2110
Ind. Eng. Chem. Res. 1992,31, 2110-2118 Nitsch, W.; Schoor, A. V. The Kinetica of Coextraction in the System Uranylnitrate, Nitric Acid, Tributylphosphate. Chem. Eng. Sci.
Baumgartner, F.; Finsterwalder, L. On the Transfer Mechanism of Uranium(V1) and Plutonium(1V) Nitrate in the System Nitric Acid-Water/Tributylphosphate-Dodecane. J. Phys. Chem. 1970,
1983,38, 1947-1957.
Nitsch, W.; Schuster, U. Liquid/Liquid-Mass Transfer at Swarms of Droplets: Kinetics of Maas Transfer in the System Uranylnitrate, Nitric Acid, Tributylphosphate. Sep. Sci. Technol. 1983,
74,108-112.
Chilton, T. H.; Colburn, A. P. Mass Transfer (Absorption) Coefficients. Prediction from Data on Heat Transfer and Fluid Friction. Ind. Eng. Chem. 1934,26,1183-1187. Cox, M.; Flett, D. S. Metal Extractant Chemistry. In Handbook of Solvent Extraction; Lo, T. C., Baird, M. H. I., Hanson, C., Eds.; Wiley: New York, 1983; Chapter 2.2, pp 53-89. Edwards, J. 0. Correlation of Relative Rates and Equilibria with a Double Basicity Scale. J . Am. Chem. SOC.1954, 76,1540-1547. Eigen, M.; Tamm, K. Sound Absorption in Electrolytes as a Consequence of Chemical Relaxation. 11. 2.Elektrochem. 1962, 66,
18, 1509-1533.
Ohtaki, H.; Tanaka, M.; Funahashi, S. Yoekihannnou no Kagaku; Gakken Shuppan: Tokyo, 1977; Chapter 12, p 189. Padday, J. F. The Measurement of Surface Tension. In Surface and Colloid Science; Matijevic, E., Ed.; Wiley: New York, 1969; Vol. 1, Chapter 2. Rabinowitch, E.; Stockmayer, W. H. Association of Ferric Ions with Chloride, Bromide and Hydroxyl Ions. J. Am. Chem. SOC.1942, 64, 335-347.
107-121.
Roddy, J. W.; Coleman, C. F.; Arai, S. Mechanism of the Slow Extraction of Iron(II1) from Acid Perchlorate Solution by Di(2ethylhexy1)phosphoric Acid in n-Octane. J . Inorg. Nucl. Chem.
Fuoss, R.M. Ionic Association. III: The Equilibrium Ion Pairs and Free Ions. J . Am. Chem. SOC.1958,80, 5059-5061. Gouger, S.; Stuehr, J. Kinetics of Iron(II1) Interactions with Phenol and o-Aminophenol. Inorg. Chem. 1974,13,379-384. Harada, M.; Miyake, Y. Solvent Extraction with Chelating Agents. In Handbook of Heat and Mass Transfer;Cheremisinoff, N. P., Ed.; Gulf: Houston, 1989; Chapter 21, pp 789-881. Horner, D.; Mailen, J.; Thiel, S.; Scott, T.; Yates, R. Interphase Transfer Kinetics of Uranium Using the Drop Method. Znd. Eng. Chem. Fundam. 1980,19, 103-109. Imura, H.; Takahashi, H.; Ueki, Y.; Suzuki, N. Liquid-liquid Extraction of Rhodium(II1) in Various Systems: Extraction Behavior of Some Rhodium(II1) Species in Aqueous Solutions. R o c . Symp. Solvent Extr. 1988, 1, 1-6. Matauyama, H.; Miyake, Y.; Izomo, Y.; Teramoto, M. Kinetics and Mechanism of Metal Extraction with Acidic Organophosphorus Extractanta (11): Extraction Mechanism of Fe(II1) with Di(2ethylhexy1)Phosphoric Acid. Hydrometallurgy 1990,24,37-51. Miyake, Y.; Takenoshita, Y.; Teramoto, M. Extraction of Copper with SME529. J . Chem. Eng. Jpn. 1983,16, 203-209. Miyake, Y.; Matauyama, H.; Nishida, M.; Nakai, M.; Nagase, N.; Teramoto, M. Kinetics and Mechanism of Metal Extraction with Acidic Organophosphorus Extractanta (I): Extraction Rate Limited by Diffusion Process. Hydrometallurgy 1990, 23, 19-35. Niitau, M.; Sekine, T. Solvent Extraction Equilibria of Acids-I. Extraction of Nitric and Perchloric Acid with Trioctylphosphine Oxide. J. Inorg. Nucl. Chem. 1975,37, 1056-1057.
1971,33, 1099-1118.
Sekine, T.; Tetauka, T. The Solvent Extraction of Iron(II1) in Perchlorate Solutions Containing Chloride or Bromide Ions with 2Thenoyltrifluoroacetone and Trioctylphosphine Oxide. Bull. Chem. SOC.Jpn. 1972,45, 1620-1625. Sekine, T.; Honda, H.; Zeniya, Y. Kinetic Studies of Solvent Extraction of Metal Complexes-VI. J3quilibrium and Rate of Solvent Extraction of Iron(II1) Perchlorate with Trioctylphosphine Oxide in Hexane. J. Inorg. Nucl. Chem. 1976,38, 1347-1350. Tarasov, V. V.; Yagodin, G. A. Interfacial Phenomena in Solvent Extraction. In Ion Exchange and Solvent Extraction; Marinsky, J. A,, Marcus, Y., Eds.;Dekker: New York, 1988; Vol. 10, Chapter 4.
Wilke, C. R.; Chang, P. Correlation of Diffusion Coefficients in Dilute Solutions. AIChE J. 1955, 1, 264-270. Yamada, S.; Tanaka, M. Softness of Some Metal Ions. J. Inorg. Nucl. Chem. 1975, 37, 587-589. Yasunaga, T.; Harada, S. Temperature-jump Studies of the Kinetics of the Formation of the Monochloro and Monobromo Complexes of Iron(II1). Bull. Chem. SOC. Jpn. 1969, 42, 2165-2169.
Received for review January 13, 1992 Revised manuscript received June 1, 1992 Accepted June 12, 1992
Mechanism and Kinetics of the Reaction between Sulfur Dioxide and Ammonia in Flue Gas Klaus Hjuler and Kim Dam-Johansen* Department of Chemical Engineering, Technical University of Denmark, Building 229, DK-2800Lyngby, Denmark
The reaction between sulfur dioxide and ammonia in flue gas is of importance in understanding scaling on process equipment and monitoring systems in a number of flue gas cleaning processes. The reaction was studied a t atmospheric pressure using simulated flue gas, 1.5-12.0-mm-internal diameter tube reactors, and flows ranging from 0.5 to 4.0 (NL)/min. The reaction temperature, sulfur dioxide concentration, and ammonia concentration were 40-80 "C,e 1 0 0 0 ppm, and 120-2000 ppm, respectively. The water content of the gas was typically 6 vol 9% and the oxygen content 5 vol ?%. The fiist step of the reaction is the formation of liquidlike ammonium salt containing sulfur in oxidation state IV on surfaces. With oxygen and small amounts (5 ppm) of nitrogen dioxide present, S(1V) is irreversibly oxidized to S(VI), and the product becomes solid and thermally stable. This product contains primarily ammonium sulfate. An overall mechanism for the reaction involving two intermediary species is proposed, and a simple three-parameter expression for the reaction rate has been developed from the mechanism. Introduction Ammonia is generally used for the selective catalytic and noncatalytic reduction of nitrogen oxides and for the removal of sulfur dioxide in wet scrubbing. Scaling on catalyst surfaces or in flue gas ducts or monitoring sample lines by solid or liquid ammonium salt is a potential problem in these processes, and more information on the 0888-5885/92/2631-2110$03.00/0
formation mechanism is needed in order to avoid it. On the other hand, the production of ammonium salt is desirable in the electron beam dry scrubbing process, where sulfur dioxide removal is due substantially to the reaction between sulfur dioxide and ammonia. Very little is in fact known about the reaction mechanism, its kinetics, or its products. Most investigations of 0 1992 American Chemical Society
Ind. Eng. Chem. Res., Vol. 31, No. 9,1992 2111 the reaction between gaseous sulfur dioxide and ammonia have been performed in the field of atmospheric chemistry. In the atmosphere, the temperature and the concentrations of sulfur dioxide and ammonia are significantlylower than in flue gas, and consequently such investigations are not necessarily relevant. As will be discussed later, some information is available on the thermodynamics of the formation of ammonium sulfite from ammonia, water vapor, and sulfur dioxide. However, under flue gas conditions it has been observed that the reaction product primarily contains sulfur in oxidation state VI @(VI)), i.e. sulfate. Thus the thermodynamics, mechanism, and kinetics of the S(IV) oxidation step are of particular interest. In this study the reaction between SO2, NH3, and H20 was investigated in the temperature range from the water dew point of the gas to the temperature where no reaction was observed. The gas composition was typical of a coal flue gas. The inlet sulfur dioxide, ammonia, and water concentrations were varied to examine the kinetics. The influence of flue gas components such as oxygen, carbon oxidea, nitrogen oxidea, sulfur trioxide, and fly ash was also investigated. S(VI)product composition was studied by X-ray diffraction.
Figure 1. Experimental setup.
0.8
N
% 0.6 L
0
Experimental Setup Preliminary experiments were carried out using an oven-thermostated, variable volume, 25-mm4.d. mixedflow reactor with a movable "piston" bottom where a 2-wm membrane filter was placed. These experiments made clear that the reaction was very sensitive to temperature. Because of problems with gas mixing and temperature control, tube reactors operated in the laminar flow range were found to be better suited. The tube reactors were various lengths of 1.5-12.0-mmi.d. tubes made of steel, Teflon, nylon, or glass, thermostated either in an oven or by a heating jacket. The tube length, gas flow rate, and residence time were typically 20-30 cm, 0.6-0.7 (N L)/min, and 0.2 s, respectively (at 20 "C and atmospheric pressure). A simulated flue gas was mixed in a manifold from cylinders of the individual components by means of reduction valves and mass flow controllers. The manifold and the gas lines up- and downstream of the reactor were heated to above 100 OC to avoid condensation of water and undesirable reactions. The inlet gas temperature had little or no influence on the observed conversion in the 40-120 OC temperature range; any reador temperatures quoted in the following are therefore surface temperatures. Water vapor was introduced by humidifying a stream of nitrogen or by evaporating water from an HPLC pump. Sulfur dioxide was measured continuously using nondispersive dual wavelength absorption (Hartman & Braun Radas 1G)and multiwavelength differential absorption (OPSIS AR 600) in the UV range. The OPSIS analyzer was used extractively and equipped with a 14 cm long measuring cell thermostated at 150 OC. The OPSIS analyzer was also used for continuous monitoring of ammonia. Nitrogen oxide and oxygen were measured respectively by means of chemiluminescence (Monitor Labs) and a paramagnetic gas analyzer (Hartman & Braun Magnos 5T). In the following, gas concentrations are given as volume fractions for the sake of convenience. Figure 1shows the experimental setup with gas-mixing manifold, water vaporizer, and jacket-thermostated tube reactor. Also shown is the setup for the OPSIS analyzer: an oven with a UV lamp and a measuring cell. The light is transmittd to the spectrometer through an optical fiber. The spectrometer was operated by means of a built-in
0 0.2 I
t 0.0 45
50
55
-
Surf ace Temperature (deg .C)
60
Figure 2. Conversion of SO2versus temperature in a 30-cm reactor (a) and a 435-cm reactor (0). The a data were taken from four similar experiments. Inlet conditions: 930-1100 ppm SO,; NH3/S02 ratio, 1.7-2.0; 5.9 vol % HzO. Flow range: 0.8-0.9 (N L)/mm.
personal computer, which was also used for data acquisition and analysis. Experimental Results (I): Formation of Sulfite The experimental study started with the simplest possible system, namely, a mixture of SO2, NH3, and H20 in NP The concentrations (volume fractions) of SO2and NH3 were in the ranges 930-1100 and 1580-2150 ppm, respectively. N2was added to balance; about 5 vol % of O2was added in some runs. All experiments were carried out at atmwphericpressure using oven-thermostated Teflon tube reactors (id. 3.2-4.0 mm). The reaction product was always found on the reactor walls, especially close to the ammonia inlet. It was liquidlike, yellowish in color, and unstable with respect to decomposition: it sublimed when the reactor was heated or one of the gaseous components was removed, indicating that sulfur in oxidation state IV was a main constituent, as discussed below. No direct influence of oxygen on the reaction rate was found, but the product oxidized slowly in the presence of oxygen, forming a solid, white ammonium salt containing S(V1). This is in agreement with results obtained by Landreth et al. (1975),Scargill (1971), and Haber et al. (1974). Figure 2 shows the conversion of SO2versus the surface temperature for tube reactors 30 cm (0) and 435 cm ( 0 ) in length. Through the transparent Teflon wall of the 435-cm reactor, it was observed that product deposita formed on the first 60 cm only, corresponding to an "active" residence time of about 0.46 s at 50 O C . Conse-
2112 Ind. Eng. Chem. Res., Vol. 31, No. 9, 1992 '9 0.6
-
&
-
c 0 0.4
\'
0
- \,'\ $
C 0
0:.
"0
?\
O
0
\\
0.0-
O
'
K
.-0 0.4
a
'
P
0.2 0
.
Eq: (NH4)rSOa
Eq: (NHI)zSO,
% ':K
a
-
*
.,"\ \'
le\, -
Pa,
0
0
C
.-0
\
o',,',
6 v,
0.8
0 Exp. equilibrium
-Y&,
\\e\ \
\ \
\\
0
L
'e
\
%
3
b
45
Figure 3. Comparison between experimental equilibrium data and data calculated using the equilibrium constanta listsd in Table I. Inlet conditions: 1070 ppm SO,; NH3/SOZratio, 1.7; 6.0 vol % H20. Table I" ammoniumpbosulfite (a): (NH,),SOS(s) + 2NH3(g) + 2SOAg) + H2O(g) ammonium sulfite (b): (NH~)ZSO~(E)2NH3(g) + SOz(g) + HzO(g) hydrated ammonium sulfite (c): ( N H ~ ~ ~ S O ~ * H+ Z O2"3(g) (S) + SOz(g) + 2H20k) ammonium bisulfite (d): NH4HSO&) + NH,(g) + SOzk) + HzOk) Ka p"Q2pS0:pH20 log K, 41.89 - 17705/T Kb = PNH$ISO&~O log Kb = 33.27 - 14171/T K, = p N H s P s O ~ H ~ ~ log ~ K, = 42.00 - 17411/T log Kd = 23.77 - 9958/T Kd = P N H ~ P S O ~ H ~ O
is equilibrium partial pressures in atmospheres and T the temperature in kelvin. O p
quently, equilibrium at the outlet was assumed in this experiment. However, in Figure 2, it is shown that the results from the equilibrium experiment do not differ from the data obtained using the 30-cm reactor, and it was concluded that the formation of sulfite took place rapidly. Notice that the scattering of the data reflects the experimental uncertainty.
Thermodynamics of Sulfite Formation Equilibrium constants of the formation of four S(1V) salts, given by Scargill (1971) and Tock et al. (1979), are listed in Table I. In Figure 3, the temperature profile of the abovementioned equilibrium experiment is compared to profiles calculated from the thermodynamics. Bearing in mind the experimental uncertainty, the composition of the S(IV)salt or salt mixture formed cannot be determined on this basis. It was found that results obtained from thermodynamic data on ammonium salt formation may vary to a great extent. For example, thermodynamic functions derived by Tock et al. (1979)based on data given by Earhart (1969) were tested but gave significant deviations from both experimental results and calculations using the expressions in Table I. In experiments carried out by Landreth et al. (1975), the enthalpy and entropy for ammonium sulfite formation was determined with about 20% uncertainty. Experimental Results (11): Formation of Sulfate The reaction product obtained under flue gas conditions primarily contained sulfur(V1). A search was therefore initiated to find one or more flue gas components which in concentrations typical of coal combustion flue gases effected a rapid oxidation of S(1V) or a direct formation of sulfate. The search was conducted adding one compo-
50
60
55
70
65
S u r f ace Tempera t u r e (deg .C) Figure 4. Temperature profiles obtained with HCI (0) and NO present ( 0 , A) compared to the case without additive (the area within the broken lines). Inlet conditions: 775-1110 ppm SO,; NH3/SOZratio, 1.7-2.4; 5.5-6.0 vol % H,O; 4.5 vol % 0,.Specific surface area of the tube reactor: 3.0-4.4 m2/((N m3)/min) at 0.74.8 (N L)/min.
6 cn 0.6 + 0
n
.-80 0.4
-
\
Pa, >
c 0 0.2 0
0.0
t'
45
' ' ' '
4
o-\
x
' 50
' ' '
O\O 0
r *' 55
'
" "
60
' '
0
'
' '0'
65
'AI 70
Su rf ace Tern peratu re (deg .C) Figure 5. Comparison of temperature profiles obtained with and without NOz. Inlet conditions: 930-1100 ppm SO,; NH3/SOZratio, 1.7-2.0; 5.5-6.0 vol % HzO 4.5 vol % OP Reactor specific surface area: 3.0-4.4 m2/((N m3)/min).
nent at a time to a gas containing SO2, NH3, H20,02,and N2. Thermodynamics indicate that NH4C1aerosols may be formed by reaction between NH3 and HC1 in flue gas at or below about 100 OC (Barin, 1989). On the basis of the hypothesis that NH&l particles may act as nuclei for the formation of other solid ammonium salts, HCl(g) was introduced. Figure 4 compares the results obtained by introducing 15 ppm HCl (0) to results obtained without additive. The broken lines define an area of experimental uncertainty beyond which any effect is assumed to be significant. Evaluated in this way, HC1 has no significant effect. Also shown in Figure 4 are the results obtained by adding 300 and 3000 ppm NO. With 300 ppm NO present (o),the temperature at which reaction was first observed (initiation temperature) was not significantly higher than it was without the additive. With 3000 ppm NO present (A),the initiation temperature was significantly higher, and the slope of the profile is not as steep. In order to check if this was due to the formation of NO2 from NO, NO2 was added. Figure 5 shows the marked effect of NO2on the observed conversion of SO2 versus temperature. The data shown are taken from a number of experiments performed under
Ind. Eng. Chem. Res., Vol. 31, No. 9, 1992 2113 0.6
5 ppm NOz
0
1 \a, 1
1
1
1
1
1
1
1
1
,
1
1
1
10 ppm NO2
10 pprn NO2
0.5 A
0" V, 0.4
, , , , , ,
20 ppm NOz
c
0 0.3
.-0 Pe, 0.2 >
C
0
00.1
55
60
65
75
70
Surface Tern perat ure (deg r.C)
-.- 55 nn
65
60
70
Surface Tern peratu re (deg .C) Figure 6. Temperature profiles obtained by varying the concentration of NO2. Inlet conditions: 920-1000 ppm SO,; NH3/SO, ratio, 2.0; 5.5-6.0 vol % HzO; 4.5 vol % 0,.Reactor specific surface area:
Figure 8. Temperature profile obtained with the addition of 15 ppm NO, and 300 ppm NO (X) compared to the 10 and 20 ppm NO, profiles shown in Figure 6. Inlet conditions: lo00 ppm SO,; NH3/SOpratio, 2.0; 6.0 vol % H20;4.5 vol % 02.Reactor specific surface area: 4.5 m2/((Nm3)/min).
4.4 mZ/((Nm3)/min).
0
-
0.5
0 " -V, 0.4 %
-
10
~01%
Hz0
0
6 VOI-S HzO
A
3
O
VOI-S
-
HzO
1.5 VOI-S Hz0
-
-
0
0
1
2
3
Oxygen Concentration (vol-n)
4
Figure 7. Influence of oxygen on the reaction at 59 "Cwith 16 ppm NO2 present. Notice the starting point of the Y-axis. Inlet conditions: 960 ppm SO,; NH3/SOz ratio, 2.0; 5.8 vol % HzO; 4.5 vol % OZ.Reactor specific surface area: 4.4 m2/((N m3)/min).
similar conditions. In the presence of only 15 ppm NO2 (and 5 vol % 02), the initiation temperature of the reaction increased about 15 "C. Moreover, the appearance of the reaction product changed from liquidlike to white and crystalline. The temperature profiles shown in Figure 6 were obtained by varying the NO2concentration from 5 to 30 ppm. It is apparent that the conversion of SO2at a given temperature increases with the addition of NO2. Sudden "jumps" in the profiles were observed in several cases as shown, which probably indicate a surface dependence of the reaction. The influence of the oxygen concentration on the reaction at 59 "C with 16 ppm NO2 present is shown in Figure 7. Above about 1 vol %, the reaction order in oxygen is apparently zero. With no oxygen present, a conversion of zero is expected because S(1V) is not stable at the actual conditions. Oxygen was not removed completely, due to the presence as an impurity of dissolved oxygen in the humidifier water and in the nitrogen cylinders. The search was continued with the addition of fly ash, HC1, NO, CO, C02, CHI, and SO3to a gas containing SO2, H20, NH3, NO2, and 02. Fly ash was introduced onto quartz wml placed in a reador t u h experiments were also conducted with quartz wool alone for comparison. Only NO was found to have a significant effect: with both NO and NO2preaent, a "tail"was observed on the temperature profile (Figure 8). This may be due to an effect of NO on the absorption of NO2in a water film layer, as discussed below. Presence of 10-20 ppm of SO3 did not influence
Figure 9. Observed conversion of SO, versus the reactor surface temperature at water contents between 1.5 and 10 vol %. The upward and downward arrows indicate cooling and heating profiles, respectively. The cooling profile with about 3 vol % water was not obtained. The reactor was equipped with a heating jacket. Inlet conditions: 500 ppm SO2;NH3/S02ratio, 2.0; 5.5-6.0 vol % H,O; 4.5 vol % 0,;10 ppm NO,. Reactor surface area: 27 cm2. Gas flow: 0.58 (N L)/min.
SO2 removal at all, but formation of ammonium sulfate via condensed sulfuric acid was observed above 100 O C .
Surface Dependence Phenomena indicating a surface dependence of the reaction have been observed in several experiments, the sudden jumps shown in Figure 6 being an example. Another example is the temperature hysteresis shown in Figure 9. These results were obtained when starting from a reactor temperature around 70 OC, cooling at a rate of 0.1 "C/min to about 55 "C, and then reheating (without removing the product formed) until no conversion was observed. No product was present initially. As shown, the cooling and heating profiles are not identical. In one experiment, the above procedure was repeated three times without cleaning the reactor in between (without removing the product formed). The heating profile remained about the same, while the gap between cooling and heating profiles diminished gradually. Finally, an influence of the reactor construction material on the initiation step of the reaction was observed. Most significant was the effect when a nylon tube (i.d. 4.5 mm/length 44 cm) was used: at 61 "C, the conversion of SO2was only 0.08 compared to 0.47when a Teflon reactor
2114 Ind. Eng. Chem. Res., Vol. 31, No. 9, 1992
-
n 4 1.4 1.7
1 .e
2.1
2.3
Inlet NH3/S02 Stoichiometry
F i r e 10. Observed N to S reaction stoichiometry with NO2 preaent calculated from measured inlet and outlet concentrations of NH, and SOI. Inlet conditions: 516560 ppm SO2;NH,/SO, ratio, 1.85-2.10: 5.4-6.1 vol % H,O 4.5 vol % 0,:10 oom NO,: reaction tempera&, 39-59 oC. &iZO cma.-'Gaaflow: 0.6-3.9 (N L)/min.
with similar dimensions, inlet conditions, and thermal history was employed. When the nylon reactor was cooled about 10 "C and then reheated to 61 'C, the conversion increased to 0.51. The choice of construction material had no effect once the surface had heen covered with reaction product. Composition of Sulfur(V1) Product The reaction product obtained in the presence of both NO2and 0,was thermally stable above 100 OC. Very little or no generation of SO, and NH3 was detected when product formed in the presence of NO, and 0,at the S ( N formation temperature was flushed with N, or heated moderately. Two samples of product, one obtained at 60 OC and one at 65 "C,were analyzed by X-ray diffractometry. The characteristic peaks of the samples were compared with reference characteristic peaks of ammonium sulfate, (NH4),S04,and ammonium pyrosulfate, (NH4),S207,and most major peaks of the samples could be accounted for. From the ratio of the height of the highest (NH&SO, peak to the height of the highest (NH4),S,0, peak, the molar N to S ratios of the samples were found to be 1.64 and 1.77, respectively. Taking the uncertainty of the method into consideration, it was concluded that the two samples were identical, with a molar mass of about 146 i 5 g/mol. The overall reaction stoichiometry, a,was determined on the basis of measured outlet concentrations of SO, and NH,. Figure 10 shows a versus the inlet N to S ratio. The reaction stoichiometry calculated from the 32 data points shown is 1.M (solid line) f 0.15. The result from the X-ray analysis agrees quite well with this result. Discussion The marked effect of NO, on the reaction rate indicates that the oxidation of SO, is not directly affected by 0,but that some intermediate step is involved. A similar conclusion has heen drawn from many years of experience with the lead-chamber process used previously for the manufacture of sulfuric acid (Duecker and West, 1959). In a theory of Berl(1935) it is proposed that the oxidation of SO, actually takes place at the gas-liquid interface and that the reaction is a cyclic process involving the alternate oxidation and reduction of a transiently formed intermediate species,written as NO-H+304. Due to decomposition of NO.H2SO4, the operation of the cycle is dependent on a continuous supply of NO,(g). A catalytic effect of NO, on sulfite oxidation by 0,was reported by Takeuchi et al. (1977) and Wittig et al. (1988) in agreement with the results obtained in this study, because as little as 5 ppm NO,
Figure 11. Mechanism proposed.
is capable of oxidizing, for example, 500 ppm SO, in the presence of 5 vol % 0,. The effect of NOz on the initiation temperature of the reaction (Figure 5) may be due to the heterogeneous reactions of NO, and HzO studied by Jenkin et al. (1988). Jenkin et al. found that adsorbed nitric acid, HNO,(ads), was formed and that the reaction NO,(adn) + H,O(ads) NOz-HzO(adn) (1) determined initial kinetics. The study was carried out at room temperature and partial pressures of NO, and HzO of about 4 ppm and 0.5 vol % HzO, respectively. With dependence on the relative humidity, NOz in probably absorbed in a water film rather than adsorbed prior to reaction. According to Schwartz and White (1981), gaseous NO2 is in equilibrium with NO; and NO3-:
-
2N02(g) + H,0(1)
+ 2H+ + NO3- + NOz-
(2)
3NOz(g) + HzO(l) % 2H+ + 2N03- + NO@
(3)
NO(g) + NO,(g)
(4)
+ HZO(1) % 2H+ + 2NO2'
Notice in eqs 2-4 that the molar ratio of NOz- to NO&) is l/,, 0, and 2, reapectively. Martin et al. (1981)concluded that NO,- oxidizes S(IV)more rapidly than diasolved NO or NO, does. Thus if NO, is actually absorbed and NOzis rate-limiting, then the rate of reaction should increase with NO, partial pressure and in the presence of NO (eq 4), as was also shown experimentally (Figures 6 and 8). Martii et al. (1981) also reported that nitrous oxide, N,O, was evolved at pH values below 2. However, no N20 was detected in this study when Samples were analyzed using gas chromatography. Oblath et al. (1981) found that the initiate product of the reaction between NO2- and HSO, wan hydroxylamine disulfonate (HADS), HON(SO3)?-, at 15-30 "C in the 4.5-7 pH range using approximately equimolar solutions: NO,-
+ 2HS03- =, HON(S03)z2-+ OH-
(5)
The HADS eventually reacts further with SO,?-or HzO to produce S(VI). On the basis of eqs 2-4, an increane in pH is expected to increase the rate of absorption of NOz, while the rate of oxidation of S(IV) decreases, eq 5. Investigations carried out by Tanaka et al. (1979) a t 50 "C showed a 300-fold increase in the rate of conversion of NOz- when the pH value was decreased from 7.5 to 5.6. Proposal for Reaction Mechanism and Kinetics It is proposed that the reaction mechanii can be simp l i e d as outlined in Figure 11. The proceas takes place on a surface, and three phases are considered gas, water film, and solid. Except for the solid, all species situated on the surface are considered as being absorbed in adsorbed water (or water film), assuming that a few or eaveral
Ind. Eng. Chem. Res., Vol. 31, No. 9, 1992 2115 moles per unit surface area of R1and & present. kl and klare defined by k 1 / k l = kl*/kl*Kso *bKN abKWaQ where Krb are absorption Constants, and KwaBis an ad: sorption constant of water. Similarly we have k 3 / k 3= k?*/k-3*KN0;bS. The absorption constants, Killb, are derived below from the pertinent wet-chemical equilibra:
' B h¶
HZO(1) + SOz(g) t;H2SO3
5 0.2
0 0.0
(9)
KlO
0
20
40
60
80
HzS03'=; H+ + HS03-
100
Spec. Surface Areo, m2/(Nm3/min)
Figure 12. Fit of k, and kox (solid line) using data obtained at 60 "C with 15 ppm of NOz present. Inlet conditions: lo00 ppm SO,; NHB/SOzratio, 2.0; 6.0 vol % H,O; 4.5 vol 7% 0,;10 ppm NO,.
layers of adsorbed water molecules can provide a medium for fast solution chemistry. The first step of the reaction is the absorption of SO2, NH3, 02, and NO2 into the water film. This step is not considered to be rate-limiting: rapid equilibrium exists between the gaseous and absorbed species. The next step is the formation of intermediary soluted species. The existence of two intermediary products (Rl and R2) is proposed, containing sulfur in oxidation states IV and VI, respectively. & is the product of the reaction between R1 and absorbed NOz. The final step of the reaction is the transformation of R1and R2 into crystalline ammonium salts denoted S(1V) and S(VI), respectively. S(1V) is unstable with respect to decomposition, while S(V1) is thermally stable under the conditions studied. S(V1) is primarily ammonium sulfate. An important feature of the mechanism is that adsorbed NO2acts as an oxygen carrier, so that & is formed only in the presence of both NOz and 02.Notice that the actual chemical compositions of R1 and Rzare not precisely known and that additional intermediary products may be formed in the course of the reaction. The rate of oxidation of S(V1) is probably influenced by film conditions such as pH and the presence of trace metals or oxidants, but this will not be considered. A simple mathematical model has been developed on the basis of the mechanism and the following assumptions: (1)No reaction takes place in the gas phase. (2) The first reaction step is an absorption of gaseous reactants in a water film on surfaces. The quantity of water adsorbed is approximated, assuming direct proportionality with the water partial pressure. (3) Direct proportionality exists between the partial pressures and the film concentrations of SOz, NH3, and NOz. All equilibra are rapid. (4) There is quasi-stationarity with respect to the intermediary products Rland & and zero-order dependence in solid S(1V) and in 02(abs) (cf. Figure 7). There is first-order dependence in H20(ads),SOz(abs),NH3(abs), and NOz(abs). (5) The direct oxidation of S(1V)to S(V1) by Oz is slow and does not influence the overall reaction rate. The effect of gaseous NO on NOz absorption is not considered. (6) The overall reaction rate is not limited by the crystallization process. Equations 6-8 are derived, where Rjis the rate of formation in number of moles of component i per unit surface area and per unit time. N R 1 and N R z are the number of RS02 klPH20PS0#NH3 - k-lNR1 (6)
= k-lNR1 - klPH20PS0#NH3 + k2NR1 - k-2 + ~~NO,N kR- 3, N ~=~0 (7) (8) R, = k - 3 N ~- ~k $ N 0 2 N R l + k4Nk = O
(10)
Kii'
H+ + S032- + HS03net:
Kmh
HzO(l) + S02(g) + S032- % 2HS03-
where Ks0;"
(11) (12)
= h&Kl,,/Kll). With respect to NH3we have
-
h13
NH3(g)9 NH3(aq) NH3(aq) + H 2 0 KIS
K14
(13)
NH4++ OH-
HSOc * H+ + S032OH- + H+
(14)
(15)
Kw-1
+ HzO
(16)
net: NH,(g)
+ HZO(1) + HSO,
Kmh 9
H20
+ 502- + NH4+ (17)
where K"sabs = h13(K14K15/KW). Making eqs 7 and 8 equal to zero due to the assumption of quasi-statiomrity and rearrangement gives
where
As long as product is formed, it is reasonable to assume that kl
