Literature Cited Anderson,T. S., Coull, J.,A.l.Ch.E. J.. 16, 542 (1970). Dassau, W.J., Wolfgang, G. H.. Chem. Eng. Progr., 59,4, 43 (1963). DeWasch, A. P., Froment, G. P., Chem. Eng. Sci.. 26, 629 (1971). Eigenberger, G.. Chem. Eng. Sci., 27, 1909 (1972). Gilles. E. D.. Chem.-lno.-Techn., 40. 469 (1968). Lunde, P. J.', Kester, F. L.. lnd. Eng. Chem.. Process Des. Develop., 13, 27 (1974). McGreavy, C.. Cresswell, D . L.. Can. J. Chem. Eng., 4 7 , 6 , 583 (1969).
Marek, M., HlavaEek, V., Chem. Eng. Sci., 21,(193 (1966). Marek, M., HlavaEek, V., Chem.-lng.-Techn., 40,1086 (1968). Paynter, J. D., Dranoff, J. S.,Bankoff, S. G., Ind. Eng. Chem., Process Des. Develop., 9,303 (1970). Sinai, J., Foss, A.S..A.l.Ch.E. J., 16, 658 (1970). Vortmeyer, D., Jahnel, W.. Chem. Eng. Sci., 27, 1485 (1972). Wicke. E., Padberg, G., Chem.-Ing.-Techn., 40, 1033 (1968).
Received for review M a y 9, 1973 Accepted M a r c h 6,1974
Nonisothermal Kinetics Studies of the Hydrodesulfurization of Coal Alfred L. Yergey,* Frederick W. Lampe,' Marvin L. Vestal,2 Alan G. Day, Gordon J. Fergusson, William H. Johnston,3 Judith S. Snyderman, Robert H. E ~ s e n h i g hand , ~ James E. Hudson Scientific Research lnstruments Corporation, Baltimore. Maryland 27207
The application of nonisothermal kinetic methods to the hydrodesulfurization of coal is described. The results indicate that with only a few exceptions, the hydrogen-induced release of sulfur as hydrogen sulfide can be described adequately, in a practical kinetic sense, by five processes. We have designated the five processes in terms of the source of sulfur as Organic I , Organic II, Pyrite, Sulfide, and Organic I I I. Preexponential factors and activation energies are determined for each. The importance of the back-reaction of H2S with partially desulfurized coal is discussed and it is shown that the most important of the back-reactions in the inhibition of H2S release is reaction with iron. It is also shown that Organic I I I is most likely formed by back-reaction of H2S with carbon in the partially desulfurized coal.
Introduction The current recognition of sulfur oxides as a major air pollutant has generated a renewed interest in an old industrial problem, namely, the desulfurization of coal and coke prior to their combustion. It is generally accepted that sulfur occurs in coal in three forms: (1) in organic chemical combinations and, therefore, as an integral part of the coal substance itself; ( 2 ) as pyrites and/or marcasites; and (3) as sulfates, which generally are of calcium and iron. In view of the nature and abundance of the organic sulfur in coal, it is readily apparent that chemical methods of sulfur removal, in which the coal substance itself is chemically altered, present the only practical means of efficient desulfurization prior to combustion. Thus knowledge of the chemical kinetics of this very complex chemical transformation is desirable from the practical-process point of view and because of any fundamental information concerning structure that may be derived. The desulfurization of coal and coke has been well-studied over the past half-century beginning with the investigations of Wibaut and Stoffel (1919) in Europe and Powell (1920) in the United States. An excellent review of work done on this subject prior to 1932 is given by Snow (1932), while a summary of work done to about 1945 is presented by Thiessen (1945). Reports of desulfurization studies have continued to appear to the present time (Batchelor, et al., 1960; Blayden, 1958; Brewer, 1949; Chapman, 1955; Chowlhury, et al., 1952; Curran, et al., 1958; Fuchs, 1951; Gray, et al., 1970; Mahmoud, et al., 1969; Mason, 1959; Zielke, et al., 1954, 1970). The general picture of hydrodesulfurization emerging
' *
Department of Chemistry, Pennsylvania State University, University Park, Pa. 16802. Department of Chemistry, University of Utah, Salt Lake City, Utah 84112. Deceased. Department of Fuel Sciences, Pennsylvania State University, University Park, Pa. 16902.
from these studies is that the organically. bound sulfur and the sulfur contained in pyrites are evolved from the coal or coke principally as hydrogen sulfide at processes that begin to be of significance at about 300°C. Presumably ferrous sulfide is a by-product of the pyrite decompositions and this sulfide is thought to decompose in the range of 400-600°C. Sulfates are thought to be reduced to sulfides a t temperatures below 500°C. Apparently a significant fraction of the original organically bound sulfur is retained in the solid product as chemically bound sulfur that is at least partially removable by hydrogen treatment a t temperatures of 700-1000°C. The reaction of the hydrogen sulfide with the partially desulfurized coal is rapid and unless this back-reaction is somehow inhibited practical desulfurizations are precluded. This back-reaction of the hydrogen sulfide may also be an important source of the strongly bound sulfur in the solid partially desulfurized product. A common feature of all the early investigations of coal desulfurization is the use of the classical, time-honored method of chemical kinetics in which the extent of conversion is measured as a function of time for a series of experiments, each series at a different but constant temperature. Unfortunately, this method is severely limited when applied to such a complex conversion as the hydrodesulfurization of coal because of the occurrence of a number of parallel chemical reactions, both at the desired reaction temperature and during the time the sample is being brought to the reaction temperature. The result is that one observes, at each temperature measurement, the overall end result of a superposition of reactions. A method in which the velocity is measured at a constant and known heating rate of the solid sample, and which circumvents this difficulty, has been developed recently by Juntgen and coworkers (Hanbaba, et al., 1968; Juntgen, 1964; Juntgen and Traenckner, 1964; Juntgen and van Heek, 1968; Peters and Juntgen, 1965; van Heek, et al., 1967a,b). The method permits evaluation of the usual kiInd. Eng. Chern., Process Des. Develop., Vol. 13,No. 3, 1974
233
netic parameters of activation energy and preexponential factor for the various parallel reactions occurring. It has been shown (van Heek, et al., 1967a), that the method yields kinetic parameters that are in good agreement with those obtained by classical methods when it is applied to the decomposition of basic magnesium carbonate. This paper describes applications of the method to the hydrodesulfurization of coal.
gas. If only the j t h process was producing HzS, the temperature dependence of I would be given by ( 5 ) and (6), which predict that a plot of I os. T is a single peak whose shape and position depend upon ko,j, E,, n,, and M . However, in actuality, a number of superimposed processes produce HzS in the hydrodesulfurizatiop of coal so that the observed temperature dependence of I is given by
Theory of the Method Let us depict a hydrodesulfurization process in coal by the chemical reaction iAJ-S)solid
+ H1
-
+ H,S
(A,)=lld
(1)
where A,-S represents one of the variety of ways in which the sulfur is bound into the solid and AJ represents the pertinent site of this bonding after removal of the sulfur. The rate of this reaction is given by expression 2, namely
-d[AJ-S1 dt
(-':fl')
= k ,~HJ[A,-SP' =
k,[A,-SI"' ( 2 )
in which [A,-SI is the concentration of solid reactant sites of type j, [Hz] is the molecular hydrogen concentration, which remains essentially constant, n, is the reaction order, and k , is the specific reaction rate of the j t h process. The reaction rate constant for the j t h process may be written in the usual Arrhenius form shown by
k , = k , ]e-El RT
(3)
Consider now the situation in which the temperature is not held constant but rather is varied linearly according to the expression in (4), where M is a constant.
dT-- M dt
(4)
Substitution of (3) and (4) into (2) and integration of the result, following Juntgen, is a valid procedure provided the elementary processes establishing thermal equilibrium are rapid compared with the heating rate M . Assuming this to be true, one obtains for the j t h process the results given by
in which V, is the volume of H2S evolved by the j t h process at time t, VO,, is the volume evolved by this process when the reaction is complete, LY is a proportionality factor relating the volume of H2S to the concentration of A,-S, i . e . , [A,-SI = LY(VO,, - V I ) ,and the other symbols are as described previously. In the derivation of (5) and (5a) it is assumed that Ej/RT >> 1. In our experiments the concentration of H2S in a stream of hydrogen gas passing through a bed of coal particles is measured mass spectrometrically as a current of ions, I, whose magnitude is related to dV/dT by the relationship
In (6), Q is the flow rate of gas, G is the weight of the coal sample, and fl is a known proportionality factor relating the intensity I to the mole fraction of H2S in the flowing 234
Ind. Eng. Chern., Process Des. Develop., Vol. 13,No. 3,1974
It may be shown that at the maxima, Tm,j, predicted by ( 7 ) , relationship 8 holds for each process. Equation 8 is
independent of order, n,, provided that EJ/2n,RT >> 1. Therefore, from ( 8 ) , provided this latter condition is realized, we may, in principle, determine ko,, and E, from the temperatures of the peak maxima measured as functions of the heating rate M . This approach is relied upon in the determination of the kinetic parameters ko,, and Ej in this work. Examples appear in Figures 5 and 6. The shape of each H2S evolution peak, described by its width at half-height, is a function of the reaction order n,. Knowledge of ko,, and E, permits calculation by ( 5 ) of the corresponding peak widths as a function of the reaction order; comparison with the experimentally observed peak widths then allows determination of the pertinent reaction order. The foregoing treatment is based on the assumption that the hydrogen sulfide produced does not react with the partially desulfurized coal. This is most likely not true, as has been mentioned in the Introduction. We may approximate the effect of this back-reaction of hydrogen sulfide on each evolution peak by imagining that there is only one HzS-producing process and by including in the reaction scheme the process
+c
H?S
C-s +H*
(9)
when C represents partially desulfurized coal and kb is the specific reaction rate of the back-reaction. In describing the kinetics of this complex situation it is necessary to include explicitly the geometry of the solid reactant and the flow velocity of the sweep gas. For the present application, we are concerned with a static bed of solid reactant having cross section A and depth d, so that for a volume flow rate of sweep gas of Q cm3/min the average residence time 7 of a volume element of gas within the bed is given by
Ad
VBed
Q
Q
7=-=-=-
d
v
In (lo), V, is the bed volume and u is the average linear flow rate of the sweep gas. The rate of change in the concentration of HzS with residence time 7 in the bed may be written
-4H1'dT
- k,[A,-SI"' -kb[H,S]
where the first term on the right-hand side is the rate of formation of HzS from the particular process under consideration and the second term is the rate of disappearance of H2S due to the back-reaction with the partially desulfurized coal. For simplicity we assume this backreaction to be first order in HzS and that the concentration of sites for back-reaction is sufficiently large that it
may be taken as constant. The term [Aj-SI is the instantaneous concentration of sulfur containing species of the j t h type and is assumed to be uniform over the bed of coal. Integration of (11) over the residence time T , with the boundry condition that [HzS] = 0, for T = 0, gives
For a nonisothermal experiment the rate constants k , and k b , being functions of temperature, are also time dependent. However, for moderate heating rates and reasonably short residence times, the change in temperature (AT = M,) is sufficiently small that the variation in the rate constants during the time T may be neglected. The concentrations [A,-SI and [HzS] are functions of real time t (just as k , and k b are) since the sweeping out of the product gas H2S depletes the sulfur content of the bed. Explicitly, this may be written
Combination of (12) and (13) yields
where we denote explicitly that the rate constants are functions of real time t. Substitution of (4) into (14) and integration gives [A - ~ ] - c n , + ~ J - [ A ~ - S & - ( " J=+ ~ ) I
where [Aj-SIo is the initial concentration of reactant Aj-S in the bed. After a series expansion of the exponential, this integral can be evaluated numerically using tables of the exponential integral or it may be evaluated analytically with the aid of the approximation to the exponential integral employed by Juntgen. The results, however, are very cumbersome and of very limited practical use. The case which is of most interest to the present experimental study is that of a very fast back-reaction or a very large k b . For this situation, we may assume k b T >> 31. Then, with the Arrhenius form of the rate constants (3) and following Juntgen, we may integrate (15) and ( E a ) to yield
Comparison of (16) and (16a) with (5) and (5a), respectively, shows that the dependence on temperature is identical, with E, - E b in (16) and (16a) replacing E, in (5) and (5a). Hence, the same data evaluation techniques, described previously for the case of an absence of back-
reaction, may be applied when there is a fast back reaction. When this is done, the "activation energy" obtained is actually the activation energy differences E, - Et,; the preexponential factor so obtained is k o , , / k b , o T , and, since T is known, one obtains the ratio k O , , / k b , o . Experimental Section Apparatus a n d Procedure. A block diagram of the nonisothermal kinetics laboratory is shown in Figure 1. Hydrogen or helium, a t flow rates selected for an experiment, are passed over a finely ground (100-200 mesh) sample of coal placed in the center of a quartz reaction tube inside the furnace. The tube is 0.5 cm i.d. and has 15 cm of its length inside the furnace. Temperatures within the furnace are increased linearly with time and are maintained to within a few degrees of a preselected increase rate by differential controls within the temperature programmer. The rates of temperature increase are continuously variable from about l"C/min to 10O0C/min. Within measurement accuracy there are no radial temperature gradient8 within the furnace, and axial temperatures are constant within the central 12 cm. Temperatures are unaffected by gas flow rates under 400 scc/min. Combinations of mesh size, samples size, and gas flow rate are selected so that carrier gas flow remains relatively constant during an experiment, even at the peak of product evolution from the coal bed. Gas residence times in the coal sec. bed are on the order of 7.5 x The sweep gases used in this work also serve as chemical ionization (Field, 1968; Munson and Field, 1971) reagent gases since their concentrations are much greater than the gaseous reaction products. Source pressures in the mass spectrometer are maintained at about 1 Torr, so that ions formed originally by electron impact undergo multiple ion-neutral collisions. In the case of hydrogen, this results in H3+ being the most abundant ion in the source, as shown in eq 17 and 18. Ions derived from the desulfurization reaction product gases are formed by a proton transfer from H3+ as shown in eq 19. HP+e-H,++2e
(17)
H, +H,-H,++H
( 18)
&++ H2S --t H,S+ 4- H?
(19)
Because the proton affinities of the product gases are higher than that of H2, the product gases are preferentially protonated, an effect which results in a great sensitivity increase compared with normal electron impact ionization. An analogous situation involving charge transfer occurs when He is used as the reagent gas. Ions leaving the source region are mass-analyzed in a quadrupole mass filter (Brubaker, 1961; Paul, 1953) and their intensities are amplified by an electron multiplier. Signals from both the multiplier and a thermocouple are input to the digitizer and mass spectrometer controller. This device allows the monitoring of the intensities of up to nine peaks in the mass spectrum by stepping the mass filter to appropriate focussing voltages. The rate at which the controller steps through the mass settings is selected to be comparable to the furnace heating rate, so that intensity us. temperature plots result. Materials. The ten bituminous coals used in this study are described in Table I. The first three columns show the agency from which the sample was obtained along with the mines and seams whioh yielded the samples. A partial summary of the proximate analyses performed, according to ASTM procedures, on each sample, including nominal sulfur level, per cent volatiles, and fixed carbon, is given in the last three columns. We wish to thank Mr. Jack Ind. Eng. Chem., Process Des. Develop., Vol. 13, No. 3, 1974
235
VENT
TUIPERANRE MEASUREMENT
GAS + IN
FURNACE REACTOR
. >-
).
CHEWCU IONIZATION QUADRUPOLE
TIM,M
DIGITIZER
*...Mn,
PAPER TAPE PUNCN
MASS SELECTOR
MASS SPECTROMETER
I COMPUTERIZED FURNACE CONTROLLER
HANDLING SYSTEM
1 Figure 1. Block diagram of the nonisothermal kinetics laboratory.
Simon of the Illinois State Geological Survey for providing the first five samples and the U. S. Bureau of Mines for the last five samples. To ensure that the lack of microscopic homogeniety in the coals did not influence the results, samples for a nonisothermal experiment were selected by riffle separation of the desired coal. Portions of the riffled sample were then pulverized and sieved to select the 100-200 mesh fraction. Coke used in the experiments to determine the kinetic parameters of the strongly bound sulfur in the partially desulfurized solid coals was prepared by placing approximately 3-g samples of a riffled coal 40-60 mesh in a furnace maintained at a temperature of 600°C with a hydrogen or helium flow. After pyrolysis, the coke was pulverized and sieved to the desired mesh fraction. Pyrite and marcasite, FeS2, samples were obtained from the U. S. Bureau of Mines. These materials were pulverized and sieved prior to their use in the characterization of these two forms of sulfur occurring in coal. The hydrogen and helium used in these studies were purchased from the Matheson Co. and were used as obtained from the cylinders. Results and Discussion HzS Evolution Curves. Typical HzS evolution curves from bituminous coal, obtained in the hydrodesulfurization of two nominally 5% sulfur Illinois coals (ISGS: 3211.60F, Little Dog Mine and ISGS: 28-1, Will Scarlet Mine), are shown in Figures 2 and 3, respectively, in which the HzS intensity, plotted as H3S+ ion ( m / e = 35), in arbitrary units, is plotted us. the absolute temperature. One may recognize in these figures that the overall HzS evolution curves are superpositions of a number of HzS evolution peaks, in accord with ( 7 ) . In Figure 4 is shown a typical HzS evolution curve obtained in the hydrodesulfurization of iron pyrite. We assign the two processes producing HzS in the hydrodesulfurization of iron pyrite to the successive reactions FeS, + H P FeS + H,S (20)
-
and
236
Ind. Eng. Chem., Process Des. Develop., Vol. 13,No. 3, 1974
1
600
..
N T E N
..
I I
T y
40
'
200
.
100
*
. . :
SW
*
wo
7w
1100
1
3
TEMPERATURE O K
Figure 2. HsS intensity SRI Coal No. 5 , Hz carrier, 100-200 mesh, 26"/min.
FeS
+ H,
Fe
+
+ Ha
(21)
with (20) exhibiting a sharp peak at -790°K and (21) a broader peak a t -890°K. Assuming that in coal the processes (20) and (21) will occur independently of other HzS-producing reactions, their occurrence in coal hydrodesulfurization is seen in Figure 2 to be responsible for the small but clearly discernible peaks at 790 and 880°K. Similarly, but not so unambiguously, we assign in Figure 3 the peak at -780°K to (20) and the shoulder at -890°K to (21). Comparison of Figures 2 and 3 then indicates that in the coal designated as SRI 4, some process not present in the coal designated as SRI 5 produces an HzS evolution curve having a peak at -810°K (Figure 3). This peak at 810°K has been observed in only one other of the coals examined, namely the coal designated as SRI 7 in Table I.
6400
:. I
.
.
H,S 48QQ
I
N
1 E N I I
1
32M
1640
300
iw
700
Po0
IlW
1300
TEMPERhTURE .1'
750
950
TEMPERATURE O K
Figure 3. H3S intensity SRI Coal No. 5, Hz carrier. 100-200
mesh, 30"/min.
Figure 4. HzS intensity from pyrite, Hz carrier, -200 mesh, loo/ min.
The low-temperature peak appearing at 685°K in Figure 2 is common to all ten coals studied (Table I), appearing always in the range of 660-740°K. The exact position of the low-temperature peak depends, of course, upon the heating rate employed, in accord with the discussion leading to eq 8. In addition, consideration of the shapes of this peak obtained from all the coal samples leads to the conclusion that it represents a superposition of at least two hydrodesulfurization processes producing H2S. The exact position of the observed single peak at a given heating rate will also necessarily depend upon the relative contributions of the unresolved processes to the overall H2S evolution. Since the relative contributions of the two or more processes depend upon the relative amounts of the pertinent sulfur types in the coal, the exact position of this peak varies somewhat from coal to coal. This variation is always small compared with the variation due to heating rate. We have not succeeded in resolving the lowtemperature H2S evolution peak into its components at the heating rates used. Since we know two or more processes are involved, and since it is generally accepted (Batchelor, et al., 1960; Blayden, 1958; Brewer and Ghosh, 1949; Chapman and Jones, 1955; Chowlhury. et al., 1952; Curran, et al., 1958; Fuchs, 1951; Gray, et al., 1970; Mahmoud, et al., 1969; Mason, 1959; Ode, 1963; Powell, 1920; Snow, 1932; Thiessen, 1945; Wandless, 1955; Wibaut and Stoffel. 1919; Zielke, et al., 1954, 1970), that organically bound sulfur and pyritic sulfur are released as H2S a t the lowest temperature in hydrodesulfurization, we designate this low temperature H2S evolution peak as Organic I and Organic 11. It has been shown by Juntgen and coworkers that application of eq 5-8 to peaks that comprised more than one process leads to kinetic parameters that are averages of the individual processes. Finally, common to all coals studied, we have observed a higher temperature H2S evolution exemplified in Figures 2 and 3 by the broad peak a t -930°K which must have an onset near 800°K. We have called this high-temperature peak, common to all the coals studied, Organic 111. In a few of the experiments, there appears to be H2S production at even higher temperatures, an example being the peak at about 1200°K in Figure 3. However, we do not observe this peak in all the coals, nor is it reproducible in successive experiments with the SRI 4 coal. Clearly, we do
not understand the appearance of this peak in Figure 3; it may be an experimental artifact, but even if not, we believe that it has no bearing on our conclusions reached relative to those evolution peaks and processes described that are common to all the coals examined. According to the foregoing discussion, the evolution of H2S in the hydrodesulfurization of the ten coals listed in Table I can be described almost completely by sulfur release from what we have called, in order of increasing temperatures, Organic I, Organic 11, pyrite, sulfide, and Organic 111. Sotable exceptions to a complete description are: (1) the evolution of H2S by a process occurring in a temperature range intermediate to the pyrite and sulfide evolution peaks, a process that appears most prominently in the coals designated as SRI 4 and SRI 7; and (2) a higher temperature H2S evolution (Tpeak= 1200°K) that is not reproducible and may be an experimental artifact. Determination of Kinetic Parameters. As discussed in the section dealing with the theory of the method, measurement of the temperature of each evolution peak as a function of the heating rate permits evaluation by (8) of the Arrhenius preexponential factor and activation energy pertinent to the process involved. In Figures 5 and 6 are shown plots according to (8) for the H2S evolution peaks arising from Organic I and I1 and pyrite. respectively. The coals employed in the study of the variation of temperature maxima with heating rate were those designated in Table I as SRI 1, SRI 4, SRI 5, and SRI 10. The exponential points shown in Figure 5 were obtained from all four coals and it is clear from this plot that the experimental error is greater than the variation of the Organic I and I1 in the coals. The rate data shown in Figure 6 for the formation of H2S from pyrite were obtained in hydrodesulfurization of pure iron pyrite. Similar plots were obtained for H2S evolution from sulfide and Organic 111. In accord with eq 8, the slopes of plots exemplified by Figures 5 and 6 lead immediately to the apparent activation energies shown in the third column of Table 11, while the intercepts yield the preexponential factors shown in the fourth column. The second column of Table I1 gives the reaction orders of the various processes obtained, as described earlier, from a study of the half-height width of the H2S evolution peaks. The true activation energies and frequency factors of Ind. Eng. Chem., Process Des. Develop., Vol. 13, No. 3, 1974
237
Table I. Description of Coal Samples Used in Hydrodesulfurization Studies SRIC No. Source Mine Bed 1 2 3
4 5 6 7
8 9 10
Illinois ISGS: 1-1 Illinois ISGS: 5-1 Illinois ISGS: 5-1 Illinois ISGS: 28-1 Illinois ISGS: 32-11.60F Ohio USBM: 205 Maryland USBM: 106 Ohio USBM: 107 Pennsylvania USBM: 109 Kentucky USBM: 110
% Nominal
%
% Fixed
sulfur
Volatiles
carbon
Crown
Seam 6
5
36.2
44.7
No. 21
Seam 6
1
33.8
54.6
Northern Illinois Will Scarlet Little Dog Stanley
Seam 2
2.5
33.7
42.5
Davis
5
37.6
43.2
Seam 6
4.5
40.8
44.1
Pittsburgh No. 8 Franklin
3
38.1
49.6
3
20.5
62.8
Pittsburgh No. 8 Lower Kittanning(B)
3.5
37.2
50.4
1
27.9
64.2
4
31.3
46.5
Royal Cravat Greenwich No. 8
No. 14
Shamrock
Table 11. Kinetic Parameters for Hydrodesulfurization Reactions
Reaction order with respect to sulfur species
+ + + + + + + + + +
Org I Org I1 Hf -+ H2S Pyrite H2 6 HzS Sulfide Hz -,H2S Organic I11 HP-* H2S Fe(s) H2S + FeS Hz C (9) H2S-,H2S (adsorbed) CaO(s) H2S --* C a s H20
ko (apparent)
Eapp
2
22.0 42.1 25.1 38.1
1 1/2
1 2 1 1 1
1.8 2.5
... ...
-16.00
2 In
.20.0(
1.i4
I.&
i.is 1/Tm (
.ll.rn
(A) Tn'
.12.w
.13.W
.I4.U 1.45 li
10
1.52 -3
1.66
1.21
1
Figure 5 . Best fit curve for Organic I and 11, E = 22.0, log ko = 6.28: 0,SRI Coal No. 1; A , SRI Coal No. 4; 0 , SRI Coal No. 5; X , SRI Coal No. 10.
the various H2S producing processes are obtained only after consideration of the effects of the back reaction of H2S with the coal, as described in the next section. Effect of Back-Reaction of H2S. As mentioned in the Introduction, it is a generally accepted fact that the back reaction of H2S with partially desulfurized coal is a rapid process. Moreover, as shown in the discussion of the theory of the nonisothermal method, the occurrence of a rapid back reaction of H2S will result in the kinetic parameters obtained by (8) being apparent parameters related to the 238
105
...
.lO.W
.12.w
x
x 107
...
...
4.W
-lo.w
2 x 108 1 . 3 X 10" 9 . 7 x 108 1 . 3 X loll 6 . 5 X 104 2 . 3 X lo8 4 . 7 x 10'J
...
.s.w
4.W
ko
22.0 42.1 43.1 56.1 18 32 38
x 106
1 . 3 X 10"
0
I"(+) 1m
E, kcal/mol
Ind. Eng. Chem.,Process Des. Develop., Vol. 13, No. 3,1974
1.26 l/Tm (
1.30
1.35
0
* 10 -1 1
Figure 6. Best fit curve for pyrite, E = 42.1. log k o = 11.11.
true values by the relationships (22) and (23), namely
E=E,,,+Ea
(22)
and
ko = koapp/kwT (23) where Eb is the activation energy of the back-reaction and kbo is the preexponential factor of the back-reaction. We have attempted to assess the relative importance of the H2S back-reaction with the carbon and iron constituents of coal by conducting nonisothermal experiments in
1.20
1.40
2.00
1.80
Figure 7. Arrhenius-type plot of the H2S absorption by iron.
which 1000 ppm of H2S in He is passed over solid samples of charcoal, iron filings, and coal chars produced from SRI 1 and SRI 2 coals with continuous monitoring of the H2S concentration in the gas. The results are as follows. (1) The onset of H2S reaction with the iron filings occurs a t -800°K and H2S removal from the gas stream is essentially complete at 800°K. (2) The onset of H2S reaction with charcoal occurs at -850°K; the reaction reduces the H2S content of the gas stream to less than 100 ppm a t -1000°K. (3) The behavior of the coal chars, both of which contain iron from pyrite desulfurization in the preparation of the chars, is intermediate to that of the iron filings and the charcoal. As may be seen in Figures 2 and 3, reduction of iron pyrite in coal to iron sulfide and subsequently to iron by (20) and (21) does not occur until temperatures are reached a t which removal of H2S from Organic I and I1 is complete. Therefore, reaction of the H2S produced from Organic I and I1 with iron will not occur; neither will it occur with carbon, since the onset of the H2S reaction with carbon lies at temperatures above which evolution from Organic I and IT is complete. We conclude, therefore, that the backreaction of H2S is unimportant in the evolution of H2S from Organic I and I1 and that the apparent kinetic parameters in Table I1 for this source of sulfur are identical with the true values. In a similar manner, we may recognize from Figures 2-4, and particularly from 4, that the onset of H2S formation from iron sulfide via (21) does not occur until the conversion of pyrite to sulfide uia (20) is essentially complete. Therefore, the back-reaction of H2S from the pyrite in coal with iron and carbon will also not be of importance. Possibly some back-reaction with FeS will occur but we are not able to assess its relative importance and we make the therefore somewhat arbitrary conclusion that the kinetic parameters for pyrite in Table I1 are not influenced by back-reaction and are the true values. The H2S evolution processes from sulfide and Organic 111 will definitely be influenced by back reactions of H2S, this back reaction being comprised of reaction of H2S with iron and carbon and other constituents in the coal. However, since the reaction with iron is so predominant at temperatures up to -950" (the temperature of the peak corresponding to Organic III), we conclude that the kinetic effect on evolution from sulfide and Organic I11 is due primarily to back-reaction with iron; the kinetic parameters obtained from the plots typified by Figures 5 and 6 are thus apparent values which need correction according to (22) and (23).
We have made this correction by measurement of the kinetic parameters kbO and E b for reaction of H2S with iron filings, in experiments in which 1000 ppm of H2S in He is passed over a sample of iron filings. Consider the reaction to be H,S
+ Fe
-
FeS + H,
(24)
and assume, for simplicity, that it is first-order in H2S and that the Fe is in vast excess. Then we may write for the rate of disappearance of H2S from the gas stream
For a given temperature of the solid sample, we may integrate (25) to obtain the concentration of the H2S in the gas leaving the bed, uit.
[H,S]
= [H$],
e-'&
e--E6'R77
(26)
where fH2S10 is the concentration of H2S in the inlet gas and 7 is the average residence time in the bed. For a nonisothermal experiment in which the heating rate is sufficiently low that the temperature change during the time T is small, k b o e ( - E b / R T ) in (26) may be considered as a constant. Assuming that this condition is met we obtain
and upon taking logarithms of both sides, we have finally
Hence, by plotting semilogarithmically the quantity In [(H2S)o/(H2S)]us. 1/T, the kinetic parameters E b and k b o may be determined from the slope and intercept, respectively. Since T is known from the flow rates and solid bed size, k b O is determined. A typical plot of the data for take-up of H2S by iron filings, according to (28), is shown in Figure 7 . Two temperature regions are clearly discernible in this figure: (1) a region from onset to about 600°K in which chemical reaction is rate-determining and, (2) the region above 600°K in which diffusion of H2S is rate-determining. Such behavior is typical of solid-gas reactions in which the gaseous reactant is brought into contact with the surface of the solid. In the back-reaction case of interest here, H2S is produced by hydrodesulfurization within the pores of the solid and, hence, diffusion to reactive sites is not the limiting factor of interest. The area of interest,is that limited I nd. Eng. Chem., Process Des. Develop., Vol. 13, No. 3, 1974
239
by chemical reaction and, from the slope and intercept of the reaction-limited region of Figure 7 , we obtain the values for the back-reaction shown in Table 11. These values were then used in eq 22 and 23 to obtain true kinetic parameters for the H2S evolution from sulfide and Organic 111. Although not used in the correction of apparent desulfurization kinetic parameters to true values, we have also determined, by the method just described, kinetic parameters for the back reaction of H2S with carbon (charcoal) and with calcium oxide. These are shown also in Table 11. The reaction orders for removal of sulfur in pyrite and Organic I11 shown in Table I1 are interesting in that they are not unity but rather are M and 2, respectively. The former can be rationalized by supposing the existence of an equilibrium between iron pyrite, iron, and adsorbed sulfur, and a rate-determining step that involves reaction of hydrogen with the adsorbed sulfur. In the latter case, a reaction order of 2 indicates that, in the rate-determining step of sulfur-evolution from Organic 111, two sulfur-containing species must be somehow involved, as, for example, in the schematic reaction below
H --H
s s
#
I
S
S
S
Further speculation is not possible since we do not know the type of carbon-sulfur binding involved. Some further comment is warranted concerning the nature of the form of sulfur in coal that we have designated as Organic 111. In our experimental studies of H2S take-up by charcoal (essentially mineral-free) described above, the solid product of the reaction was obviously a sulfur-containing solid in which sulfur was bound to carbon. Nonisothermal hydrodesulfurization studies were conducted in this solid with the H2S evolution exhibiting an onset at -700°K and a broad peak at -900°K. This behavior is strikingly similar to that of Organic I11 from coal hydrodesulfurization and suggests that the Organic I11 form of sulfur may not be present in the original, but rather is
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Ind. Eng. Chern., Process
Des. Develop., Vol. 13, No. 3, 1974
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Received f o r reuieu: June 21, 1973 Accepted March 6, 1974 This research was supported by the Environmental Protection Agency under Contract No. P H 86-68-65, CPA 70-50, and 68-020206. In addition, reports of this work are available from the National Technical Information Service as P B 185882 or from Scientific Research Instruments as reports SRIC 68-13. SRIC 50-14, and SRIC 71-15.
