ppm NO, 10 ppm NO2, 12.5% CO2, 5% H2O, and N2 (balance).

flue gas. Its concentration depends on combustion conditions. Experiments on the reactions of sulfur dioxide with ammonia in simulated flue gas containing nitrogen dioxide showed interesting results. Figure 9 shows an FTIR spectrum of the reaction products formed under the condition of 215 ppm NO and 10 ppm NO2 without oxygen. Double strong absorption peaks (II-a′ and -b′) around 1270 cm-1 and two absorption (IIc′ and -d′) around 800 and 600 cm-1 on the spectrum are attributable to sulfamic acid and ammonium sulfamate as described above. The relative abundance ratio (I-b′/II-b′) was found to be 1.5 from this spectrum. This value was lower than that of 2.6 obtained from Figure 8 under the condition of 225 ppm NO and 0% oxygen. This shows that nitrogen dioxide enhances the formation of sulfamic acid. However, the presence of oxygen depressed the formation: the depth ratio increased from 1.5 in the absence of oxygen to 8.2 in the presence of 10% oxygen. Aerosol Formation. Many studies on the aerosol formation from the NH3-SO2 reaction have been conducted in the absence of oxygen or nitric oxide (e.g. Christensen et al. (1992) and Bai et al. (1992)). The size and the number concentration of aerosols formed by the reactions were measured in the presence of oxygen and nitric oxide in our experiments. The results are shown in Figures 10a and b obtained in the gas mixture containing 600 ppm SO2, 1200 ppm NH3, 10% oxygen, 225 ppm NO, 12.5% carbon dioxide, 10% water, and nitrogen (balance) at reaction temperatures of 48.4 and 47.2 °C, respectively. The total number concentration of particles increased dramatically with only a 1.2 °C decrease in reaction temperature from 48.4 to 47.2 °C. The number concentration of medium (0.6-1.2 µm) and larger size (above 1.2 µm) particles increased substantially at lower temperature with relative lower number concentration of smaller size (below 0.6 µm). The particles became larger at lower reaction temperature. Growth of particles at lower temperature suggests that the particle surface is an another possible place of homogeneous NH3-SO2 reaction as well as a reactor surface. Actually, particles were not detected in sulfur dioxide removal lower than about 20% in all experiments. The diameter of particles formed initially by the NH3-SO2 reaction was probably less than 0.33 µm, which is the minimum size measurable for the detector. The particles became larger as sulfur dioxide removal increased at lower temperature. Figure 11 shows a semilog plot of total particle volume concentration versus sulfur dioxide removal. Assuming that all particles are spherical ammonium sulfate, a couple of

3366 Ind. Eng. Chem. Res., Vol. 35, No. 10, 1996 Table 1. Equilibrium Constants for NH3-SO2 Reactionsa reaction formulas

equilibrium expressions

(NH4)2SO3‚H2O ) 2NH3 + SO2 + 2H2O (NH4)2SO3 ) 2NH3 + SO2 + H2O NH4HSO3 ) NH3 + SO2 + H2O (NH4)2S2O5 ) 2NH3 + 2SO2 + H2O

log Ka ) 42.00 - 17411/T log Kb ) 33.27 - 14171/T log Kc ) 23.77 - 9958/T log Kd ) 41.89 - 17705/T

a

Figure 10. Particle size distribution of reaction products formed at reaction temperature of (a, top) 48.4 and (b, bottom) 47.2 °C. Inlet gas: 600 ppm SO2, 1200 ppm NH3, 225 ppm NO, 10% O2, 12.5% CO2, 10% H2O, and N2 (balance).

Hjuler and Dam-Johansen (1992).

absence of oxygen and nitric oxide for thermodynamic research (e.g. Landreth et al. (1975), Scargill (1971), and Scott and Lamb (1970)). The reactions of sulfur dioxide with ammonia produce those assuming compounds, listed in Table 1. Sulfur dioxide removal can be plotted against temperature using the equilibrium constants listed in Table 1. On the basis of the Bai et al. (1994) study, sulfur dioxide removal was calculated from the equilibrim constants for the four assuming products. Figure 12 shows the calculated removal of sulfur dioxide with our experimental results obtained in the absence of oxygen and nitric oxide. The calculations show the different profile of sulfur dioxide removal and the onset of removal. No agreement in the profile of sulfur dioxide removal was obtained among our results and the calculations. This indicates that sulfur dioxide reacted with ammonia via a combination of the reactions listed in Table 1 or other reaction paths. Our reaction products formed in the absence of oxygen and nitric oxide were analyzed with FTIR. It revealed that the reaction products contained ammonium sulfate with a relatively large amount of sulfamic acid and ammonium sulfamate (the relative abundance ratio ) 1.8). The formation of ammonium sulfate indicates that the oxidation of sulfur dioxide actually occurred under our experimental condition. It is considered that a trace amount of oxygen oxidized sulfur dioxide dissolved in a water layer on the reactor surface:

SO2 + H2O f HSO3- + H+

(1)

HSO3- f SO32- + H+

(2)

SO32- + 1/2O2 f SO42-

(3)

NH3 + H2O f NH4+ + OH-

(4)

2NH4+ + SO42- f (NH4)2SO4

(5)

Combining reactions 1-5, the overall reaction is

2NH3 + SO2 + 1/2O2 + H2O f (NH4)2SO4 Figure 11. Total particle volume concentration as a function of SO2 removal. Inlet gas: 600 ppm SO2, 1200 ppm NH3, 225 ppm NO, 10% O2, 12.5% CO2, 10% H2O, and N2 (balance).

percent of removed sulfur dioxide formed aerosols when the sulfur dioxide removal was 40%. Almost all removed sulfur dioxide may be deposited on the surface of the reactor. The fraction of removed sulfur dioxide forming aerosols increased with sulfur dioxide removal. For example, about 5% of removed sulfur dioxide formed aerosols when the sulfur dioxide removal was 50%. Discussion Oxygen Effect on Reactions. A number of previous studies of the NH3-SO2 reaction were done in the

(6)

while sulfamic and ammonium sulfamate are formed as follows (Hartley and Matteson, 1975; Kiang et al., 1973) H 2O

NH3 + SO2 98 NH3SO2 H 2O

(7)

NH3SO2 + NH3 98 (NH3)2SO2

(8)

NH3SO2 + 1/2O2 f HOSO2NH2

(9)

(NH3)2SO2 + 1/2O2 f NH4SO3NH2

(10)

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N-sulfonate (NHAS, -ON(NO)SO3-) (Nunes and Powell, 1970; Littlejohn et al., 1986):

2NO + SO32- f -ON(NO)SO3-

(15)

2NO + HSO3- f -ON(NO)SO3- + H+

(16)

The formed NHAS is readily hydrolyzed in acidic solution to yield sulfate ions: -

Figure 12. Comparison of SO2 removal efficiencies as a function of reaction temperature obtained from experiments and equilibrium constants listed in Table 1. Inlet gas: 600 ppm SO2, 1200 ppm NH3, 12.5% CO2, 5% H2O, and N2 (balance).

The overall reactions 7 and 9 are

NH3 + SO2 + 1/2O2 f HOSO2NH2

(11)

The overall reactions 7, 8, and 10 are

2NH3 + SO2 + 1/2O2 f NH4SO3NH2

(12)

Sulfamic acid (Chang and Lee, 1992) and ammonium sulfamate can be hydrolyzed to form sulfate and ammonium ions:

HOSO2NH2 + H2O f NH4+ + SO42- + H+

(13)

NH4SO3NH2 + H2O f 2NH4+ + SO42-

(14)

Schwartz and Freiberg (1981) have studied the oxidation reaction of sulfur dioxide in aqueous solution, where hydrogen sulfite and sulfite ions formed by the dissolution of gaseous sulfur dioxide were oxidized by dissolved oxygen to form hydrogen sulfate and sulfate ions. They proposed that there were some steps to complete the oxidation: transfer of sulfur dioxide and oxygen across the gas-liquid interface, establishment of their solubility equilibrium, establishment of equilibria between hydrogen sulfite and sulfite ions, and so on. From their discussion, a lot of sulfate ions are not formed under the trace amount of oxygen condition because Henry’s law coefficient of oxygen dissolving in liquid water is much lower than that of sulfur dioxide. Sulfur dioxide is oxidized by oxygen via reactions 6, 11, and 12 in the presence of oxygen. In fact, the relative abundance ratio of ammonium sulfate to sulfamic acid was found to increase from 1.8 to 3.2 by the addition of 10% oxygen from the FTIR spectra. From our results and discussion, it is found that oxygen enhanced the oxidation of sulfur dioxide to form ammonium sulfate. Reactions of 6, 11, and 12 simultaneously occurred in our experiments. Nitric Oxide Effect on Reactions. The FTIR analyses revealed that a higher content of ammonium sulfate was yielded in the presence of nitric oxide. Nitric oxide also had an effect on the oxidation of sulfur dioxide in the water layer on the reactor. Previous investigations have showed that the reactions of hydrogen sulfite and sulfite ions with nitric oxide in aqueous solution led to the formation of N-nitrosohydroxylamine-

ON(NO)SO3- f N2O + SO42-

(17)

In a solution of pH > 7, reaction 15 predominates in the NHAS formation reactions, and less hydrolysis of NHAS occurs (Ackermann and Powell, 1967). A water layer on the surface of the reactor in this study is probably alkaline, since Henry’s law coefficient and the gas phase concentration of ammonia are larger than those of sulfur dioxide. Hence, the alkaline water layer on the reactor surface leads to reaction 15 for NHAS formation with less NHAS hydrolysis in our experiments. However, NHAS is continuously hydrolyzed to form sulfate ions through reaction 17 in order to establish equilibrium between NHAS and sulfate because sulfate is consumed by ammonium ions. As a result, hydrogen sulfite and sulfite ions are oxidized to sulfate ions through reactions 15-17 in the presence of nitric oxide. Some NHAS may directly react with ammonium ions to form the compounds which have absorption bands in the range of 1200-1350 cm-1, as seen in the FTIR absorption spectrum of Figure 7. Reactions in the Presence of Oxygen and Nitric Oxide. The NHAS produced through reactions 15 and 16 is slowly oxidized to form sulfate and nitrite ions when exposed to gaseous oxygen in alkaline solution (Ackermann and Powell, 1967). Actually, ammonium sulfate was formed in a large amount with an increase in oxygen concentration shown in Figure 8. From the discussion of Ackermann and Powell, and our results, it is considered that the slow oxidation of NHAS occurred under our experimental conditions, i.e., in an alkaline water layer in contact with gaseous oxygen:

ON(NO)SO3- + O2 + 2OH- f

-

2NO2- + SO42- + H2O (18) Oxygen enhances the oxidation of NHAS to form sulfate ions. It is found that oxygen oxidized not only hydrogen sulfite and sulfite ions in the aqueous layer, but also NHAS formed through the reactions of these ions with nitric oxide. Oxygen and nitric oxide are very important contributors to the oxidation of sulfur dioxide in the aqueous layer. Hjuler and Dam-Johansen (1992) have observed the temperature at which the onset of removal was obtained. The temperature was shifted higher by the addition of 3000 ppm nitric oxide in the study of the reactions of sulfur dioxide with ammonia. This is consistent with our results. An increase in oxygen and nitric oxide shifted the onset temperature higher, as shown in Figures 3 and 5. The removal of sulfur dioxide was observed at higher temperature in the presence of abundant oxygen and nitric oxide. The removal of sulfur dioxide originates from its dissolution into the water layer formed on the reactor surface. The resulting S(IV) ions are oxidized to sulfate ions by oxygen and nitric oxide as described above. The removal is constant

3368 Ind. Eng. Chem. Res., Vol. 35, No. 10, 1996

under gas-liquid equilibrium of sulfur dioxide. However, consumption of sulfate ions through the reactions with ammonium ions breaks the equilibrium and causes more dissolution of sulfur dioxide. Consequently, the oxidation of sulfur dioxide in the aqueous layer by oxygen and nitric oxide shifted the onset temperature. Influence of Nitrogen Dioxide. Sulfamic acid was formed in the reaction products when 10 ppm nitrogen dioxide was added to the simulated flue gas in the absence of oxygen. Strong contribution of nitrogen dioxide was confirmed for the formation of sulfamic acid. Nitrogen dioxide was found to have something with the oxidation of sulfur dioxide in the aqueous layer. The reactions of nitrogen dioxide with hydrogen sulfite and sulfite ions produce sulfite radicals (Chang and Lee, 1992):

NO2 + SO32- f NO2- + SO3•-

(19)

NO2 + HSO3- f NO2- + HSO3•

(20)

Reactivity of sulfite radicals is considerably high, and the reactions of these radicals with ammonium ions will yield the compounds such as sulfamic acid. Another resulting NO2- in reactions 19 and 20 reacts with hydrogen sulfite ions formed by the dissolution of sulfur dioxide. It gives hydroxyamidosulfonate (HADS, HON(SO3-)2) (Littlejohn and Chang, 1991). Further sulfonation and/or hydrolysis of the compound may finally produce sulfamic acids and sulfate ions under acidic conditions (Chang and Lee, 1992). From the above discussion and our FTIR results, it is considered that reactions 19 and 20 may occur under our conditions and the formation of sulfamic acid is markedly depressed in the presence of oxygen. Conclusions The reactions of sulfur dioxide with ammonia were studied using a simulated flue gas for a variety of oxygen and nitric oxide concentrations at temperatures in the range of 40-60 °C. The reactions in the lack of oxygen and nitric oxide formed sulfamic acid and ammonium sulfamate. The reactions in the presence of the two species increased removal efficiency of sulfur dioxide with the formation of ammonium sulfate by the oxidation of sulfamic acid and ammonium sulfamate. The oxidation may occur in an aqueous layer on the reactor surface. Qualitative analyses of reaction products by FTIR showed that nitrogen dioxide enhanced the formation of sulfamic acid. Literature Cited Ackermann, M. N.; Powell, R. E. Air Oxidation of HydroxylamineN-sulfonate. Inorg. Chem. 1967, 6, 1718-1720. Bai, H.; Biswas, P.; Keener, T. C. Particle Formation by NH3SO2 Reactions at Trace Water Conditions. Ind. Eng. Chem. Res. 1992, 31, 88-94. Bai, H.; Biswas, P.; Keener, T. C. SO2 Removal by NH3 Gas Injection: Effects of Temperature and Moisture Content. Ind. Eng. Chem. Res. 1994, 33, 1231-1236. Chang, S. G.; Lee, G. C. LBL PhoSNOX Process for Combined Removal of SO2 and NOx from Flue Gas. Environ. Prog. 1992, 11, 66-73.

Christensen, P. S.; Madsen, N. M.; Livbjerg, H. The Formation of Aerosols from SO2 and NH3 in Humid Air. J. Aerosol Sci. 1992, 23, S261-S264. Hartley, E. M.; Matteson, M. J. Sulfur Dioxide Reactions with Ammonia in Humid Air. Ind. Eng. Chem. Fundam. 1975, 14, 67-72. Hirota, K.; Niina, T.; Anwar, E.; Namba, H.; Tokunaga, O.; Tabata, Y. Reactions of Sulfur Dioxide with Ammonia. J. Environ. Sci. 1993, 6, 143-150. Hirota, K.; Tokunaga, O.; Miyata, T.; Sato, S.; Osada, Y.; Sudo, M.; Doi, T.; Shibuya, E.; Baba, S.; Hatomi, T.; Komiya, M.; Miyajima, K. Pilot-scale Test for Electron Beam Purification of Flue Gas from a Municipal Waste Incinerator with Slaked-Lime. Radiat. Phys. Chem. 1995, 46, 1089-1092. Hjuler, K.; Dam-Johansen, K. Mechanism and Kinetics of the Reaction between Sulfur Dioxide and Ammonia in Flue Gas. Ind. Eng. Chem. Res. 1992, 31, 2110-2118. Kawamura, K.; Aoki, S.; Kimura, H.; Adachi, K.; Kawamura, K.; Katayama, T.; Kengaku, K.; Sawada, Y. Pilot Plant Experiment on the Treatment of Exhaust Gas from a Sintering Machine by Electron Beam Irradiation. Environ. Sci. Technol. 1980, 14, 288-293. Kiang, C. S.; Stauffer, D.; Mohnen, V. A. Possibilities for Atmospheric Aerosol Formation involving NH3. Nature Phys. Sci. 1973, 244, 53-54. Landreth, R.; de Pena, R. G.; Heicklen, J. Thermodynamics of the Reaction of Ammonia and Sulfur Dioxide in the Presence of Water Vapor. J. Phys. Chem. 1975, 79, 1785-1788. Littlejohn, D.; Chang, S. G. Modeling of the Chemistry of Wet Limestone Flue Gas Desulfurization Systems. Energy Fuels 1991, 5, 249-254. Littlejohn, D.; Hu, K. Y.; Chang, S. G. Kinetics of the Reaction of Nitric Oxide with Sulfite and Bisulfite Ions in Aqueous Solution. Inorg. Chem. 1986, 25, 3131-3135. Namba, H.; Tokunaga, O.; Hashimoto, S.; Tanaka, T.; Ogura, Y.; Doi, Y.; Aoki, S.; Izutsu, M. Pilot-scale Test for Electron Beam Purification of Flue Gas from Coal-combustion Boiler. Radiat. Phys. Chem. 1995, 46, 1103-1106. Nunes, T. L.; Powell, R. E. Kinetics of the Reaction of Nitric Oxide with Sulfite. Inorg. Chem. 1970, 9, 1916-1917. Osada, Y.; Hirota, K.; Sudo, M.; Baba, S.; Shibuya, E.; Doi, T.; Nakajima, M.; Komiya, M.; Miyajima, K.; Miyata, T.; Tokunaga, O. Pilot-scale Test on Electron Beam Treatment of Municipal Solid Waste Flue Gas with Spraying Slaked-Lime Slurry. Radiat. Phys. Chem. 1995, 45, 1021-1027. Scargill, D. Dissociation Constants of Anhydrous Ammonium Sulphite and Ammonium Pyrosulphite Prepared by Gas-phase Reactions. J. Chem. Soc. A 1971, 2461-2466. Schwartz, S. E.; Freiberg, J. E. Mass-Transport Limitation to the Rate of Reaction of Gases in Liquid Droplets: Application to Oxidation of SO2 in Aqueous Solutions. Atmos. Environ. 1981, 15, 1129-1144. Scott, W. D.; Lamb, D. The Solid Compounds Which Decompose into a Common Vapor. Anhydrous Reactions of Ammonia and Sulfur Dioxide. J. Am. Chem. Soc. 1970, 92, 3943-3946. Tokunaga, O.; Namba, H.; Hirota, K. Experiments on Chemical Reactions in Electron-Beam-Induced NOx/SO2 Removal. In Non-Thermal Plasma Techniques for Pollution Control; NATO ASI Series Vol. 34, Part B; Penetrante, B. M., Schultheis, S. E., Eds.; Springer-Verlag: Berlin, Heidelberg, 1993; pp 55-62.

Received for review September 6, 1995 Accepted May 24, 1996X IE950560K

X Abstract published in Advance ACS Abstracts, July 15, 1996.