Reaction of Sulfur Oxides with Alumina and Platinum/Alumina Jack C. Summers Physical Chemistry Department, General Motors Research Laboratories, Warren, Mich. 48090
~~~~~
~~
At high temperatures (>500 "C) very little SO2 is adsorbed on A1203 in the absence of 0 2 and noble metals. However, in the presence of 0 2 significant quantities of SOz, above that occurring independently in the gas phase, are oxidized to SO3 and subsequently stored on the AlzO3. Platinum greatly enhances the rate of SO2storage on Alz03.In the absence of 0 2 , Pt appears to catalyze the disproportionation of SOz. With 0 2 present in the feedstream, Pt catalyzes the oxidation of SOz. SO2 and SO3 adsorb on different A1203 sites. The adsorption of SO:] does not affect the rate of SO2 adsorption. These findings are not surprising since SOz is a Lewis base and SO3 is a Lewis acid. The surface coverage of A1203 by sulfur ranged from a few tenths to a few percent of a monolayer depending on the reaction conditions employed. Gasoline contains small quantities of organosulfur compounds, e.g., thiophene. These compounds are oxidized to SO2 in the automobile engine. If the automobile is equipped with a catalytic converter, some of the SO2 may be converted to SO3. The SO3 readily reacts with water vapor in the exhaust to form H2S04, small quantities of which may then be emitted from the converter as an aerosol ( I , 2 ) . With the discovery that catalyst-equipped vehicles emit traces of H2SOI, considerable effort was expended in characterizing the sulfur emission products ( 3 )and in determining under what operating conditions they were formed (4-6). I t was soon realized that the formation and emission of H2S04 from catalyst-equipped vehicles were complicated by the storage of sulfur species on the catalyst support. The alumina support moderates the quantity of Has04 that is emitted, since large quantities of H2S04 can be stored when the vehicle is running under oxidizing conditions and released as either , 9 0 2 or H2S under certain reducing conditions (7). In spite of the intensive efforts to understand the mechanisms of SO2 oxidation and storage on noble metal catalysts, much still remains obscure. For example, the role of the bare support in SO2 storage has not been fully elucidated, nor has the effect of noble metals on SO2 storage in the absence of 0 2 . In this study we try to answer these questions by studying the reactions of SO2 and SOz/O2 mixtures over a bare A1203support and Pt/A1203.
Experimental The experiments described in this report were conducted in a reactor system constructed so that the inlet and effluent gas streams did not come into contact with metal: the gas lines were all made of polypropylene tubing up to the SO2 analyzer, and the reactor was made of quartz. The flow rates of the component gases were controlled by rotameters and differential flow controllers. The gases were mixed in a gas blending manifold and then passed into the reactor. In those experiments in which water was added to the feed, a variable speed metering pump was used to control the water injection rate. The water was pumped through a capillary tube into the reactor, and dripped onto preheated silicon carbide (Sic) pellets (20 cm3). The silicon carbide was placed on top of the catalyst sample. I t served to vaporize the water and to mix the water vapor with the SOZ-containing feed gas. Any H2S04 (SO3)made during the course of an experiment
was condensed in a Goksoyr-Ross coil (8) which was placed after the reactor. This was done not for the purpose of sulfate analysis but to protect the analytical train from corrosive H2SO4. The Goksoyr-Ross coil consisted of a helical condensing coil a t the end of which was placed a medium pore (10-15 km) glass frit filter. The condensing coil and the glass frit filter were enclosed in a water jacket which was maintained a t 68 "C, a temperature which is below the dew point of sulfuric acid. Hence, the sulfuric acid condenses in the coil, while the other gases pass through. The SO2 content of the inlet and the effluent gases were measured by a pulsed fluorescence analyzer. The inlet SO2 level was monitored periodically while the effluent SO2 levels were measured continuously. The 0 2 levels were measured by a paramagnetic oxygen analyzer. Ahead of the analytical train was a permeation dryer that was used to remove the water from the effluent stream. For the SO2 adsorption experiments the inlet gas was passed through 10 cm3 of either A1203 or a Pt catalyst supported on the same type of A1203. The properties of these materials are listed in Table I. The reaction time for the experiments was 300 min. The gas hourly space velocity was -32 000 h-' (STP).The temperature was measured and controlled by a thermocouple placed a t approximately the center of the catalyst bed. A prepurified Nz (99.998% Nz) and a 1.07 vol % SOz/Nz blend were used in the experiments. The SO2 blend contained less than 30 ppm (vol) of 0 2 . The sulfur pickup on the A1203 and on the Pt/A1203 catalyst was measured a t the conclusion of each experiment. The sulfur stored on the catalysts was measured by a combustion method. The samples were heated in 0 2 to convert the stored sulfur to SOz, which was subsequently measured by a LECO Model 32 IR analyzer. Two SO3 saturation experiments were run in which the A1203 and the Pt/A1203 were presaturated with SO3 prior to adsorbing SO2. The presaturation by SOBwas accomplished by first passing a feed stream containing 200 ppm (vol) of SO2 and 5.1 vol % 0 2 a t 400 "C and a space velocity of -10 000 h-l (STP) over a 1.0 wt % Pt/A1203 catalyst (20 cm3) to convert the SO2 to SO3 ( 4 ) .The 1.0 wt % Pt/A1203 catalyst was separated from the A1203 and Pt/A1203 samples (15 cm3) by Sic (30 cm3). After 750 min, the SO2 was turned off and the SO3 presaturated samples were heated to 630 "C for 120 min in order to desorb any SO2 that might have been adsorbed on the support during the exposure to SO3. SO2 is desorbed from A1203a t 600 "C, but SO3 is not desorbed a t temperatures less than 800 "C (IO). The effect of 0 2 on the nature of the stored sulfur species was studied by passing either a S02/N2 or a SOn/air mixture (1:l volume) over 75 cm3 of A1203and 0.1 wt % Pt/A1203 a t 500 "C. Nitrogen was first passed (1h) over the materials prior to exposure to the SO2 in order to remove adsorbed water and oxygen. The SO2 mixture was then passed over the A1203 or the Pt/A1203 for 1 h. Finally, the sample was cooled to room temperature in flowing nitrogen. The reactor system consisted of two quartz tubes connected in series and heated by tube furnaces. The first tube served as a preheater (to 500 "C) and the second contained the catalyst. Some samples were analyzed by ESCA in order to identify the oxidation state(s) of the adsorbed sulfur species. An
0013-936X/79/0913-0321$01 .OO/O @ 1979 American Chemical Society
Volume 13, Number 3, March 1979
321
r
24
Inlet
so2
/
540'C
occ 0:c
20
-
5 16
p'
-f
5 12 ru c
-
z
I
f f 8
iF O l O l
4
0
60
120
180
\
300
240
R e a c f ~ o nT ~ m efmanut?sI
Figure 1. Effect of temperature on the adsorption of SOa on A1203 in the absence of 0 2
Oo2
t
I 0 6 x 10131*
I
I
200
300
I
I
400
A1203
BET surface area, m2/g total pore volume, cm3/g
PVA1203
96 0.595
600
lcCl
Figure 2. Effect of temperature on SOp accumulation on Al2O3. The asterisk indicates atoms of sulfur/cm2 BET area
91
-
0.055
Pt distribution, ,urna
edge
Pt dispersion, %
73.8 0.05
+
0.085
+
0.012
a Determined by the SnCI2 method (9). tion.
10131.
I
+
wt % R
wt % Fe wt % Na wt % Ca
(0 2
500
Ternprrature
Table 1. Catalyst Properties
\
62 & 30
+ Determined by CO chemisorp-
ESCA-3 instrument was employed for all the measurements. Because the intensity of the sulfur peaks was small, it was necessary to do signal averaging. The sulfur 2P peak was used to identify the chemical shift of the sulfur species. Sulfur was present as either S4+(SO2 or S032-) or S6+ (SO3or S042-).The S4+ species were identified by a low-energy shoulder (Pl/p state) on the S6+ band (P3/2 state). The average binding energy of the S4+ species was 167.4 eV and the average binding energy of the S6+ species was 169.0 eV.
Results and Discussion SO2 Adsorption on A1203 Hammerle and Truex ( 5 ) re-
ported that in automobile exhaust oxidation catalysis, so3 is primarily responsible for sulfur storage. However, other investigators have found that significant quantities of SO2 can also be stored on A1203 (10, 11).As a first step toward understanding the role of the support in sulfur storage, we studied the adsorption of SO2 on the bare A1203. The study of SO2 adsorption on A1203was performed by passing a feedstream containing about 22 ppm (vol) of SO2 and varying quantities of 0 2 (0.0-5.1 vol %) in N2 over the A1203. The rate of SO2 removal from the inlet stream was measured as a function of reaction time. SO2 removal can occur by either the adsorption of SO2 on the A1203,or by the oxidation of SO2 and SO3. Characteristically, the initial adsorption of SO2 was rapid followed by a more gradual uptake (5,6). Temperature was the first variable studied. The rate of SO2 adsorption and the quantity of sulfur stored on the A1203are strong functions of temperature ( 4 , 5, 10): both decrease markedly with increasing temperature (Figures 1and 2). The 322 Environmental Science & Technology
4t
0
0
I
1
I
60
120
180
R < I on
Figure 3. Effect of "C
02
1
I *
I TI n u i i
I 240
I 300
s
level on the adsorption of SO2 on AI203 at 540
decline in SO2 uptake with increasing temperature is probably associated with an increase in the rate of desorption with increasing temperature. Chang (IO)has shown by IR that SO2 adsorbs on A1203 in the absence of 0 2 to give two distinct species. The first is chemisorbed SOz. In vacuo this species is almost completely desorbed from A1203 a t temperatures as low as 200 OC. The second is a surface sulfite ( S 0 3 2 - )species that is chemically stable up to about 600 "C. Since all of our experiments were done at temperatures of 200 "C or above, it is presumably the surface sulfite species that constitutes the bulk of sulfur stored on the A1203. ESCA analysis of A1203exposed to SO2 revealed the presence of S4+species. When heated in vacuo a t 500 "C, the S4+species can be desorbed from the A1203. The surface coverage by sulfur decreased from 2.9 X 1013 to 0.20 X atoms of sulfur/cm2 BET area as the temperature was increased from 200 to 560 "C. This coverage represents from 2.9 to 0.2% of a monolayer coverage assuming an alumina site density of 1 X 1015 sites/cm2 BET area (12). The effect of 0 2 level on SO2 adsorption was investigated a t 540 "C. One-tenth volume percent 0 2 had little effect on either the rate of SO2 uptake or the quantity of sulfur stored on the A1203 (Figures 3 and 4). Hammerle and Truex ( 5 )reported a similar finding. However, while Hammerle and Truex found no enhancement in sulfur storage at 0 2 levels less than
"It
SI
0
1
1
1
1
20
40
60
80
0 2 ~n the Feed l v o l %I
Figure 4.
Figure 6.
OC).
O C
Effect of O2 level on sulfur storage on AI2O3 and R/A1203(540 The asterisk indicates atoms of sulfur/cm2BET area
1
1
1
1
1
1
1
I
1
1
1
1
100 120 140 160 180 200 220 240 260 280 300 320 React,on Time (minutes)
Adsorption of SO2 on fresh and SO3 pretreated A1203 at 540
U
I 70
t
Lewis Base Lewis Acid Molecular structures of SO2 and SO3
Figure 5.
1 vol %, we found that the level of sulfur stored was nearly a linear function of the inlet 0 2 level from 0.1 to 5.1 vol % 0 2 (Figure 4). Since the support used in this study contains iron impurities (see Table I) and since iron oxides are known to catalyze the oxidation of SO2 (13), it may be that the iron is responsible for the increased adsorption of sulfur oxides in the presence of 0 2 . ESCA analysis of A1203 exposed to a S02/02 mixture revealed only the presence of S6+ species which could not be desorbed upon heating to 500 "C in vacuo. Presumably the SOz was converted to S6+ over the iron impurities contained in the A1203 I t is desirable to calculate the extent of conversion of SO2 to SO3 during the course of the reaction. Such a calculation can be done if two assumptions are made: (a) The difference in the rate of SO2 removal from the feed when 0 2 is not present and when 0 2 is present is due to SO2 oxidation to SO3. (b) SO2 and SO3 adsorb on the A1203independently of each other. The second assumption is based on the difference in the chemical nature of the sulfur oxides (Figure 5). SO2 is a Lewis base with a lone pair of electrons available to bond with a Lewis acid site (12). SOB,on the other hand, is a Lewis acid with an empty p orbital available for bonding with a Lewis base site. Thus, the electron structure of the adsorbing sulfur oxide should have a major influence on the nature of its adsorption site. To test the validity of the second assumption, the A1203 and the Pt/A1203 catalyst were presaturated with SO3 and then exposed to a feedstream containing 22 ppm (vol) of SO2 in Nz. The SO2 uptake curves and quantities of sulfur adsorbed (as S02) on the A1203 and Pt/A1203 catalyst were virtually identical with those obtained for the fresh materials (Figure 6). Thus, the SO2 adsorption sites (both for chemisorbed SO2 and SOB*-)were not affected by adsorbed SOB,and the assumption that SO:! and SO3 adsorb independently on the A1203is verified. Having established the validity of this key assumption, we can now calculate the conversion of SO2 to SO3 a t a given 0 2
Inlet
Figure 7.
O2 Content
ivol %I
Effect of O2 level on SO2 conversion (300 min) at 540 O C
level from the following equation:
where [S02],,letis the inlet SO2 concentration, [SO,] is the SO2 in the effluent with no 0 2 in the feed, and [S02]o2is the SO2 in the effluent a t a given 0 2 level. The conversion of SO2 to SO3 a t t = 300 min was calculated for each 0 2 level studied (Figure 7 ) .As can be seen from Figure 7, this conversion can be significant over bare A1203 a t higher 0 2 levels. For example, a t 5.1 vol % 0 2 , 14.5%of the SO2 is converted to SO3. A t 540 "C and with 5.1 vol YO0 2 in the feed the relative quantities of adsorbed SO2 and SO3 can be calculated since we know that these oxides adsorb independently of one another and adsorb on different sites. The ratio of adsorbed SO3 Volume 13, Number 3, March 1979
323
24k/ ' -- - Inlet SO2
24
u
20
p
1E
P >
E
k
7 12
x
Y N
$ E
Y 4
I
0 0
60
I
I
240
300
I
I
120 180 Reaction Time (minutes1
Figure 8. SO2 removal over A1203 and Pt/AI303 at 540 " C
in the absence of O2
-16 P >
E
-
12
pi
-
I
m
$ 8
4
0 0
60
120 Rt3LIlO"
TI,,,?
180 (,"#""lesi
240
300
Figure 9. Effect of O2level on the adsorption of SOz on Pt/A1203 at 540 OC
to SO2 under these conditions as calculated from the data in Figure 4 is approximately 10. SO2 Adsorption on Pt/A1203. The adsorption of SO2 on Pt/A1203 was studied at 540 "C in the absence of 0 2 . The general characteristics of this adsorption experiment were similar to those observed for A1203 (Figure 3). However, the rate of SO2 disappearance over Pt/A1203 was significantly greater than that observed for the A1203(Figure 8). This observation is consistent with the finding that larger quantities of sulfur were stored on Pt/A1203 than were stored on the bare A1203 (Figure 4):in the absence of 02, Pt/A1203 stores approximately 10 times as much sulfur as the bare A1203after 300 min of reaction. ESCA reveals the presence of both S4+and S6+species for a Pt/A1203 catalyst exposed to SO2. Like the A1203sample that was exposed to SO2, the S4+species can be removed upon heating at 500 "C in vacuo. It is apparent that Pt catalyzes SO2 adsorption. Chang has studied the adsorption of SO2 on A1203and Pt/A1203 by infrared spectroscopy ( 1 4 ) .He has observed that at 400 "C in the absence of 02, SO2 adsorbs on Pt/A1203 but not A1203to form surface sulfate species. He has postulated that Pt catalyzes the disproportionation of SO2 to yield the surface sulfate species. In a separate experiment, it was shown that Pd also catalyzes the rate of adsorption on S02. The significance of this finding is obvious: by studying just SO? uptake on the corresponding bare A1203support, one fails to get a real measure of how much sulfur may be adsorbed on Pt/A1203 at elevated temperatures in the absence of 0 2 . 324
Environmental Science & Technology
C
I 60
1
I
120 180 Reaction Time (minutes1
I
J
240
300
Figure 10. Effect of temperature on SOn removal by Pt/A1203 at 5.1 % O2and 540 " C
Effect of 0 2 on SO2 Storage on Pt/A1203. Next, the effects of 0 2 level on SO2 removal and sulfur storage were investigated. In accord with the kinetics of SO2 oxidation ( 1 5 ) , SO2 removal is a sensitive function of the 0 2 level. At 540 "C the rate of SO2 removal (Figure 9), the rate of SO2 oxidation (Figure 7 ) , and the quantity of sulfur stored on the catalyst (Figure 4) increase markedly with increasing 0 2 content. Finally, the effect of temperature on the reaction of SO2 and 5.1 vol % 0 2 was studied (Figure 10).At 200 "C, the SO2 uptake curve over Pt/AlzO3 was identical with that obtained for bare A1203. This indicates that at 200 "C Pt catalyzes neither the oxidation nor the disproportionation of S02. At 350 "C two opposing processes occur that affect SO2 removal (Figure 10): thermal desorption of SO2 and SO2 oxidation. From 200 and 350 "C there is a decrease in the quantity of SO2 adsorbed on the bare A1203 support. The rate of oxidation, however, increases from 200 to 350 "C ( 3 ) .The net result of these processes on sulfur storage after 300 min of reaction is that the rate of SO2 removal at 350 "C is slightly greater than at 200 "C (Figure 10). At 540 "C, the oxidation of SO2 clearly dominates and this results in a rate of SO2 removal that is much greater than that observed at 350 "C. Effect of H20 on SO2 Removal. Because automotive exhaust contains appreciable quantities of H2O (-10 to 13 vol %) and because, up to this point, our experiments were performed with dry feedstreams, a few experiments were done to determine the effect of water on the sulfur oxide removal reactions. In all the experiments the results were the same: water slightly inhibits sulfur storage. The uptake curves as illustrated in Figure 11 are typical of those obtained in the presence and absence of H2O. Prior to running the SO2 adsorption and oxidation experiments, it was determined that H20 (10 ~ 0 1 %had ) no effect on the gas-phase oxidation of SO2 at 540 "C and with 5.1 vol % 0 2 in the feedstream. The removal of SO2 from the feedstream by A1203 and Pt/A1203was studied at two 0 2 levels (0 and 5.1 ~ 0 1 %and ) with each of these at two H2O levels (0 and 10%vol). Under all conditions studied, H2O inhibited sulfur storage as was shown by a decrease in the rate of SO2 removal from the feedstream on both the A1203 (16) and the Pt/A1203. The absolute differences of sulfur stored on the A1203and Pt/A1203 were independent of the reaction conditions. Conclusions This study has supported previous work that has found that SO3 is the primary sulfur species stored on oxidation emission control catalysts. However, we have shown that the mechanisms of storage are more complicated than previously
Literature C i t e d (1) Pierson, W. R., Hammerle, R. H., Kummer, J. T., presented to
II 60
I
120 180 I f o n T ~ m iri i ~ n u i i . s I
I
I
240
300
Rt,c+c
Figure 11. Effect of H20 on SO2 by Pt/AI2O3 at 5.1% O2 and 540 OC
thought. Impurities in the A1203 support can catalyze SO2 oxidation, and Pt and. Pd apparently can catalyze the disproportionation of SO2 to form surface sulfate species. SO2 and SO3 adsorb on different A1203sites, and they adsorb independently of each other. Consequently, it was possible to determine the relative quantities of sulfur stored on A1203 as SO2 and as SOs. Acknowledgments
The experiments were conducted by J. Ulicny and D. Fournier.
Society of Automotive Engineers, Detroit, Mich., Feb 1974, Paper No. 740287. (2) Beltzer, M., Campion, R. J., Peterson, W. L., presented to Society of Automotive Engineers, Detroit, Mich., Feb 1974, Paper No. 740286. (3) Mikkor, M., Hammerle, R. H., Truex, T. J.,Ind. Eng. Chem. Prod. Res. Deu., 16,217 (1977). (4) Taylor, K. C., Ind. Eng. Chem. Prod. Res. Deu., 15, 264 (1976). (5) Hammerle, R. H., Truex, T. J., presented at the Division of Petroleum Chemistry, 172nd National Meeting of the American Chemical Society, San Francisco, Calif., Aug 1976, PETR-034. (6) Hammerle, R. H., Mikkor, M., presented to Society of Automotive Engineers, Detroit, Mich., Feb 1975, Paper 750097. (7) Barnes, G. J., Summers, J. C., presented to Society of Automotive Engineers, Detroit, Mich., Feb 1975, Paper 750093. (8) Goksoyr, H., Ross, K., J . Inst. Fuel, 35,177 (1962). (9) Michalko. E., U.S.Patents 3 259 454 and 3 259 589. (10) Chang, C. C., J . Catal., 53,374 (1978). (11) Deo, A. V., DallaLana, I. G., Habeood, - H. W.,J. Catal.. 21.270 (1971). (12) Peri, J. B., J. Phys. Chem., 69, 211 (1965). (13) Chun, K. C., Quon, J. E., Enuiron. Sci. Technol., 7, 532 (1973). (14) Chang, C. C., Infrared Studies of SO2 on Pt-Alumina, General Motors CorD.. Warren. Mich.. 1978. Drivate communications. (15) Olson, R: W., Schuler, R. W., Smith, J. M., Chem. Eng. Prag., 46,614 (1950). (16) Glass, R. W., Ross, R. A., Can. J. Chem., 50,2537 (1972).
Received for review M a y 17, 1978. Accepted October 6, 1978. This paper was presented at the 176th National Meeting of the American Chemical Society (Division of Colloid and Surface Chemistry), Miami Beach, Fla., Sept 10-15, 1978.
Chloroform and Chlorophenol Production by Decarboxylation of Natural Acids during Aqueous Chlorination Richard A. Larson* and Arlene L. Rockwell Stroud Water Research Center of the Academy of Natural Sciences of Philadelphia, R.D. 1, Box 512, Avondale, Pa. 19311
Naturally occurring carboxylic acids of several structural types reacted in dilute solution with aqueous hypochlorite to afford decarboxylation products. Incorporation of chlorine into the residual organic molecule occurred. Citric acid was efficiently converted to chloroform a t pH 7 by a pathway probably involving 3-ketoglutaric acid as an intermediate; in acidic or alkaline solution, yields of CHCl3 were lower. Several other enolizable keto acids (including all three isomers of resorcylic acid) were likewise precursors of CHC13 or of substances which could thermally be converted to CHC1,; yields varied widely. Two substituted benzoic acids common to natural waters, p-hydroxybenzoic acid and vanillic acid, were decarboxylated by hypochlorite with the production of chlorophenols. Although the association of chlorophenols with unpleasant odors and tastes in drinking waters was recognized many years ago, fundamental studies on the kinetics and mechanisms of aqueous chlorination reactions have been sparse until recent years. It has now become evident that treatment of drinking and wastewaters with chlorine leads to the formation of a variety of organochlorine compounds. Chloroform (CHC13), other haloforms, chlorinated phenols and phenolic acids, and chlorinated quinones, benzoic acids, and heterocyclic com0013-936X/79/0913-0325$01.00/0 @ 1979 American
pounds are only a few of the many structural types which have been identified (1-4). Much of the organically bound chlorine has not been fully characterized; a large fraction is associated with macromolecular organics (5). Chlorine dissolved in water (“aqueous chlorine”) exists principally as hypochlorous acid (HOCl) between pH 3.4 and 7.5, and as the hypochlorite anion OC1- a t higher pH values. The reactions of aqueous chlorine with organic molecules fall into three categories, addition, substitution, and oxidation. Chlorohydrins are produced by the addition of HOCl to olefinic double bonds, and chlorine can also be incorporated by substitution reactions into activated aromatic nuclei, amines, and enolizable ketones. Many oxidative reactions of hypochlorite are known, but until recently they have been neglected in discussions of water chlorination since it has been assumed that chlorine was not incorporated into organic molecules during their oxidation. There are a few references to oxidative decarboxylation in chlorination reactions in aqueous solution. The presence of chlorophenols (4, 6) and chlorinated quinones (7) in pulp bleaching liquors has been demonstrated. These compounds are largely produced through oxidative cleavage of lignin side chains (4-substituted guaiacols and catechols) by aqueous chlorine in strongly acid solution. Among the model compounds studied was p-hydroxybenzoic acid, which under these conditions was rapidly converted to 2,4,6-trichlorophenol
Chemical Society
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