Reaction of Sulfur Dioxide and Hydrogen Sulfide with Porous


Girard A. Simons, and Wilson T. Rawlins. Ind. Eng. Chem. Process Des. Dev. , 1980, 19 (4), pp 565–572. DOI: 10.1021/i260076a011. Publication Date: O...
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Ind. Eng.

Chem. Process Des. Dev. 1980,

g = acceleration of gravity, cm sW2 Ho = bed height, cm H,f = bed height at incipient fluidization, cm U = superficial gas velocity, cm s-l U d = superficial gas vellocity at incipient fluidization, cm s-l Literature Cited

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Adkins, H., Peterson, W. P., J . Am. Chem. Soc., 53, 1512 (1931). Alessandrini, G., Cairati, L., Foczatti, P., Viiia, P. L., Trifir6, F., J. Less-Common Met., 54, 373 (1977). Boreskov, G. K., Kolovertnov, 13. D., Kefeli, L. M., Plyasova, L. M., Karachiev, L. G., Matikhin, V. N., Popov, M. I., Dzis'kd, V. A,, Tarasova. D. V., Kinet. Ketal., 7, 144 (1966). Bossi, A., Leofanti, G., Moretti, E., Giordano, N., J . Mater. Sci., 8, 1101 (1973). Botton, R. J., Chem. Eng. Prcy. Symp. Ser. No. 101, 66, 8 (1970). Cairati, L., Di Fiore, L. (to Euteco Spa), Italian Patent 21426 (Mar 21, 1977). Cairati, L., TrifirB, F. (to Euteoo Spa), Italian Patent 27409 (Sept 9, 1977). Cairati, L., Carbucicchio, M., Fluggeri, O., Trifir6, F., Stud. Surf. Sci. Cafal., 3, 279 (1979). Carbucicchio, M., J . Chem. Phys., 70(2), 784 (1979). Carbucicchio, M., Trifir6, F., J . Catal., 62, 13 (1980).

565

79,565-572

Dente, M., Poppi, R., Pasquon, I., Chim. Ind. (Milan), 46, 1326 (1964). Dente, M., Coliina, A., Chim. Ind. (Milan), 47, 821 (1965). Forsythe, W. L., Hertwig, W. R., Ind. Eng. Chem., 41, 1200 (1949). Jiru, P., Wichteriova, B., Tichy, J., Proc. 3rdInt. Congr. Catal. Amsterdam, 1, 199 (1964). Le Page, J. F., Cosyns, J., Courty, P., Freund, E., Frank, J. P., Jacquin, Y., Juguin, E., Marclily, C., Martino, G., Miquel, J., Montarnai, R., Sugier, A,, Van Landeghem, H., "Catalyse de Contact", Technip, Paris, 1978. Matsen, J. M.. Tarmy, B. L., Chem. Eng. Prog. Symp. Ser. No. 101, 66. 1 (1970). May, W. J., Chem. Eng. Prog., 55, 49 (1959). Pernicone, N., J . Less-Common Met., 36, 289 (1974). Stewart, P. S. B., Ph.D. Thesis, University of Cambridge, Cambridge, England, 1965.

Received f o r review April Accepted June

18, 1979 18, 1980

This work has been supported by Euteco Spa and by Consiglio Nazionale delle Ricerche (Roma).

Reaction of Sulfur Dioxide and Hydrogen Sulfide with Porous Calcined Limestone Girard A. Slmons' and Wilson T. Rawlins Physical Sciences Incorporated, Woburn, Massachusetts 0 180 7

A simple theory is developed to describe the mass transport and heterogeneous chemistry which occurs when either SO2 or HzS reacts with calcined limestone. The reactant gas diffuses into the porous calcine and is consumed on the interior surface. The basic heterogeneous rate constants for HzS and SOz with CaO have been inferred from existing laboratory data. Our results indicate that the reaction of HS , with CaO proceeds almost as fast as that of :SOzwith CaO. Hence, any procedure which utilizes limestone removal of SOz is potentially capable of removing HzS at approximately the same rate.

I. Introduction The high-temperature removal of HzS by calcium-based sorbents such as limestone and dolomite may provide a useful cleanup technique for fluidized bed coal gasification processes (Keairns et al., 1976). A similar method for SOz removal in fluidized bed coal combustion processes has already been shown to be a promising technique (Case et al., 1978). In both processes, high-temperature capability of the sorbent permits in situ cleanup of the producer gas and the resulting long gas-solid contact time in the fluidized bed allows high sorbent utilization. Limestone is primarily calcium carbonate (CaCOJ, or calcite. Upon heating, the calcite decomposes into calcium oxide (CaO) and C02. The CaO is referred to as calcined limestone or simply calcine. The HzS reacts with calcine to form calcium sulfide (Cas) whereas SOz reacts with calcine to form calcium sulfate (CaS04). The ultimate level to which sulfur can be removed depends upon the equilibrium properties of the gas mixture; however, the efficiency of sorbent utilization is determined by the overall rate of the gas-solid readion. This work is concerned with evaluating the kinetic rates with which SO2 and H2Sreact with calcine. The practical utility (of the limestone technique cannot be evaluated on the baisis of the chemical kinetics alone. In the case of SOz removal, there is clear evidence (Borgwardt and Harvey, 1972; Hartman and Coughlin, 1974) that the diffusion of SO2 through porous CaO may be rate controlling. In addition, it has been proposed oi96-43051a0111 i9-0565$oi.o010

(Hartman and Coughlin, 1976,1978) that the diffusion of the reactant gas through the solid sulfur deposits within the porous structure is, under some circumstances, rate limiting. The whole process is complicated by the fact that the porous structure itself is dependent upon the degree of calcination and sulfation. These complicated processes must be described in order to predict the sulfur removal rate in a diffusion-controlled environment. Transport theories have been developed to describe diffusion-limited processes in SO2 cleanup. Hartman and Coughlin (1976, 1978) represent the porous CaO by a pattern of small spherical grains surrounded by pores. The SOz diffuses through the pores to the grain where it reacts with the CaO. As the CaO reacts, CaS04 is formed on the outside of the grain and further reaction requires that the SOz must diffuse through the solid phase CaS04. The net reaction is then limited by both the diffusion of SOz through solid CaS04 and by the diffusion of SOz through the pores. The reaction ultimately terminates because the porosity of the stone is reduced to zero and prevents SOz from reaching the grain. The grain theory possesses four parameters: the grain size, the gas phase diffusion coefficient, the activated diffusion coefficient, and the heterogeneous rate constant. The grain radius was obtained by X-ray diffraction analysis and the gas phase diffusion coefficient (D) was taken from the experimental results of Campbell et al. (1970). The heterogeneous rate constant ( K ) was evaluated from experimental data in the limit of low fractional conversion 0 1980

American Chemical Society

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( x = 0.04) and the activated diffusion coefficient (D,) was adjusted until the theoretical SO2 removal rates agreed with measured values. Basic measurements by Borgwardt and Harvey (1972) were reduced in a much more direct manner. The size of the calcine sample was reduced until the SO2removal rate became independent of particle size. This resulted in a kinetically controlled situation in which the rate constant could be measured directly. Unfortunately, the rate constant measured by Borgwardt and Harvey (1972) at 1033 and 1253 K is 30 times slower than that inferred by Hartman and Coughlin (1976) at 1123 K. This implies that either the grain theory does not adequately describe the structure of porous calcine or the four parameters introduced into the grain theory are not uniquely determined. A single pore structure model, transport theory, and heterogeneous rate constant are needed which can reproduce both the kinetic and diffusion-controlled data reported by both observers. Such a theory is described and verified in this paper. In order to develop a theory capable of predicting the H2S or SO2 removal by fully calcined limestone, we must fully understand the porous structure, the kinetics of the reaction, the diffusion of the reactant gas through the porous CaO, and the diffusion of the reactant gas through the sulfur deposits. As a first step in the development of such a theory, we have concentrated on the kinetics and gas phase diffusion through the porous CaO. The theory is valid only at early time, before sulfur deposits become appreciable. A clear evaluation of the kinetics for the reaction of CaO with both SO2 and H2S is established. Future theoretical efforts must incorporate the diffusion through the sulfur deposits, thereby describing the late time reactivity and the conversion efficiency of the CaO. 11. Kinetics It is generally agreed (Borgwardt and Harvey, 1972; Hartman and Coughlin, 1976) that the rate (kj) of the reaction CaO + SO2 + 1/202 CaS04

-

is directly proportional to the SO2 concentration in an oxygen-rich environment. Similarly, we assume that the rate of the reaction CaO + H2S CaS + H20

-

is proportional to the H2S concentration. Hence, we express the rate of the reaction of species j with CaO as k I . = k~ I. P. I where pi is the partial pressure of the species j in atm, kj is the kinetic rate in grams of CaO removed per second per square centimeter of surface, and k , is the heterogeneous rate constant in grams of CaO per second per square centimeter per atmosphere of species j . This rate constant may also be expressed as Ka, in units of grams of species J' per second per square centimeter per atmosphere of species j . Hence, we note K,,j = kaj/Zj where Z j is the solid mass released per reactant gas consumed. For the particular reactions discussed above Zj = 56/Mj where Mj is the molecular weight of species j . A third common form for the rate constant is in units of centimeter per second, K , where

0 . 5 ~

0

n

E

n

0.4-

0.3-

5

g

0.2-

t

?

2

0.1-

0 2.0

I

I

I

2.6

3.2

3.8

Pore Radius, log rp

I

4.4

5.0

[81

Figure 1. Volume distribution in calcine before SO2 reaction.

is the density of species j in grams per cubic centimeter and p . is in atmospheres. (Note that K does not represent a surCace recession rate.) The measurements of Borgwardt and Harvey (1972) indicate that the rate constant for the SO2reaction at 1250 K is K = 0.22 f 0.05 cm/s pj

Hence, it follows that

Kaj = 1.4

X

g of S02/s cm2 atm SO2

and

kaj = 1.2 X

g of CaO/s cm2 atm SO2

Throughout this paper, the heterogeneous rate constant is expressed as kaj in units of grams of CaO per second per square centimeter per atmosphere of species j . A pore structure and transport theory is described which, when used with the rate constant quoted above, is consistent with the early time SO2 removal rates of Borgwardt and Harvey (1972) and of Hartman and Coughlin (1974) in both the kinetic and diffusion-controlled limits. The theory is also applied to describe H2S removal and the basic heterogeneous rate constant for the H2S reaction is inferred from the H2S data base. 111. Pore Structure The removal of H2S and/or SO2 by CaO is, in part, controlled by the diffusion of the reactant species through the porous structure. Since the porous structure does influence the rate of diffusion (Thiele, 1939), it is necessary to adequately describe that structure. Hartman and Coughlin (1976, 1978) treat the porous structure as a cluster of identical grains (spheres) of CaO. This is equivalent to assuming that all of the pores are identical in size. However, the porosity data of Hartman and Coughlin (1974), illustrated in Figure 1,indicate a much broader distribution of pore sizes. We will approximate the data of Hartman and Coughlin (1974) with a linear profile integrated porosity a In r p and proceed to demonstrate that the structure of porous CaO can be treated in the same way as that of porous carbon char. The pore structure theory has been derived in detail elsewhere (Simons and Finson, 1979; Simons 1979a); the salient features of the final formulation are summarized here. The pores are assumed to be cylindrical tubes of length I , and radius r,. Each pore that reaches the exterior surface of the calcine sample is depicted as the trunk of a tree. The number of tree trunks whose radius is between

Ind. Eng. Chem. Process Des. Dev., Vol. 19, No. 4, 1980

rp and r, + dr, is denoted by 4aa2g(r,)dr, where a is the radius of the calcine sample and g ( r ) is the number of pores per unit surface area. The pore chstribution function g(r,) is given by

material. The quantity c j denotes the gaseous mass fraction of the species j . The net mass diffusion of the species j into a single pore tree of the structure is denoted by mpj and is given by mpj = - p S j ( a r

where 0 is the porosity and p is In (rmar/rmin). The minimum pore radius is a few hundred angstroms (see Figure 1)and the maximum pore radius is related to the radius of the calcine sample by where KOis approximately the pore aspect ratio (KOi= 5). Each tree trunk of radius rp is associated with an internal structure whose surface area is given by

2)dcjl dx,

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(5)

61/3rmin

where the internal surf,aceof the pore tree, St(rp)is greater than the surface of the trunk by the factor rp/rmin. Hence, each trunk is associated with a very large network of pores whose structure is depicted as an ordinary tree or river system. While this pore structure has been well substantiated for porous char (Simons and Finson, 1979; Simons, 1979a), there is very little information about the structure of porous calcine. The theory predicts that the total porous volume between rmhaind an arbitrary value of rp is given by integrated porosity

0:

11

rp2g(rp)drp0: In

(4)

where the quantity p G is the total gas density within the pore, Djis the self-diffusion coefficient, and the gradient of the species j is measured at the surface of the char sample. This gradient is a function of the heterogeneous reaction within the pore structure and will adjust itself such that mpjrepresents the mass depletion rate of species j within a pore tree whose trunk radius is rp' The total mass depletion rate Mp, of species j is obtained by integrating nipj over the pore distribution function

27rKor,3(1 - 0)

St(rP)=

587

(rp/rmin)

This fundamental scaling law has been compared to data reported by Hartman and Coughlin (1974) in Figure 1. While the l/r: distribution does not yield perfect agreement, it does suggest ithat, to a first approximation, the theory of char structure may be applied to calcine. One of the key paraimeters in the pore structure is the minimum pore radius, rmin. The value of rmhis related to the internal surface area [The internal surface area is the integral of rg(rp)drp(Simons and Finson, 1979; Simons, 1979a)l by 20 'S, = (3) PPsrmin

where ps is the density of the nonporous calcine (3.32 g/ cm3) and spis in units of area/mass. The data illustrated in Figure 1 indicate that the smallest pores are approximately 100 to 200 A in radius. Equation 3 predicts that these values of rminarle consistent with internal surface areas of the order of 2 to 5 m2/g. These predicted values of the internal surface area are similar to the values measured by Borgwardt and Harvey (1972). Hence, all indications are that the theory of char structure may be applied to calcined limestone. This structure is now used to develop a transport theory describing the flux of a reactant species into the porous calcine. IV. Transport Theory Following the transport theory of Simons (1979b),which describes the species dif'fusion and heterogeneous reactions which occur in porous char, consider the porous structure described in section I11 to be placed in a gaseous environment in which one of the gaseous species will undergo a heterogeneous reaction with the porous material. In particular, consider the reaction of species j with the

where 4aa2 is the exterior surface of the porous sample. The total mass removal rate of the porous structure, Mpc, is obtained by multiplying Mpj by the ratio of the weight of solid to reactant gas consumed (2,). To obtain a diffusive solution for m the diffusion of species j into a single tree trunk is baflnced by the heterogeneous reaction a t the walls of the trunk. Following Thiele (1939), the conservation of species j is expressed as d2Cj

d ~ =2 (2*rp)Kj

P&j(*r,2)

(6)

where x is measured from the open end of the pore and Kj is the heterogeneous rate measured in mass of reactant gas consumed per time-area. Before integrating the pore diffusion equation, it is necessary to specify the heterogeneous rate Kj. The heterogeneous reaction rate is assumed to be first order in the reactant species concentration. Following the definition of the rate constant in section I1 (7) Kj = K aJpJ . . = k a1. pI . / z j Denoting the total gas pressure by pG and the species mole fraction by X,, p j is expressed as Pj = PGX,

The species mass fraction is defined as XjMj cj

=

7

where M is the average molecular weight of the gas. Hence Kj may be written as Kj = kajcj

( 8)

kaj = kajp&/ZjMj

(9)

where While

La, is a weak function of cj through M , we treat

Raj as invariant with species concentration. Hence, eq 6 is a linear, second-order differential equation with constant coefficients. Equation 6 may be integrated to obtain dcj/dx which, when evaluated a t x = 0, yields m j . The details of the integration are given by Simons (1979&).The mass depletion rate within a single pore tree may be limited by the diffusion of the reactant species through the trunk of the tree, or by the kinetics occurring on the surface St(rp). These solutions are, respectively

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Des. Dev., Vol. 19, No. 4,

1980

Table I. ProDerties of Limestone Calcined at 1253 K Borgwardt and Harvey

and mpj

=

R,jCojSt

(11)

where coj is the value of c j at the exterior surface of the calcine. There is a critical pore radius rcjat which the diffusion limited solution for mpjis identical with that in the kinetic limit. Solving eq 10 and 11 for rcj,we note that mpj for r p I rcj is given by eq 11 and mp,for rp 2 rcj is given by eq 10. The integration of eq 5 then follows The integration of eq 5 is further complicated by the fact that Dj is controlled by Knudsen diffusion in the smaller pores and by continuum diffusion in the larger pores. The diffusion coefficient is given by elementary kinetic theory (Jeans, 1954) and is expressed as 1 D . = -),.V. I

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theoretical predictions

(1972)

3 1 J

where Aj is the mean free path and V jis the mean thermal speed. The mean thermal speed is expressed as

where kB is the Boltzmann constant and T G represents both the particle temperature and the temperature of the gas within the pore, which, by assumption, is in thermal equilibrium with the porous material. For sufficiently small radii, the relevant mean free path is the pore diameter and the diffusion equation is governed by the Knudsen diffusion coefficient D , 2Dkj = -Vjrp 3 In the continuum limit, Xi is the mean free path of species j undergoing collisions with all species within the system. Hence, the continuum diffusion coefficient Dcj, is a complicated function of the molecular weight and concentrations of all species. However, an approximate expression for Dcj is used for the gas mixtures considered here Dcj = [10(TG)1,7 g cm/s3]/p~

(12)

where TG is the gas temperature in K. This expression yields 0.14 cm2/s at STP. Since Dcj is basically viscosity/density, it is clear that this temperature dependence law for viscosity. is inferred from a The value of the pore radius at which Dkjis equal to Dcj is denoted by rDj 3Dci rDj = 2 vj

For all rp greater than rDj,the continuum diffusion coefficient is appropriate whereas the Knudsen coefficient is used for all r less than rDj. This physical limitation leads to additionaf expressions for mpjand further complicates the integration in eq 5. The detai!s are given by Simons (1979b) and the final solutions for Mpjand MpGare utilized without further discussion. To calculate the rate at which species J’ is removed by the porous calcine, we assume a value of the heterogeneous rate constant in units of grams CaO per second per square centimeter per atmosphere of species j . In addition, we specify the density of the nonporous calcine (3.32 g/cm3), the initial porosity, the initial internal surface area, and the concentration and molecular weight of species j .

UP,

S p

type

cm3/g

m /g

1 2 3 4 5 6 8 9 10 11

0.25 0.30 0.29 0.32 0.42 0.40 0.36 0.36 0.032 1.21

10.2 0.7 1.8 2.4 3.9 4.1 0.8 3.7 0.6 2.2

rmin,

A

40 1000 320 245 155 140 950 150 195 440

P,

e

g/cm3

0.45 0.50 0.49 0.52 0.58 0.57 0.54 0.54 0.10 0.80

1.81 1.66 1.69 1.61 1.39 1.43 1.51 1.51 3.00 0.66

Measured values o .e initial porositv and internal surface area are utilized where avaiiable. if unavailable, these properties are either calculated from known quantities or they are estimated. The heterogeneous rate constant is the only “free parameter” in the theory. Its value is adjusted to agree with measured values of the SO2 and H2S consumption rate. It must be emphasized that the theory is valid only at early time, before sulfur deposits within the porous calcine become appreciable. The method of analysis described above is applied to the SO2 data base in section V and to the H2S data base in section VI. V. SO2 Analysis Since the literature possesses a wider and better documented data base for SOz removal than for H2S removal, the theory is first compared to the SOz data. A series of well controlled experiments have been conducted by Borgwardt and Harvey (1972). They examined ten different types of calcined limestone over a wide range of porosity and average pore diameter. The porous volume (up) and internal surface area (sp)of each stone were measured. The porous volume is interpreted as a porosity (0) and density (p) via (Simons and Finson, 1979) P =

Ps(1 - 6)

and up

= 0/P

where ps is the density of the nonporous sample (3.32 g/cm3). The internal surface area and the porosity are uniquely related to the minimum pore radius rminby eq 3. Calculated values of r-, 0, and p are illustrated in Table I, together with the measured values of u and sp. Thus, the properties of the calcines are completefy characterized and the only undetermined parameter is the rate constant kaj. The SOz removal rates are quoted (Borgwardt and absorbed Harvey, 1972) in milligrams of SO3 (SOz + 1/202) by a 30-mg sample of calcine. The data were obtained for SOz removal a t 1253 K in flue gas at 1 atm. The particle size was 96 pm in diameter and the SOz concentration was 3000 ppm. Under these conditions, the reaction CaO

-

+ SO2 + 1/202 CaS04

is oxygen rich and the rate is directly proportional to the SO2 concentration. We have calculated the initial SO3 removal rate for all 10 samples for rate constants equal to g of CaO/s cm2 atm of SOz. 2X and 4 X Sample results are compared to data in Figures 2-4. The reactivity of Type 1 calcine is compared to data in Figure 2. The theory is in good agreement with data over the first 20 s of the reaction. At later time, the sulfur deposits clearly decrease the calcine reactivity and must

Ind. Eng. Chem. Process Des. Dev., Vol. 19, No. 4, 1980 - 1

I

589

8

/ :: 5 m'

I-

Theory must include Sulfur Deposition

[SO,]

TG

= 3000 ppm

=

1253' K

PG = 1 atrn TG

1253' K

1

PG = 1 atm a =48um

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0 ; 0

10

20 30 40 Exposure Time -Seconds

50

1

0

p

d

12-

e

:: 10m' 8-

B

Q m

% E

c

.= 5

4Data: Borgwardt & Harvey ( 1 9 7 2 ) Type 11 Calcined at 1253' K

10 0

Exposure Time -Seconds

Figure 3. Reactivity of type 11 calcine a t 1253 K.

be described in order to predict the late time sulfur sorption capability. However, the early time data suggests the following value for the kinetic rate constant

kaj

-

1.0 X

I

I

20 30 40 Exposure Time -Seconds

I

50

I I

Figure 4. Reactivity of type 4 calcine a t 1253 K.

surface (Borgwardt and Harvey, 1972). The test of the validity of our theory in this limit is illustrated in Figure 3. Again, the theory agrees with the data for a rate constant of about 1.5 X g of CaO/s cm2 atm of SO2. From the remaining eight types of calcine studied, the value of the kinetic rate constant inferred from the data varies between and 3 X g of CaO/s cm2 atm of SO2. An example of the higher value of the rate constant is given in Figure 4 for Type 4 calcine. Hence, our best approximation for the rate constant in the SO2 reaction with CaO is kaj = (2 f 1) X g of CaO/s cm2 atm of SOz or, in the units used by Borgwardt and Harvey K = 0.36(1 f 0.5) cm/s

6-

0

10

i

60

Figure 2. Reactivity of type 1 calcine a t 1253 K.

14

Data: Borqwardt & Harvey (1972) Type Calcined at 1253' K

'11

Data: Borgwardt & Harvey (1972) Type 1 Calcined at 1253OK

0

r/ /

to 1.5 X

g of CaO/s cm2 atm of SO2

which is within f25% of that obtained by Borgwardt and Harvey (1972) and discussed in section 11. Perhaps the most significant aspect of this data comparison is the fact that the reactivity of Type 1 calcine is diffusion limited (limited by SOz diffusion through the porous structure) (Borgwardt and Harvey, 1972). Hence, we have properly duplicated the diffusion limited reactivity of Type 1calcine using the rate constant obtained by Borgwardt and Harvey from the kinetically limited data. This offers further credibility to our pore structure and pore transport theory (sulfur deposition excluded). Just as the reactivity of Type 1 calcine was limited by the diffusion of SO2 through the porous structure, Type 11is limited by the chemical kinetics acting on the internal

which is reasonably consistent with that quoted by Borgwardt and Harvey (1972) ( K = 0.22 cm/s). It is not, at this time, our intention to determine the rate constant to a higher degree of accuracy. We wish only to emphasize that the present theory accurately simulates the early time data of Borgwardt and Harvey in both the kinetic and diffusion-controlled environment and the rate constant measured herein is consistent with that which Borgwardt and Harvey determined from the kinetically controlled data. Borgwardt and Harvey also measured SO2removal rates by Type 4 limestone a t 1033 K. The gas was again 1 atm flue gas with 3000 ppm of SO2. The calcination temperature was lowered to 1063 K and the subsequent internal surface area was 7.3 m2/g. The porosity of the sample was not quoted. We assume that the porosity was the same as that quoted in Table I for higher calcination temperatures. The corresponding theoretical predictions are illustrated in Figure 5. The rate constant appears to be between lo4 and 2 X lo4 g of CaO/s cm2atm SO2at 1033 K. Since Figure 4 illustrates that the corresponding rate constant for Type 4 calcine at 1253 K is about 3 X g of CaO/s cm2 atm of SO2, there appears to be about a factor of 2 difference between the rate constant a t 1033 K and that at 1253 K. This implies that the activation energy for the SO2reaction is approximately 10 kcal/mol. This value for the activation energy represents a relatively weak temperature dependence for the chemical kinetics. Equations 10 and 12 indicate that the temperature dependence of the diffusivity Dj is almost as important as that of the chemical kinetics. Hence, empirical efforts to

570

Ind. Eng. Chem. Process Des. Dev., Vol. 19, No. 4, 1980

a

8

sz

7

E 1 6

6 2

a

4

8

3

E z

f

2 1

,

o l

0

Data Borgwardt & Harvey (1972) Type 4 Calcined at 10630 K -

20 40 60 80 Exposure Time -Seconds

0

,

20

10

I

I

1

I

1

30 40 50 60 70 Exposure Time - Minutes

EO

90

100

Figure 7. Comparison with data of Keairns et al. a t 0.005 atm HzS.

I

I

Data: Keairns et al. (1973) 0 Denbighshire Limestone A Rockland Limestone 0 Lee Limestone 0 Pfizer Limestone ( A d a m )

100 [HzS] = 5000 ppm

Figure 5. Reactivity of type 4 calcine a t 1033 K. Downloaded by UNIV OF NEBRASKA-LINCOLN on August 26, 2015 | http://pubs.acs.org Publication Date: October 1, 1980 | doi: 10.1021/i260076a011

PG = 10 atm 0.5 .. I

0.4L

E

/

0.3-

0

.0565cm Data : Keairns et al. (1973)

Y)

a

Glasshouse Stone

i

s 0.2- P 24 E E

Data' Hartman & Coughlin (1976)

[SO,]

= 3000 ppm T G =1123OK PC = l a t m

---- 1144OK 1060°K 01 0

I

I

1

1

2 3 4 5 Exposure Time - Minutes

I

I

I

6

Figure 8. Comparison with data of Keairns et al. a t 0.05 atm HzS. -

1

0

5

lb

1k Exposure Time

io

25

3b

- Minutes Figure 6. Comparison of present theory to SOz data of Hartman and Coughlin (1976).

deduce the activation energy from laboratory data must be careful to separate the diffusive properties from the kinetic properties. An additional data base for SOz removal was obtained by Hartman and Coughlin (1974), which they later analyzed by grain theory (Hartman and Coughlin, 1976,1978). The derived value of the rate constant was 6.6 cm/s at 1123 K. This is a factor of roughly 20 higher than we obtain from Borgwardt and Harvey's data at 1253 and 1033 K and it is a factor of 30 higher than that quoted by Borgwardt and Harvey (1972). This discrepancy suggests that the parameters introduced into the grain theory may not be uniquely determined. For example, our eq 10 indicates that the sulfur sorption rate in a diffusion-limited environment scales as (k; Equation 12 predicts the value of D, at 1123 K to e 1.5 cm2/s, whereas Hartman and Coughlin use a value of D j that is a factor of 20 lower. (Figure 1 illustrates that half of the porosity is contained in pores whose diameter is sufficiently large to describe the gas diffusion coefficient via the continuum value.) Hence, it is possible that Hartman and Coughlin could have obtained the correct product of k , and Djwithout uniquely determining the value of the heterogeneous rate constant. The present theory, while valid only at early time, possesses only one parameter (the rate constant) which should be uniquely determined. The test of this statement is our ability to duplicate the early time data of Hartman and Coughlin using the rate constant 2 x g of CaO/s cm2 atm of SOz. To characterize the calcine, we use the quoted values of 8 = 0.54 and rmin;= 150 A (sp = 3 m2/g). The theory is compared to data in Figure 6 for particles of diameter 0.0565,0.0900, and 0.112 cm. The

t

Data: Keairns et ai. (1973) Tvmochtee Dolomite

1

.9-

u .7-

P P

2

.5-

+

P .3-

10: 0

10 20 30 40 Exposure Time -Seconds

I

50

Figure 9. Comparison with data of Keairns et al. a t 0.15 atm HzS.

predicted CaO conversion rates are in good agreement with the early time data. Hence, the present approach has uniquely determined the rate constant and it has been shown to be in agreement with a large body of data obtained by two independent observers.

VI. H2S Analysis Calcined limestone will react with HzS to form Cas and HzO, via the reaction CaO + HzS Cas + HzO

-

The kinetic rate constant controlling this reaction may also be obtained using the present theory. The data of Keairns et al. (1973) offer sufficient data base for this exercise. Since the properties of the stone are not characterized, we will assume 8 = 0.54 and rmin= 150 A. These properties are consistent with those measured for calcined limestones prior to SO2 reaction (Borgwardt and Harvey, 1972; Hartman and Coughlin, 1974). The theory is compared

Ind. Eng.

to data in Figures 7-9. Figure 7 illustrates the reactivity of four different stones in 0.005 atm of HzS at 1144 K. The initial rate of sulfur scrption suggests that the heterogeneous rate constant is

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k,,

-

g of CaO/s cmz atm of HzS

Figure 8 demonstrates the same behavior for Glasshouse stone in 0.05 atm of Hz!3 at 1144 K. Since Glasshouse stone is only 50% calcite (Keairns et al., 1973),we assumed that the porosity was reduced to 0.27 but r,,, was still 150 A. This has the net effect of reducing the internal surface area from 3 m2/g to 1.5 m2/,g. However, the internal areas were not measured and theae characteristics of the stones are not confirmed. The rate constant inferred by the 0.05 atm of HzS data base is consistent with that determined at 0.005 atm of HzS. ThLe data comparisons are carried to 0.15 atm of HzS at 1118 K in Figure 9 and, in this case, a slightly lower value of the rate constant is inferred. The sensitivity of thle HzS reaction to the temperature is best illustrated in Figure 8. If we compare reactivities at a sulfur loading of 2 mmol/g of limestone, the reactivity a t 1144 K is 1.2 times the reactivity a t 1060 K. The theoretical calculations predict that the reactivity scales as k , 0 6 for these conditions. Hence, the derived kinetic rate constant a t 1144 K is 1.4 times the kinetic rate constant at 1060 K. This corresponds to an activation energy of approximately 10 kcal,lmol. Thus, the rate constants and the activation energies for the SOz and HzS reactions with CaO are roughly the s,ame.

VII. Conclusions A detailed kinetic investigation has shown that the heterogeneous rate constant for CaO with HzS is of the same magnitude as that with SOz and that the activation energy for both reactions is approximately 10 kcal/mol. The derived value of the rate constant for the SOzreaction is within a factor of 2 of that quoted by Borgwardt and Harvey (1972) but approximately 20 times slower than that obtained by Hartman and Coughlin (1976). For the SOz and HzS reactions with CaO at 1200 K, we obtain

It,, = (2 f 1) X

g of CaO/s cmz atm of SOz

k,, =

g of CaO/s cm2 atm of HzS

respectively. The pore transport theory of section IV, together with the SOz rate constant quoted above, is consistent with the SOz sorption data of bloth Hartman and Coughlin (1976) and Borgwardt and Harvey (1972). The theory was verified in both the limit that the sulfur sorption is controlled by SOz diffusion through the porous structure, and in the limit of kinetic control. We therefore conclude that the rate constant obtained by Hartman and Coughlin (1976) is approximately 20 times too fast and we suggest that this may be due to the use of a diffusion coefficient that is a factor of 20 too low. The present theory is valid only at early time because we do not consider the sulfur deposition process. The sulfur deposits prevent the reactant gas from reaching the CaO, thereby reducing the late time reactivity. The late time reactivity of H2S is higher than that of SOz because the formation of CaS does not completely obstruct the porous structure of CaO, as does the formation of CaS04. Empirical evidence of this is demonstrated by the increased CaO utilization with H2S as opposed to SOz. See, for example, Figures 6 imd 9 for 500-pm diameter particles. The CaO conversion with SOz is 4070, whereas that with

1980

571

HzS is 90%. This offers clear evidence that the sulfur deposits play a role in the late time reactivity. Hence, the next step in the theoretical description of sulfur sorption by limestone must include the effects of the sulfur deposits. Such a theory would not only yield the late time reactivity, but it would indicate the operating conditions under which the rate of CaO utilization could be maximized. While grain theory potentially possesses this capability, unique, physically consistent values are required for all parameters. In addition to developing a theory which includes the effects of sulfur deposits, we must determine the effects of the calcination temperature and COPbackground on the pore structure of the calcine. The enhancement of CaO reactivity by lower calcination temperature (Borgwardt and Harvey, 1972), and higher COPbackground level (Ruth et al., 1972), have been experimentally observled, although very little quantitative work has been reported. In particular, the work of Ruth et al., (1972) on the reaction of HzS with CaC03 calcined in situ clearly shows the effects of these conditions on the transition from high reactivity a t early times to lower reactivity at late times. These effects are important for CaO utilization and additional experimental work is required before a viable theory can be constructed. Acknowledgment The authors are grateful to R. H. Borgwardt for providing his data in the format used in this paper. This research was supported by the US.Department of Energy under Contract No. DE-AC21-78MC08450. Nomenclature a = particle radius c = species mass fraction D = species self-diffusion coefficient D, = continuum value of D Dk= Knudsen value of D g(r,) = pore distribution function kB = Boltzmann constant k, = kinetic rate: mass of solid/time-area k , = rate constant: mass of solid/time-area-atm species j K = kinetic rate: mass of species j/time-area = rate constant: mass of species j/time-area-atm species

d,. J

and

Chem. Process Des. Dev., Vol. 19, No. 4,

K = rate constant: cm/s KO= constant = ratio of pore length to radius = 5 1, = pore length = consumption of species j per pore tree 2 1 = total porous consumption of species j h & = total porous consumption of solid M = molecular weight of species j = average molecular weight p = pressure rp = pore radius rmax= maximum value of rp r,,, = minimum value of rp r D = vdue of rp separating continuum and Knudsen diffusion rc = value of rp separating diffusion and kinetically limited pores s = internal surface area (area/mass) = internal surface area of a pore tree T = temperature u = porous volume (volume/mass) d = mean thermal speed x = coordinate along pore tree trunk X = species mole fraction 2 = solid mass released per reactant gas consumed Greek Symbols P = In (rmJrmln) 0 = porosity X = mean free path p = mass density

id

&

572

Ind. Eng. Chem. Process Des. Dev. 1980, 19, 572-580

Subscripts G = gas in pore j = speciesj o = external surface p = pore tree s = nonporous Literature Cited Borgwardt, R. H., Harvey, R. D., Environ. Sci. Techno/. 6, 350 (1972). Case, G. D., et al., "Chemlstry of Hot Gas Cleanup in Coal Gasification and Combustion", Morgantown Energy Research Center, MERC/SP-78/2, 1978. Campbell, F. R., Hills, A. W. D., Pauiin, A., Chem. Eng. Sci., 25, 929 (1970). Hartman, M., Coughlln, R. W., Ind. Eng. Chem. Process Des. Dev., 13, 248 (1974). Hartman, M., Coughlln, R. W., AIChE J., 22, 490 (1976).

Hartman, M., Coughlln, R. W., Ind. Eng. Chem. ProcessDes. Dev., 17, 411 (1978). Jeans, J. H., "The Dynamical Theory of Gases", 4th ed, p 310, Dover Publications, 1954. Keairns, D. L., Archer, D. H., Newby, R. A,, ONeiil, E. P., V U , E. J., "Evaluation of the FluMized-Bed Cornbustion Process, Vol. IV, Fluidized-Bed Oil Gasification/Desulfurizatlon", EPA-650/2-73-048d, NTIS PB 233-101, 1973. Keairns, E. L., Newby, R. A., O'Nelll, E. P., Archer, D. H., Am. Chem. Soc., Div. FuelChem., Prepr., 21, 91 (1976). Ruth, L. A., Squires, A. M., Graff, R. A., Envkon. Sci. Techno/.,6, 1009 (1972). Simons, G. A., Finson, M. L., Combust. Sci. Techno/., 10, 217 (1979). Simons, G. A., Combust. Sci. Techno/., 10, 227 (1979a). Simons, G. A., Combust. Scl. Techno/., 20, 107 (1979b). Thiele, E. W., Ind. Eng. Chem., 31, 916 (1939).

Received for review May 7, 1979 Accepted June 19,1980

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Selective Action of Hematite in Coal Desulfurization Diwakar Garg, Arthur R. Tarrer," James A. Guin, Christlne W. Curtis, and James H. Clinton Chemical Engineering Department, Auburn University, Auburn, Alabama 36849

The effects of hematite (Fe203)on dissolution, hydrogenation, and hydrodesulfurizationreactions in the liquefaction of Western Kentucky No. 9/14 coal have been studied in a 300-cm3 autoclave. The ranges of process variables studied are reaction temperature, 385 to 420 'C; reaction time, 15 to 120 min; hematite loading, 0 to 6% of maf coal feed; and hydrogen partial pressure, 7.0 to 20.8 MPa. The selectivity of hematite for desulfurization over hydrogenation is determined and is used as a measure of the relative performance attained at different levels of the process variables. The amount of hematite required for satisfactory removal of the organic sulfur from the Western Kentucky No. 9/14 coal/solvent system is roughly equal to the stoichiometric amount. These results are verified by examination of the sulfur removal ability of hematite on a model sulfur compound. The effect of varying the hematite surface area from 0.15 to 8.9 m2/g on desulfurization activity is also studied. The effect of mass transfer on the hydrogenationand hydrodesulfurizationreactions of Western Kentucky No. 9114 coal is evaluated.

\

Introduction

The purpose of most coal liquefaction processes, including the Solvent Refined Coal (SRC) process, is to produce an environmentally acceptable solid or liquid fuel or chemical feedstock from coal. Almost all processes rely on direct or indirect (e.g., from a donor solvent) hydrogenation of coal to accomplish this goal. In the SRC process, coal is dissolved in a coal-derived solvent in the presence of hydrogen gas. The organic insolubles, inorganic sulfur, and other minerals are then removed by physical methods while some of the organic sulfur in the coal is removed as hydrogen sulfide gas. Hydrogen economy is a major concern in SRC processing. The stoichiometric amount of hydrogen required solely for removal of an acceptable amount of sulfur is an order of magnitude less than that actually consumed. Actual consumption of the SRC process is about 2 w t % of moisture-ash free (maf) coal feed. Excess hydrogenation of liquid product and formation of gases account for most of the hydrogen consumed. The overall objective of this work is to develop a methodology for utilizing the beneficial effect of certain mineral additives on the rate of desulfurization during coal liquefaction. The addition of some minerals to the dissolver feed may allow sulfur removal requirements to be met with shorter reaction times, which could result in lower hydrogen consumption. The present results demonstrate that a specific mineral additive, hematite, is effective in 0196-4305/ao/i i19-0572$01.0010

sulfur scavenging, provided that the hematite has sufficient surface area and is present in a t least a stoichiometric amount. The selectivity for hydrodesulfurization over hydrogenation, Se, is a measure used to evaluate the performance of the additive as well as the effect of various reaction conditions such as reaction time, amount of the additive, reaction temperature, and hydrogen partial pressure. It is proportional to the fraction of sulfur removed per unit fraction hydrogen consumed and is defined as follows (Garg et al., 1979)

where Sfis the sulfur content of the total liquid product, Sois the original sulfur content calculated from the solvent sulfur and the organic sulfur in the coal, 0.816% (Garg et al., 1979), Hf is the final hydrogen partial pressure, and Ho is the initial hydrogen partial pressure. Previous W o r k

Clays were the first cracking catalyst used in the petroleum industry. Given et al. (1975) studied the catalytic properties of clay minerals and found that certain clay minerals catalyze hydrogen transfer from tetralin to phenanthrene and from partly hydrogenated phenanthrenes to naphthalene. Wright and Severson (1972) showed a linear relationship between percentage of hy0 1980 American Chemical Society