Influence of Nitrogen Monoxide on the Complex Phase and Chemical

Faculty of Engineering, Department of Thermodynamics, Gerhard-Mercator-University Duisburg, Lotharstrasse 1, D-47057 Duisburg, Germany. Ind. Eng. Chem...
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Ind. Eng. Chem. Res. 2003, 42, 1406-1413

Influence of Nitrogen Monoxide on the Complex Phase and Chemical Equilibria in Wet Flue Gas Cleaning Processes M. A. Siddiqi,* J. Petersen, and K. Lucas† Faculty of Engineering, Department of Thermodynamics, Gerhard-Mercator-University Duisburg, Lotharstrasse 1, D-47057 Duisburg, Germany

The present paper reports the absorption behavior of sulfur dioxide in the presence of nitrogen oxides (NO and NO2). The absorption process has been studied at 298.65 and 318.45 K and ambient pressure. The results show that after a few hours NO2 vanishes from the gas phase. This leads to a reproducible frozen (stable) state in which tetravalent sulfur exists in phase and chemical equilibrium. The gas phase contains SO2, N2O, and NO. The liquid phase contains a number of nitrogen-sulfur compounds besides tetravalent and hexavalent sulfur and nitrate. The absorption of NO2 is accelerated by the presence of NO, and the solubility of SO2 increases with the increasing amount of NO. The results of model calculations that give the concentration profile of the species produced in this system as a function of time are presented. The effects of temperature and the initial concentration of the gas mixture on the absorption behavior of SO2 are discussed. Introduction The studies on the desulfurization process have shown that the chemistry of the wet scrubbing is strongly influenced by the concentration of nitrogen oxides present. In a recent paper1 from this laboratory, the effect of nitrogen dioxide (NO2) on the absorption of sulfur dioxide (SO2) in wet flue gas cleaning processes was reported. The study has been extended by substituting a part of nitrogen dioxide with nitrogen monoxide (NO), which is believed to be less reactive but forms a large proportion of the total nitrogen content in the real flue gases2 and is thus of great technical importance. The simultaneous absorption of NO, NO2, and SO2 in water is studied. The oxides of S(IV) and N(II,III,IV) undergo a series of concurrent and consecutive chemical reactions in aqueous solution. This leads to the formation of a series of mixed nitrogen-sulfur compounds, which are believed to inhibit the oxidation of sulfite in aqueous solutions.3 The possible compounds which may be formed during the process are analyzed and quantified. The initial concentrations of SO2 in the gas mixture are pertinent to its concentration in flue gases and varied from 2000 mg mN-3 ()volume fraction of 700 ppm) to 10 000 mg mN-3 ()volume fraction of 3500 ppm), the concentration of NO2 varied from 100 mg mN-3 ()volume fraction of 49 ppm) to 3000 mg mN-3 ()volume fraction of 1462 ppm), and the concentration of NO varied from 500 to 2000 mg mN-3 ()volume fraction of 373.5-1494 ppm). The absorption process has been studied at 298.65 and 318.45 K. The gas and liquid phases could be specified completely. The scheme for the reactions that can take place in an aqueous solution proposed earlier1 has been extended to include the influence of NO on the simultaneous absorption of SO2 and NO2 in aqueous solutions. The results of model calculations that give the concentration profile of species * To whom correspondence should be addressed. E-mail: [email protected]. Telephone: +203 379 3353. Fax: +203 379 1594. † Chair of Technical Thermodynamics, RWTH Aachen, Schinkelstrasse 8, D-52062 Aachen, Germany.

produced in this system as a function of time are presented. The effects of the temperature and initial concentration of the gas mixture are discussed. Experimental Section The measurements have been carried out using the apparatus described in previous papers.4,5 The whole apparatus was placed in an air thermostat. The temperature within the thermostat was checked by measuring the temperature at different reference points within the thermostat with the help of precalibrated thermocouples and did not vary more than (0.5 K. An ancillary part of the apparatus, not shown in the diagram,4 consisted of a FTIR spectrometer from Bio-Rad Laboratories (Richmond, CA) equipped with a long-path gas cell. This assembly was used to measure the concentration of nitrous oxide (N2O) in the gas phase. Materials. The certified standard mixtures (sulfur dioxide + nitrogen, nitrogen dioxide + nitrogen, and nitrogen monoxide + nitrogen) were supplied by MesserGriesheim (Krefeld, Germany). The standard mixtures having the volume fraction of 10 000 ppm SO2, 1000 mg mN-3 NO2, 10 000 mg mN-3 NO2, and 5000 mg mN-3 NO were used to prepare the desired gas mixtures. The sulfur dioxide in the standard mixtures had a mole purity x(SO2) g 0.9998 [impurities: x(CO2) e 30 × 10-6; x(H2O) e 50 × 10-6]. The nitrogen dioxide standard mixtures had a mole purity x(NO2) g 0.98 [impurities: x(CO2) e 30 × 10-6; x(H2O) e 50 × 10-6]. The nitrogen monoxide standard mixtures had a mole purity x(NO) g 0.98 [impurities: x(CO2) e 30 × 10-6; x(H2O) e 50 × 10-6]. Pure nitrogen and the nitrogen in the standard mixtures had a purity x(N2) g 0.999 99 [impurities: x(O2) e 0.5 × 10-6; x(H2O) e 0.5 × 10-6]. Purified and deionized water that had a conductivity 97%). The other nitrogen-sulfur compounds, viz., hydroxylamine-

10.1021/ie020739c CCC: $25.00 © 2003 American Chemical Society Published on Web 02/22/2003

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trisulfonic acid (HATS), hydroxylaminedisulfonic acid (HADS), hydroxylaminemonosulfonic acid (HAMS), aminetrisulfonic acid (nitrilotrisulfonic acid; NTS), aminedisulfonic acid (imidodisulfonic acid; IDS), and hydroxylamine-N,O-disulfonic acid (HAODS), used for calibration purposes in ion chromatography were not available commercially and were synthesized.1 All of the other chemicals used for the ion chromatography measurements were “pro-anlysis”. All of the chemicals were dried before use. Procedure. The experimental procedure was already described in previous papers,4,5 and so only a brief description of the modifications will be given here. The basic procedure for filling the gas mixtures is the same except that first a SO2 standard gas mixture was filled to a precalculated pressure in the gas vessel, then the NO2 standard gas mixture was pressurized up to its precalculated pressure, then the NO standard gas mixture was pressurized up to its precalculated pressure, and last the mixture was diluted with pure nitrogen up to the experimental pressure to obtain the mixture of desired composition. The initial pressure in the system depended (sulfur dioxide + nitrogen dioxide + nitrogen monoxide + nitrogen) on the desired experimental temperature. It was chosen to furnish a total pressure of nearly 104 kPa after making allowances for the saturated vapor pressure of the liquid phase. It was kept at 101 or 95 kPa, respectively, for studies at 298.65 or 318.45 K. The volume of the gas phase was 22.25 dm3. After attainment of the experimental pressure and temperature, the gas mixture was brought into contact with water (about 0.4 kg; accurately weighed) in the reaction cell by opening the appropriate valves and starting a circulatory pump. The spectra of the gaseous phase and of the liquid were taken by scanning through the wavelength range of 220-450 nm with the help of the respective spectrophotometers at different time intervals. It was found that after about 2.5 h all of the nitrogen dioxide disappeared from the gas phase (no absorption band in the 390-410 nm region). This stable state was considered as the state of quasi-equilibrium. The gas- and liquid-phase spectra were taken in the UV-vis region and evaluated to determine the composition of the gas phase. Only sulfur dioxide could be detected in the gas phase. The liquid-phase spectra were evaluated by taking the molar absorption coefficient for SO2(aq) from the literature4 at 260 and 276 nm. In this way the concentration of molecular dissolved sulfur dioxide SO2(aq) could be determined. The samples of the liquid phase were taken from the bottom of the reaction cell and analyzed by ion chromatography for NO3-, NO2-, SO32-, and SO42- as well as for HAMS, HADS, HAODS, HAOMS, HATS, IDS, AS, and NTS separately. For the ion-chromatographic determination, a Dionex DX-100 ion chromatograph equipped with a membrane suppressor and a conductivity detector was used. The details for the ion-chromatographic measurements are given in an earlier paper.1 At the end of each experiment, the gas phase was sent through the long-path gas cell in the FTIR spectrometer and the IR spectra were taken in the region 2180-2260 cm-1 for the analysis of nitrous oxide (N2O). The spectra were evaluated in the same manner as UV spectra using the multivariate analysis method described in a previous publication.4 Before the spectroscopic measurements were performed, a calibration was made by taking

Figure 1. Measured and calculated SO2 concentrations in the gas phase as a function of the initial NO concentration for some selected SO2 initial concentrations at 298.65 K (initial NO2 concentration ) 1000 mg mN-3).

spectra of (nitrous oxide + nitrogen) gas mixtures of known compositions. The initial concentrations for SO2 varied from 2000 to 10 000 mg mN-3, those for NO2 from 500 to 3000 mg mN-3, and those for NO from 500 to 2000 mg mN-3. The studies were performed in an acidic solution, and the pH was varied between 2.1 and 2.6. Results Measurements for the SO2/NO2/NO/N2/H2O system at 298.65 and 318.45 K show that after 2.5 h all NO2 vanishes from the gas phase; i.e., the mass transfer of NO2 from the gas phase to the liquid phase is completed. This time for the gas mixtures in the absence of NO was 4-6 h, and so it is concluded that the absorption of NO2 is accelerated by the addition of NO. The specification of the system has been undertaken after attaining this reproducible state when the mass transfer of NO2 from the gas phase to the liquid phase is completed. A comprehensive analysis of the liquid phase and the gas phase at 298.65 K showed the presence of the same components as those for the SO2/NO2/N2/H2O system; i.e., the gas phase consists of SO2, N2O, and N2 besides a small amount of NO, and the liquid phase consists of molecular dissolved SO2, NO3-, HADS, and HAMS as well as tetravalent sulfur [S(IV)] {HSO3-} and hexavalent sulfur [S(VI)] {HSO4- + SO42-}. It may be remarked at this point that a small proportion (2-5%) of sulfur measured as S(VI) in solution results from the oxidation of HSO3- through atmospheric oxygen during sampling and analysis with ion chromatograph. At 318.45 K, another nitrogen-sulfur compound, HAODS, was found in the liquid phase and no HADS could be detected. This means that at this temperature and in the presence of NO the hydrolysis of HADS to HAMS is completed within 2.5 h. The results show that the addition of NO to the system leads to an increase in the solubility of SO2; i.e., with increasing amount of NO in the initial gas mixture, the final concentration of SO2 in the gas phase decreases. For the sake of brevity, only some typical results of the analysis of the gas phase and of the liquid phase after 2.5 h are presented in Tables 1 and 2. Some

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Table 1. Experimental Results for the Absorption of SO2 + NO2 + NO in Water at 298.65 K after a Mixing Period of 2.5 h initial cSO2/cNO2/cNO(g) [mg mN-3]

cSO2(g) [mg m-3]

cNO(g) [mg m-3]

8000/500/0 8000/500/500 8000/500/1000 8000/500/2000 8000/2000/0 8000/2000/500 8000/2000/1000 8000/2000/2000 10000/500/0 10000/500/500 10000/500/1000 10000/500/2000 10000/2000/0 10000/2000/500 10000/2000/1000 10000/2000/2000 10000/3000/0 10000/3000/500 10000/3000/1000

1865 1647 1621 1577 1403 1106 908 689 2514 2356 2284 2200 1997 1646 1565 1404 1672 1422 1021

0 334 584 1177 0 233 423 1099 0 320 554 1068 0 278 412 1092 0 226 414

cN2O(g) [mg m-3]

40 49 112 141 205 31 44 123 161 211 141 214 366

mNO3[mol kg-1]

mHADS [mol kg-1]

mHAMS [mol kg-1]

pH

6.8 × 10-5 2.3 × 10-5 1.9 × 10-5 0 3.79 × 10-4 1.82 × 10-4 1.60 × 10-4 1.10 × 10-4 7.7 × 10-5 2.1 × 10-5 2.1 × 10-5 1.9 × 10-5 3.64 × 10-4 1.73 × 10-4 1.55 × 10-4 8.4 × 10-5 5.53 × 10-4 3.50 × 10-4 3.24 × 10-4

7.60 × 10-5 3.46 × 10-4 3.72 × 10-4 4.03 × 10-4 2.06 × 10-4 6.40 × 10-4 9.72 × 10-4 8.22 × 10-4 6.90 × 10-5 2.73 × 10-4 3.34 × 10-4 4.38 × 10-4 1.38 × 10-4 7.04 × 10-4 7.40 × 10-4 7.75 × 10-4 1.77 × 10-4 7.42 × 10-4 8.17 × 10-4

2.17 × 10-4 2.61 × 10-4 3.37 × 10-4 3.60 × 10-4 8.62 × 10-4 8.19 × 10-4 7.49 × 10-4 1.08 × 10-3 2.72 × 10-4 2.47 × 10-4 3.73 × 10-4 3.93 × 10-4 8.93 × 10-4 8.72 × 10-4 1.03 × 10-3 1.17 × 10-3 1.46 × 10-3 1.11 × 10-3 1.36 × 10-3

2.43 2.38 2.37 2.36 2.31 2.27 2.24 2.21 2.36 2.32 2.31 2.30 2.27 2.23 2.21 2.17 2.20 2.17 2.13

Table 2. Experimental Results for the Absorption of SO2 + NO2 + NO in Water at 318.45 K after a Mixing Period of 2.5 h initial cSO2/cNO2/cNO(g) [mg mN-3]

cSO2(g) [mg m-3]

cNO(g) [mg m-3]

8000/500/0 8000/500/500 8000/500/1000 8000/500/2000 8000/2000/0 8000/2000/500 8000/2000/1000 8000/2000/2000 8000/2000/2605 10000/1000/0 10000/1000/500 10000/1000/1000 10000/1000/2000 10000/1000/2754

2445 2439 2356 2304 1916 1592 1448 1192 1170 3168 2922 2885 2712 2361

0 264 574 1091 0 246 571 1157 1608 0 266 520 1112 1474

cN2O(g) [mg m-3] 50 70 311

mNO3[mol kg-1]

mHAMS [mol kg-1]

pH

7.3 × 10-5 3.4 × 10-5 1.8 × 10-5 1.3 × 10-5 2.9 × 10-4 2.3 × 10-4 9.2 × 10-5 8.6 × 10-5 6.9 × 10-5 1.4 × 10-4 6.8 × 10-5 4.2 × 10-5 2.3 × 10-5 2.4 × 10-5

2.29 × 10-4 3.18 × 10-4 3.79 × 10-4 4.94 × 10-4 6.22 × 10-4 7.68 × 10-4 8.51 × 10-4 9.14 × 10-4 9.14 × 10-4 4.12 × 10-4 6.13 × 10-4 6.96 × 10-4 9.05 × 10-4 9.63 × 10-4

2.58 2.53 2.52 2.49 2.43 2.31 2.36 2.31 2.31 2.48 2.41 2.39 2.41 2.39

phase with increasing NO concentration is almost the same for all initial SO2 concentrations. Figure 2 shows that at a constant initial SO2 concentration an increase in the initial NO2 concentration results in an increase in SO2 absorption. This observation is supported by the reaction of NO with NO2 to form HNO2(aq) or NO2according to the reaction:6,7

NO(aq) + NO2(aq) + H2O(l) f 2HNO2(aq)

Figure 2. Measured and calculated SO2 concentrations in the gas phase as a function of the initial NO concentration for some selected NO2 initial concentrations at 298.65 K (initial SO2 concentration ) 8000 mg mN-3).

more results are displayed in the figures which follow. Figures 1 and 2 show the equilibrium concentration of sulfur dioxide in the gas phase for a series of experiments at different SO2, NO2, and NO initial concentrations. The measured values are shown as points. It may be seen from Figure 1 that at a constant initial NO2 concentration the extent of decrease in SO2 in the gas

(1)

As HNO2(aq) reacts with HSO3- to produce nitrogensulfur compounds, more and more SO2 goes into solution [see reactions (6)-(12)]. The variation of the total molality of nitrogen-sulfur compounds (mHADS + mHAMS) is shown in Figure 3. It is seen that the molality increases with an increase in the NO and NO2 concentration. The formation of a nitrogen-sulfur compound is accompanied by an increase in the formation of N2O by the side reaction (12). The observed decrease in the molality of NO3- with increasing initial NO concentration indicates that reaction (1) is very fast and pushes the concurrent reaction (2) for NO2 species in the background. The observed faster mass transfer of NO2 from the gas phase to the liquid phase in the presence of NO supports this view. Model Calculations The physicochemical behavior of aqueous solutions of sulfur dioxide and nitrogen oxides is governed by the

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NOSO3(SO3)23- + H2O(l) f HATS SO3NHOSO32- + HSO4 (11) HAODS HONHSO3- + HNO2(aq) f HAMS N2O(aq) + HSO4- + H2O(l) (12)

Figure 3. Measured and calculated total molalities of HADS + HAMS as a function of the initial NO concentration for some selected initial NO2 concentrations at 298.65 K (initial SO2 concentration ) 8000 mg mN-3).

kinetics of a number of concurrent and consecutive reactions taking place in the liquid state. The system can be modeled on the basis of the reaction scheme proposed in the previous paper.1 For the sake of ready reference, the reactions are summarized here:

NO(aq) + NO2(aq) + H2O(l) f 2HNO2(aq) 2NO2(aq) + H2O(l) f HNO2(aq) + HNO3(aq)

(2)

2HNO2(aq) + HSO4- (3) 2NO2(aq) + 2HSO3- f 1 1 HON(SO3)22- + /2N2O(aq) + /2H2O(l) + O2(aq) (4) HADS

3HNO2(aq) f HNO3(aq) + 2NO(aq) + H2O(l)

(5)

HNO2(aq) + HSO3- f ONSO3- + H2O(l) NSS

(6)

1 1 ONSO3- + /2H2O(l) f /2N2O(aq) + HSO4 (7) NSS

(8)

-

HON(SO3)22- + H2O(l) f HONHSO3- + HSO4 HADS HAMS

(13)

SO2(aq) + H2O(l) T HSO3- + H+

(14)

HSO4- T H+ + SO42-

(15)

HNO2(aq) T H+ + NO2-

(16)

HNO3(aq) T H+ + NO3-

(17)

NO(g) T NO(aq)

(18)

NO2(g) T NO2(aq)

(19)

N2O(g) T N2O(aq)

(20)

(1)

2NO2(aq) + HSO3- + H2O(l) f

ONSO3- + HSO3 f HON(SO3)22NSS HADS

SO2(g) T SO2(aq)

(9)

HON(SO3)22- + 2HNO2(aq) + HSO3 f HADS NOSO3(SO3)23- + 2NO(aq) + 2H2O(l) (10) HATS

Reactions (1)-(12) run completely in the direction of the products. For reactions (1), (3), and (5), it is indicated by the thermodynamic data.8,9 For the other reactions which involve nitrogen-sulfur compounds, it may be inferred from the available literature.3,10 The significance of each individual reaction depends on its kinetics and thus on its activation energy. Equations (13)-(20) describe the equilibrium reactions which occur spontaneously and are to be considered at all times. Reactions (1)-(4) describe the dissolution of NO and NO2 and are responsible for an increase in SO2 absorption with an increase in the initial NO2 concentration. Reactions (1)-(3) are frequently mentioned in the literature.10-15 Reaction (4) was postulated for the formation of HADS directly from NO2(aq) and HSO3- on the basis of experimental evidences.1 Reaction (5) gives the formation of nitric acid and NO from nitrous acid in solution. Reactions (6) and (7) show the formation and the hydrolysis of NSS to produce N2O. The reaction products N2, N2O, NO3-/HNO3-, HAODS, HAMS, and SO42-/HSO4- are regarded as end products because they do not react further in the considered time period.16-18 The eventually possible oxidation reactions of NO, NO2, HNO2, and HSO3- with O2 produced in reaction (4) have been neglected because their concentrations are very low. The amount of nitrogen which could not be accounted for as NO3- or nitrogen-sulfur compounds was taken as N2O and NO for the sake of simplicity and to keep the number of reactions to a minimum. The gas-phase reactions

NO2(g) + SO2(g) f SO3(g) + NO(g)

(21)

NO2(g) + NO(g) + H2O(g) f 2HNO2(g)

(22)

are principally possible19 but will not influence the results because of their low concentrations. They are, therefore, neglected.

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The temporal change of the partial pressure of NO2 in the gas phase may be written as

dpNO2

)

dt

dpNO2(1) dt

+

dpNO2(2)

dpNO2(3)

+

dt

dt

+

dpNO2(4) dt

(23)

with

dpNO2(1) dt dpNO2(2) dt dpNO2(3) dt

( )

KNO2 m ˆ Liq ) -RT axk1mNODNO2 pNO2 V - VLiq p0 (24)

m ˆ Liq ) -RT a V - VLiq

( )

KNO2 4 k2DNO2 3 p0

x

x43k m

m ˆ Liq a ) -RT V - VLiq

3

dt

(37)

r10 ) k10mHADSmN(III)

(38)

r11 ) k11mHATSmH+

(39)

r12 ) k12mHAMSmN(III)

(40)

and the equilibrium constants are

K13 ) KSO2 )

1.5

pNO21.5 (25)

K14 )

( ) KNO2

K15 )

1.5

pNO21.5 (26) K16 )

)

m ˆ Liq a -RT V - VLiq

( )

KNO2 4 k4mHSO3-0.5DNO2 3 p0

x

pNO21.5 (27)

r1 ) k1mNO2mNO

(28)

r2 ) k2mNO22

(29)

r3 ) k3mNO22mHSO3-

(30)

r4 ) k4mNO22mHSO3-0.5

(31)

mHNO24

(32)

mNO2

r6 ) k6mH+0.5mN(III)mS(IV) r6 ) k6mH+mN(III)mS(IV)

m*pSO2

mSO2m* / / γSO 2- γH+mSO 2-mH+ 4 4 / γHSO -mHSO -m* 4 4 / / γNO - γH+mNO -mH+ 2 2

mHNO2m*

(41)

(42)

(43)

(44)

1.5

where m ˆ Liq is the mass of liquid water, a is the masstransfer surface area, V is the total volume of the gas mixing vessel, and VLiq is the volume occupied by the aqueous solution. Equation (23) is formulated for the transfer of NO2 into the liquid phase using the film model for the material transport for a fast reaction running completely in the direction of products under the consideration of reactions (1)-(4). The equations for the rate of reactions are then

r5 ) k5

mSO2p0

/ / γHSO - γH+mHSO -mH+ 3 3

HSO3-DNO2

p0

dpNO2(4)

r9 ) k9mH+mHADS

(at 298.65 K) (33) (at 318.45 K)

(34)

r7 ) k7mH+mNSS

(35)

r8 ) k8mNSSmHSO3-

(36)

K17 )

/ / γNO - γH+mNO -mH+ 3 3

mHNO3m*

mNOp0 K18 ) KNO ) m*pNO

K19 ) KNO2 )

K20 ) KN2O )

mNO2p0 m*pNO2 mN2Op0 m*pN2O

(45)

(46)

(47)

(48)

The above equations for the reaction rate and the equilibrium reactions were written in terms of the degree of reaction. The molalities and the partial pressures of the concerned components were then calculated with the help of the degree of reaction for all of the reactions in which the concerned component participates (for details, see the previous paper1). The equilibrium constants were calculated from the thermodynamic data given in the literature as cited before.1 A modified Pitzer equation20,21 with only long-range interactions was used to calculate the activity coefficient of any ionic species. For the modeling of the SO2/NO2/NO/N2/H2O system at 298.65 K, only one additional reaction, reaction (1), is considered as compared to those taken for the SO2/ NO2/N2/H2O system. Reactions (10)-(12) were not considered at 298.65 K because the experimental results showed that these were relevant only at higher temperatures. The values of the rate constant k1 reviewed in the literature10 vary from 1 × 107 to 5 × 109 kg mol-1 s-1. Its value was, therefore, fitted to obtain the best agreement with the observed experimental values of

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Figure 4. Time dependence of the concentration of SO2 and NO2 in the gas phase for a typical SO2/NO2/NO gas mixture (cSO2,ini ) 4000 mg mN-3, cNO2,ini ) 500 mg mN-3, and cNO,ini ) 2000 mg mN-3) at 298.65 K.

cSO2, mNO3-, and mHADS + mHAMS after a period of 2.5 h. The complete set of rate constants at 298.65 K is thus

Figure 5. Time dependence of the concentration NO2 in the gas phase for a typical SO2/NO2/NO gas mixture (cSO2,ini ) 8000 mg mN-3 and cNO2,ini ) 2000 mg mN-3) for various NO concentrations at 298.65 K.

k4 ) 2.0 × 109 kg1.5 mol-1.5 s-1

on the absorption of NO2. It is noted that the model also predicts correctly the observed increase in the absorption of NO2 in aqueous solutions in the presence of NO. The modeling at 318.45 K is done by considering reactions (10)-(12) in addition to those considered for 298.65 K. The calculations for the system SO2/NO2/NO/ N2/H2O were performed by using the same constants as those given in the previous paper1 for the SO2/NO2/ N2/H2O system, namely,

k5 ) 9.4 × 10-7 kg mol-1 s-1

k1 ) 2.2 × 109 kg mol-1 s-1

k6 ) 142 kg1.5 mol-1.5 s-1

k2 ) 3.33 × 108 kg mol-1 s-1

k1 ) 5.4 × 108 kg mol-1 s-1 k2 ) 7.0 × 107 kg mol-1 s-1 k3 ) 8.5 × 109 kg2 mol-2 s-1

k8/k7 ) 7 × 109/1 × 109 ) 7 k9 ) 1.75 × 10-2 kg mol-1 s-1 k10 ) k11 ) k12 ) 0 Some typical model calculations are shown below. The calculated values for the concentration of SO2 in the gas phase as a function of the initial NO concentration are shown as drawn lines in Figure 1 for an initial NO2 concentration of 1000 mg mN-3 for various SO2 concentrations. Similar results for an initial SO2 concentration of 8000 mg mN-3 for various NO2 concentrations are shown in Figure 2. Figure 3 shows the measured and calculated molalities of HADS + HAMS in the liquid phase for various initial concentrations of NO2 and NO. The average mean deviation of about 5% for cSO2 and mHADS+HAMS may be regarded as quite satisfactory for such a complex system. The molalities of nitrate show large deviations (up to 50%), which is acceptable considering their small values (on the order of 10-5 mol kg-1). To test if the model can also predict the time dependence of the system, the concentrations of NO2(g) and SO2(g) were calculated at definite time intervals and compared with the measured values. Figure 4 shows the measured and calculated time profiles of the concentrations of NO2(g) and SO2(g) for typical initial concentrations of SO2, NO, and NO2. It is seen that the model calculations are in reasonable agreement with the experimental values. Figure 5 shows the effect of NO

k3 ) 1.5 × 1011 kg2 mol-2 s-1 k4 ) 1.0 × 1010 kg1.5 mol-1.5 s-1 k5 ) 9.4 × 10-7 kg mol-1 s-1 k6 ) 17431 kg1.5 mol-1.5 s-1 k8/k7 ) 7 × 109/1 × 109 ) 7 k9 ) 0.1775 kg mol-1 s-1 k10 ) 15 kg mol-1 s-1 k11 ) 0.289 kg mol-1 s-1 k12 ) 350 kg mol-1 s-1 A comparison of the calculated values with the experimental values at 318.45 K for some typical mixtures is shown as full lines in Figure 6. Figure 7 shows the measured and calculated time profiles of the concentrations of NO2(g) for typical initial concentrations of SO2, NO2, and NO. The results may be seen as satisfactory

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Ind. Eng. Chem. Res., Vol. 42, No. 7, 2003 Greek Letters γ ) activity coefficient λ ) wavelength [m] Subscripts HADS ) hydroxylaminedisulfonic acid HAMS ) hydroxylaminemonosulfonic acid ini ) initial i, j ) species i, j or reaction i, j Liq ) bulk liquid N ) normal state (TN ) 273.15 K; pN ) 101.3 kPa) Superscripts * ) standard state of the hypothetical ideal aqueous solution of unit molality 0 ) standard state at standard pressure (p0 ) 100 kPa)

Literature Cited Figure 6. Measured and calculated SO2 concentrations in the gas phase as a function of the initial NO concentration for some selected SO2 initial concentrations at 318.45 K (initial NO2 concentration ) 1000 mg mN-3).

Figure 7. Time dependence of the concentration NO2 in the gas phase for a typical SO2/NO2/NO gas mixture (cSO2,ini ) 6000 mg mN-3 and cNO2,ini ) 1000 mg mN-3) for various NO concentrations at 318.45 K.

considering the approximations made in obtaining basic thermodynamic and kinetic data at 318.45 K. Acknowledgment This work was supported by a grant from the Deutsche Forschungsgemeinschaft (DFG), which is gratefully acknowledged. Nomenclature a ) specific mass-transfer surface area [m2 m-3] aq ) aqueous c ) molarity [mol L-1] c ) concentration in the gas phase [mg m-3] D ) diffusion coefficient [m2 s-1] g ) gas K ) K value, equilibrium constant k ) rate constant [mol, kg, s] l ) liquid m ) molality [mol kg-1] m ˆ ) mass [kg] M ) molecular weight [g mol-1] p ) pressure, partial pressure [Pa] r ) rate of the reaction [mol kg-1 s-1] R ) molar gas constant () 8.3145) [J mol-1 K-1] t ) time [s] T ) temperature [K] x ) mole fraction

(1) Siddiqi, M. A.; Petersen, J.; Lucas, K. A Study of the Effect of Nitrogen Dioxide on the Absorption of Sulfur Dioxide in Wet Flue Gas Cleaning Processes. Ind. Eng. Chem. Res. 2001, 40, 2116. (2) Gutberlet, H.; Dieckmann, H.-J.; Schallert, B. Auswirkung en von SCR-DENOX-Anlagen auf nachgeschltete Kraftwerkskomponenten. VGB Kraftwerkstech. 1991, 71, 584. (3) Gutberlet, H.; Finkler, S.; Pa¨tsch, B.; van Eldik, R.; Prinsloo, F. Bildung von Schwefel- und Schwefel-Stickstoff-Verbindungen in Rauchgasentschwefelungs-anlagen und ihr Einfluss auf die Oxidationskinetik von Sulfit. VGB Kraftwerkstech. 1996, 76, 139. (4) Siddiqi, M. A.; Krissmann, J.; Peters-Gerth, P.; Luckas, M.; Lucas, K. Spectrophotometric measurement of the vapour-liquid equilibria of (sulphur dioxide + water). J. Chem. Thermodyn. 1996, 28, 685. (5) Siddiqi, M. A.; Krissmann, J.; Lucas, K. A novel fibre-optic based technique for the spectrophotometric measurement of phase and chemical equilibrium of (sulphur dioxide + water) at elevated temperatures. J. Chem. Thermodyn. 1997, 29, 395. (6) Komiyama, H.; Inoue, H. Absorption of Nitrogen Oxides into Water. Chem. Eng. Sci. 1980, 35, 154. (7) Shadid, F. T.; Handley, D. Nitrous Acid Formation and Decomposition during the Absorption of Nitrogen Oxides. Chem. Eng. J. 1990, 43, 75. (8) Wagman, D. D.; Evans, W. H.; Parker, V. B.; Schumm, R. H.; Halow, L.; Bailey, S. M.; Churney, K. L.; Nuttals, R. L. The NBS Tables of Chemical Thermodynamic Properties. J. Phys. Chem. Ref. Data 1982, 11, Suppl. 2. (9) Schwartz, S. E.; White, W. H. Kinetics of Reactive Dissolution of Nitrogen Oxides into Aqueous Solution. Adv. Environ. Sci. Technol. 1983, 12, 1. (10) Aspen, Aqueous Datenbank, Aspen Plus, Version 9.3; Aspen Technology Inc.: Cambridge, 1997. (11) Schwartz, S. E.; White, W. H. Solubility Equilibria of the Nitrogen Oxides and Oxyacids in dilute aqueous Solution. Advances in Environmental Science and Engineering; Gordon and Breach Science Publishers: New York, 1981; Vol. 4, p 1. (12) Clifton, C.; Altstein, N.; Huie, R. E. Rate Constant for the Reaction of NO2 with Sulfur(IV) over the pH Range 5.3-13. Environ. Sci. Technol. 1988, 22, 586. (13) Sato, T.; Matani, S.; Okabe, T. The Oxidation of Sodium Sulfite with Nitrogen Dioxide. With Special Reference to Analytical Methods for Nitrogen-Sulfur Compounds produced in the Reaction System. Nippon Kaguku Kaishi 1979, 7, 869. (14) Lee, Y. N.; Schwartz, S. E. Kinetics of Oxidation of Aqueous Sulfur(IV) by Nitrogen Dioxide. Fourth International Conference on Precipitation Scavenging, Dry Deposition, and Resuspension, St. Monica, CA; Elsevier: New York, 1982. (15) Littlejohn, D.; Wang, Y.; Chang, S.-G. Oxidation of Aqueous Sulfite Ion by Nitrogen Dioxide. Environ. Sci. Technol. 1993, 27, 2162. (16) Ellison, T. K.; Eckert, C. A. The Oxidation of Aqueous SO2. 4. The Influence of Nitrogen Dioxide at Low pH. J. Phys. Chem. 1984, 88, 2335. (17) Martin, L. R.; Damschen, D. E.; Judeikis, H. S. The Reactions of Nitrogen Oxides with SO2 in Aqueous Aerosols. Atmos. Environ. 1981, 15, 191. (18) Candlin, J. P.; Wilkins, R. G. Sulphur-Nitrogen Compounds. Part II. The Hydrolysis of Hydroxylaminetrisulfonate and

Ind. Eng. Chem. Res., Vol. 42, No. 7, 2003 1413 Hydroxylamine-N,O-disulphonate Ions in Perchloric Acid. J. Chem. Soc. 1961, 3625. (19) Sander, S. P.; Seinfeld, J. H. Chemical Kinetics of Homogenous Atmospheric Oxidation of Sulfur Dioxide. Environ. Sci. Technol. 1976, 10, 1114. (20) Pitzer, K. S.; Kim, J. J. Thermodynamics of Electrolytes. IV. Activity and Osmotic Coefficients for Mixed Electrolytes. J. Am. Chem. Soc. 1974, 96, 5701.

(21) Edwards, T. J.; Maurer, J.; Newman, J.; Prausnitz, J. M. Vapor-Liquid Equilibria in Multicomponent Aqueous Solutions of Volatile Weak Electrolytes. AIChE J. 1978, 24, 966.

Received for review September 19, 2002 Revised manuscript received December 16, 2002 Accepted December 30, 2002 IE020739C