Ind. Eng. Chem. Res. 1988,27, 1344-1348
1344
results of the present study demonstrate that this mechanism must be common to cobalt manganese oxide catalysts with Co/(Co + Mn) ratios < 0.5 since these formulations demonstrate the same 01 value for the Schulz-Flory distribution in addition to showing a linear dependence of C3 production on CO conversion. The nature of the C1 and C2 reactive species is as yet not defined, but the operation of such a mechanism with cobalt manganese oxide catalyst explains why methane yields lower than those expected from a Schulz-Flory distribution are obtained. However, the production of high yields of C3 and C4 hydrocarbons compared to the current commercial catalysts is of significant technological and economic importance. For optimal unpromoted formulations, the C3and C4 hydrocarbons comprise ca. 28 mass % with a methane selectivity of only 3.5%. The C3 and C4 hydrocarbons contain >80% alkenes and hence would be a suitable feedstock either for oligomerization to produce diesel-range hydrocarbons, e.g., via the Mobil MOGD process, or for alkylation under acid conditions to produce gasoline-range hydrocarbons. Hence, use of a cobalt manganese oxide CO hydrogenation catalyst could provide a synfuel process with flexibility in the diesel/gasoline ratio produced. Conclusions The results of this study indicate that cobalt manganese catalysts can give high activity, long-lived catalysts for CO hydrogenationwith high propene and low C1 and C2yields. The optimized unpromoted catalyst formulation has a Co/(Co + Mn) ratio in the range 0.15-0.25 and must be operated at ca. 220 "C. While upper limits for the optimum conditions for reaction pressure and reactant space velocity have yet to be determined, the catalysts should be operated at >600 kPa and GHSV >300 h-l. Further studies are now required to determine if this catalyst performance can be enhanced by addition of suitable promoters, and these studies are now in progress in our laboratories. Acknowledgment We thank Mark Dry, Egbert Gritz, and Dirk Vermaire for useful discussions. We thank Sasol Technology, the University of the Witwatersrand, and the Foundation for Research Development for financial assistance. We also thank Peter Loggenberg and Mark Betts for obtaining the X-ray photoelectron spectra.
115-07-1;C3H3,74-98-6;CH,OH, 67-56-1; CZHSOH, 64-17-5; l-C,OH, 71-23-8; 1-C40H, 71-36-3; l-C,OH, 71-41-0; l-C60H, 111-27-3; l-CVOH, 111-70-6.
C3&,
Literature Cited Barrault, J. Metal Support and Metal Additive Effects in Catalysis; Imelik, B., Ed.; Elsevier: Amsterdam, 1982; p 225. Barrault, J.; Renard, C. French Patent 2 571 274, 1984. Barrault, J.; Renard, C. Appl. Catal. 1985, 14, 133. Barrault, J.; Renard, C.; Yu, L. T.; Gal, J. Proc. 8th Znt. Congr. Catal., Berlin 1984, 2, 101. Betts, M. J.; Copperthwaite, R. G.; Hutchings, G. J.; Loggenberg, P. M., submitted for publication in Surf. Sci. 1988. Biloen, P.; Sachter, W. M. H. Adu. Catal. 1981, 30, 165. Blanchard, M.; Vanhove, D.; Laine, R. M.; Petit, F.; Mortreux, A. J. Chem. SOC., Chem. Commun. 1982, 570. Bussemeier, B.; Frohning, C. D.; Horn, G.; Kluy, W. German Patent 2 518 964, 1976. Copperthwaite, R. G.; Hack, H.; Hutchings, G. J.; Sellschop, J. P. F. Surf. Sci. 1985, 164, L827. Copperthwaite, R. G.; Hutchings, G. J.; van der Riet, M. Znd. Eng. Chem. Res. 1987,26, 869. Copperthwite, R. G.; Derry, T. E.; Loggenberg, P. M.; Sellschop, J. P. F. S. Vacuum 1988, in press. Cornils, B.; Frohning, C. D.; Morow, K. Proc. 8th Int. Congr. Catal., Berlin 1984, 2, 23. Dry, M. E. Catalysis-Science and Technology; Springer Verlag: Berlin, 1981; Vol. 1, p 159. 1980, 102, 2478. Fraenkel, D.; Gates, B. C. J . Am. Chem. SOC. Friedel, R. A.; Anderson, R. B. J . Am. Chem. SOC.1950, 72, 1212, 2307. Frost, D. C.; McDowell, C. A.; Woolsey, I. S. Molec. Phys. 1974,27, 1473. Henrici-Oliv6, G.; Oliv6, S.; Angew. Chem., Znt. Ed. Engl. 1976, 15, 136. Hutchings, G. J. S. Afr. J. Chem. 1986, 39, 65. Hutchings, G. J.; Boeyens, J. C. A. J. Catal. 1986, 100, 507. Jensen, K. 5.;Mossoth, F. E. J . Catal. 1985, 92, 98, 109. Kolbel, H.; Tillmetz, T. US Patent 4 177 203, 1979. Kugler, E. L. Prep.-Am. Chem. SOC.,Diu. Pet. Chem. 1980, 564. Lochner, U.; Papp, H.; Baerns, M. Appl. Catal. 1986,23, 339. Maiti, G. C.; Melessa, R.; Baerns, M. Appl. Catal. 1983, 5, 151. Maiti, G. C.; Melessa, R.; Lochner, U.; Papp, H.; Baearns, M. Appl. Catal. 1985, 16, 215. Pichler, H.; Schulz, H.; Elstner, M. Brennst. Chem. 1967, 48, 78. Ponec, V. Catalysis-Specialist Periodical Report; Royal Society of Chemistry: London, 1982; Vol. 5, p 48. Scofield, J. H. J. Electron Spectrosc. 1976, 8, 129. Van der Riet, M.; Hutchings, G. J.; Copperthwaite, R. G. J . Chem. SOC.,Chem. Commun. 1986, 798. Van der Riet, M.; Hutchings, G. J.; Copperthwaite, R. G. J . Chem. SOC.,Faraday Trans. 1 1987, 83, 2963.
Registry No. Co/MnO, 51845-85-3; CO, 630-08-0; Co, 744048-4; MnO, 1344-43-0;CHI, 74-82-8; CzH4,74-85-1; C5Hs, 74-84-0;
Received for reuiew December 23, 1987 Accepted March 30, 1988
Oxidation of HS03- by O2 David Littlejohn, Ke-Yuan Hu,+and Shih-Ger Chang* Applied Science Division, Lawrence Berkeley Laboratory, University of California, Berkeley, California 94720
The oxidation of bisulfite ion by dissolved oxygen has been studied using a high-pressure flow system and Raman spectroscopy. We have detected an intermediate in the reaction which appears to be formed from 2HS03- and lo2. We propose that the intermediate is disulfate ion, S 2 0 T 2 - which , hydrolyzes into SO2-, H f , and/or HSOd2-with a hydrolysis rate constant of 1.33 f 0.26 X s-l a t 25 "C. An improved understanding of the kinetics of the oxidation of dissolved SO2 by oxygen in aqueous solutions would benefit flue gas desulfurization programs in which On leave from the Research Center for Eco-Environmental Sciences, Academia Sinica, Beijing, People's Republic of China.
0888-5885/8812627- 1344$01.50/0
control of the oxidation of SO2is desirable (Hudson, 1980). This oxidation reaction has been studied for many years, yet the oxidation mechanism is unclear (Alyea and Backstrom, 1929; Backstrom, 1934; Braga and Connick, 1982). Previous studies of this reaction have reported conflicting results because the reacton is very sensitive to 0 1988 American Chemical Society
Ind. Eng. Chem. Res., Vol. 27, No. 8, 1988 1345 Table I comvd
Raman veak, cm-'
RSE"
235 934 967 981 1023 1050 1090 1092 1152 1550
0.15 1.06 0.10 1.00 0.12 0.19 ca. 0.33 1.48 0.94 0.38
~
"\ A
a
D
Figure 1. Schematic diagram of the experimental apparatus: (A) gas tanks, (B)regulators, (C) tanks for liquid reactants, (D) flow regulation valves, (E) flow meters, (F) mixer, (G) Raman cell, and (H) pressure reduction valve.
many kinds of impurities. Some of these impurities act as catalysts (Huss et al., 1982), while others are inhibitors at very low concentrations (Backstrom, 1934; Lim et al., 1982). Studies of this reaction have generally monitored the rate of disappearance of one of the reactants or the rate of formation of sulfate ion by using wet analytical techniques. The reaction stoichiometry, 2HSO,-
+ O2
-
2S042-+ 2H+
(1)
was used to deduce the concentrations of the other species. We have studied this reaction by using laser Raman spectroscopy (LRS) to monitor the behavior of all observable species during the reaction. Our results indicated that a mass balance for sulfur or oxygen could not be obtained from just the products and reactants. A signal due to an intermediate in the reaction was subsequently discovered. This intermediate has not been reported in earlier studies.
Experimental Section A high-pressure stopped-flow/continuous-flowmixing system was constructed for use with the Raman spectrometer and is shown in Figure 1. Two 20-L cylindrical stainless steel vessels (C) were used for preparing and storing the reactant solution under pressure from gas tanks (A). The system is capable of being pressurized to 100 atm to obtain a sufficiently high dissolved oxygen concentration for detection by LRS. Besides dissolved oxygen, use of LRS permits the observation of the following species in the reaction system: SO,(aq), HS03-, HS04-, S042-, S2052-, and S20g2-(Littlejohn and Chang, 1984). The reactants flowed through separate flow meters (E) into a mixer (F). The mixer was designed to efficiently mix the two solutions at flow rates above 2 mL/s. This was confirmed by observing the mixing of Fe"(EDTA) with dissolved nitric oxide in a spectrophotometer. The Fe"(EDTA) + NO reaction is sufficiently rapid so that the rate of product formation was limited by mixing. After mixing, the solution flowed through a quartz cell (G) where it was illuminated by the 514-nm line from an argon ion laser. The light scattered by the solution was collected and passed through a Spex 1403 double spectrometer. A Spex Datamate computer was used to acquire and analyze the data. The solution then passed through a metering valve (H), where the pressure dropped to ambient levels, and past a pH electrode.
Relative scattering efficiency-compared
to SO:-
The bisulfite solution was prepared by adding reagent-grade sodium metabisulfite to deionized water that had been deoxygenated by passing nitrogen through it. After it was prepared, it was sealed in one of the tanks and pressurized with nitrogen. Potassium perchlorate was added as a reference so that quantitative measurements could be made. The oxygen solution was prepared by passing oxygen through a potassium perchlorate solution under pressure. The oxygen concentration was determined by comparing the oxygen Raman peak with the perchlorate Raman peak. The ionic strengths of the solutions were adjusted to the desired levels using NaCl. The species observed in these experiments, along with their Raman peaks and relative scattering efficiencies (RSE's), are listed in Table I. The RSE for a compound is defined as compound's peak height /compound's molarity RSE = 981-cm-I S042- peak height/S042- molarity (2) Two types of experiments were carried out: continuous flow and stopped flow. In the former type of experiment, the mixed solution flowed through the cell at a steady rate while the spectrometer was scanned over the desired region. The volume between the mixer and the cell could be changed to vary the time between mixing and observation. This allowed us to obtain a continuous spectrum of the reacting solution, although a larger quantity of the reactant solutions was consumed than in the stopped-flow experiments. In the stopped-flow experiments, the desired flow through the cell was established, and the spectrometer was set at the desired wavenumber setting. The flow of solution through the cell was stopped under pressure, and the computer recorded the signal change due to the reaction. Afterwards, the spectrometer was set to another position, and the process was repeated under the same conditions until all the desired species had been observed. These measurements permitted the observation of the behavior of the species of interest during the reaction. A number of experiments have been done with dissolved oxygen concentrations ranging from 8 X lod3to 4 X M (after mixing) and bisulfite concentrations ranging from 6 X loW2 to 5 X lo-' M (after mixing). Initial O2to HSO< concentration ratios ranged from 0.02 to 0.4. The pH of the solution immediately after mixing was in the range 3.5-5. All runs were done at 25 "C and at ionic strength p = 0.5 M.
Results and Discussion Most experiments were carried out under stopped-flow conditions. In these experiments, the signals due to S042-, HSOf, SOz(aq),and O2were observed. After the solution had reacted, the Raman spectra at 900-1200 and
1346 Ind. Eng. Chem. Res., Vol. 27, No. 8, 1988 1000'
-
Time (sec)
Figure 2. Relative concentrations of oxygen (solid line) and total S(1V) (dashed line) with time after mixing. 13ocp------
I
,
i
Raman shift (cm-')
Figure 3. Raman spectra of the intermediate (left peak) and S202(right peak).
1540-1560 cm-l were collected. From these data, the concentration vs time profiles and the final concentrations of HSO,-, 02,SO:-, and S02(aq)could be determined by comparing their peaks with the C104- peak. The H+ concentration could be determined from the equilibrium (Huss and Eckert, 1977) [H+][HS03-]/ [S02(aq)]= 0.0139 M (3) and the HS04- concentration could be estimated from (Nair and Nancollas, 1958) [H+][SO:-] /[HSO,-] = 0.0102 M (4) With the concentration vs time profiles established, mass balances for sulfur and oxygen could be obtained. For sulfur, the curves for HS03-, S02(aq),S042-,and HS04were added together. For oxygen, the curves for SO:- and HS04- were added, multiplied by two, and then added to the oxygen curve. The concentrations of other species, such as S2052-, were negligible and were not included in the mass balances. The mass balance plots for both sulfur and oxygen generally showed a dip shortly after mixing, which implied the presence of an intermediate. Figure 2 shows the rapid decrease in the O2 reactant and the slow + HS04-) after mixing. While atbuildup in Sv1(S042tempting to observe S206*-a t 1092 cm-l, we obtained a curve with a profile in time close to that expected for the intermediate. Subsequent study showed that the peak of the intermediate was centered at 1090 cm-' and was somewhat broader than S2062-(Figure 3). Scans taken during continuous-flow experiments indicated the presence of a second, weaker peak of the intermediate at about 740 cm-'. We attribute the Raman peaks to S-0 stretches.
I 0
50
'00
,
150
1
200
250
33C
Time (seconds)
Figure 4. Typical decomposition plots of the intermediate.
Other oxysulfur compounds with Raman peaks near 1090 cm-' include S20e2-(1092 cm-l) and SzOs2-(1074 cm-'). Dithionate ion is very stable and does not decay in the manner of the intermediate. No peaks were observed that would suggest a peroxy linkage, as in H202(875 cm-') or S202- (835 cm-'). In runs where the intermediate was rapidly formed and then decayed away, the changes in concentrations of 02, HSOB-,and S(V1) could be used to determine the stoichimetry of the intermediate. Both the ratios of the change in HS0,- (after correction for the change in pH) to the change in O2 and of the change in S(V1) to the change in O2 were about two to one, indicating that the intermediate is formed from two HSO, and one 02.This ratio has been obtained by many other studies. We have studied the reaction mixture by ESR spectroscopy and found no signal which could be ascribed to the intermediate, indicating that it is not a free radical. We propose the intermediate has the formula of S2072-.The decomposition of the intermediate does not appear to be significantly influenced by the experimental conditions. It appears that the decomposition occurs by
S2072-+ H20
-
2H+ + SO4*-
(5)
Plots of log, [intermediate] vs time (Figure 4) are linear, indicating first-order decomposition. A series of stopped-flow experiments were conducted over a range of reactant concentrations. Values of the rate constant for the decay of the intermediate were obtained from the slopes of log, (concentration) vs time plots. These values are tabulated in Table I1 with the initial reactant concentrations. The data in the column labeled "from intermediate" are from the intermediate signal a t 1090 cm-'. The data in the column "from total S(V1)" are from the growth of the normalized signals a t 981 cm-' (SO?-) and 1050 cm-l (HS04-). The data in the final column "from total H+" was obtained by calculating the concentration of free H+from eq 3 and adding to it the concentrations of HS04- (1050 cm-l) and S02(aq)(1152 cm-l). The initial pH of the mixed reactants for the runs listed varied from 3.6 to 4.3. The values obtained from the decay of the intermediate agree well with the values obtained from the buildup of S042-and H+. The average value of
Ind. Eng. Chem. Res., Vol. 27, No. 8, 1988 1347 3 5 0 0 1 : :
Table I1 initial IHSOs-l, -.
M
0.19 0.22 0.22 0.22 0.22 0.22 0.23 0.25 0.27 0.30 0.34 0.34 0.34 0.34 0.475 0.68
~
:
I
:
:
.
:
t
decay rate constant, s-l from from intermedtotal from iate S(V1) total H+
initial M
[02l,
0.014 0.012 0.016 0.032 0.038 0.0385 0.014 0.036 0.014 0.024 0.016 0.017 0.017 0.017 0.014 0.014
0.0135 0.0141 0.0139 0.0112 0.0133 0.0178 0.0139 0.0141 0.0099 0.0180 0.0126 0.0183 0.0125 0.0122 0.0150 0.0144
0.010 0.0124 0.0130 0.0104 0.0153 0.0132 0.0111 0.0138 0.0141 0.0168 0.0091 0.0116 0.0115 0.0086 0.0161 0.0203
0.0115 0.0130 0.0183 0.0116 0.0126 0.0127 0.0114 0.0148 0.0141 0.0118 0.0104 0.0119 0.0099 0.0112 0.0113 0.0177
all determinations is k = 0.0133 f 0.0026 s-' at 25 "C and p = 0.5 M. The decomposition rate did not appear to be influenced by the pH of the solution over the range of pH 2-5 nor was it influenced by the presence of Fe3+or Mn2+ a t concentrations up to about lo-* M. The properties and behavior of the intermediate are very similar to those of the disulfate ion, S2072-. To make a confirmation that the intermediate is indeed disulfate ion, we prepared sodium disulfate by heating sodium bisulfate to 180 OC for 5 days (Hofmeister and Van Wazer, 1962). The hydrolysis of the disulfate ion was observed by monitoring the decay of the 1090-cm-' Raman peak with time. Values of the hydrolysis constant similar to those in Table I1 were obtained ( k 0.012 s). There have been three previous studies (Hofmeister and Van Wazer, 1962; Thilo and von Lampe, 1963; Ryss and Drabkina, 1973) of the hydrolysis rate of the disulfate ion in aqueous solution. All of the studies found an activation energy of about 11 kcal/mol, although Hofmeister and Van Wazer obtained a lower rate constant at 25 "C ( k = 0.0067 s-l) than Thilo and von Lampe ( k = 0.0126 s-l) and Ryss and Drabkina (k = 0.0110 s-l). The hydrolysis rate constant we obtained in this study agrees well with these values. Thilo and von Lampe (1963) reported that sodium ions accelerate the rate somewhat. This effect is sufficiently small so that our value of the rate constant should require no correction. To obtain the Raman spectrum of S20:- at 25 OC, about 0.2 g of the salt was mixed with about 5 mL of water for 5 s, and a sample was quickly collected. A Raman spectrum was collected over a 40-cm-' band for about 25 s. This was done for bands at 305-345,400-440, 720-760, 760-800, 800-840, and 1070-1110 cm-'. A composite spectrum of these results is shown in Figure 5A, after correcting for the background. The approximate peak maxima are located at 319,739, and 1090 cm-l. The only previously reported Raman spectrum of disulfate ion in aqueous solution was obtained by Millen (1950), who was uncertain of the peak at 735 cm-'. Frequency assignments for solid disulfate salts were made by Simon and Wagner (1961). The spectrum confirms that the peak at about 740 cm-' in spectra of the reaction mixtures is due to disulfate ion and not another intermediate. To obtain a mass balance for sulfur and oxygen, the relative scattering efficiency (RSE) for S2072-is needed. To obtain a measure of the relative scattering efficiency of S20$'-,about 2 g of the sodium disulfate salt was mixed with 30 mL of water at 0 OC. After mixing for 30 s, it was added to a reservoir, in a 0 "C bath, which fed the solution through the spectrometer cell. Fresh aliquots of the so-
-
j 300
C 750
1200
Raman shift (cm-')
Figure 5. (A) Composite Raman spectrum of disulfate ion in solution a t 25 "C. (B) Raman spectrum of flowing disulfate ion in solution at 0 "C. (C) Difference between Raman spectra of the reaction mixture during and after reaction.
lution were added every 90 s while the spectrum was collected. The spectrum is shown in Figure 5B. Some decomposition had occurred by the time the solution flowed through the cell, as indicated by the S042-peak at 981 cm-'. By the change of height of the SO:- and HS04peaks between the flowing solution and the completely decomposed solution, the RSE of the S20:1090-cm-l peak was calculated to be about 0.33. For comparison with the spectra of S2072-,Figure 5C shows the difference between Raman spectra of the O2 HS03- reaction during and after reaction. SO:- and HSO, are formed by the reaction and have negative peaks (not shown). HS03- and the intermediate decrease as the reaction goes to completion and have positive peaks. The spectra in Figure 5 suggest that the peak at about 740 cm-' has two components. It is possible that other intermediates are present in the reaction mixture. If they are present, they are at lower concentrations or exist for shorter times than S2072-.No other sulfur oxyanion we know of has Raman peaks at 740 and 1090 cm-' and decays in aqueous solutions with a rate constant of k = 0.013 s-'. Connick and co-workers (1987) have been studying the oxygen-bisulfite ion reaction and have found indirect evidence of the intermediate. They are studying the reaction at more dilute conditions than those used in this study, with reactant concentrations of about M oxygen and M HS03-. Their measurements of the rate constant for the decay of the intermediate are in good agreement with our results. The formation of the observed intermediate shows complicated dependence on reactant concentrations and trace amounts of metal ions such as Fe2+/Fe3+and Mn2+. As the pH of the reaction mixture is raised above 4.5, the relative amount of intermediate formed decreases, suggesting the presence of another reaction channel. The fraction of S(1V) in the form of S032-increases rapidly as the pH increases. It may be that another oxidation reaction channel proceeds through the involvement of SO3*at a faster rate than the reaction involving HS03-. The observed intermediate does not seem to be involved in this process. The factors that govern the formation of the intermediate are under investigation.
+
Registry No. HS03-, 15181-46-1; 02,7782-44-7.
Literature Cited Alyea, H. N.; Backstrom, H. J. Am. Chem. SOC.1929,51, 90. Backstrom, H. 2.Phys. Chem. 1934, 258, 122. Braga, T. G.; Connick, R. E. In Flue Gas Desulfurization; Hudson, J. L., Rochelle, G. T., Eds.; ACS Symposium Series 188; American Chemical Society: Washington, D.C., 1982; pp 153-171.
1348
Ind. Eng. Chem. Res. 1988, 27, 1348-1356
Connick, R. E., private communication, 1987. Hofmeister, H. Z.; Van Wazer, J. R. Inorg. Chem. 1962, 1, 811. Hudson, J. L. "Sulfur Dioxide Oxidation in Scrubber Systems". EPA-600/7-80-083, 1980; EPA, Washington, D.C. Huss, A., Jr.; Eckert, C. A. J. Phys. Chem. 1977,81, 2268. Huss, A., Jr.; Lim, P. K.; Eckert, C. A. J. Phys. Chem. 1982,86,4224, 4229. Lim, P. K.; Hws, A., Jr.; Eckert, C. A. J. Phys. Chem. 1982,86,4233. Littlejohn, D.; Chang, S. G. Enuiron. Sci. Technol. 1984, 18, 305.
Millen, D. J. J. Chem. Soc. 1950, 2589. Nair, V. S. K.; Nancollas, G. H. J. Chem. Soc. 1958, 4144. Ryss, I. G.; Drabkina, A. Kh. Kinet. Katal. 1973, 14, 242. Simon, A.; Wagner, H. Z. Anorg. Allg. Chem. 1961, 311, 102. Thilo, E.; von Lampe, F. Zn. Anorg. Allg. Chem. 1963, 319, 387. Received for review November 30, 1987 Revised manuscript received March 9, 1988 Accepted March 21, 1988
Asphaltene Reaction Pathways. 4. Pyrolysis of Tridecylcyclohexane and 2-Et hyltetralin Phillip E. Savaget and Michael T. Klein* Department of Chemical Engineering and Center for Catalytic Science and Technology, University of Delaware, Newark, Delaware 19716
The thermal reactions of the alkylnaphthenic and alkylhydroaromatic moieties likely present in petroleum asphaltenes were investigated Via pyrolysis of the model compounds tridecylcyclohexane (TDC) and 2-ethyltetralin (2ET), respectively. TDC pyrolyzed to the major product pairs cyclohexane plus l-tridecene and methylenecyclohexane plus n-dodecane, both in initial selectivities of about 0.08. Cyclohexene, methylcyclohexane, and n-tridecane also formed but in lower selectivities of about 0.0388,0.035, and 0.030, respectively. Arrhenius parameters of [log A (s-'), E * (kcal/mol)] = [14.9, 59.41 were determined over the temperature range 400-450 "C. 2ET pyrolyzed to 2-ethylnaphthalene, 2-ethyldialins, naphthalene, dialin, and tetralin as major products in yields of 1.73%, 1.19%, 1.13%, 1.09%, and 0.17%, respectively, a t low (=5%) conversions and 3.79%,0, 30.2%, 2.19%, and 6.49%, respectively, at high (=65%) conversions. The temporal variation of these products permitted deduction of the reaction network. It includes pathways for primary reaction of 2ET to dialin and the 2-ethyldialins and secondary reactions accounting for the formation of the other major products. The Arrhenius parameters for the disappearance of 2ET were [log A @I), E * (kcal/mol)J = [12.7, 53.51. The pyrolyses of both compounds were consistent with free-radical reaction mechanisms. The implications of the results of this model compound study to the reactions of the analogous moieties in asphaltenes are suggested. Petroleum asphaltenes, operationally defined as the heptane-insoluble, benzene-soluble fraction of heavy oils and residua, are a physically and chemically complex mixture of polar and aromatic compounds. Their ubiquity in heavy oils and their propensity for deactivating catalysts have rendered asphaltenes, and their reactions under hydroprocessing conditions, to be of considerable practical significance. Unfortunately, however, the complexity of both asphaltenes and their reaction products frustrates the deduction of reaction engineering fundamentals from direct experiments with precipiated asphaltenes. Resolved information is typically limited to the apparent kinetics of asphaltene disappearance and the yields of solubilitybased product fractions (Schucker and Keweshan, 1980; Savage et al., 1985; Ternan, 1983). This complexity motivates complementary studies wherein model compounds mimicking key reactive moieties in asphaltenes are reacted to probe asphaltene reaction pathways. Experiments with model compounds can provide intrinsic kinetics, operative reaction pathways, and, in favorable instances, reaction mechanisms. The literature suggests (Yen et al., 1984a,b; Takegami et al., 1980; Speight and Moschopedis, 1981; Speight and Pancirov, 1983) that condensed aromatic and naphthenic ring systems with peripheral aliphatic substituents are important moieties in petroleum asphaltenes. Therefore,
* Corresponding author. +Present address: Department of Chemical Engineering, University of Michigan, Ann Arbor, MI 48109. 0888-5885/88/2627-1348$01.50/0
model compounds of the basic asphaltene hydrocarbon framework should contain alkylaromatic, alkylhydroaromatic, and alkylnaphthenic moieties. We previously (Savage and Klein, 1987a,b) reported on the pyrolysis of n-pentadecylbenzene and l-phenyldodecane, prototypical alkylaromatics, and herein report on the pyrolyses of a mode alkylnaphthene, n-tridecylcyclohexane (TDC), and a model alkylhydroaromatic, 2-ethyltetralin (2ET). Few reports of the pyrolysis of singly substituted alkylcyclohexaneshave appeared in the literature. Fabuss et al. (1964) pyrolyzed several saturated cyclic hydrocarbons including cyclohexane and methyl-, ethyl-, propyl-, and n-butylcyclohexane,and they developed an empirical correlation between the first-order rate constant for substrate disappearance at 427 "C and molecular structure. Mushrush and Hazlett (1984) pyrolyzed tridecylcylohexane at 450 O C and observed n-alkanes, l-alkenes, alkylcyclohexanes, toluene, benzene, and methylcyclohexane as reaction products. However, this latter investigation provides only limited kinetics results and little information regarding the operative reaction pathways. The literature concerning the pyrolysis of alkyltetralins is likewise sparse, and only methyl-substituted tetralins have been previously pyrolyzed. For instance, Trahanovsky and Swenson (1981) subjected l-methyl- and 2methyltetralin to flash vacuum pyrolysis conditions (700-900 "C, 0.1 Torr), and they concluded that cleavage of alkyl groups was facile because the major products were naphthalene and 1,2-dihydronaphthalene. No reports of 2-ethyltetralin pyrolysis were found. 0 1988 American Chemical Society
