- Email: [email protected]
Thermolysis of hydrogen sulphide in an open tubular reactor
hr. J. H-vdrogen Energy, Vol. 20, No. 10, pp. 777 783, 1995 Copyright @ 1995 International Association for Hydrogen Energy Elsevier Science Ltd
Pergamon
Printed in Great Britain. All rights reserved 036&3199/95 $9.50 + 0.00
0360-3199(95)00001---l
THERMOLYSIS
OF HYDROGEN SULPHIDE REACTOR A. A. ADESINA,*t
V. MEEYOOt
IN AN OPEN TUBULAR
and G. FOULDSZ
tSchool of Chemical Engineering and Industrial Chemistry, University of New South Wales, Sydney, NSW, Australia 2052 ZChemical Engineering Division James Cook University Townsville, Queensland, Australia 4811
(Received
for publication 10 Junuary 1995)
Abstract--This paper addresses the kinetic aspects of the thermal decomposition of H,S in a flow reactor. Experiments conducted in a quartz tube with argon/H,S feed over a wide composition spectrum (ZO-100% H,S) at four temperatures (1030.-1070 K) show that the reaction is essentially first order in H,S partial pressure. Theoretical models based on a free radical mechanism involving the abstraction of hydrogen from H,S as the rate-determining step show the linear dependency of thermolysis rate on reactant composition. Hydrogen yield also increases monotonically with feed composition at all temperatures. Interestingly, the activation energy for H, production of 200 kJ mall’ is lower than that for the global decomposition reaction of about 241 kJ mall ‘, consistent with the view that the initiation step requiring the breaking of the H-S bond is the rate-controlling step rather than the termination involving the combination of two hydrogen radicals.
2H,S + 0, + 2H,O + S,
NOMENCLATURE a, . . . a5
F
Coefficients in equation (3) Feed molar flow rate (mol s-l) Experiments in equation (3) Measured temperature (“C) Set temperature (“C) Reactor volume (cm3) Inlet H,S composition Exit H,S composition Distance measured from reactor inlet (cm)
INTRODUCTION Hydrogen sulphide is an unwanted by-product of many treatment operations in the petroleum, petrochemical and mineral processing industries. It is normally produced during hydrodesulphurization of fossil fuels and in the ore reduction stages during metal extraction. Increasingly, stringent environmental regulations require that this toxic gas must be completely removed from industrial waste gases before atmospheric venting. In particular, acid rain resulting from H,S pollution produces severe deleterious effects on buildings, aquatic and plant ecosystems. Conventionally, H,S is removed through catalytic oxidation to elemental sulphur and water in the Claus process
*Author
to whom correspondence
(1)
or via absorption in caustic solution. It is obvious that the large quantities of H, used in the upstream pretreatment operations are simply jettisoned as water in the Claus oxidation process. In a typical refinery producing about 50 tons day-’ of H,S, this translates to a loss of the order of a million dollars per year. While sulphur is recovered as a saleable product, the simultaneous recovery of H, will further improve the economics of material utilization in plant operation since H, can be recycled for further re-use, hence the need to develop a technology for the dissociation of H,S into H, and sulphur. Moreover, natural gas contains large quantities of H,S and thus the development of an efficient process for hydrogen recovery from H,S will be a welcome contribution to current
efforts
on the production
of clean techniques
for
energy generation from natural gas. Indeed, there is a discernible shift towards HZ, as a source of energy for the future (e.g. H, fuel cells) since the combustion product is environmentally harmless. Clearly, the development and improvement of H, production technologies will remain a major area of industrial and government research activity.
Within the last 2 decades, a number of investigators Cl-93 have undertaken studies on the decomposition of H,S to hydrogen and elemental sulphur. These attempts can be classified into the following categories: (a) non-oxidative thermal cracking; (b) thermochemical decomposition by open and closed loop processes;
should be addressed. 771
A. A. ADESINA
118
(c) electrochemical decomposition; (d) photolytic activation (using solar radiation); (e) catalytic decomposition. In general, the thermal decomposition of H,S is the simplest dissociative route and must be well understood if other process options [(b)(e)] are to be meaningfully evaluated. Unfortunately, there is a paucity of information on the thermolytic kinetics of H,S and even among these few studies, agreement is poor. The thermal decomposition of H,S, written 2H,S ti 2H, + S,
AG& = 116.3 kJ mol-‘,
et al.
prevailed above 875 K [3]) and the analytical technique used by each research group. Furthermore, while most workers favoured quartz asthe reactor material, Kaloidas and Papayannakos [7] opined that this material possessed catalytic activity for H,S decomposition, and hence they employed an cc-Al,O, tube to avoid wall effects.Intriguingly, in a solar-aided decomposition study, Kappauf and Fletcher [ 1l] observed enhanced catalysis by the alumina walls of the reactor. This brief review suggeststhe need for greater clarification on the thermolytic reaction.
(2)
is an endothermal reaction, favoured at high temperatures, typically 973-1273 K. Equilibrium conversions increase with temperature, attaining 14% at about 1200K and 1 atm [7]. From the stoichiometry, it is easily seen that conversions will be favoured at low pressure. Raymont [l] observed that the decomposition rate could be adequately described by an irreversible first order equation in H,S partial pressure. However, earlier work by Darwent and Roberts [3] indicated a second order dependency.Kaloidas and Papayannakos [7] found that at 1.3-3atm and within a temperaturespreadof 873-l 133K the decomposition reaction rate follows a reversible first order (in H,S partial pressure) expression. Tesner et al. [lo] also proposed a reversible model whose forward rate has a second order kinetics in H,S concentration. It is obvious from these various studies that unanimity on the kinetics of this apparently simple reaction is lacking. The diverse views could be due to the differences in the choice of process conditions (for example, below 875 K and subambient, the reaction shows a first order dependency on H,S pressure while a second order kinetics
EXPERIMENTAL Figure 1 is a schematic representation of the experimental rig housed in a ventilated cabinet. Argon and H,S were metered and regulated by rotameters and mixed in a stainless steel chamber before being fed to the quartz tubular reactor (inner diameter, 5.4 mm). The 300 mm long quartz tube was placed in a cylindrical electrical furnace equipped with a temperature controller. Preliminary tests indicated no change in decomposition activity even with a fivefold increasein reactor surface area (using a 30 mm inner diameter tube) at about 1073 K. Thus, contributions from possible catalytic wall effects were probably non-existent. Sulphur vapours from the reactor were trapped in a U-tube packed with inert glass beads immersed in an ice bath. The sulphur-free gaseswere then sent to a Shimadzu 8A gaschromatograph (GC) equipped with thermal conductivity detector (TCD) for analysis. A Haysep Q column gave good peak separation and resolution for H, and H,S at the column temperature of 363 K and carrier flow rate of 40 cm3 min- ’ [standard temperature and pressure (STP)]. The GC effluent was
rd
ir
H;S
Fig. 1. Experimentalrig for H,S decomposition.
THERMOLYSIS
further bubbled through a NaOH solution before venting. Blank experiments with a telescopically arranged thermocouple showed an axial temperature variation which is a function of the set temperature and gas flow rate as shown in equation (3): T = (a0 + a,z + u*z2 + a3z3 + a4z4 + u5z5)FnTz,,
(3)
where T is the measured temperature (“C), T,,, is the set temperature (‘C), z is the distance measured from the reactor inlet (cm), F is the feed molar flow rate (mol s- ‘), 11= 6.1538 x 10-4, m = 1.00023, a, = 0.7310, a, = O.l0715,cr, = -1.211 x lo-*& = 4.60387 x 10m4,uq = -8.78 x lo-’ and LJ~= - 1.66 x IO-‘. This temperature distribution is primarily an artifact of the furnace design. The average reaction temperature was determined from
L
779
20-
0 1029 K + 1049 K CT 2 2
n 1059 K o 1068 K
O0-m
1.0
H2S partial pressure (atm)
‘j T(z)dz T,, = L
OF H,S
(4)
’
Fig. 3. Influence of H,S partial pressure on decomposition
rate.
whereupon T,,, = 0.998 F”TE,.
(5)
Clearly, the temperature difference between T,, and the maximum temperature, T,,,, over the reactor length is relatively small, rarely exceeding 3 K even at the highest temperature used in this study. Figure 2 shows the relationship between conversion and residence time in the empty reactor. As may be seen from this plot, conversion varies linearly with residence time up to about 0.2 min for 1030 and 1070 K at 1 atm. Beyond this point, flow rates were sufficiently low to
permit the attainment of reaction equilibrium within the reactor length. Consequently, all kinetic measurements were made with total feed rates in excess of 100 cm3 min-’ (stp). Under these conditions, the initial decomposition rate, r,,,, was calculated as F(Y;,“~,” - Y”,k) r rx” = _ v ’
where F is the feed molar flow rate, y$’ is the inlet H,S composltlon, yr’, is the exit H,S composition and I/ is the reactor volu’me. RESULTS
0 1068 K l 1029 K
AND DISCUSSION
All kinetic data were collected over the composition range0.20 < y uZS< 1.OOat 103&1070 K using high space velocities (about 7 cm s-l). Figure 3 shows that the decomposition rates increased monotonically with H,S partial pressure. Therefore the rate data was modelled by the empirical law: -THIS = kc,e-EIRTp\ 2s.
0.1
(6)
0.2
0.3
Residence time (min) Fig. 2. Conversion-residence
time plot.
(
(7)
Non-linear regression analysis using PROC NLIN (from the SAS package) gave n = 1.02, an indication that the overall reaction is first order in H,S partial pressure. The associated activation energy of 240 kJ molt ’ compares with the literature value of 230 kJ mol-’ [l, 7, 121. The kinetics of H,S thermal decomposition has been highly debated. While one school of thought favours an irreversible rate expression [l, 43, Kaloidas and Papayannakos [7] correlated their results with the reversible kinetics law described by:
780
A. A. ADESINA et al.
- rH,s =
k PHz, - -
1
PHI P,112
Ke,
’
(8)
As discussed in a later section, the data obtained in this study could not be adequately described by equation (8). Indeed, data regression gave a negative parameter estimate for the equilibrium constant, K,,. In practically every investigation reported in the literature, it was difficult to carry out direct on-line sulphur vapour composition measurement and as a result it had to be estimated from cumulative gravimetric analysis or from H, concentration assuming only hydrogen and S, are produced in the reaction. Nevertheless, the nature of sulphur vapour remains a contentious issue. Chivers et al. [ 121 observed that the predominant allotropic form of sulphur in their system was the cycle-octasulphur species while others [13, 141 have demonstrated that beyond 873 K sulphur vapours contain mostly diatomic molecules. An extensive thermodynamic analysis by Berk et al. [lS] also lends credence to this view. Clearly, in agreement with Raymont [l] and Chivers et aI. [12], our data show that the reaction has a first order irreversible kinetics. Since many of the previous kinetic laws were empirical expressions, it is instructive to investigate rate equations derived from formal mechanisms which accommodate experimental observations reported in the literature. This will allow a more unified interpretation of the data. In particular, as will be evident in the forthcoming analysis, these mechanisms may also give rise to rate laws which are structurally similar to the empirical equations. In the following, we consider mechanistic propositions for the thermolytic reaction. Type A mechanisms assumed that the free radical reaction steps are irreversible while type B allows reversibility in the elementary steps.
H’+
k, H’ + H,.
Writing the rate law for each elementary step, we obtain -rH,S
= k,Ptp
-
U=&
+ k,Pn2sPn
r HS= k,Pn,s - k&is 'H = k,P,zs
rs = k,Pi,
-
+ k,PnzsPn - k,PsPns
k, pn2spn
+ k,PnsPs
-
V':,
- k,P,,P,
and using the concept of the pseudo steady-state hypothesis, the production rate of the intermediate species rHS, r, and rH may be set to zero, thus allowing the computation of P,,, P, and P, in terms of PHz, and the kinetic parameters, k;. This exercise yields -'H,S
=
k H,SPH,S~
(9)
where kHz, = 2k, as shown in the Appendix. Obviously, this mechanism is consistent with a first order rate law for the irreversible decomposition of H,S. Indeed, linear regression analysis of the data gave a correlation coefficient of 0.99-0.998 at all temperatures with an activation energy of 241 kJ mol-i. Mechanism A.2. Berk et al. [lS] have provided evidence for the thermodynamic feasibility of sulphane (H,S,, 2 < i < 8) species during the decomposition of H,S. To accommodate this possibility, a reaction mechanism in which H,S, is an intermediate is proposed:
Type A mechanism Mechanism A.I. This proposal stipulates that chain initiation involves the unimolecular dissociation of H,S to HS’ and H’ radicals. Bradley and Dobson [16, 171 and Merryman and Levy [ 181 had provided evidence for the existence of these radicals using spectroscopic techniques. The propagation steps consist of the interaction between two HS’ species to yield another H,S and S followed by the reaction between the nascent H,S and a H’ radical to produce a H, molecule and another HS’. Sulphur is then produced from a combination of s’ and HS’ with a concomitant release of H’ radicals. The termination step is the production of H, from two H’ species. Thus, the mechanism may be written as: H,S 2 HS’ + H HS+HS1:H,S+S
H,S 2 HS + H’ HS + H,S 2 H, + HS’
HS+HS+H,S+S H1S+S+H2S2+
+H,+S,.
As may be seen, the initiation step still involves the thermal abstraction of a hydrogen radical from H,S to give HS and H’. The propagation steps also remain the same. However, the termination step involves the production of an intermediate sulphane species, H,S, which further decomposes to H, and S,. Following the same procedure as in A.l, the formal rate expression for this mechanism was derived as
H,S + H’ 3 H, + HS (10)
THERMOLYSIS
where - rH s = k, ,P, s and k, s = 2k, Clearly, there is no distinction, at ~east’mathematicallv. between the rate equation from either A.1 or A.2. Observe, however, that the overall reaction requires two molecules of H,S, unlike A.1 in which the global reaction is seen as unimolecular decomposition. It has been proposed [11] that since the overall stoichiometry involves two molecules of H,S then the kinetics is probably second order in H,S partial pressure (mass action kinetics) while others [3, 12, 181 maintained that the first order dependency on H,S observed recommends the reaction as a unimolecular one. Our analysis shows that both unimolecular and bimolecular reactions could give rise to first order rate laws. For a system involving increase in number of moles upon reaction, increased pressure will aid the reverse reactionand thus the global rate law must account for this effect. Additionally, conversion drops with increase in system pressure [S] and the rate begins to exhibit non-linear dependence on H,S partial pressure. Thus, reversibility and non-linear kinetics are both favoured by high pressure. Indeed, the assumption that P, + 1 (used in the formal derivation of the kinetic equation) can only be valid if total system pressure is atmospheric. In addition, at the relatively low temperatures employed, decomposition is slow enough to prevent the attainment of chemical equilibrium [ll]. It would seem that workers who operated close to the “plateau” portion of the conversionresidence time plot would obtain data describable by reversible kinetic laws and this may be further compounded by high pressure operation. Researchers who used high pressures have generally obtained data with a second order fit while those who studied at about 1 atm reported first order H,S decomposition kinetics. Future process commercialization, however, requires that the kinetics be investigated at about atmospheric pressure. Type B mechanisms The principal difference between this mechanism and the previous ones is the reversibility permitted in each elementary step and the assumed termination step requiring sulphur production. Mechanism B.l.
This may be written as H,S > HS’ + H’ HS’ + HS’ 2 H,S + S H,S + H’ 3 H, + HS’ s’ s l/2 s,
After going through the formal motions, the final rate equation, structurally identical to equation (8), emerges as
781
OF H,S
-rH,S
=
kl
pH2s
- & P”*P;j2 E
,
(11)
where Ki = ki/k-i and K, is the equilibrium constant for the overall unimolecular decomposition of H,S. Mechanism B.2. This is the bimolecular model B.l. The mechanism
version of
2H,S 3 21~s’ + 2H’ HS+HS+H,S+S H,S + H’ 2 H, + HS s’ 3 1/2s, yields a rate expression -r
H,S =
k,
p;,s !
-
f
W)
p;ips2 e
>
which is second order in H,S and H, partial pressures and first order in sulphur vapour pressure. As pointed out earlier, reversible rate equations could not reasonably describe the rate-composition behaviour observed in this investigation. Thus, models arising from type B mechanisms failed to fit the data. In the present investigation, we believe that since low conversions far from equilibrium were used, the associated low partial pressure for H, in concert with a correspondingly low PSi (1 d i 6 8) means that the second terms on the right-hand sides ofequations (11) and (12) contribute very little to the rate. In fact, using the appropriate value,s for K,, at the experimental temperature, these equations still failed to describe adequately the data in this study. It would therefore seem that the mechanism for the thermolysis of H,S consists of irreversible elementary steps in which the abstraction of hydrogen from the H,S molecule constitutes the rate-determining step. In principle, any supplementary “aid” for the abstraction of hydrogen from the H,S (such as the use of a catalyst, electrolytic or photochemical activation) may further enhance the production of hydrogen from H,S. This will be to subject of a future article. Hydrogen production In spite of its obvious importance, analysis of the H, yield has received very little attention from other workers. Figure 4 shows that the production of H, increases linearly with the reactant partial pressure at all temperatures used in this study. From this plot, first order rate constants were determined and utilized further to obtain the activation energy for H:! production as may be seen
A. A. ADESINA et al.
782
0 1029K + 1049 K n 1059 K 0 1068 K
2
0.4
1
0.8
0.6
H#Y partial pressure (atm) Fig. 4. Influence of H,S partial pressure on H, production
rate.
in Fig. 5. The resulting activation energy of about 200 kJ mol-’ for H, production is lower than the global thermolytic activation energy of 241 kJ mol-‘. This suggeststhat the overall decomposition reaction is not controlled by the elementary step involving H, release (a termination or propagation step). As alluded to previously, the decomposition of H,S is probably governed by the initiation step requiring the abstraction of the hydrogen radical from H,S. CONCLUSIONS This study reveals that at the low pressures(about 1
-13.4
c
-13.6
-13.8 3 5 z
-14.0
-14.2
-14.4
-14.6 9.3
9.4
9.5
9.6
9.1
l/T x 104 Fig. 5. Arrhenius plot for H, production
rate.
atm) appropriate for H,S decomposition, a first order irreversible rate expression in reactant partial pressureis an adequate representation of the reaction kinetics. Mechanisms which permit reversibility in the elementary stepsfailed to describe the data (even though such could lead to a first order dependency in H,S partial pressure). However, the two mechanisms which assumedirreversibility of elementary steps gave first order rate equations irrespective of the overall reaction stoichiometry (unior bi-molecular). Hydrogen production rates were also linear in H,S partial pressure at all four temperatures studied. The estimated activation energy was about 40 kJ mol- ’ lower than that for the overall decomposition reaction (about 241 kJ mol-‘) suggesting that the elementary step resulting in H, production cannot be the rate-determining step for the global reaction. Thus, the abstraction of hydrogen from H,S appears to be the most likely rate-controlling step in the free radical mechanism for the thermolysis of H,S and non-thermal methods (such as catalytic, photochemical and electrolytic processes)may enhance the activation and breakage of the H-S bond with a consequent increase in hydrogen yield. Acknowledgements-Partial support for this work by the CSIRO Division of Coal and Energy Technology, Lucas Heights, New South Wales, is appreciated. We are also indebted to John Starling, Philip McAuley and Dean Benke for technical back-up.
REFERENCES 1. M. E. D. Raymont, Hydrocarbon Processing 54, 139 (1975). 2. K. Fukuda, M. Dokiya, T. Kamayama and Y. Kotera, Ind. Engng Chem. Fundnm. 17, 243 (1978). 3. B. Darwent and R. Roberts, Proc. R. Sot. 216, 344 (1953). 4. R. C. Kainthla and J. 0. M. Bockris, Inc. J. Hydrogen Energy 12, 629 (1987). 5. T. N. Veziroglu and W. Seifritz (eds), Proc. 2nd World Hydrogen Energy Conf, Zurich, Switzerland (1978). 6. T. Cl. Gregory, D. L. Feke, J. C. Angus, C. B. Borsilow and U. Landau, J. Appl. Electrochem. 10, 405 (1980). 7. V. Kaloidas and N. Cl. Papayannakos, Gem. Engng Sci. 44, 2493 (1987). 8. F. Bandermann and K. B. Harder, Int. J. Hydrogen Energy 7, 471 (1982). 9. V. Kaloidas and N. G. Papayannakos, I & EC Res. 30,345 (1991). 10. P. A. Tesner, M. S. Nemirovskii and D. N. Motyl, Kinet. Katal. 31, 1232 (1990). 11. T. Kappauf and E. A. Fletcher, Energy 14,443 (1989). 12. T. Chivers, J. B. Hynes and C. Lau, Int. J. Hydrogen Energy 5, 499 (1980). 13. H. Rau, T. R. N. Kutty and G. R. F. Guedes de Carvalho, J. Chem. Thermodyn. 5, 833 (1973). 14. V. Kaloidas and N. G. Papayannakos, Int. J. Hydrogen Energy 12, 403 (1987). 15. D. Berk, R. A. Heidemann, W. Y. Svreck and L. Behie, Can. J. Chem. Engng 69,944 (1991). 16. J. N. Bradley and D. C. Dobson, J. Chem. Phys. 46, 2865 (1967). 17. J. N. Bradley and D. C. Dobson, J. Chem. Phys. 46, 2872 (1967). 18. E. L. Merryman and A. Levy, J. Phys. Chem. 76,1925 (1972).
THERMOLYSIS
783
OF H,S
APPENDIX
p, = b:,,p,,,
Derivation of mechanism-based rate equation for H,S thermolysis The sequence of elementary steps for Mechanism A.1 are:
where k,, = k,/k,,
645)
and from equation (A2), we obtain
k,PH2s + k,P,$‘,
= (4 + k,k,,)Pih
whence (A7)
H,S 2 HS’ + H
where k,, = k, + k,,. For atmospheric operation and H being an intermediate, P,, < 1 and thus Pf, would be very close to zero. Using this approximation and introducing equation (A7) into equation (A3) yields, upon reshuffling,
HS’+HS’%H,S+S H,S + H’ $ H, + HS’ S’+HSaS,+H’ H’+H’*H,
646)
k,f’,zs + kA&s
4
= W-‘frs - k,Pn2s
648)
from which, so that the rate expression for each species is: (Al)
-rnas = klPHJS - k2P& + k3f’HJ’H rHS = klPH2s - k2P& + W’H2SPH - k4PsPHs ‘H
=~,PHJ
rs = k,P&
-
k,f’&-‘,
- k,P,,Ps.
+
k,f’,sPs
P,, = k,P,$
-
k,P:,
and k, = 2k,lk,.
Consequently,
(A4 P, =
(A3) (A4)
Assuming that all steps except the rate-determining one [cf. equation (Al)] obey the pseudo steady-state hypothesis, then rs = rHS = rH = 0. From equation (Al), we obtain:
(A9)
h&,2 - k,
(A101
k,
Substituting equations (A9) and (AlO) into equation (Al) leads to -rH,S
=
kf(b
where k,,>, = 2k,.
+
h&-t
-
kPHzs
=
kH,SPH,S,
(All)
Pergamon
Printed in Great Britain. All rights reserved 036&3199/95 $9.50 + 0.00
0360-3199(95)00001---l
THERMOLYSIS
OF HYDROGEN SULPHIDE REACTOR A. A. ADESINA,*t
V. MEEYOOt
IN AN OPEN TUBULAR
and G. FOULDSZ
tSchool of Chemical Engineering and Industrial Chemistry, University of New South Wales, Sydney, NSW, Australia 2052 ZChemical Engineering Division James Cook University Townsville, Queensland, Australia 4811
(Received
for publication 10 Junuary 1995)
Abstract--This paper addresses the kinetic aspects of the thermal decomposition of H,S in a flow reactor. Experiments conducted in a quartz tube with argon/H,S feed over a wide composition spectrum (ZO-100% H,S) at four temperatures (1030.-1070 K) show that the reaction is essentially first order in H,S partial pressure. Theoretical models based on a free radical mechanism involving the abstraction of hydrogen from H,S as the rate-determining step show the linear dependency of thermolysis rate on reactant composition. Hydrogen yield also increases monotonically with feed composition at all temperatures. Interestingly, the activation energy for H, production of 200 kJ mall’ is lower than that for the global decomposition reaction of about 241 kJ mall ‘, consistent with the view that the initiation step requiring the breaking of the H-S bond is the rate-controlling step rather than the termination involving the combination of two hydrogen radicals.
2H,S + 0, + 2H,O + S,
NOMENCLATURE a, . . . a5
F
Coefficients in equation (3) Feed molar flow rate (mol s-l) Experiments in equation (3) Measured temperature (“C) Set temperature (“C) Reactor volume (cm3) Inlet H,S composition Exit H,S composition Distance measured from reactor inlet (cm)
INTRODUCTION Hydrogen sulphide is an unwanted by-product of many treatment operations in the petroleum, petrochemical and mineral processing industries. It is normally produced during hydrodesulphurization of fossil fuels and in the ore reduction stages during metal extraction. Increasingly, stringent environmental regulations require that this toxic gas must be completely removed from industrial waste gases before atmospheric venting. In particular, acid rain resulting from H,S pollution produces severe deleterious effects on buildings, aquatic and plant ecosystems. Conventionally, H,S is removed through catalytic oxidation to elemental sulphur and water in the Claus process
*Author
to whom correspondence
(1)
or via absorption in caustic solution. It is obvious that the large quantities of H, used in the upstream pretreatment operations are simply jettisoned as water in the Claus oxidation process. In a typical refinery producing about 50 tons day-’ of H,S, this translates to a loss of the order of a million dollars per year. While sulphur is recovered as a saleable product, the simultaneous recovery of H, will further improve the economics of material utilization in plant operation since H, can be recycled for further re-use, hence the need to develop a technology for the dissociation of H,S into H, and sulphur. Moreover, natural gas contains large quantities of H,S and thus the development of an efficient process for hydrogen recovery from H,S will be a welcome contribution to current
efforts
on the production
of clean techniques
for
energy generation from natural gas. Indeed, there is a discernible shift towards HZ, as a source of energy for the future (e.g. H, fuel cells) since the combustion product is environmentally harmless. Clearly, the development and improvement of H, production technologies will remain a major area of industrial and government research activity.
Within the last 2 decades, a number of investigators Cl-93 have undertaken studies on the decomposition of H,S to hydrogen and elemental sulphur. These attempts can be classified into the following categories: (a) non-oxidative thermal cracking; (b) thermochemical decomposition by open and closed loop processes;
should be addressed. 771
A. A. ADESINA
118
(c) electrochemical decomposition; (d) photolytic activation (using solar radiation); (e) catalytic decomposition. In general, the thermal decomposition of H,S is the simplest dissociative route and must be well understood if other process options [(b)(e)] are to be meaningfully evaluated. Unfortunately, there is a paucity of information on the thermolytic kinetics of H,S and even among these few studies, agreement is poor. The thermal decomposition of H,S, written 2H,S ti 2H, + S,
AG& = 116.3 kJ mol-‘,
et al.
prevailed above 875 K [3]) and the analytical technique used by each research group. Furthermore, while most workers favoured quartz asthe reactor material, Kaloidas and Papayannakos [7] opined that this material possessed catalytic activity for H,S decomposition, and hence they employed an cc-Al,O, tube to avoid wall effects.Intriguingly, in a solar-aided decomposition study, Kappauf and Fletcher [ 1l] observed enhanced catalysis by the alumina walls of the reactor. This brief review suggeststhe need for greater clarification on the thermolytic reaction.
(2)
is an endothermal reaction, favoured at high temperatures, typically 973-1273 K. Equilibrium conversions increase with temperature, attaining 14% at about 1200K and 1 atm [7]. From the stoichiometry, it is easily seen that conversions will be favoured at low pressure. Raymont [l] observed that the decomposition rate could be adequately described by an irreversible first order equation in H,S partial pressure. However, earlier work by Darwent and Roberts [3] indicated a second order dependency.Kaloidas and Papayannakos [7] found that at 1.3-3atm and within a temperaturespreadof 873-l 133K the decomposition reaction rate follows a reversible first order (in H,S partial pressure) expression. Tesner et al. [lo] also proposed a reversible model whose forward rate has a second order kinetics in H,S concentration. It is obvious from these various studies that unanimity on the kinetics of this apparently simple reaction is lacking. The diverse views could be due to the differences in the choice of process conditions (for example, below 875 K and subambient, the reaction shows a first order dependency on H,S pressure while a second order kinetics
EXPERIMENTAL Figure 1 is a schematic representation of the experimental rig housed in a ventilated cabinet. Argon and H,S were metered and regulated by rotameters and mixed in a stainless steel chamber before being fed to the quartz tubular reactor (inner diameter, 5.4 mm). The 300 mm long quartz tube was placed in a cylindrical electrical furnace equipped with a temperature controller. Preliminary tests indicated no change in decomposition activity even with a fivefold increasein reactor surface area (using a 30 mm inner diameter tube) at about 1073 K. Thus, contributions from possible catalytic wall effects were probably non-existent. Sulphur vapours from the reactor were trapped in a U-tube packed with inert glass beads immersed in an ice bath. The sulphur-free gaseswere then sent to a Shimadzu 8A gaschromatograph (GC) equipped with thermal conductivity detector (TCD) for analysis. A Haysep Q column gave good peak separation and resolution for H, and H,S at the column temperature of 363 K and carrier flow rate of 40 cm3 min- ’ [standard temperature and pressure (STP)]. The GC effluent was
rd
ir
H;S
Fig. 1. Experimentalrig for H,S decomposition.
THERMOLYSIS
further bubbled through a NaOH solution before venting. Blank experiments with a telescopically arranged thermocouple showed an axial temperature variation which is a function of the set temperature and gas flow rate as shown in equation (3): T = (a0 + a,z + u*z2 + a3z3 + a4z4 + u5z5)FnTz,,
(3)
where T is the measured temperature (“C), T,,, is the set temperature (‘C), z is the distance measured from the reactor inlet (cm), F is the feed molar flow rate (mol s- ‘), 11= 6.1538 x 10-4, m = 1.00023, a, = 0.7310, a, = O.l0715,cr, = -1.211 x lo-*& = 4.60387 x 10m4,uq = -8.78 x lo-’ and LJ~= - 1.66 x IO-‘. This temperature distribution is primarily an artifact of the furnace design. The average reaction temperature was determined from
L
779
20-
0 1029 K + 1049 K CT 2 2
n 1059 K o 1068 K
O0-m
1.0
H2S partial pressure (atm)
‘j T(z)dz T,, = L
OF H,S
(4)
’
Fig. 3. Influence of H,S partial pressure on decomposition
rate.
whereupon T,,, = 0.998 F”TE,.
(5)
Clearly, the temperature difference between T,, and the maximum temperature, T,,,, over the reactor length is relatively small, rarely exceeding 3 K even at the highest temperature used in this study. Figure 2 shows the relationship between conversion and residence time in the empty reactor. As may be seen from this plot, conversion varies linearly with residence time up to about 0.2 min for 1030 and 1070 K at 1 atm. Beyond this point, flow rates were sufficiently low to
permit the attainment of reaction equilibrium within the reactor length. Consequently, all kinetic measurements were made with total feed rates in excess of 100 cm3 min-’ (stp). Under these conditions, the initial decomposition rate, r,,,, was calculated as F(Y;,“~,” - Y”,k) r rx” = _ v ’
where F is the feed molar flow rate, y$’ is the inlet H,S composltlon, yr’, is the exit H,S composition and I/ is the reactor volu’me. RESULTS
0 1068 K l 1029 K
AND DISCUSSION
All kinetic data were collected over the composition range0.20 < y uZS< 1.OOat 103&1070 K using high space velocities (about 7 cm s-l). Figure 3 shows that the decomposition rates increased monotonically with H,S partial pressure. Therefore the rate data was modelled by the empirical law: -THIS = kc,e-EIRTp\ 2s.
0.1
(6)
0.2
0.3
Residence time (min) Fig. 2. Conversion-residence
time plot.
(
(7)
Non-linear regression analysis using PROC NLIN (from the SAS package) gave n = 1.02, an indication that the overall reaction is first order in H,S partial pressure. The associated activation energy of 240 kJ molt ’ compares with the literature value of 230 kJ mol-’ [l, 7, 121. The kinetics of H,S thermal decomposition has been highly debated. While one school of thought favours an irreversible rate expression [l, 43, Kaloidas and Papayannakos [7] correlated their results with the reversible kinetics law described by:
780
A. A. ADESINA et al.
- rH,s =
k PHz, - -
1
PHI P,112
Ke,
’
(8)
As discussed in a later section, the data obtained in this study could not be adequately described by equation (8). Indeed, data regression gave a negative parameter estimate for the equilibrium constant, K,,. In practically every investigation reported in the literature, it was difficult to carry out direct on-line sulphur vapour composition measurement and as a result it had to be estimated from cumulative gravimetric analysis or from H, concentration assuming only hydrogen and S, are produced in the reaction. Nevertheless, the nature of sulphur vapour remains a contentious issue. Chivers et al. [ 121 observed that the predominant allotropic form of sulphur in their system was the cycle-octasulphur species while others [13, 141 have demonstrated that beyond 873 K sulphur vapours contain mostly diatomic molecules. An extensive thermodynamic analysis by Berk et al. [lS] also lends credence to this view. Clearly, in agreement with Raymont [l] and Chivers et aI. [12], our data show that the reaction has a first order irreversible kinetics. Since many of the previous kinetic laws were empirical expressions, it is instructive to investigate rate equations derived from formal mechanisms which accommodate experimental observations reported in the literature. This will allow a more unified interpretation of the data. In particular, as will be evident in the forthcoming analysis, these mechanisms may also give rise to rate laws which are structurally similar to the empirical equations. In the following, we consider mechanistic propositions for the thermolytic reaction. Type A mechanisms assumed that the free radical reaction steps are irreversible while type B allows reversibility in the elementary steps.
H’+
k, H’ + H,.
Writing the rate law for each elementary step, we obtain -rH,S
= k,Ptp
-
U=&
+ k,Pn2sPn
r HS= k,Pn,s - k&is 'H = k,P,zs
rs = k,Pi,
-
+ k,PnzsPn - k,PsPns
k, pn2spn
+ k,PnsPs
-
V':,
- k,P,,P,
and using the concept of the pseudo steady-state hypothesis, the production rate of the intermediate species rHS, r, and rH may be set to zero, thus allowing the computation of P,,, P, and P, in terms of PHz, and the kinetic parameters, k;. This exercise yields -'H,S
=
k H,SPH,S~
(9)
where kHz, = 2k, as shown in the Appendix. Obviously, this mechanism is consistent with a first order rate law for the irreversible decomposition of H,S. Indeed, linear regression analysis of the data gave a correlation coefficient of 0.99-0.998 at all temperatures with an activation energy of 241 kJ mol-i. Mechanism A.2. Berk et al. [lS] have provided evidence for the thermodynamic feasibility of sulphane (H,S,, 2 < i < 8) species during the decomposition of H,S. To accommodate this possibility, a reaction mechanism in which H,S, is an intermediate is proposed:
Type A mechanism Mechanism A.I. This proposal stipulates that chain initiation involves the unimolecular dissociation of H,S to HS’ and H’ radicals. Bradley and Dobson [16, 171 and Merryman and Levy [ 181 had provided evidence for the existence of these radicals using spectroscopic techniques. The propagation steps consist of the interaction between two HS’ species to yield another H,S and S followed by the reaction between the nascent H,S and a H’ radical to produce a H, molecule and another HS’. Sulphur is then produced from a combination of s’ and HS’ with a concomitant release of H’ radicals. The termination step is the production of H, from two H’ species. Thus, the mechanism may be written as: H,S 2 HS’ + H HS+HS1:H,S+S
H,S 2 HS + H’ HS + H,S 2 H, + HS’
HS+HS+H,S+S H1S+S+H2S2+
+H,+S,.
As may be seen, the initiation step still involves the thermal abstraction of a hydrogen radical from H,S to give HS and H’. The propagation steps also remain the same. However, the termination step involves the production of an intermediate sulphane species, H,S, which further decomposes to H, and S,. Following the same procedure as in A.l, the formal rate expression for this mechanism was derived as
H,S + H’ 3 H, + HS (10)
THERMOLYSIS
where - rH s = k, ,P, s and k, s = 2k, Clearly, there is no distinction, at ~east’mathematicallv. between the rate equation from either A.1 or A.2. Observe, however, that the overall reaction requires two molecules of H,S, unlike A.1 in which the global reaction is seen as unimolecular decomposition. It has been proposed [11] that since the overall stoichiometry involves two molecules of H,S then the kinetics is probably second order in H,S partial pressure (mass action kinetics) while others [3, 12, 181 maintained that the first order dependency on H,S observed recommends the reaction as a unimolecular one. Our analysis shows that both unimolecular and bimolecular reactions could give rise to first order rate laws. For a system involving increase in number of moles upon reaction, increased pressure will aid the reverse reactionand thus the global rate law must account for this effect. Additionally, conversion drops with increase in system pressure [S] and the rate begins to exhibit non-linear dependence on H,S partial pressure. Thus, reversibility and non-linear kinetics are both favoured by high pressure. Indeed, the assumption that P, + 1 (used in the formal derivation of the kinetic equation) can only be valid if total system pressure is atmospheric. In addition, at the relatively low temperatures employed, decomposition is slow enough to prevent the attainment of chemical equilibrium [ll]. It would seem that workers who operated close to the “plateau” portion of the conversionresidence time plot would obtain data describable by reversible kinetic laws and this may be further compounded by high pressure operation. Researchers who used high pressures have generally obtained data with a second order fit while those who studied at about 1 atm reported first order H,S decomposition kinetics. Future process commercialization, however, requires that the kinetics be investigated at about atmospheric pressure. Type B mechanisms The principal difference between this mechanism and the previous ones is the reversibility permitted in each elementary step and the assumed termination step requiring sulphur production. Mechanism B.l.
This may be written as H,S > HS’ + H’ HS’ + HS’ 2 H,S + S H,S + H’ 3 H, + HS’ s’ s l/2 s,
After going through the formal motions, the final rate equation, structurally identical to equation (8), emerges as
781
OF H,S
-rH,S
=
kl
pH2s
- & P”*P;j2 E
,
(11)
where Ki = ki/k-i and K, is the equilibrium constant for the overall unimolecular decomposition of H,S. Mechanism B.2. This is the bimolecular model B.l. The mechanism
version of
2H,S 3 21~s’ + 2H’ HS+HS+H,S+S H,S + H’ 2 H, + HS s’ 3 1/2s, yields a rate expression -r
H,S =
k,
p;,s !
-
f
W)
p;ips2 e
>
which is second order in H,S and H, partial pressures and first order in sulphur vapour pressure. As pointed out earlier, reversible rate equations could not reasonably describe the rate-composition behaviour observed in this investigation. Thus, models arising from type B mechanisms failed to fit the data. In the present investigation, we believe that since low conversions far from equilibrium were used, the associated low partial pressure for H, in concert with a correspondingly low PSi (1 d i 6 8) means that the second terms on the right-hand sides ofequations (11) and (12) contribute very little to the rate. In fact, using the appropriate value,s for K,, at the experimental temperature, these equations still failed to describe adequately the data in this study. It would therefore seem that the mechanism for the thermolysis of H,S consists of irreversible elementary steps in which the abstraction of hydrogen from the H,S molecule constitutes the rate-determining step. In principle, any supplementary “aid” for the abstraction of hydrogen from the H,S (such as the use of a catalyst, electrolytic or photochemical activation) may further enhance the production of hydrogen from H,S. This will be to subject of a future article. Hydrogen production In spite of its obvious importance, analysis of the H, yield has received very little attention from other workers. Figure 4 shows that the production of H, increases linearly with the reactant partial pressure at all temperatures used in this study. From this plot, first order rate constants were determined and utilized further to obtain the activation energy for H:! production as may be seen
A. A. ADESINA et al.
782
0 1029K + 1049 K n 1059 K 0 1068 K
2
0.4
1
0.8
0.6
H#Y partial pressure (atm) Fig. 4. Influence of H,S partial pressure on H, production
rate.
in Fig. 5. The resulting activation energy of about 200 kJ mol-’ for H, production is lower than the global thermolytic activation energy of 241 kJ mol-‘. This suggeststhat the overall decomposition reaction is not controlled by the elementary step involving H, release (a termination or propagation step). As alluded to previously, the decomposition of H,S is probably governed by the initiation step requiring the abstraction of the hydrogen radical from H,S. CONCLUSIONS This study reveals that at the low pressures(about 1
-13.4
c
-13.6
-13.8 3 5 z
-14.0
-14.2
-14.4
-14.6 9.3
9.4
9.5
9.6
9.1
l/T x 104 Fig. 5. Arrhenius plot for H, production
rate.
atm) appropriate for H,S decomposition, a first order irreversible rate expression in reactant partial pressureis an adequate representation of the reaction kinetics. Mechanisms which permit reversibility in the elementary stepsfailed to describe the data (even though such could lead to a first order dependency in H,S partial pressure). However, the two mechanisms which assumedirreversibility of elementary steps gave first order rate equations irrespective of the overall reaction stoichiometry (unior bi-molecular). Hydrogen production rates were also linear in H,S partial pressure at all four temperatures studied. The estimated activation energy was about 40 kJ mol- ’ lower than that for the overall decomposition reaction (about 241 kJ mol-‘) suggesting that the elementary step resulting in H, production cannot be the rate-determining step for the global reaction. Thus, the abstraction of hydrogen from H,S appears to be the most likely rate-controlling step in the free radical mechanism for the thermolysis of H,S and non-thermal methods (such as catalytic, photochemical and electrolytic processes)may enhance the activation and breakage of the H-S bond with a consequent increase in hydrogen yield. Acknowledgements-Partial support for this work by the CSIRO Division of Coal and Energy Technology, Lucas Heights, New South Wales, is appreciated. We are also indebted to John Starling, Philip McAuley and Dean Benke for technical back-up.
REFERENCES 1. M. E. D. Raymont, Hydrocarbon Processing 54, 139 (1975). 2. K. Fukuda, M. Dokiya, T. Kamayama and Y. Kotera, Ind. Engng Chem. Fundnm. 17, 243 (1978). 3. B. Darwent and R. Roberts, Proc. R. Sot. 216, 344 (1953). 4. R. C. Kainthla and J. 0. M. Bockris, Inc. J. Hydrogen Energy 12, 629 (1987). 5. T. N. Veziroglu and W. Seifritz (eds), Proc. 2nd World Hydrogen Energy Conf, Zurich, Switzerland (1978). 6. T. Cl. Gregory, D. L. Feke, J. C. Angus, C. B. Borsilow and U. Landau, J. Appl. Electrochem. 10, 405 (1980). 7. V. Kaloidas and N. Cl. Papayannakos, Gem. Engng Sci. 44, 2493 (1987). 8. F. Bandermann and K. B. Harder, Int. J. Hydrogen Energy 7, 471 (1982). 9. V. Kaloidas and N. G. Papayannakos, I & EC Res. 30,345 (1991). 10. P. A. Tesner, M. S. Nemirovskii and D. N. Motyl, Kinet. Katal. 31, 1232 (1990). 11. T. Kappauf and E. A. Fletcher, Energy 14,443 (1989). 12. T. Chivers, J. B. Hynes and C. Lau, Int. J. Hydrogen Energy 5, 499 (1980). 13. H. Rau, T. R. N. Kutty and G. R. F. Guedes de Carvalho, J. Chem. Thermodyn. 5, 833 (1973). 14. V. Kaloidas and N. G. Papayannakos, Int. J. Hydrogen Energy 12, 403 (1987). 15. D. Berk, R. A. Heidemann, W. Y. Svreck and L. Behie, Can. J. Chem. Engng 69,944 (1991). 16. J. N. Bradley and D. C. Dobson, J. Chem. Phys. 46, 2865 (1967). 17. J. N. Bradley and D. C. Dobson, J. Chem. Phys. 46, 2872 (1967). 18. E. L. Merryman and A. Levy, J. Phys. Chem. 76,1925 (1972).
THERMOLYSIS
783
OF H,S
APPENDIX
p, = b:,,p,,,
Derivation of mechanism-based rate equation for H,S thermolysis The sequence of elementary steps for Mechanism A.1 are:
where k,, = k,/k,,
645)
and from equation (A2), we obtain
k,PH2s + k,P,$‘,
= (4 + k,k,,)Pih
whence (A7)
H,S 2 HS’ + H
where k,, = k, + k,,. For atmospheric operation and H being an intermediate, P,, < 1 and thus Pf, would be very close to zero. Using this approximation and introducing equation (A7) into equation (A3) yields, upon reshuffling,
HS’+HS’%H,S+S H,S + H’ $ H, + HS’ S’+HSaS,+H’ H’+H’*H,
646)
k,f’,zs + kA&s
4
= W-‘frs - k,Pn2s
648)
from which, so that the rate expression for each species is: (Al)
-rnas = klPHJS - k2P& + k3f’HJ’H rHS = klPH2s - k2P& + W’H2SPH - k4PsPHs ‘H
=~,PHJ
rs = k,P&
-
k,f’&-‘,
- k,P,,Ps.
+
k,f’,sPs
P,, = k,P,$
-
k,P:,
and k, = 2k,lk,.
Consequently,
(A4 P, =
(A3) (A4)
Assuming that all steps except the rate-determining one [cf. equation (Al)] obey the pseudo steady-state hypothesis, then rs = rHS = rH = 0. From equation (Al), we obtain:
(A9)
h&,2 - k,
(A101
k,
Substituting equations (A9) and (AlO) into equation (Al) leads to -rH,S
=
kf(b
where k,,>, = 2k,.
+
h&-t
-
kPHzs
=
kH,SPH,S,
(All)