Energy & Fuels 1992,6, 526-532
526
Microcalorimetric Study of Oxygen Adsorption on Catalytically Promoted Gasification Chars: Mechanistic Evidence for Alkali- and Alkaline-Earth-Metal Carbonate Catalyzed Reactions Arthur S.Gow,I11 Department of Chemical Engineering, University of New Haven, 300 Orange Avenue, West Haven, Connecticut 06516-1999
Jonathan Phillips* Department of Chemical Engineering, 120 Fenske Laboratory, Penn State University, University Park, Pennsylvania 16802-4400 Received March 4, 1992. Revised Manuscript Received May 20, 1992
A series of microcalorimetric studies have been undertaken in order to better understand the mechanism of catalytic coal char gasification by alkali-metaland alkaline-earth-metalcatalysts. The present work is a report on a study of the interaction between oxygen and coal chars promoted with either sodium or magnesium carbonate. Results suggest that the active phase of sodium is a highly dispersed metallic phase. This contrasts with earlier studies on the behavior of alkali-metalcatalysts in which it was found that potassium is in a partially oxidized state and even more highly dispersed. This indicates that not all alkali metals promote coal char gasification through the same mechanism. It was also found that magnesium did not wet the char surface and failed to adsorb a significant quantity of oxygen. This is consistent with earlier studies on the behavior of alkali-earth-metalchar catalysts. carbon surface sites using this method. It is ale0 impoesible to get quantitative data on the difference in chemical During the past decade, there has been considerable character of the various catalytic metals. Thua, it has not research on all phases of catalytic carbon gasification. been possible to determine if increased reaction rates are Studies of catalytic gasification as reviewed else~herel-~ a function of the amount of oxygen adsorbed or the have focused on the catalytic agent, method of addition, character of the adsorbing sites. pyrolysis conditions, gasifying agent, and conditions. The present calorimetric study was designed to yield However debate continues regarding the mechanism of quantitative data regarding the number and chemical gasification and identification of the catalytically active character of oxygen adsorbing (active) sites on chars made species. from bituminous coal (Pitt no. 8) to which sodium and One current model of metal enhanced coal char remagnesium carbonates have been added. This represents activity, the so-called oxygen-transfer mechanism, suggests a continuation of an earlier study of untreated, calciumthe reaction takes place in three steps: (i) adsorption onto promoted and potassium-promoted cham8 One major metal or carbon active sites, (ii) gas "activation", and (iii) finding of the present study is that there are considerable activated atom diffusion to the reaction ~ i t e . ~ ~The ~ - ' char qualitative differencesbetween the structure and chemistry reactivity for this scheme should be proportional to the of potassium carbonate versus sodium carbonate catalysts. number of active (both catalyst and carbon) sites on the surface. It is necessary to obtain accurate information Experimental Section regarding the number and nature of the active sites in Calorimeter. A Calvet-type calorimeter of novel design was order to properly evaluate the above model of reactivity used in this work. The design and operation of this calorimeter are described in detail e l s e ~ h e r e . ~ ~ "The J ~ reliability of this for catalyzed chars. Recent work8p9indicates that differtechnique is confirmed by the results of previous studies in this ential microcalorimetry can provide information regarding lab~ratory.~~~ The ' ~ .standard ~' operating procedure is simple. catalyst dispersion and the identity of the active phase. These studies showed that potassium formed from de(1)Wen, W. Y.Catal. Rev. Sci. Eng. 1980,22,1. composition of potassium carbonate was highly dispersed (2)Wood,B.J.; Sancier, K. M. Catal. Reu. Sci. Eng. 1984,26,233. (a. 50%) and that the active phase was a form of partially (3) Walker,Jr., P. L.; Shelef, M.; Anderson, R. A. Chemistry and oxidized potassium. In contrast, methods previously used Physics of Carbon;Walker, P. L., Jr., Ed.; Dekker: New York, 1968,Vol. 4,pp 287-383. to characterize gasification catalysts including O2chemi(4)McKee, D. W.; Chatterji, D. Carbon 1978,16, 53. sorption and various desorption studies for investigation (5)Veraa, M.J.; Bell, A. T. Fuel 1978,57,194. of surface active sites6Johave not provided information (6)Radovic, L. R.;Walker, Jr., P. L. Fuel Processing Technol. 1984, 8,149. about both the number and identity of active sites. For (7)McKee, D. W.;Chatterji, D. Carbon 1976,13,381. example, oxygen chemisorption has been the principle (8)Gow, A. S.; Phillips, J. J. Catal. 1991,132,388. method used for obtaining an approximate value for the (9)Gow, A. S. Ph.D. Thesis, The Pennsylvania State University, University Park, PA, 1991. total catalytic surface area in carbon-catalyst systems. Yet, (10)Featea, F.S.;Harris, P. S.; Rueben, B. G. J. Chem. SOC.,Faraday it is impossible to distinguish active catalyst sites from Trans. 1 1974,70W,2011.
Introduction
* To whom correspondence should be addressed.
(11)ONeil, M.; Lovrien, R.; Phillips, J. Rev. Sci. Instrum. 1986,56, 2312. (12)ONeil, M.; Phillips, J. J. Phys. Chem. 1987,91, 2867.
0887-0624/92/2506-0526$03.00/00 1992 American Chemical Society
Catalytically Promoted Gasification Chars Table I. Coal Char Sample Preparation Specifications" OUtgaesin$ aample added by physical mixing temp,O C 3
500
4
900
1
.I-+
1
/II
"Demineralized Pitt no. 8 bituminous. *All samples were outgassed for 3 h at a pressure of 1 X loJ Torr. The sample is exposed to a measured amount of gas. Heat generated by the adsorption of the gas flows out of the sample chamber throughthermopiles, which generate a signal proportional to the instantaneous heat flow. The integrated signal is proportional to the total heat generated. The shape of the heat peak yields qualitative information regarding the kinetics of the adsorption process. The qualitative kinetic parameter used in this study and earlier work from this lab is the normalized cooling width at half-maximum (NCWHM). This nondimemional parameter is determined by dividing the time required for the heat peak to go from ita maximum value to its half-maximum value by the time required for the same process to take place for the first gas dose on that sample. The procedure is repeated once the system has returned to equilibrium. This is done repeatedly (generally more than 20 times) until it is clear from the signal shape and intensity that very little heat is being produced and that gas is only physically adsorbing on the sample. BET a n d Polanyi-Dubinin (C02) Measurements. The BET and COz measurements were conducted in a standard glass high-vacuum system with an ultimate pressure of 1 X lo4 Torr. After char production, described below, each of these samples was exposed to the ambient laboratory atmosphere for several days before surface area measurements were conducted. The procedures used in making the surface area measurements are standard and are described in detail elsewhere.ls X-rayDiffraction. XRD characterizationof the samples was conducted in air using a Rigaku Model DMAX-IA Dfiactometer. This diffradometer employs a copper target X-ray source, curved crystal graphite monochromator, and Ti-driftsdNaI scintillation detectors. The distance between sample and detector is 185mm and a divergent slit of lo and a receiving slit of 0.3O were used. Sample Preparation. Experiments were conducted on samples made from a single batch of ground Pitt no. 8 bituminous coal received from the DOE/PSU Coal Bank. Demineralization was carried out, using the standard method, as follows: (1)6 g of coal,ground under nitrogen to pass 200 mesh, were mixed with 40 mL of 5 N HCl in a plastic beaker (125mL). The slurry was then agitated at 60 OC for 1 h. (2)The coal was fiitered, washed, and bathed in concentrated HF at 60 "C for 1h. (3) Step 2 was repeated except with HC1. (4)The coal was then filtered and continuously washed with distilled water. (5)A slurry was prepared with an excess of COS-freedistilled water and refluxed for several hours. It was necessary to repeat step 4 until no chlorine was detected in the fitrate with a silver nitrate indicator. The coal was thoroughly dried in air at 50 OC for 24 h. Altogether four samples were studied. Char samples 1 and 3 were prepared by physically mixing the demineralized parent coal with metal carbonates (NazC03,or 3M&O&g(OH)2.3H20) and then outgassing at 500 OC. Samples 2 and 4 were prepared by physically mixing the demineralized parent coal and metal carbonates and then devolatilizingat 900 OC. Relevant parameters for each sample are given in Table I. To produce a fine powdered material and avoid interparticle diffusion limitations, all of the coal underwent a three-stage treatment before measurementswere made. That is, the coal was f i t pyrolyzed, then crushed, and fiially heat treated. This sequence was chosen since high-rank, low-volatile coals form a plasticized material upon ~ o o l i n g .The ~ exact procedure is de~~
~~
~
(13)Gatte,KR.; Phillipa, J. Langmuir 1989,5, 758. (14)Gatte, R. R.;Phillipa, J. Thermochim. Acta 1989,154,13. (15)Mahajas, 0.P.; Walker,Jr., P. L. Anolytical Methods for Coal and Coal Products;Karr, C., Ed.; Academic Press: New York, 1978;Vol. 1, p 125.
0
10
20
30
40
50
1.50
60
n (micrwnoles 0 2 ads& sample)
Figure 1. Differential heat of oxygen adsorption (A)and NCWHM (m) versus the quantity of gas adsorbed for the 500 OC 3MgCO3.Mg(OH),3H20 pretreated char. Relatively little oxygen adsorbs on this sample, and in fact the total amount adsorbed is similar to that o b ~ e ~ on e dchar to which no catalyst has been added. However, the kinetics of adsorption and to some extent the observed heats of adsorption are different than that o b e e ~ e d on unloaded samples. scribed below. Note, 10% by weight of a fine powder of alkaline-earth- or alkali-metal carbonate was physically mixed into the coal before any heat treatments. Demineralized coal and powdered carbonatewere heated under flowing Nz in a tube furnace to 500 "C, held at that temperature for 1h, and then cooled gradually (ca. 0.5 h) to room temperature. This pmcess resulted in the formation of an agglomerated materid This material was crushed to a fine powder using a mortar and pestle. From this powder two samples were made for each metal carbonate catalyst. A "high-temperature" sample was prepared by heating the powder slowly (1.5h) to 900 O C in a vacuum furnace (1x lo-' Torr) holding it there for 3 h, then cooling it gradually (1.5h) to room temperature. The aample was held under vacuum for an additional 10h before being transferred under vacuum into the calorimeter sample cell. A "low-temperature" sample was prepared in a nearly identical manner; however, this sample was heated only to 500 OC. Again, catalysts were added by physically mixing metal carbonates with the demineralized coal prior to preparation. This method was wqd since prior studieszJ6have indicated that this will leave highly dispersed catalyst particles on the coal char in certain cases (e.g., alkali-metal carbonates) and poorly dispersed particles in others (e.g., alkaline-earth-metal carbonates). Some insight regarding this contrasting behavior was desired here.
Rssults Oxygen Adsorption on Magnesium (Alkaline-Earth Metal) Carbonate Promoted Chars. All of the results for the work done in this study are presented graphically in terms of regions in which a particular type of adsorption behavior (kinetics/thermdynamica)is observed. In Figure 1 the differential heats of oxygen adsorption and the NCWHM are plotted versus the amount of oxygen adsorbed on the low-temperature treated magnesium carbonate promoted char (sample 3). There is considerable scatter in the data particularly in region I. This suggests that either the numbers are not reliable or that there are competing adsorption processes, not in equilibrium, taking place. However, the same gas volumetric data used to determine the amount adsorbed were also used to plot standard isotherms (Figure 2). The amount adsorbed increased smoothly and steadily with pressure. This is one indication that the data are reliable and that the scatter is due to a competition between adsorption sites which differ in activation energy. (16)Mime,C.A.;Pabet, J. K.Am. Chem. Soc., Fuel Chem. Diu. Prepr. 1980,25,258.
Gow and Phillips
528 Energy & Fuels, Vol. 6,No. 4, 1992 90
I
I
,
I.
I
b
4
b
0 m
1E b
>
0
30
b 0 0
20
b 0
lo 0
0
0
0
0
0
,
r
0
50
100
150
200
250
n (micromoles 0 2 addg sample)
Figure 4. Differential heat of oxygen adsorption (A) and NCWHM (m) v e m the quantity of gas adsorbed for the 500 O C Na2C03pretreated Pitt no. 8 bituminous char. Relative to unloaded and alkali-earth-metal-loadedchar this sample adsorbs a great deal of oxygen. The pattern of kinetics is distinct as well.
I
14
I
. , P =A P A a
1'
Thus, it is suggested that in this case the carbonate does not contribute significantlyto oxygen adsorption and that the little adsorption which does occur takes place primarily at unblocked carbonlike active sites on the char surface. It is interesting to note that calcium carbonate addition also has little impact on the amount or nature of oxygen adsorption.8 Standard oxygen adsorption isotherms for the 900 "C magnesium carbonate promoted char also show that the amount adsorbed increases smoothly and steadily with pressure, which suggest that the data are reliable. Oxygen Adsorption on Alkali-Metal Carbonate Promoted Chars. The differential heats of oxygen adsorption and the NCHWHM are plotted in Figure 4 for the sample which was loaded with Na2C03and outgassed at 500 OC. This sample clearly adsorbed significantly more oxygen than the magnesium carbonate treated samples. The amount of oxygen chemically adsorbed was nearly 5 times as much per gram as for the magnesium carbonate loaded or unloaded char samples discussed thus far. The high heats of adsorption and the relatively large amount adsorbed suggest that the oxygen is adsorbing on reduced metal sites on this sample. This sample hae three distinct regions of adsorption. In region I there are high heats (70-100 kcal/mol adsorbed) and, after an initial increasing region, steadily declining rates of adsorption. The heats drop steadily in region I1 while the rate of adsorption increases dramatically. The low heats in region I11 clearly signifies that physical adsorption is dominant. The differential heat and rate of oxygen adsorption versus the amount adsorbed are plotted in Figure 5 for the high-temperature treated sample to which Na2CO9had been added. There are several important features to these data which should be noted. Again, a great deal of oxygen adsorbs on this sample. It is about 3 times more than the amount which adsorbs on the unpromoted8or alkalineearth metal carbonate promoted char, and about 70% of the amount which was found to adsorb on the low-temperature treated sample to which Na2CO3had been added. Second, there are also three distinct regions to the kinetics of adsorption for this sample. The first 11gas doses (region I) all occurred at nearly the same rate. The next 5 doses showed a considerable decline in the rate of adsorption. As the heat begins to fall (region 111, there is a marked increase in the rate of adsorption, until physical adsorption begins to dominate and the rate once again becomes very stable and rapid (region 111). Third, the heats of chemisorption in region I are all higher (ca. & B O kcal/mol) than
Catalytically Promoted Gasification Chars
Energy & Fuels, Val. 6, No. 4, 1992 529
Table 11. Oxygen Chemisorption Capacities of Coal Chars at Room Temperature catalytic surface area, m2/g .from AH plota from single total from dual total temp, (AH c 10 O2 uptake O2 uptake sample "C kcal/mol) isotherm isotherms MgC03A4g(OH)2-3H20promoted Pitt no. 8 bituminoua 3.4 4.6 4.2 MgC03.Mg(OH)2.3H20promoted Pitt no. 8 bituminous 6.0 7.7 7.3 Na2C03 promoted Pitt no. 8 bituminous 23.6 23.6 23.4 22.9" 15.2 17.0"
Na2C03promoted Pitt no. 8 bituminous a
21.0 15.6 16.6
20.0 15.6 16.1
Rerun.
I 1
8
* 40
y" 0
,
; 1., ' I
I
I
I
i 2
,
P"""J-
20
0 0
30
60
90
120
150
o 0
40
I
0 I
0 0
30
60
90
120
0
*- ' 150
n (micromoles 0 2 adslp sample)
Figure 5. Differential heat of oxygen adsorption (A) and NCWHM (m) versus the quantity of gas adsorbed for the 900 O C Na2C03pretreated Pitt no. 8 bituminous char. This sample adsorbs oxygen with about the same heat as the low-temperature prepared sodium loaded sample,but adsorbs about 30% less, with a distinctly different pattern to the kinetics. those found on raw char samples. There is also some scatter in these data. The structure of these data in region I is somewhat different from that of the low-temperature treated, sodium carbonate loaded sample. The values of the NCWHM for that sample increased steadily with coverage in region I and there was more scatter in the AH& values. To explain this phenomenon, nonequilibrium surface adsorption was postulated. For the high-temperature treated samples, in contrast, the NCWHM values are remarkably steady for most of region I. It is interesting to note that most of the "scatter" in the region I heata of adsorption occurs at very early coverage where the rates of adsorption actually slightly increase, and thus it is probable that nonequilibrium competitive adsorption, which can lead to scatter in the values of AH, commences very early. This is followed by an equilibrium proteas for the remainder of region I. In contrast, the heats and rates of adsorption in region I1 have the "signature" of uncovered carbonlike active sites. This suggests that the sodium does not completely cover the active carbontype sites. Again, as was the case with the 500 "C treated sodium carbonate char, the reliability of the data is confirmed by the regular behavior of the NCWHM and the adsorption isotherms, and by a repeat calorimetric experiment on another 900 "C treated sodium carbonate promoted char sample (Figure 6). Adsorption Isotherms and Catalytic Surface Area. Operation of the calorimeter also produces "apparent" equilibrium adsorption data, that is, the final pressure versus cumulative amount adsorbed is obtained for each dose (for example, see Figure 2). This information can be used to estimate catalytically active surface area which can be directly compared with calculations based on differ-
n (miuwnoles 02 adslp samle)
Figure 6. Differential heat of oxygen adsorption versus the quantity of gas adsorbed for two identical samples of the 500 O C Na&03 pretreated Pitt no. 8 bituminous char. This figure demonstrates that the values for heat and amount adsorbed determined with different samples of the same material treated in the same fashion are very similar. The NCWHM values (not shown) are virtually identical. ential heat data. Second isotherms were collected for each sample as well. That is, after the first isotherms were complete, the sample cell was outgassed and held under vacuum (1X lo4 Torr) for 10 h, and then the procedure described earlier for determining heat production spectra and isotherms was repeated. The heata produced for these second isotherms were low (less than 10 kcal/mol adsorbed) and the values were consistent with mainly physical adsorption occurring. It is not surprising that an evacuation at room temperature does not remove a significant amount of strongly chemisorbed gas. Two methods for computing the amount chemisorbed from the isotherm data were employed. First, the intercept was obtained by extending the high-pressure linear part of the curve to the zero-pressure axis. Second, the difference in the amount adsorbed between first and second isotherms at 200 Torr was measured. The values for all four samples using both of these methods are presented in Table 11. While isotherm data, which are typically the only data available, are valuable for determining the amount adsorbed, they do not provide any information regarding the possible distribution of chemisorbing sites and the strength or kinetics of adsorption at these sites. In contrast, the calorimetric heat of adsorption data provide all of this information. Integral Heat of Adsorption. The integral heat of adsorption (Table 111) is another important byproduct of calorimetric operation. The integral heat is obtained by dividing the total area under the differential heat profiie by the cumulative amount of gas chemisorbed on the sample. The integration is carried out to the point where the measured AH values fall below 10 kcalfmol adsorbed.
Gow and Phillips
630 Energy & Fuels, Vol. 6, No. 4, 1992 1.o
Table 111. Integral Heats of Oxygen Adsorption on Coal Chars sample 500 "C MgC03-Mg(OH)2.3H20promoted Pitt no. 8 bituminous 900 "C MgC03.Mg(OH)z.3H20promoted Pitt no. 8 bituminous 500 "C Na2C03promoted Pitt no. 8
Qht,kcallmol
45.7 39.6 75.3 68.3" 74.0 73.2"
bituminous 900 "C Na2C03promoted Pitt no. 8 bituminous
.-b cn
-I C
0.5
.->
L
a!
U
"Rerun. Table IV. Surface Areas of Coal Chars
BET-SA, sample 500 "C MgC03.Mg(OH2).3H20promoted Pitt no. 8 bituminous 900 "C MgC03.Mg(OH2).3H20promoted Pitt no. 8 bituminous 500 "C Na2C03promoted Pitt no. 8
bituminous 900 "C Na2C03promoted Pitt no. 8 bituminous
m2/g 53
COZ-SA, m2/g 291
I
0.0, 30
35
40
45
50
Degree (28) 21
255
10.0 2' 3
273 258" 7 1.o
"Rerun.
Heats lower than this were assumed to result from physical adsorption. BET and P-D Surface Area Estimates. Values of BET and P-D surface area for all chars made from the same coal (Pitt no. 8 bituminous) studied here and in previously published work are given in Table IV. These data exhibit several interesting characteristics. First, the Polanyi-Dubinin surface areas are greater than the BET (N2) surface areas for all chars studied. In most cases the COPsurface area is between 1and 2 orders of magnitude in excess of the corresponding BET surface area. This is indicative that most chars possess a great deal of microporosity (high COz surface area) while having a limited macroporous (feederpore) structure8Js (low BET surface area). Note that P-D surface areas of the alkali metal carbonate promoted chars are lower than those of the other chars. This suggests that these catalysts achieve and maintain good dispersion and block gas access to the micropores. The BET measurements indicate that all of the chars have little or no macropore surface area. X-rayDiffraction Analysis. It was shown in an earlier study that the diffraction pattern for the uncatalyzecl char pretreated at 500 "C has no strong peaks other than those attributable to the holder. Hence, it is likely that the demineralizationremoves much of the mineral matter from the coal and that any residual material is well dispersed. Diffraction patterns for the two 3MgCO3&fg(OH&3HZO promoted samples (not shown) indicate that low-temperature (500 "C) heat treatment is insufficient to decompose the carbonate. All diffraction peaks for this sample are consistent with the original mineral precursor 3MgC03. Mg(OH2)-3H20.In contrast, high-temperature (900 "C) heat treatment apparently decomposes the carbonate to form MgO. The diffraction patterns for the two sodium carbonate catalyzed samples following oxidation in the calorimeter and air exposure are presented in Figure 7. (Note: broad reflections from the Grafoil sample holder are found at angles of 13" and 25"17 and serve as an internal standard.) The presence of Na2C03is indicated on the 500 "C treated sample (Figure 7a). This is anticipated on thermodynamic
-
.-
5,
u)
-e
a!
L
.-
0.5
P)
L
a!
K
0.0, 30
35
40
45
Degree (28)
Figure 7. X-ray diffraction spectra of sodium carbonate loaded char: (a) Following low-temperature treatment; (b) following high-temperature treatment.
grounds, because decomposition of the carbonate at 1 X 10-4 Torr (oxygen pressure) is not predicted to occur until 650 "C. The diffraction pattern for the 900 "C sodium metal as well as treated sample (Figure 7b) is very different, apparently, all of the sodium carbonate has decomposed to form NazO and NazOz.
Discussion The data collected in this work permit a fairly precise description of the number and chemistry of the active sites of each of the samples studied. The contrast in the adsorption behavior of the different systems also correlates with previously obeerved differences in reactivity.'41B This suggests that the differentid calorimeter may be a valuable tool in predicting the suitability of various materiala aa coal char catalysts. Alkaline-Earth Metal Carbonate Promoted Chars. On the basis of the present study and prior work? it is clear that there are only subtle differences in the character of oxygen adsorption between unpromoted chars and the (18) Spiro, C. L.; McKea, D. W.; Karky, P. G.; Lamby, E. J. Fuel 198.3,
---.
63: --, I n n
(17) Wu, N. L.; Phillips, J. Surf. Sci. 1987, 184, 463.
!I
(19) Huhn, F.; Klein, J.; Juntgen, H.Fuel 1983, 62, 196. (20) Otto, K.; Bartoeiewicz, L.; Shelef, M. Fuel 1979,58, 566.
Energy & Fuels, Vol. 6, No. 4, 1992 531
Catalytically Promoted Gasification Chars
group IIA metal carbonate promoted chars. The major difference is that over a wide range of coverage there are few fluctuations in the heat of adsorption on magnesium carbonate (present work) or calcium carbonate8promoted chars. The similarities suggest that physically mixed magnesium or calcium carbonate catalysts do not create a large number of new active sites. Much of the oxygen adsorption occurs on a variety of carbon-type sites; little actually occurs on the catalyst particles. Competition between site types could explain the initial fluctuations just as was the case for the uncatalyzed chars. The conclusion that "magnesium or calcium" catalysts are not creating a significant number of new active sites is also supported by the X-ray studies. For example, it is clear that on the low-temperature MgC03*Mg(OH2)3H20 sample most of the carbonate is not fully decomposed and is present as poorly dispersed particles. The high-temperature treatment converts a great deal of the catalyst metal to alkaline-earthmetal sulfides or to poorly dispersed oxides (i.e., MgO or CaO). It is also noteworthy that the surface areas (BET and PD) are similar for unpromoted and group IIA metal promoted chars. This model of alkaline-earth metal salt behavior is in disegreement with findings from other studia21-ain which it was shown that an increase in gasification rate resulted from new active sites associated with the catalyst. However, the enhancement in reactivity was witnessed for chars to which the catalyst was applied via incipient wetness, impregnation or ion exchange. This suggests that physical mixtures of chars with alkaline-earth-metal-containing catalysts are not effective in creating additional sites. Indeed, the dynamic mobility of alkaline-earth-metal catalysts during gasification is poor in this case. This is significant from a processing standpoint where physical mixing of char and catalyst would be most practical. Alkali-Metal Carbonate Promoted Chars. Sodium carbonate evidently has a pronounced effect on the oxygen adsorption behavior of Pitt no. 8 bituminous chars. The characteristic heat of adsorption on both samples is between 80 and 100 kcal/mol O2adsorbed over a broad range of coverage, and much more oxygen adsorbs on these samples than on uncatalyzed or alkaline-earth-metal carbonate catalyzed chars. Moreover, the adsorption process appears to be kinetically controlled on the low temperature samples while it is apparently an equilibrium process on the high-temperature chars. Various mechanisms have been proposed to explain sodium salt catalyzed char ga~ i f i c a t i o n , 2 and , ~ , ~the ~ ~experimental ~ observations within this work give additional insight into which is correct. In general, mechanisms of alkali metal salt catalyzed gasification may be classified as either oxygen-transfer mechanisms or electrochemical mechanisms., An oxygen-transfer mechanism involves donation of oxygen from the reactant gas (e.g., H20 or C02)to the catalyst which is in at least a partially reduced state. Here the catalyst salt is envisioned to be. reduced by the carbonaceous substrate whereby oxygen is transferred to the carbon in activated form. Thus the entire process repmnta a redox cycle. McKee and c o - ~ o r k e r s ~ suggest ~ ~ J a carbothermic reduction scheme (consistent with oxygen transfer) such as: Na2C03+ C + O2 Na20 + 2C02 Na20 + Y2O2 Na202 Na202+ C Na20 + CO
--
(21) Hippo, E.; Walker, Jr., P. L. Fuel 1975,54, 245. (22) Jenkins, R. G.;Nandi,S.P.; Walker, Jr., P. L. Fuel 1973,52,288. (23) Dunks,G. B.; Stelman, D.; Yosim, S. J. Carbon 1980, 18, 365.
Table V. Oxidation Reactions for Na and Na20 reaction AHm, kcal/mol
---
4NaW + Oz(g) 2Na(s) + Oak)
2 Na20(s)
-199.6 -122.5 -63.0 -46.4 -3.5
NazOkd
Na(4 + O M NaOz(s) 2NazO(s) + Oz(g) 2NazOz(s) 2/aNazO(e) + O h ) 4/3 NaOz(s)
The first reaction produces the catalyst Na20 and the second and third reactions constitute the redox cycle. As a modification, it has been proposed that the following mechanism is operative in typical H20 gasification conditions: Na2C03+ 2C 2Na + 3CO 2Na + 2H20 2NaOH
--
+ CO
2NaOH
+ H,
+
Na2C03 H,
The first of these models involves Na20 and Na202(Le., in the vicinity of lo00 "C) as 'reduced" intermediates while the second scheme proposes that completely reduced sodium metal is the reduced intermediate in the cycle. Both the calorimetric and X-ray diffraction data indicate that sodium metal is present following the pyrolysis step. Thermochemical data alone for the decomposition of sodium carbonate at 500 "C suggest that little or no decomposition should occur during pyrolysis. The partial pressure of CO, predicted for the reduction to Na20 is only about 7 X Torr at 500 "C, while the partial pressure of CO predicted from the formation of reduced metallic sodium from the carbonate is 2.5 X lo4 Torr. However, two other factors may account for Na metal formation during devolitization at 500 "C. First, the equilibrium of both the proposed decomposition reactions is strongly influenced by the continuous removal of products. These reactions are pulled to the right by the constant flushing of the system with N2during pyrolysis and by keeping the total pressure low (i.e., ca. X lo4 Torr) during outgassing pretreatment. The effect on the second reaction is more pronounced due to a significantly higher CO equilibrium partial pressure at 500 "C. There is also kinetic evidence that carbothermic reduction of Na2C03to Na20 and Na is feasible under pretreatment conditions. Wigmans et al.24 employed temperature-programmed desorption (TPD) to study the effects of heating carbon/Na2C03mixtures in vacuo. They witnessed a significant evolution of C02 and CO commencing at 300 "C. Also, in agreement with these results are the experimental findings of Hughes et al.25regarding carbothermic reduction of sodium carbonate. These workers observed a fairly high CO/CO2product ratio. This evidence suggests that a great deal of Na2C03is reduced to metallic Na during both pyrolysis and outgassing pretreatment steps. Thus the data suggest that the second model in which carbon acts to reduce sodium to the metallic state and not merely to a partially oxidized state is a more plausible general mechanism. Table V lists the possible oxidation reactions (on a per mole of O2 basis) involving Na and Na20 which are the most likely species formed by the sodium carbonate precursor. Clearly, the observed heats of oxygen adsorption in region I of both sodium samples are too high to be explained by the complete oxidation of any single partially (24) Wigmans, T.; Haringa, H.; Moulijn, J. A. Fuel 1983, 62, 185. (25) Hughes, R. L.; Smith, 1. C.; Lawless, E. W. Production of the Boranes and Rekrted Research; Holtzmann, R. T., Ed.; Academic Press: New York, 1967.
532
Energy & Fuels 1992,6, 532-534
Table VI. Oxide Film Dispersion on Sodium Carbonate Promoted Chars % Na oxidized sample wt % Na to NazOza 500 "C 17.6 20.6 500 "C 20.3 17.4b 900 "C 11.0 21.2 900 " C 10.5 24.ab a See
text for explanation of assumptions employed. bRerun.
oxidized species. The heats observed are more consistent with a mixture of processes, all involving the oxidation of sodium metal. Again as has been pointed out, it is thermodynamicallyreasonable for sodium metal to form during the decomposition of the carbonate species above 500 "C. The X-ray diffraction data are also in accord with the formation of completely reduced sodium metal during pretreatment. Diffraction patterns for both the high- and low-temperaturesamples show evidence that some sodium metal still remains following calorimetric oxidation and air exposure. This suggests that a significant amount of sodium metal was present following the initial high temperature treatments. Some insight regarding morphological behavior is also available. Indeed, the moderately high dispersion (Table VI) suggests that the reduced sodium metal melts and spreads on the char surface to a considerableextent. This is supported by the normal melting point of sodium, 98 OC.= Moreover, there is ample evidence that sodium wets the surface of carbonaceous materials, especially coal (26) Barin, I. Thermochemical Data of Pure Substances; Verlagsgesellschaft: Weiskin, FRG, 1989.
cham2 It is also possible that some of the sodium on the high-temperature char evaporates during the final heat treatment step since the normal boiling point of sodium is about 900 "C. The sodium dispersion was eatimated as follows. First, the sodium content in each sample was determined by quantitative analysis. Oxidation stoichiometrythen provided the maximum amount of NazOzpossible. That is, it was assumed that all oxygen adsorbed by the samples participated in oxidation of reduced metallic sodium to Na202. Finally, the ratio of the amount of oxide formed to the maximum amount possible provided the fractional dispersion of sodium on the char surface. Table VI summarizes these results for each of the sodium promoted chars. Clearly, the dispersion is relatively high for all samples. However, it must be noted that the sodium dispersion is not as high as that found for potassium on the same char.*pg AB a final note, the results of this study suggest that the mechanism of char catalysis for sodium and potassium have both similarities and differences. They are similar in that in both cases gasification occurs via an oxidation/reduction process. They are different in that in the potassium case the metal is never fully reduced, whereas in the sodium case most of the metal apparently is fully reduced following the carbothermic reduction part of the cycle. The work done here suggests that additional work should be undertaken using different sodium and potassium precursors and perhaps the pure metals in order to more fully understand the role of alkali metal catalysts in gasification. Registry No. Na2C03,497-19-8; MgC03, 546-93-0;02,778244-7.
. . Communications N
Solvent Quality Assurance in Porphyrin Research Sir: Questions concerning the chemical stability of porphyrins in various solvents are important to porphyrin research in geochemistryand related analysis.'-7 Sporadic measurements of nickel porphyrin at low concentrations in dichloromethane suggested a question of solvent-related instability.6 Spectrophotometer source irradiation was found to cause a second problem, namely, photodegradation of nickel etioporphyrin-I. The effect is rapid in carbon tetrachloride: 119% loss during a 2-min scan. Under similar conditions, corresponding free-base porphyrin photodecomposes at about the same rate, while the vanadyl complex was 40-fold slower. Given the extent to which chlorinated solvents can enhance porphyrin research-as well as porphyrin decomposition-the present study was made to minimize such difficulties. (1) Symposium issue: Energy Fuels 1990, 4, no. 6. (2) Barwise, A. J. G.; Wolff, G. A.; Eglinton, G.;Maxwell, J. R. J. Chromatogr. 1986,368, 1-9. (3) Flynn, J. S.; Freeman, D. H. J. Chromatogr. 1987,386, 111-121. (4) Freeman, D. H.; Angeles, R. M.; Keller, S. Prepr.-Am. Chem. SOC.,Diu. Pet. Chem. 1988, 33, 231-8. (5) Boreham, C. J.; Fookes, C. R. J. Prep.-Am. Chem. SOC.,Diu. Pet. Chem. 1989,467, 195-208. (6) Freeman, D. H.; Swahn,I. D. Energy Fuels 1990,4,699-704. (7) Freeman, D. H.; OHaver, T. C. Energy Fuels 1990, 4, 695-99.
0887-0624/92/2506-0532$03.00/0
The presence of impurities in dichloromethanehas been the subject of numerous reports. To illustrate, chlorocarbons, phthalic acid and its esters? cyanogen chloride: hydrogen chloride, methyl chloride and chloroform,1°and tetrachloroethane" impurities have been identified. We have observed several instances where porphyrin losses were due to reactive impurities. These were traced to special lots of HPLC-grade solvent containing cyclohexane (C8H12:150 ppm) stabilizer where rapid partial degradation of nickel etioporphyrin-I, as well as free base, was found to occur in the absence of light. This is a plausible explanation for previously observed erratic spectral data? The use of dichloromethane containing cyclopentene was linked to certain HPLC difficulties12that were not found (8) Bowers, M. L.; Parsons, M. L.; Clement, R. E.; Eiceman, G.A,; Karasek, Jr. F. W. J. Chromatogr. 1981,206,279-288. (9) Franklyn, R. A.; Heatherington, K.; Morrison, B. J.; Sherren, P., Word, T. J. Analyst (London) 1978, 103, 662. (10) Sedivec, V.; Flek, J. Handbook of Annlysis of Organic Soluents; Halsted Press (Wiley): New York, 1976; p 146. (11) Bouis, P. A., J. T. Baker, Inc. Private Communication, 1991. (12) Pentene-stabilized dichloromethane caused anomalous HPLC peaks when an alcohol, either methanol or isopropyl alcohol, was introduced as the coeolvent, but not with hexane or ethyl acetate. The anomaly worsened with increasing air exposure. Cyclohexane-stabilized dichloromethane did not exhibit this property (Jarvis, B. B. Private Communication 1991).
0 1992 American Chemical Society
