(7) Grieves, R. B., Wilson, T. E., Shih, K. Y., A.Z.Ch.E. J. 11, 820 (1965). (8) Haas, P. A., Ph.D. thesis, University ofTennessee, 1965. ( 9 ) Haas, P. A., Johnson, H. F. A.Z.Ch.E. J . 11,319 (1965). (10) Karger, B. L., Poncha, R. P., Miller, M. M., Anal. Chem. 38, 764 (1966). (11) Kishimoto, H., Kolloid-Z. 192, 66 (1963). (12) Metzner, A. B., Brown, L. F., Znd. Eng. Chem. 45, 2040 (1956). (13) Poncha, R. P., Karger, B. L., Anal. Chem. 37,422 (1965). (14) Rubin, E., Ph.D. thesis, Columbia University, 1962. (15) Rubin, E., Everett, R., Jr., Znd. Eng. Chem. 55,44 (1?63). (16) Rubin, E., Gaden, E. L., “Foam Separation” in “New
Chemical Engineering Separation Techniques,” H. M. Schoen, ed., Interscience, New York, 1962. (17) Rubin, E., Schonfeld, E., Everett, R., Jr., Oak Ridge National Laboratory, R.A.I. 104 (October 1962). (18) Shinoda, K., Mashio, K., J . Phys. Chem. 64, 54 (1960). (19) Walling, C., Ruff, E. E., Thornton, J. L., Jr., Zbid., 56, 989 (1952). RECEIVED for review May 5, 1966 ACCEPTEDNovember 7, 1966 Work performed as part of a research project toward the Ph.D.
degree of Morris Goldberg.
OXIDATION OF COKED SILICA-ALUMINA CATALYST F.
E. M A S S O T H
Rejnery Processes Division, Gurf Research C3 Development Co., Pittsburgh, Pa. 15230
The kinetics of oxidation of artificially coked silica-alumina catalysts were studied a t several temperatures between 425” and 480” C. and oxygen concentrations of 4 to 21% a t atmospheric pressure. A flow microbalance in conjunction with product gas analysis enabled the course of the oxidation to b e followed semicontinuously. A rapid, initial temperature rise occurred, followed by an essentially isothermal, carbon oxidation which was not influenced by mass-transport phenomena under the reaction conditions employed. Oxidation of hydrogen from the coke proceeded more rapidly than that of carbon. Appreciable oxygen uptake attended oxidation. These phenomena were interpreted as an initial adsorption of oxygen to form a surface carbon complex with rapid liberation of heat; a sustained, chemically controlled carbon oxidation a t the outer surface of the coke particles; a simultaneous oxidation of the hydrogen in the coke, limited by diffusion of oxygen through the outer carbon layer to an inner coke boundary; and oxygen uptake mainly due to the sorption of oxygen in the outer carbon layer.
employed in the hydrocarbon processing industry deactivated, with the concomitant deposition of a condensed carbonaceous polymer called coke. T h e catalyst is normally reactivated by burning off the coke a t elevated temperature in a n air-nitrogen stream. The object of the present work was to study the kinetics of the oxidation of a coked commercial catalyst. Considerable information exists on the oxidation of pure carbon and graphite (78). Research devoted to carbon blacks and cokes is less extensive. Only a few definitive studies of the oxidation of coke on porous supports have been reported. I n particular, the reaction of the hydrogen associated with the coke seems not to have been investigated in any detail. T h e processes involved in oxidation of coke on a porous support are akin to but more complex than those involved in general gas-solid reactions. For example, not only must the relative contributions of gas phase diffusion, surface adsorption, and chemical reaction be considered, but intraparticle mass transport, coke distribution, and the presence of hydrogen in the coke must be added. In this paper, we show qualitatively the relative importance of each of these conjunct steps and define quantitatively the kinetics of the controlling steps for the carbon and hydrogen oxidation reactions. ATALYSTS
C eventually become
Experimental
Catalyst. The catalyst used was a fluid-bed cracking catalvst consisting of 759% by weight silica and 25% by weight alumina ( C y a n a k d , Tdple A),-which was tableted’ using a binder [5y0poly(viny1 acetate) and 2% Acrawax-C], calcined in a muffle furnace at 540” C. overnight, and broken down to 200
I & E C PROCESS D E S I G N A N D DEVELOPMENT
10- to 20-mesh particle size. Analysis before coking showed less than 0.2% by weight carbon and a BET area of 436 sq. meters per gram. Coking was accomplished by first immersing the catalyst in a prepurified furnace oil, followed by a 4-hour treatment a t 425’ C. under a hydrogen pressure of 750 p.s.i.g. The coked catalyst was washed with n-pentane and oven-dried a t 120’ C. overnight; it had a BET area of 349 sq. meters per gram. Chemical analysis showed 9.6% by weight carbon and 1.1yo by weight hydrogen. T h e latter value includes support water driven off the catalyst during analysis besides the hydrogen associated with the coke. Table I shows a typical determination of carbon, hydrogen, and water on the coked catalyst. A small amount of volatile carbon evolved on heating the coked catalyst in nitrogen to 500” C. All oxidation runs were preceded by a 480’ C. prevolatilization in a nitrogen stream for 2 hours. The coke had an average composition a t this point corresponding to CHo.4. Another batch of coked catalyst containing 7.0% C was caustic-treated to isolate the coke according to a leaching
Table 1. Source
Water 500° C. Coke Volatile Nonvolatile
Coke Composition of Catalyst
c, wt. 5%
... ...
H,
wt. %
0.3-0.65 0.35
0.5 0.05 9.1 0.25 Total 9.6 0.95-1 ,25 Anal. 9 . 6 & 0 . 2 1.15 & 0 . 1 a Varies with sample exposure to room atmosphere and prior thermal history.
method described by Haldeman and Botty (8). Analysis of the extracted coke phase for two samples, one the 10- to 20-mesh and the other ground up prior to leaching, yielded : Surf. H, H / C Mole Area, Extracted C, Ratio Sq. M./G. Coke wt.yo wt.yo 10-20 88.39 3.65 0.50 237 Ground 84.68 3.45 0.49 102 T h e carbon gradient throughout the bulk particle was determined by a n attrition sieving of the coked catalyst. T h e catalyst was systematically reduced in size by carefully forcing it through the next smaller screen size, and determining carbon in the batch through and the fines produced a t each step reduction. T h e results, obtained on another batch of catalyst coked in the same manner as described above but a t 375' C., are shown in Table 11. No appreciable carbon variation was noted on reducing the particle diameter by a factor of 7. T h e fines did show slightly higher carbon values than the respective onscreen samples, indicating that some fine coke particles may have been dislodged in the screening process. Eberly et al. (5) also found the coke to be evenly distributed throughout a coked silica-alumina cracking catalyst. Equipment. The basic apparatus used in the oxidation runs was a flow, McBain-type quartz spring balance. The reactor consisted of three basic parts joined together by 29/42 standard-taper joints: a lower tube of l'/Z-inch 0.d. Vycor, surrounded by a high temperature electrical furnace; a n upper tube of Il/n-inch borosilicate glass, containing a circulating water jacket for providing a constant spring temperature; and a borosilicate glass cap. A glass yoke was affixed to the cap, from which hung a quartz spiral spring located in the upper section. T h e catalyst sample was contained in a n 80mesh platinum screen bucket suspended from the spring by a fine platinum wire. Catalyst weight changes were monitored by measuring the vertical displacement of the spring with a IO-cm. Gaertner cathetometer. Two glass fibers attached to the yoke and extending to near the bottom of the spring provided a n internal fiducial reference mark; attached to the spring bottom was a glass fiber containing a fine crosspiece, permitting easy reading of the spring extension. The spring sensitivity a t 30' C. was 5.45 mg. per mm., extension being proportional to load. T h e spring reading reproducibility was 1 0 . 0 2 mm., corresponding to a weight change error of 1 0 . 1 1 mg. T h e cap and spring assembly could be replaced by a n auxiliary thermocouple which permitted separate but not concurrent determination of sample temperature changes during the initial stages of oxidation. Because of the small samples employed (150 mg.), a microthermocouple was used, positioned in the middle of the sample. Air and nitrogen for the oxidation runs were predried through Type A molecular sieves and mixed by flowing through calibrated rotameters. A preheater containing glass spheres brought the gas up to about 320' C. before entering the reactor section, Reactor temperature was achieved by a metal-block heater, controlled to within 1 2 ' C. by means of a thermoregulator. Sample temperature was about 4' C. lower than furnace temperature. Flow through the reactor was normally downward. Product gases were successively passed through a Gilbarco moisture analyzer having a range up to 1500 p,p.m.
of water and a Greenbrier process analyzer, which measured C O z and C O separately u p to 1000 p.p.m. with a 10-minute analysis time. Cumulative amounts of CO, and H 2 0 released were obtained by measuring the areas of the respective concentration us. time plots. All runs were made a t atmospheric pressure (average 745 mm.) plus the pressure drop through the water and gas process analyzers. For the standard flow rate of 400 cc. S T P per minute used, the pressure drop amounted to about 55 mm., giving a total average pressure for most of the runs of 800 =t10 mm. Procedure. The general procedure adopted for most of the oxidation runs consisted of three steps: prevolatilization, isothermal oxidation, and final burnoff. The pretreatment step involved a 2-hour heatup to reaction temperature followed by a 2-hour heating at temperature, all under a dry nitrogen flow. By the end of this treatment, the concentration of desorbed water was close to the water analyzer base line. The desired air-nitrogen mix was then substituted for the nitrogen, usually a t a flow of 400 cc. STP per minute. After 150 seconds, from start of the oxidation, the process analyzer was activated. Meanwhile, spring changes lvere continually recorded. After oxidation for a sufficient period of time, complete carbon burnoff was obtained by switching to straight air and raising the furnace temperature to 600' C. Both visual inspection and chemical analysis showed no carbon left on the catalyst a t this temperature. T h e final burnoff provided a material balance check on the run. T h e catalyst before and after the run was weighed on a n analytical balance.
con,
Results
Preliminary Runs. A sample of the data obtained in a typical oxidation of the coked catalyst is shown in Figure 1. T h e top line represents the change in spring readings with time, which are converted to milligram weight changes by previous spring calibration. The apparent weight gain noted a t the beginning of reaction is real. The lower lines show product concentration profiles during reaction. Both COz and CO were usually observed in all oxidation runs, their ratio varying with the reaction conditions, COz always being greater than CO. T h e water arises from oxidation of the hydrogen in the coke. I t is significant that the water decays more rapidly than the carbon oxide concentrations. Before attempting to define the kinetics of the oxidation reactions, it was necessary to assess the contribution of physical transport processes relative to chemical reaction rates. OFving to the low thermal conductivity of porous supports, severe temperature gradients are possible ( 7 4 , especially for highly exothermic reactions. I n order to check sample temperature during oxidation, several runs were made using the thermo-
14
800
600
n
\Li -1.5
0
Table II. Carbon Distribution in Coke Catalyst Median c, wt.70 Mesh Size Diam., Mm. On mesh Through 100 10-20 1.42 6 . 9 =k 0 . 2 6.6 20-40 0.63 6.3 7.35 40-60 0.33 6.7 8.2 60-100 0.20 7.0 Av. 6 . 7 iz 0 . 3 7 . 4 f.0 . 5
zoc
f
TIME, 1000 SEC.
Figure 1.
Oxidation product concentration VOL. 6
NO. 2 A P R I L 1 9 6 7
201
couple probe described earlier. A rapid rise in sample temperature was immediately obtained upon introduction of the oxidizing gas, followed by a moderately rapid decrease in temperature to a constant value for the duration of the oxidation. T h e initial temperature peak was about 70' C. for the most severe oxidation conditions used (air a t 480' C.) and considerably less for milder conditions. Sample temperature returned to essentially isothermal conditions within several hundred seconds, whereas oxidation runs usually lasted to 10,000 seconds or more. Thus, the reaction can be treated as isothermal except in the earliest stage of oxidation. The effect of gas mass transfer on oxidation rates was checked by varying the total flow rate during a run between 100 and 800 cc. per minute. T h e concentration of CO2 C O showed no deviation from the normal concentration decay profile when corrected to a basis of 400 cc. per minute, showing gasphase mass transfer to be rapid a t our reaction conditions. T h e rate of oxygen supply was always greater than 50 times the estimated initial rate of oxygen consumption by the sample. Interparticle diffusion, or bed effect, was checked in two ways-sample size and bucket geometry. Identical oxidation results were obtained when the standard 150-mg. sample size was reduced to 50 mg. Similar results were obtained in using a quartz bucket in place of the platinum mesh bucket. I n the latter case, however, a slight retarding effect was observed in the very initial stages of oxidation using a quartz bucket. Therefore, the platinum wire bucket was employed in the kinetic studies. According to Weisz (27),our reaction temperature of 475' C. is close to the upper limit for chemical control of coke oxidation in pellets of 2-mm. diameter; above this limit, reaction rates will be strongly modified by intraparticle diffusion. T o ensure that this was not a factor, oxidation runs were made on smaller particles. A sample of 30- to 50-mesh size gave CO gas product concentrations as the standidentical COz ard 10- to 20-mesh, while a sample of 100- to 140-mesh (contained in a quartz bucket) showed a slightly faster initial reaction rate and a slightly slower sustained rate. T h e latter effect may have been due to the somewhat higher amount of coke fines arising from the screening process (see Experimental), causing an initial, enhanced reactivity, rather than a true diffusional effect. Comparison of the Weisz criterion of diffusional effect with our reaction indicates that intraparticle diffusion might be important only a t the early stages of oxidation, in agreement with our experimental findings. For example, in the most severe case of oxidation in air a t 480' C., the maximum rate of oxygen diffusion in 10- to 20-mesh particles is about 3 X 10-6 mole of 0 2 per cubic centimeter per second (27) ; under the same conditions, we estimate an initial reaction rate of about 5 X 10-6 mole of carbon per cubic centimeter per second, and a sustained rate a t 30% converted of about 1 x 10-6 mole of carbon per cubic centimeter per second. Additional evidence that intraparticle diffusion was not ratelimiting under our oxidation conditions came from inspections of partially oxidized samples. If oxygen diffusion were controlling, oxidation of the coke would occur preferentially on the outside of the particle and move toward the center. Yet our catalysts visually appeared black even u p to 90% coke burnoff. Further, an attrition sieving performed on a 10- to 20-mesh sample of the same catalyst reported in Table I1 but after partial oxidation, gave 4.3% carbon in 60- to 100-mesh and 5.5y0 carbon in the fines through 100 mesh. These data show that the outside of the particle (fines) still contained appreciable carbon and the inside of the particle (60 to 100)
+
+
202
l&EC PROCESS DESIGN A N D DEVELOPMENT
was reduced in carbon content. Thus, we conclude that the carbon oxidation is occurring equally throughout the entire catalyst particle, and is not intraparticle diffusion-limited. Oxidation Runs, Water given off by the reaction decayed faster than C o t and C O . Cumulative moles of hydrogen and carbon reacted were obtained by measuring the areas under concentration-time curves similar to those shown in Figure 1. A plot of the mole fractions converted with time is given in Figure 2. Clearly, the hydrogen in the coke is being converted faster than the carbon. Significantly, almost all the hydrogen reacts in 10,000 seconds under these run conditions, while about 25% of the carbon still remains. The varying slopes of these curves imply a changing ratio of hydrogen to carbon being consumed with time. Calculations reveal that the first material reacted has a n average formula near CH2; as reaction continues a more carbon-rich coke is left. Haldeman and Botty ( 8 ) , oxidizing a 3% coked silica-alumina catalyst, also found the hydrogen to react faster than the carbon and a n initial ratio corresponding to CHZ. Dart et ~ l .( 4 ) also report a higher hydrogen oxidation for oxidation of coked cracking catalysts; their data in the same temperature range show about 0.5 mole fraction H converted a t 0.1 mole fraction C converted and 0.9 H converted a t 0.5 C converted, in close agreement with results given in Figure 2. I n comparing catalyst weight changes during oxidation with cumulative product gases evolved, a serious discrepancy was noted in material balance. This is shown in Figure 3, where the solid curve is the actual data and the dashed line represents a material weight balance. Clearly, the sample weight was systematically high. T h e difference in weight is attributed to a n uptake of oxygen by the sample during reaction. This is a considerable amount and varies with the progress of the reaction and the reaction conditions, showing a maximum a t about one-third carbon conversion. Snow (76) has reported a similar effect with carbon blacks, which he describes as a residual solid enriched in oxygen. T h e nature of this oxygen addition is discussed below. I t is not simple, reversible oxygen adsorption, as proved in the following special experiment. A sample was partially oxidized a t 425' C., then nitrogen was substituted for the air; no further weight change nor oxidation ensued. Upon heating the sample in the nitrogen flotv up to 650' C., Con, CO, and HzO were liberated. (Mole fractions were 0.23, 0.51, and 0.26, respectively.) T h e catalyst weight
0 W
l-
a W > Z
0 0
O0 V '
2'
'
'
4
'
'
6
'
8'
'
'
IO
'
T I M E , 1000 S E C .
Figure 2.
Carbon and hydrogen conversion with time
'
I2
(3
I (D
m
0
J v)
a (3
I-
o
SPRING
Figure 3.
'
WT. LOSS, M G .
T I M E , 1000 SEC.
Product loss vs. spring loss
Figure 4.
change exactly corresponded to the evolved gases, showing no oxygen was desorbed. O n carbons, chemisorbed oxygen can only be removed as oxides of carbon (19). A very rapid oxygen adsorption could account for the initial sample temperature rise observed. Kinetics of Carbon Oxidation. We have established that mass-transport effects can be ignored, except for the very beginning of the reaction. Also, the carbon associated with the coke reacts more slowly than the hydrogen. We must, consequently, deal with two separate and parallel, gas-solid reactions, which may or may not interfere with each other. Considering first the carbon oxidation reaction, it seems reasonable that this would involve a surface reaction (77). We will make use of a model of coke deposit in silica-alumina catalyst proposed by Haldeman and Botty (8), who picture the coke as a discrete, finely divided, randomly dispersed phase, located within the interparticle spacing between the ultimate support particles. If we assume the basic coke particles to be spherical, the following kinetic relationships apply a t constant temperature (see Appendix for details) :
+
where m c is the carbon present a t time t; ( C o r CO) is the total carbon oxide concentration; 0 2 is the oxygen partial pressure which is raised to the nth power; So is the original coke surface area; a, is the mole fraction of carbon converted; k , is the intrinsic surface reaction rate constant; and p is a dimensional conversion factor. Equation 1 states that the surface reaction is proportional to the product ( C o r CO) concentration, which is an actual point rate. T h e rate is also proportional to the carbon surface area left and the oxygen pressure to the nth power. Since the carbon surface diminishes during reaction by virtue of the carbon being consumed, the reaction rate also decreases with time, which is reflected in C O ) concentrations. Integration of Equation lower (COz 1 in terms of carbon conversion and time a t constant oxygen pressure yields Equation 2, sometimes referred to as the contracting sphere model.
+
+
'/3
k,t = 1
-
(1
- aJ1I3= fs
(a,)
(2)
Here, k , includes k,, the original surface area, the total moles of coke, and the oxygen pressure term. Thus, a test of surface reaction control for spherical particles requires a linear relationship between fa(a,)and time. Such a plot is shown in Figure 4, where it can be seen that a satisfactory fit is obtained,
Surface reaction correlation
except a t very small times. We are not concerned about the nonconformity of the correlation a t the beginning of the reaction because of the nonisothermal nature of the reaction and possible transitory mass-transfer effects here. An empirical factor has often been employed to correct for this initial accelerating effect (73). I n order to evaluate the effects of oxygen partial pressure, coke level, and temperature on the carbon reaction rates without recourse to tedious integration summations, Equations 1 and 2 were combined to obtain an integrated expression in terms of (CO2 CO) concentration and time (cf. Appendix) :
+
(cor
+ C0)'/2 = I - st
(3)
where I = constant
(4)
(5) w ois the initial carbon weight in the sample, and r , is the initial
+
coke particle radius. Thus, if the square root of the (COz CO) concentration is plotted against time, the slope of the
resulting line, s, will allow correlation of the oxygen partial pressure ( 0 2 term), and the temperature ( k , term). An example of such a plot is shown in Figure 5 ; agreement with theory was generally good except a t early times and the latest stage of oxidation, where oxidation of the final 10% of carbon seemed to be unusually slow. T o correlate oxygen pressure, the slopes of the best straight lines through the data as plotted in Figure 5 were calculated; the slopes were then raised to the 2/3 power and plotted against oxygen partial pressure. This assumes that n is identical to I-Le., first-order in oxygen. T h e straight line shown in Figure 6, which passes through zero, shows that the assumption of first order is justified-if n does not equal 1, a curve would result. Therefore, our results confirm that the reaction rate is first-order in oxygen partial pressure, in agreement with earlier published work (20). T h e majority of the runs were made a t 480' C. Including the effects of oxygen pressure, sample size, and gas flow rate, values of ( k c / r o ) were calculated from the slopes of the data CO)l/z us. t and use of Equation 5. T h e replots of (Con sults are summarized in Table 111. Assuming a coke particle radius of 50 A. (see Discussion), the avcrage value for the intrinsic surface reaction rate constant for the carbon oxidation reaction a t 480' C. is about 4 X mole per sq. cm. sec. atm. This is in fair agreement with a value of 1.4 x 10-11 mole per sq. cm. sec. atm. obtained by Effron and Hoelscher (6) a t the same temperature for oxidation of graphite samples
+
VOL. 6
NO. 2
APRIL
1967
203
I
I
I
I
I
Table 111. Oxidation Data Conditions. 13.3 mg. C sample (9.67, by wt. C on catalyst) 400 cc. STP/min. gas flow 800 nim. approximate total pressure s X 703, (kJr,,) X los, Run Temp., 02 Press., P.P.M.'IZ/ MoleslCc. NO. c. Atm. Sec.u Sec. Atm.
'0
P,\
O
71* 37 38 13c 69 81 48 46d 57
7.5 8 5 480 7 1 48 0 8 0 480 7.5 480 8.1 480 8.4 7.9 480 480 0.040 6.8 Av. 7 . 8 k O . 5 52 450 0.222 0.90 2.38 53 425 0.222 0.32 0.82 a s is slope of a (Con CO)'!2 us. t plot (Figure 5). * Sample contained '/s normal charge. c Gas flow rate of 840 cc./min. d Glass bucket used.
0
\
OO
2
0.219 0 216 0 216 0 135 0.164 0.112 0.062 0.061
2.84 5 80 4 40 3 14 3.16 2.02 0.87 0.77 0.32
+
I
I
I
I
4
6
8
I O
I
480 480
TIME, 1000 SEC.
Figure 5. Carbon reaction correlation with gas concentration ,041
'
,
,
,
,
,
,
,
'i
.03 -
,M
,
,
.02-
N
u)
.05
-0
.IO
02
Figure 6.
.I 5
.20
.25
PRESS., A T M .
Carbon reaction correlation with
I .3
0 2
1.35
1.40
1000 /
1.45
1.50
T
Numbers in blocks refer to run numbers
Figure 7. Carbon activation energy
in oxygen. However, recent data published by Essenhigh et al. (7) for pure carbon samples indicate a calculated rate constant about 10-fold larger under comparable conditions. Finally, the temperature variation of the reaction was determined by a plot of the logarithm of the slope to the 2//3 power (equivalent to log k,) against the reciprocal of the absolute temperature, for runs a t constant oxygen pressure. As shown in Figure 7, a straight-line Arrhenius plot was obtained between 425' and 480' C. Reaction rates a t lower temperatures were too slow, and a t higher temperatures too fast to follow accurately with the experimental setup employed. T h e activation energy for the reaction was about 40 kcal. per mole of carbon, in agreement with other workers. For example, Weisz (27) found 38, Effron and Hoelscher ( 6 ) quote values of 35 to 40 for carbon oxidation, and Essenhigh et al. (7) employ a value of 40. Kinetics of Hydrogen Oxidation. The kinetics of the hydrogen reaction is more complex; it does not fit a surface reaction model analogous to that of the carbon. In analyzing the data, it was noted that the hydrogen reaction exhibited characteristics of diffusion kinctics, but the data did not fit
well any of the known standard forms. Therefore, a diffusion model specific for the reactions under consideration was developed from the salient features already detailed. T h e proposed model, involving coupled reactions of both carbon and hydrogen in the coke, is presented in Figure 8. A uniform carbon and hydrogen distribution throughout the basic coke particle is assumed a t the start of oxidation. Now, since the hydrogen reacts faster than the carbon, two interfaces will develop as reaction proceeds, the outer being the carbon interface and the inner the hydrogen interface. As a consequence, two distinct solid phases coexist, an outer layer of carbon only and an inner unreacted core of coke, still containing carbon and hydrogen. Because of the intrinsically faster hydrogen reaction, it will be completely consumed a t some point, leaving a residual carbon core only, which in turn will eventually disappear. Further, we assume the hydrogen reaction is controlled by diffusion of oxygen through the outer carbon layer to the inner coke core. Thus, we have a diffusion process with a double moving boundary. Based on Fick's laws of diffusion, the rate of transfer of material through a spherical surface layer is ( 2 )
204
I&EC PROCESS DESIGN A N D DEVELOPMENT
. 4 , ,
CARBON -AIR INTERFACE (C REACTION)
/
,
,
,
,
,
~
m
-
/COKE-CARBON
REA
0
2
4
8
6
10
T I M E , 1000 S E C .
Figure 8. reaction
Model for simultaneous carbon and hydrogen
Figure 9. Hydrogen diffusion correlation A.
B. (TcTH)
Rate = kl rc
ti
rc = ro(l
(7)
z d
TH
(8)
w
where ro is the original coke particle radius. Since diffusion is assumed to be rate-controlling, the hydrogen conversion is proportional to the diffusion rate; thus, Equation 6 becomes
a
- -dt
kz
[
(1 (1
- ac)I/3 = ro(1 - CYH)1’3
]
- CYc)1J3 * (1 - aH)1/3 - CYC)1J3 - (1 - aH)1’3
=
3/2
[(I)
-
(1
-
a H ) 213
1-
‘“s
(1
-
P m 0
cn
(9)
This differential equation is not readily solvable in terms of hydrogen conversion due to the coupled reactions-Le., the carbon reaction alters the boundary conditions of the hydrogen reaction. However, a partial integration yields,
kd
0.94 0.37
(6)
- TH
where T , and rH are the respective radii of the carbon and hydrogen interfaces, and kl is a constant which includes the diffusion coefficient. T h e interface radii can be expressed in terms of the mole fraction converted :
daH
CYH = CYH =
-W
x
C
Figure 10.
REMAINING, I - a c
Sorbed oxygen vs. carbon remaining
Discussion
0
daH
=
fD(aH)
(lo)
T h e integral in Equation 10 was evaluated by numerical integration and subtracted from the first term of the righthand side to yield f D ( a H ) . A plot of the latter against time should yield a straight line if the proposed diffusion model holds. T h e correlation obtained is illustrated in Figure 9. An excellent fit results, except again for the initial stages of reaction. We may now inquire whether the added oxygen detected during reaction is related to the oxygen diffusion through the carbon layer according to the model discussed above. T h e relative volume of the carbon layer present a t any time is simply equal to the fraction of carbon left minus that of coke left-i.e., (1 - a,) - (1 - C Y H ) or C Y H - cyc. When this value is calculated and compared with the sorbed oxygen, the result is very striking (Figure 10). T h e dashed line represents the carbon layer volume normalized to the milligrams of oxygen scale. T h e agreement with the experimental data curve both in general shape and position of the maximum supports the proposed model-the observed weight gain is predominantly due to the oxygen diffusing in the carbon layer. T h e total oxygen in the carbon layer (luring oxidation corresponded to an average of 1 oxygen atom per 3.7 carbon atoms.
An important difference should be mentioned in comparing the oxidation rates of coke deposited on porous supports and carbon black or graphite (unsupported). I n the latter cases, an appreciable increase in surface area is observed (77, 76) during the early stages of oxidation, and the oxidation rate proportionally increases (7 7). [Haldeman and Botty report a similar phenomenon on “extracted” coke ( g ) . ] O n the other hand, we have not observed this to occur for the coke deposited on our support, a t least as far as the carbon oxidation rates are concerned. O u r rates continue to diminish during the progress of oxidation. I t is not known a t present whether this difference is due to the nature of the coke deposited or the influence of the support on the oxidation, although our kinetic data correlation seems to argue against the latter alternative. Perhaps the already highly dispersed coke particles are not capable of developing further surface area. I t may be considered axiomatic that a carbon-oxygen complex is a prerequisite to carbon oxidation. Walker and coworkers (72, 78) have detailed several types of carbon-oxygen complexes on graphite. We cannot presume that such a requirement would not be necessary in coke oxidation. That we can account for the oxygen uptake during oxidation as residing predominantly in the carbon layer does not preclude the possibility of a carbon-oxygen complex. Indeed, the fact VOL. 6 NO. 2
APRIL 1967
205
that this oxygen is desorbed mainly as C 0 2 and CO and not as molecular oxygen as discussed earlier, would seem to suggest that the mechanism of oxygen diffusion may involve exchange with carbon-oxygen complexes within the carbon skeletal layer. I n addition to internal oxygen, there is evidence of adsorbed oxygen on the external surfaces of the basic coke particles before oxidation (70, 77). This was revealed in a special treatment with hydrogen a t reaction temperature after the volatilization period, whence a small weight loss was observed and water was evolved. The weight loss corresponded to about 0.07 mg. of oxygen per mg. of carbon, corrected for support loss (the latter being almost within experimental error). Assuming the hydrogen reacted with a monolayer of oxygen atoms adsorbed on the initial coke surface, a n average coke particle size can be estimated. Using a value for the oxygen atomic area of 16 A. (76) and a coke density of 1.6 grams per cc. (8), an average diameter of about 100 A. was calculated. This would correspond to a specific surface area of 350 sq. meters per gram of coke for the 9y0coke compared to surface areas of 102 and 237 obtained on the 770 extracted coke (see Experimental Section). T h e latter may have aggregated somewhat because of the severe caustic treatment. Coke deposit on silica-alumina catalyst has a turbostraticgraphic type character (75), similar to that exhibited by carbon blacks; the latter can exist in clusters a few hundred angstroms in size (3). Haldeman and Botty (8) estimate particle sizes of less than 100 A. for a 37, coke deposit on silica-alumina cracking catalyst. X-ray studies (3) of carbon blacks have shown the basic particles to consist of small, randomly orientated crystallites; the latter are made up of hexagon layer planes stacked parallel to each other. Similar results were reported for coke on a silica-alumina cracking catalyst (7), the major component being associated with condensed systems of fused aromatic rings. For such a configuration, we may envision the internally sorbed oxygen as being positioned between the carbon skeletal layers (the hydrogen having been previously reacted away) of the crystallite-1 oxygen atom interacting with 1 hexagon ring of each layer. This arrangement would give a C/O ratio of 4 in the carbon shell in fair agreement with the average value of 3.7 found experimentally. Then, diffusion may occur by an exchange j u m p mechanism between neighboring oxygen atoms with a net inward flux of oxygen atoms. Since reaction of hydrogen a t the coke-carbon interface boundary is limited by the oxygen diffusion rate through the carbon layer, water will be generated a t a rate equal to the oxygen diffusion rate. T h e water molecules produced must, of course, counterdiffuse to the exterior of the particle to be removed, thus competing for adsorption sites with oxygen. Water in the carbon layer cannot be distinguished from adsorbed oxygen and is included in the latter value. Water migration may also involve a jump mechanism between neighboring adsorbed water molecules, or more simply, hydrogen atoms. T h e countercurrent fluxes of oxygen and water will set up natural concentration gradients through the carbon shell layer, a basic premise of Fick's laws of diffusion. Appendix.
where m, is the carbon presei~t,moles. l&EC
CY,,
the following
For an average coke formula of CHo.b, w, is equivalent to 0.96 W,,and by substituting into Equation A3, differentiating Equation A2 with respect to time, and combining with Equation A l , we obtain dac _ dt
--
ks(l
- CY,)~'~
where k, =
3.13 k,M, (O?)n ToPC
Integration of A4 yields where A , is a n integration constant which is not zero owing to the initial deviation of the reaction from isothermicity, and CY,) represents the function of CY, shown. I n order to analyze the rate data without resort to laborious integrations of the carbon oxide concentration profiles with time to obtain CY^, we incorporate the concentration values directly, thus,
+ C0)dt
m, = Po 'J(C0. 0
or
dm - = - - --dmc dt
P(CO2
+ CO)
(-47)
dt
Combination of Equation A7 and the derivative of A2 yields
and elimination of dol,/dt by incorporation of A4 and A5 gives
(C02
+ CO)
=
woks ~
(1
MCP
- CYp3
Finally, taking the square root of Equation A9 and combining with A6 to eliminate a,, and inserting Equation A5 for k,, gives
(COz
+ CO)"2
=
1.77
Thus, a plot of (COZ having a slope, s, of
(02)n'Z (1
- A,)
-
+ CO)1'2 us. t should give a straight line
Carbon Surface Reaction
T h e rate of oxidation of carbon is assumed to be controlled by a surface reaction ; thus
206
I n terms of fraction of carbon converted, relations obtain for spherical coke particles,
PROCESS DESIGN A N D DEVELOPMENT
from which the following relations between the reaction variables and measured slopes are: Oxygen pressure. Carbon charge.
0 2
Coke particle size.
Y,
w,
= cc s2
(A121 (A12B)
cc
(A12C)
s-2'3
Rate constant.
k , cc s2I3
(A1 2D)
For a flow rate of 400 cc. (STP) per minute, P equals 3.07 X 10-10 mole/p.p.m. sec., and using a coke density of 1.6 grams per cc. (8), the slope becomes s =
5.78 X 106 ( z u , ) ” ~ (!?>”:‘
t
=
ffH
= mole fraction carbon converted = mole fraction hydrogen converted
P Pc
reaction time, sec.
dimensional conversion factor, moles/p.p.m. sec. = coke density, g./cc. =
(02)3n12
Literature Cited Acknowledgment
T h e author gratefully acknowledges the technical dexterity of J. Tabacek and W. Faust, who carried out the experimental portion of the investigation. Nomenclature
(Cot
+ CO) == function gas concentration, p.p.m. of given in Equation 10 CYH
fD(ffH) fd(ffC)
F
Z kc
= = = = = = = = = = = = = = = = = = = = =
kD
k,
ki k2 MC m, m, n (02)
rc TH
ro S
S S O
W.3
wo
function of a , given in Equation 2 gas flow rate. cc. STP/min. intercept of Equation 3, p.p.m.l/z carbon surface rate constant, mole/sq. cm. sec. atm. rate constant for oxygen diffusion, set.-' constant in Equation 2, set.-' constant in Equation 6, moles/cm. sec. constant in Equation 9, set.-' molecular weight of carbon, g./mole carbon remaining, moles total carbon reacted, moles order of rate with respect to 0 2 oxygen partial pressure, atm. coke particle radius, cm. radius a t coke-carbon interface, cm. initial particle radius, cm. slope of Equation 3, p.p.m.1/2/sec. carbon particle surface area, sq. cm. initial coke surface area, sq. cm. initial weight of carbon, g. initial weight of coke, g.
(1) Appleby, W. G., Gibson, J. W., Good, G. M., Division of Petroleum Chemistry, ACS Preprint 5, No. 4, B-71 (September 1960). (2) Barrer, R. M., Phil. Mag. 35, 802 (1944). (3) Biscoe, J.,Warren, B. E., J. Appl. P l y . 13, 364 (1942). (4) Dart, J. C., Savage, R. T., Kirkbridge, C. G., Chem. Eng. Progr. 45, No. 2, 102 (1949). (5) Eberly, P. E., Jr., Kimberlin, C. N., Jr., Miller, W. H., Drushel, H. V., IND.ENG.CHEM.PROCESS DESIGN DEVELOP. 5,193 (1966). (6) Effron, E., Hoelscher, H. E., A.I.CI2.E. J . 10, 388 (1964). (7) Essenhigh, R. H., Froberg, R., Howard, J. B., Ind. Eng. Chem. 57, No. 9, 32 (1965). (8) Haldeman, R. G., Botty, M. C., J. Phys. Chem. 63, 489 (1959). (9) Haldeman, R. G., Botty, M. C., private communication. (10) Hall, J. W., Rase, H. F., IND.ENG.CHEM.PROCESS DESIGN DEVELOP. 2, 25 (1963). (11) Hoynant, G., Comfit. Rend. 259, 2827 (1964). (12) Laine, PI’. R., Vastola, F. J., Walker, P. L., Jr., J. Phys. Chem. 67, 2030 (1963). (13) Massoth, F. E., Hensel, W. E., Jr., Zbid., 63, 697 (1959). (14) Mische, R. A., Smith, J. M., Ind. Eng. Chem. Fundamentals 1, 288 (1962). (15) Shiraski, T., Ozaki, A,, Shokubai ( T o k y o ) 4, 88 (1962). (16) Snow, C. W., Division of Rubber Chemistry, 148th Meeting ACS, Chicago, September 1964. (17) Snow, C. W., Wallace, D. R., Lyon, L. L., Crocker, G. R., Proceedings of Third Conference on Carbon,” p. 279, Pergamon Press, New York, 1959. (18) Walker, P. L., Jr., Rusinko, F., Jr., Austin, L. G., Aduan. Catalysis 11, 134 (1959). (19) Ibid., p. 140. (20) Zbid., p. 155. (21) Weisz, P. B., Goodwin, R. D., J . Catalysis 2, 397 (1963).
RECEIVED for review May 16, 1966 ACCEPTED October 21, 1966 Division of Petroleum Chemistry, 151st Meeting, ACS, Pittsburgh, Pa., March 1966.
UPGRADING OF RESIDUES BY COMBINATIONS
OF RESIDUE DESULFURIZATION AND VIS BREAK1NG B. K. S C H M I D AND HAROLD B E U T H E R
Gulf Research €3 Development Co., Pittsburgh, Pa. 75230
ROWING
emphasis on air pollution from sulfur oxides has
G led to restrictions on the sulfur content of residual fuels in some areas, and future trends may mean more such limitations ( 3 ) . Since much of the present residual fuel will not meet probable future expectations of limits on sulfur content, it will become necessary to remove sulfur from these fuels. A recent survey of various methods of removing sulfur indicates that catalytic hydrodesulfurization has the most promise ( 4 ) . Much of the current research in hydrodesulfurization of residual fuels has been directed toward decreasing process costs. One potential method for decreasing the over-
all costs or improving economics of residual fuel desulfurization is to increase the production of higher-value distillate fuels as by-products. Thus, process schemes which include a residual hydrodesulfurization step and a mild thermal cracking (visbreaking) step have been investigated. Residue hydrodesulfurization, as exemplified by the Gulf HDS process, upgrades residues by decreasing sulfur content, viscosity, and pour point, and increasing API gravity ( 2 ) . T h e extent of the decrease in viscosity and pour point is a function of the sulfur removed. I n the processing of many vacuum residues, the sulfur content must be decreased to a very VOL. 6
NO. 2
APRIL 1967
207
