Ind. Eng. Chem. Res. 1990,29, 1+14
10
desired ethylene is reduced in favor of the primary product ethane. The ratio of methane to oxygen should not be less than 5-7 since the otherwise nonselective reaction paths become very important. From a process point of view, final optimization also has to take into account separation of the various products. Conclusions for t h e Catalyst Design. From the reaction scheme and its application for predicting the various reactive radicals and products in the gas phase, some suggestions are derived for an optimum catalyst design that might result in better coupling selectivities than in the homogeneous gas phase. When pursuing this goal, it should be remembered that the present knowledge on catalytic methane coupling assumes that methyl radicals are formed on the catalytic surface by dissociation of methane and continue to react partly or fully in the gas phase after desorption. The primary coupling reaction to ethane is second order with respect to the methyl radical concentration, while the nonselective reactions are first order with respect to the concentration of methyl radicals and the various oxygencontaining species involved. For this reason, the concentration of methyl radicals should be kept high; this is presumably most easily achieved on or close to the catalytic surface. This requirement of a high concentration of CH,' radicals can be realized when the rates of methane adsorption and subsequent dissociation are fast and when the adsorption capacity of the catalytic surface for methyl radicals is high. The rate constant for the coupling step, if it occurs on the surface, should be enhanced by the catalyst. Simultaneously the catalyst surface should act as a scavenger for the gas-phase radicals containing reactive oxygen (HOP', OH', 0:)and it should preferably catalyze the reaction between these species and the hydrogen dissociated from the adsorbed methane molecule.
Warnatz of the University of Stuttgart. Registry No. CH4, 74-82-8; C2Ha,74-84-0; C2H4,74-85-1; CO, 630-08-0;Cop, 124-38-9; C2H2, 74-86-2. Literature Cited Asami, K.; Omata, K.; Fujimoto, K.; Tominaga, H. J . Chem. SOC., Chem. Commun. 1987a,1287. Asami, K.; Shikada, T.; Fujimoto, K.; Tominaga, H. Ind. Eng. Chem. Res. 1987b,26, 2348. Bos, J. Diploma Thesis, Technical University Twente, The Netherlands, 1986. Cathonnet, H.; James, H. J. Chim. Phys. 1975,72,247. CODATA. J . Phys. Chem. Ref. Data 1984,13 (4), 1259. Coffee, T. P. Combust. Flame 1984,55,161-170. Cohen, N.; Westberg, K. R. J . Phys. Chem. Ref. Data 1983,12(3), 531. Deuflhard, P.; Bader, G. Tech. Report TUM-MATH-7821; Technische Universitfit Munchen: Miinchen, 1978. Frenklach, M.; Bornside, D. E. Combust. Flame 1984,56, 1-27. Gear, W. C. Numerical Initial Value Problems in Ordinary Differential Equations; Prentice-Hall: Englewood Cliffs, NJ, 1971. Hinsen, W.; Baerns, M. Chem.-Ztg. 1983, 107,223-226. Hinsen, W.; Bytyn, W.; Baerns, M. Proc. 8th Int. Congr. Catal., Berlin, 1984; Vol. I11 pp 581-592. Ito, T.; Wang, J.-X.; Lin, C.-H.; Lunsford, J. H. J. Am. Chem. SOC. 1985,107,5062-5068. Lane, G. S.; Wolf, E. E. J . Catal. 1988,113,144-1637. Lee, d. S.; Oyama, S. T. Catal. Rev.-Sci. Eng. 1988,30,249-280. Levy, J. M.; Sarofim, A. F. Combust. Flame 1983,53,1-15. Lunsford, J. H.; Campbell, K. D.; Morales, E. J . Am. Chem. Soc. 1987,109,7900-7901. Onsager, 0. T.; Lodeng, R.; Soraker, P.; Anundskaas, A.; Helleborg, B. Catal. Today 1989,4, 355-363. Pitchai, R.; Klier, K. Catal. Rev.-Sci. Eng. 1986,28, 13. PreuB, U. Unpublished data, Ruhr-Universitat Bochum, 1986. Sofranko, J. A.; Leonard, J . J.; Jones, C. A. J . Catal. 1987, 103, 302-310. Warnatz, J. Eer. Bunsenges Phys. Chem. 1983,87,1008-1022. Warnatz, J . In Combustion Chemistry; Gardiner, W. C., Ed.; Springer Verlag: New York, 1984; p 197. Westbrook, C. K. Combust. Sci. Technol. 1929,20,5. Westbrook, C.K.; Dryer, F. L.; Shug, K. P. Combust. Flame 1983,
Acknowledgment
52, 299.
This work has partly been supported by the Commission of the European Communities under Contract EN3C/ 0023-D (M.B.). The computer program used for the kinetic simulations has been kindly provided by Professor J.
Received for review May 9, 1989 Revised manuscript received May 17, 1989 Accepted August 29, 1989
Reaction of Ferrous Chelate Nitrosyl Complexes with Sulfite and Bisulfite Ions David Littlejohn and Shih-Ger Chang* Applied Science Division, Lawrence Berkeley Laboratory, 1 Cyclotron Road, Berkeley, California 94720
The reaction of ferrous chelate nitrosyl complexes with sulfite and bisulfite ions has been studied, and a mechanism for the reaction has been developed. The mechanism has three components: (a) the reaction of coordinated nitric oxide with sulfite ion to form (NO).-$O,2-; (b) the oxidation of Fen(L) (L = ligand) by (N0)2S032-; (c) the redox reaction of Fe"(L)NOHS03- with Fe"(L)NO to generate NOS03-, which subsequently reacts with HS03- to produce HON(SO3)2-. The reaction products observed are Fe'", N20, N2, S042-,S20G2-, and HON(SO3)?-. An empirical reaction rate expression for the Fe'I(NTA)NO and Fe"(EDTA)NO complexes is reported. The ability of ferrous chelates such as Fe"(EDTA) and Fe"(NTA) (where EDTA = ethylenediaminetetraacetic acid and NTA = nitrilotriacetic acid) to strongly and reversibly bind nitric oxide to form ferrous chelate nitrosyl complexes has made them of interest in flue gas clean-up studies (Chang et al., 1983; Teramoto et al., 1978). Ap0888-5885/90/2629-0010$02.50/0
proximately 95% of the nitrogen oxides in flue gas is in the form of relatively insoluble nitric oxide. SO2 is generally also present in flue gases of coal-fired power plants, and after dissolving in aqueous solution, it will form S02H20,bisulfite ion (HSOJ, or sulfite ion (S032-),depending on the pH of the solution. It is highly desirable 0 1990 American Chemical Society
Ind. Eng. Chem. Res., Vol. 29, No. 1, 1990 11 for a scrubbing system to simultaneously remove NO, and SO2from the flue gas stream and be capable of regenerating the active compounds. There have been a number of studies (Sada et al., 1980, 1984, 1987a,b; Miyadera et al., 1978; Narita et al., 1984) done recently on the reaction of ferrous nitrosyl chelate complexes with sulfite and bisulfite ions, primarily with EDTA as the chelate. The reports of these studies indicate that the reaction is complicated and have not provided a complete understanding of the reaction process. Sada et al. (1987b) have determined the reaction products and have studied the reaction kinetics over a period of hours. However, no mechanism based on fundamental reactions has been proposed for the generation of the observed products. The kinetics of the initial reaction of Fen(L)NO (L = ligand) with HS03- and S032-have not been determined. Also, the effect of pH on the nitrogen-containing products has not been reported. We report the results of our studies on the kinetics and reaction products of the reaction of Fe"(L)NO with HSOC and SO:-. We present a reaction mechanism capable of explaining our observations.
Experimental Section The preparation of the ferrous chelate nitrosyl solutions has been described elsewhere (Littlejohn and Chang, 1982). Chelates used in this study include EDTA, NTA, iminodiacetic acid (IDA), and citric acid (cit). This work was divided into two categories-kinetic studies and reaction product studies. The kinetics of the disappearance of the ferrous nitrosyl complexes were monitored by observing one of the visible absorption bands of the nitrosyl complex with a Cary 219 UV-visible spectrophotometer (Littlejohn and Chang, 1982). A degassed solution of sodium sulfite and/or sodium metabisulfite was added to a 10-cm path length cell, followed by the ferrous chelate nitrosyl solution. After mixing, the absorption decay was monitored. Reaction product studies were performed by mixing solid sodium sulfite and/or metabisulfite with ferrous chelate nitrosyl solutions under an Ar or SF, atmosphere. Gaseous products were analyzed by gas chromatography and by mass spectrometry. A Varian Aerograph 700 gas chromatograph with a Porapak-Q column and a thermal conductivity detector was used for the gas chromatography measurements. The gas chromatograph was calibrated with mixtures of small amounts of N2and NzO in SF,. Gas samples were also collected in evacuated bulbs and analyzed on an AEI MS12 mass spectrometer. The mass spectrometer was also calibrated with mixtures of Nz,N20, and SF,. Total ferrous ion concentrations were determined by the 1,lO-phenanthroliie method. Solutions were diluted and acidified to pH 2.5 to avoid interference from the chelates used to form the complex. Other species in solution were studied by laser Raman spectroscopy (Littlejohn and Chang, 1984). Solution samples were transferred to 1-mm-i.d.capillary tubes under Ar and sealed. Quantitative measurements of the species present could be made by the addition of a known amount of inert reference compound. One of the compounds previously observed in solution, hydroxyimidodisulfate (abbreviated as HIDS), HON(S0)3)22-,was also studied colorimetrically by oxidizing it to nitrosodisulfonate ion.
-
Results and Discussion Kinetics. The studies of the absorption decay of ferrous nitrosyl chelate complexes in the presence of sulfite and bisulfite ions (referred to collectively as Sw) were done
5 0
0
I
I
I
I
J
I
2
3
4
5
(STX I O ~ M
+ S"
Figure 1. Fe"(L)NO IO
dddt
I
rate dependence on S" concentration. I
1
1
I
PH
Figure 2. Fe"(L)NO
+ SIv rate dependence on pH.
under pseudo-first-order conditions, although the Fe":SW ratio was varied. Most runs were done at 25 "C. The majority of the work was done with NTA and EDTA complexes. Fe"(H20),N0, Fe"(cit)NO, and Fe"(1DA)NO did not bind NO as strongly as Fe"(NTA)NO and Fe"(EDTA)NO. At low pH, the apparent reaction rates of the former compounds were considerably faster than those containing NTA and EDTA. The disappearance of the Fen(L)NOcomplex was found to have first-order dependence on the SIv concentration for all conditions studied, as shown in Figure 1. The rate dependence on the ferrous nitrosyl complex concentration was more complicated. The observed rate was very nearly independent of the ferrous nitrosyl complex concentration from pH 2 to 5. No definite dependence was observed between pH 5 and 7. Above pH 7, the rate was found to have variable dependence on Fen(L)NO concentration. At nitrosyl complex concentrations of about lo4 M, the rate appeared to have second-order dependence. The dependence decreased to first order at nitrosyl complex concentrations of about M and approached zero order M. above The reaction rate was found to have no dependence on ionic strength in the pH range 6-7. A plot of dc/dt vs pH is shown in Figure 2, where c = [Fe"(NTA)NO]. A plot of d(l/c)/dt shows similar behavior. The drop-off at low pH may be due to the increasing fraction of SIv as SO2H20. The rate increase starting at pH 5 can be attributed to the increasing fraction of SIv as S032-. However, the rate should continue to increase until pH 9. The decrease in the rate above pH 7 may be due to the stabilization by the
12 Ind. Eng. Chem. Res., Vol. 29, No. 1. 1990
addition of a hydroxyl ion to the nitrosyl complex, making it less reactive (Griffiths and Chang, 1986). The observed rate behavior indicates the reaction is not a simple one. However, empirical rate expressions can be obtained for a range of conditions for Fe"(NTA)NO and Fe"(EDTA)NO in the form of -
d[Fe"(L)NO] - = k,[HSO,-] df
F E I I I I INTRINO + rS03-
---t
+
4
[Fe"(L)NOI2 kb'S032-1
+
1 + a[Fe"(L)NO] + O[Fe"(L)N0l2
where cy = 1 x lo5 and fl = 1 x IO7. For L = NTA at 25 "C and pH 3-8, k , = 1.2 x s-l and kb = 1.4 X lo4M-l s-l for Fen(NTA)NOconcentrations M and for SIv concentrations bebetween and tween and lo-' M. For L = EDTA a t 25 "C and pH s-l and k b = 1.9 X lo4 M-' ss1. For L 4-8, k, = 5.6 X = EDTA at 55 "C and pH 4-8, k , = 9.0 X and kb = 6.45 x IO4 M-' s-l. These values were established for M Fe"(EDTA)NO concentrations between and and SIv concentrations between and lo-' M. At low Fe"(L)NO concentrations and at moderate pH conditions where the reaction with dominates, the expression can be approximated as d[Fe"(L)NO] -= k,[S0,2'-][Fe11(L)N0]2 dt On integration, this can be expressed as 1 1 - = k,t [S032-] [Fe"(L)NO], [Fe"(L)NO], which is similar in form to the data in Figure 1. It should be noted that the kinetics discussed here are for the initial reaction of Fe"(L)NO with SIv when there is essentially no Fe"' present. It is difficult to compare these results with those of Sada et al. (1987a),because they studied the reaction over a much longer time scale (on the order of hours rather than seconds). Their results indicate that the behavior of the initial reaction is different from that occurring at later times. They obtain first-order dependence on both SIv and Fe"(L)NO and inverse firstorder dependence on Fe"'(L). The activation energy that they obtain (7.3 kcal/mol) is considerably larger than that derived from the empirical rate constants listed here (2.0 kcal/mol). We did not study the effect of Fe"' on the kinetics of the initial reaction. Reaction Products. As the reaction between the ferrous nitrosyl complex and Sw proceeds, the solution color changes from dark green to orange or red, suggesting the presence of Fe"'. This was confirmed by measuring the total Fen in the solution by the 1,lO-phenantholinemethod. At pH 8, the fraction of Fe" oxidized increased linearly from 0.5 to 0.7 as the [SN]/[Fen(NTA)NO]ratio increased from 0.5 to 200. Runs at pH 4 showed less Fe" oxidation than similar runs at pH 8. The rate of disappearance of Fe" was slower than the rate of disappearance of the nitrosyl complex, suggesting separate reactions are causing the two processes. An increase in pH was also observed as the reaction proceeds, indicating that H+ or OH- ions are involved in the chemistry. The majority of SIv consumed by the reaction is converted to SO:-, as shown by the growth of the 981-cm-' S042-peak in the Raman spectrum. Another product is dithionate ion, S202-,which has also been reported by other workers (Sada et al., 1984). We have observed the 1092-cm-' dithionate peak as a shoulder on the larger 1084-cm-l HIDS peak in Raman spectra of reaction mix-
€50.00
1158. e0
900. 00
Raman shift Icm-!) Figure 3. Raman spectra of reaction mixture a t pH 4 (upper trace) and of S20e2-and HON(S03)?- (lower traces).
1:) 0.4
I
0.21 I I
I
I
I
I
1
I
3
5
7
9
I1
i
1
1315
[SO~]/[Fe(U)(NTA)NO]
p
-
IO.
c
iN
06
Z
\
10
= IO
o.2- [so;]/[Fe(IU(NTA)NO]
: LL
I
I
I
-3
1.0
2
I
4
J
I
-2 log [Fe(II)(NTA) NO]
ZN
-I
I
8
6
IO
PH
Figure 4. (a, upper) Effect of [S0,2-]/[Feu(NTA)NO] on the production of N 2 0 and N2. (b, middle) Effect of Fe"(NTA)NO concentration on the production of N 2 0 and N2. (c, lower) Effect of pH on the production of N20and N2.
tures at low pH, as shown in Figure 3. The peak at 875 cm-' is &,BO,, used as a reference. Spectra of HIDS and S2062-are shown below the spectrum of the reaction mixture. Dithionate was observed only in acidic conditions. A probable formation mechanism is the reaction of ferric ion with bisulfite ion (Huie and Peterson, 1982). Some of the SIv is also converted into nitrogen-sulfur compounds such as HIDS and related species. A large fraction of the nitrogen in the NO bound to the ferrous chelate is converted to NzO and N2,particularly at high pH. Figure 4a shows the effect of the [S032-]/ [Fen(NTA)NO]ratio on conversion of NO to N20 and N,.
Ind. Eng. Chem. Res., Vol. 29, No. 1, 1990 13
7
P
Fraction Present
I
0 ~ . - - - - - ~
2
'
8
6
4
10
12
PH Figure 5. Effect of pH on the production of HIDS (solid c k l e s ) , N20 (open circles), and N2 (solid squares). The sum of the three products is shown as open squares.
N 2 0 production is unaffected by this ratio above a value of 2:1, whereas N2 production increases up to a ratio of about 81. Figure 4b illustrates the effect of Fen(NTA)NO concentration on N20 and N2 production. Figure 4c illustrates the effect of pH on conversion of NO to N20 and N2 a t a [S032-]/[Fe11(NTA)NO]ratio of 10. Production of both N20 and N2 increases with pH. The threshold for N2 generation is pH 6. N,O and N2 build up somewhat more slowly than FeII', presumably because of the delay in diffusing out of the solution. Some of the remaining nitrogen from NO is converted into HIDS. The fraction of NO bound to ferrous chelate converted to HIDS, N20, and N2 is shown in Figure 5, along with the sum of these products. HIDS reacts with HSO; to produce several other nitrogen-sulfur compounds (Chang et al., 1982). The amount of conversion will depend on the time after mixing, as well as the HS03- concentration. Nitrogen-sulfur compounds such as hydroxysulfamic acid, nitridotrisulfate, and imidodisulfate are not as easily determined as HIDS, and the interconversion of the nitrogen-sulfur compounds adds uncertainty to the measurement process. The reacted nitric oxide not observed as HIDS, N20, or N2 is believed to be primarily in the form of other nitrogen-sulfur compounds. Mechanism. The observed kinetic behavior and number of products indicate that the reaction is mechanistically complicated. The mechanism can be divided into three components. The proposed first series of reaction steps involve the conversion of NO to N,O and Fe" to Fe"'.
-
Fe"(L)NO
+ S032-= Fe11(L)(NOS032-)
-
(1)
Fe11(L)(NOS032-)+ Fe"(L)NO 2FeYL) + (NO)2S032-(2) 2Fe"(L)
-
+ (NO)2S032-
2Fe1"(L)
(3)
N2022- + H2O -* N2O
-
(NO)2S032-
+ N2022-+ S032-
+ 20HN 2 0 + SO,2-
(4) (5)
These reactions are capable of explaining the observed chemistry at pH 1 8, where SO:- is the dominant Srvform. They are compatible with the reaction kinetics observed a t these conditions. By use of the steady-state approximation for Fe"(L)(NOSO,2-), a rate expression can be derived from these reactions. The rate exhibits secondorder dependence on Fe"(L)NO at low nitrosyl concentrations where k2(Fe"(L)NO) < k-, and first-order dependence a t higher nitrosyl concentrations where k2(Fen(L)NO) > k-,. The Fen(L)(NOSO& intermediate has been mentioned in Bonner and Pearsall (1982) and Buchholz and Powell (1963). The chemistry of (NO),S0,2-
has been studied by Powell and co-workers (Nunes and Powell, 1970; Ackermann and Powell, 1967). We have observed that (NO)2S032-reacts rapidly with Fe" to produce N20 and Fe'q Reactions involving unbound NO are not likely to be important because it is present only in low concentrations. The chemistry a t low pH conditions is somewhat more complicated than at high pH conditions. In addition to the reactions listed above, there are reactions producing nitrogen-sulfur compounds and dithionate ion. The mechanism below is proposed for the production of HIDS and would operate over a pH range of about 3-8. Other nitrogen-sulfur compounds can be produced from HIDS (Chang et al., 1982).
+ HS03- = Fe11(L)NOHS03Fe"(L)NOHS03- + Fe"(L)NO 2Fe"L + NOS03- + HNO HS03- + NOS03- HON(S03)22H2N202 HNO + HNO Fe"(L)NO
H2N202
-
-
+
NZO
+ HzO
(6) (7) (8)
(9) (10)
The rate of loss of the ferrous nitrosyl complex is slower at low pH, and no definite dependence on the concentration of Fe"(L)NO is observed. The oxidation state of nitrogen in NO is +2, while that in HIDS is +3. In order for one molecule of NO to be converted to HIDS, another molecule of NO has to be reduced to conserve the charge in the system. All other chemical species such as sulfite, EDTA, and Fe" are not reduced. Reactions 7 and 8 are, therefore, proposed to account for the production mechanism of HIDS. Ackermann and Powell (1967) mention the possibility of NOS03- acting as an intermediate in reactions of nitrogen-sulfur compounds. Reaction 8 is identical with that proposed for the preparation of HIDS from bisulfite ion and nitrous acid (Oblath et al., 1982). We believe that these reactions are the most feasible mechanism for the production of nitrogen-sulfur compounds. Since significant amounts of SO?- will be present in solutions down to about pH 4, reactions 1-5 will occur along with reactions 6-10 at low pH conditions. Thus, the observed kinetics cannot be easily compared with a simple rate expression. Hyponitrous acid is believed to be formed from the nitrosyl radical, HNO (Latimer and Hildebrand, 1959). The hydrolysis of hyponitrous acid has been studied by Buchholz and Powell (1963). S202- could be produced from the reaction of Fe"'(EDTA) with HS03-. Sato et al. (1978) have studied this reaction and suggested the following mechanism: Fe"'(EDTA)
+ HS03- = FeS03++ H+ + EDTA FeS03+ = Fe"
+ SO3-
+ EDTA = Fe"(EDTA) SO3- + H+ HS03 HS03 + HS03 SzOs2-+ 2H+ Fe"
-
(11)
(12) (13) (14) (15)
The three parts of the mechanism explain the observed kinetic behavior and the observed products, with the exception of N2. N2 is a minor reaction product, except at very high pH conditions. It has also been observed as a product of the autodecomposition of Fe"(L)NO at high pHs (Bonner and Pearsall, 1982). While no mechanism has been proposed for this process, it may be responsible
Ind. Eng. Chem. Res. 1990, 29, 14-21
14
for the formation of N2 in the Fen(L)NO + Srvsystem also.
Acknowledgment This work was supported by the Assistant Secretary for Fossil Energy, Office of Coal Utilization Systems, U.S. Department of Energy under Contract DE-ACOS76SF00098 through the Pittsburgh Energy Technology Center, Pittsburg, PA.
Literature Cited Ackermann, M. N.; Powell, R. E. Air Oxidation of HydroxylamineN-Sulfonate. Inorg. Chem. 1967, 6 , 1718-1720. Bonner, F. T.; Pearsall, K. A. Aqueous Nitrosyliron(I1) Chemistry. 1. Reduction of Nitrite and Nitric Oxide by Iron(I1) and (Trioxodinitrato)iron(II) in Acetate Buffer. Intermediacy of Nitrosyl Hydride. Inorg. Chem. 1982, 21, 1973-1978. Buchholz, J. R.; Powell, R. E. The Decomposition of Hyponitrous Acid. I. The Non-chain Reaction. J . Am. Chem. Soc. 1963, 85. 509-5 11. Chang, S. G.; Littlejohn, D.; Lin, N. H. Flue Gas Desulfurization; Hudson. J. L.; Rochelle, G. T., Eds.; ACS Symposium Series 188; American Chemical Society: Washington, DC, 1982; pp 127-152. Chang, S. G.; Littlejohn, D.; Lynn, S. Effects of Metal Chelates on Wet Flue Gas Scrubbing Chemistry. Enuiron. Sci. Technol. 1983. 17, 649-653. Griffiths, E. A,; Chang, S. G. Effect of Citrate Buffer Additive on the Absorption of NO by Solutions of Ferrous Chelates. Ind. Eng. Chem. Fundam. 1986, 25, 356-359. Huie, R. E.; Peterson, N. C. Trace Atmospheric Constituents; Schwartz, S.E., Ed. Advances in Environmental Science Techfiology 12; Wiley: New York, 1982; pp 117-146. Latimer, W. M.; Hildebrand, J. H. Reference Book of Inorganic Chemistry; Macmillan Co.: New York, 1959; pp 204-205. Littlejohn, D.; Chang, S. G. Kinetic Study of Ferrous Nitrosyl Complexes. J . Phys. Chem. 1982,86, 537-540. Littlejohn, D.; Chang, S. G. Identification of Species in a Wet Flue Gas Desulfurization and Denitrification System by Laser Raman
Spectroscopy. Enuiron. Sci. Technol. 1984, 18, 305-310. Miyadera, T.; Hiramine, S.; Shimada, Y.; Sugimoto, Y.; Teranishi, H. Absorption of Dilute Nitric Monoxide in Aqueous Solutions of Fe(I1)-EDTA and Mixed Solutions of Fe(I1)-EDTA and Na2SOR. J . Chem. Eng. Jpn. 1978, 11, 450-457. Narita, E.; Sato, T.; Shioya, T.; Ikari, M.; Okabe, T. Formation of Hydroxylamidobis(su1fate) Ion by the Absorption of NO into Aqueous Solutions of Na2S03Containing F e h d t a Complex. Ind. Eng. Chem. Prod. Res. Deu. 1984,23, 262-265. Nunes. T. L.: Powell. R. E. Kinetics of the Reaction of Nitric Oxide with Sulfite. Inorg. Chem. 1970, 9, 1916-1917. Oblath, S. B.; Markowitz, S. S.; Novakov, T.; Chang, S. G. Kinetics of the Initial Reaction of Nitrite Ion in Bisulfite Solutions. J . Phys. Chem. 1982,86, 4853-4857. Sada, E.; Kumazawa, H.; Kudo, I.; Kondo, T. Individual and Sim&aneous Absorption of Dilute NO and SO2 in Aqueous Slurries of MgS03 with Fe"-EDTA. Ind. Eng. Chem. Process Des. Deu. 1980, 19, 377-382. Sada, E.; Kumazawa, H.; Takada, Y. Chemical Reactions Accompanying Absorption of NO into Aqueous Mixed Solutions of Fe'L edta and Na2S03. Ind. Eng. Chem. Fundam. 1984, 23, 60-64. Sada, E.; Kumazawa, H.; Hikosaka, H. A Kinetic Study of Absorption of NO into Aqueous Solutions of Na2S03with Added Fe'L edta Chelate. Ind. Eng. Chem. Fundam. 1987a, 26, 386-390. Sada, E.; Kumazawa, H.; Machida, H. Absorption of Dilute NO into Aqueous Solutions of Na2S03with Added FeI'NTA and Reduction Kinetics of FeII'NTA by Na2S03. Ind. Eng. Chem. Fundam. 1987b, 26, 2016-2019. Sato, T.; Simizu, T.; Okabe, T. Studies of the Formation and Decomposition of Dithionate. 11. The Formation of Dithionate by the Reaction of Iron(II1)-edta with Sodium Sulfite. Nippon Kagaku K a ~ h 1978, i 361-366. Teramoto, M.; Hiramine, S.; Shimada, Y.; Sugimoto, Y.; Teranishi, H. Absorption of Dilute Nitric Monoxide in Aqueous Solutions of Fe(I1)-EDTA and Mixed Solutions of Fe(I1)-EDTA and Na2S03. J . Chem. Eng. Jpn. 1978, 11, 450-457. Receiued for review April 5, 1989 Accepted September 6, 1989
Physicochemical Aspects of the Leaching of Molybdenum from Co-Mo/y-A120s Hydrodesulfurization Catalyst Waste Using DMSO-SO2 Mixed Solvent Prafulla R. Raisoni and Sharad G. Dixit* Department of Chemical Technology, University of Bombay, Matunga, Bombay 400 019, India
Molybdenum can be selectively leached from a spent HDS catalyst using a solution of SOz in DMSO. The spent catalyst contains CoMoO,, CoA1204,Cogs8,and MoSz, and Mo(IV), MOW),and Mo(V1) are present after calcination. The solution contains S032-and S2052-.SEM reveals the formation of a layer of product on the surface of reacting particles. The rate measurements are consistent with control by diffusion through this layer. The Co-Mo/y-Al20, catalyst consisting of molybdenum deposited on alumina together with a promoter, cobalt, is currently being used for the hydrodesulfurization (HDS) process in the petroleum refining. A large quantity of catalyst waste is generated. The spent Co-Mo/Al,O, HDS catalyst is a rich secondary source of cobalt and molybdenum. Therefore, several efforts have been made to recover the valuable metals from the waste catalyst (Berkesi et al., 1985; Georgescu et al., 1980; Haehn et al., 1985; Vicol et al., 1986). In recent times, solutions containing sulfur dioxide have attracted the attention of researchers as lixiviants to extract metallic values. Dimethyl sulfoxidesulfur dioxde (DMSO-SOZ), dimethylformamide-sulfur dioxide (DMF-S02), and acetonitrile-sulfur dioxide (CH3CN-S02) have been shown to be very good leaching *To whom all correspondence should be addressed.
0888-5885/90/2629-0014$02.50/0
agents (Gill et ai., 1984). Also, aqueous solutions of sulfur dioxide have been studied as leachants (Miller and Wan, 1983; Khalafalla and Pohlman, 1981). We have investigated the leaching of cobalt and molybdenum from calcined spent catalyst using aqueous SOz solutions (Raisoni and Dixit, 1988a,b). Both cobalt and molybdenum could be easily leached out, and no selectivity was observed. The DMSO-SO, solvent system has also been studied with respect to Co-Mo/Al,03 spent catalyst, since it offers the possibility of selective leaching. Furthermore, DMSO is not an expensive solvent, and it can be easily recovered and reused. In the present case, the use of DMSO-SOZ may be economically justified since high value products are obtained. 'IJntil now, the dissolution chemistry, reaction mechanism, and kinetic factors affecting the rate of leaching with the DMSO-SO, mixed solvent have not been studied. C 1990 American Chemical Society
