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Ind. Eng. Chem. Res. 2004, 43, 3304-3313
CS2 Formation in the Claus Reaction Furnace: A Kinetic Study of Methane-Sulfur and Methane-Hydrogen Sulfide Reactions Kunal Karan*,† and Leo A. Behie Department of Chemical and Petroleum Engineering, The University of Calgary, 2500 University Drive NW, Calgary, Alberta, Canada T2N 1N4
Improving the understanding of reaction kinetics of CS2 formation in the Claus plant front-end reaction furnace (RF) is a key step in developing strategies to reduce CS2 formation and, consequently, the environmental impact of Claus plants. Specifically, experiments were carried out in a high-temperature flow reactor with pressures of 101-150 kPa, temperatures of 8001250 °C, and residence times of 90-1400 ms to study the kinetics of CH4-S2 and CH4-H2S reactions; these conditions are typical of those encountered in the Claus RF. The reaction between methane and sulfur was found to be very rapid, resulting in complete consumption of sulfur in less than 100 ms at 1100 °C with formation of CS2 and H2S as the primary sulfur-containing products. At higher temperature (>1000 °C), the produced H2S decomposes with a proportional increase in CS2 formation. A simple rate expression for CS2 formation was obtained, and a kinetic model was developed to describe H2S formation/consumption in the CH4-S2 system. In the CH4H2S reacting system, H2S thermal decomposition appears to be the rate-limiting step for CS2 formation. The consumption of H2S in the CH4-H2S system proceeds at a rate characteristic of thermal decomposition of H2S, i.e., at a rate independent of any reaction of H2S with methane. Introduction Background. The reactions of methane with sulfur and hydrogen sulfide are thought to be primary causes for the production of undesirable CS2 in the thermal step of the modified Claus process.1-3 The main purpose of the thermal step of the modified Claus process is to partially oxidize H2S in the acid gas feed stream to SO2, which is then reacted with unreacted H2S in the lowtemperature catalytic step carried out in a series of fixed-bed catalytic converters.4 The thermal step is achieved in the Claus plant front-end units comprising a reaction furnace (RF) and its associated waste heat boiler (WHB). In the RF, the partial oxidation of H2S is achieved in the diffusive flame region where temperatures may exceed 1500 °C and reaction times are on the order of a few milliseconds. The product gases from the flame region, however, continue to react in the postflame region of the RF, where average temperatures are in the range of 900-1300 °C and gases have a residence time of 0.5-2.0 s. The interaction of impurities such as CO2 and hydrocarbons with sulfur-containing species leads to the occurrence of numerous side reactions and to the formation of at least two important sulfurcontaining species: COS and CS2.1,4 Both COS and CS2 are formed in the front end of the plant and can be hydrolyzed to CO2 and H2S in the downstream catalytic reactors provided the converters are operated at higher temperatures (>280 °C). However, the catalytic converters are operated at lower temperatures in order to maximize sulfur recovery from the main reaction between SO2 and H2S. Typically, only the first catalytic converter is operated at relatively high temperatures, but even these conditions are insufficient * To whom correspondence should be addressed. Fax: (613) 533-6637. E-mail:
[email protected]. † Current address: Department of Chemical Engineering, Queen’s University, Kingston, Canada K7L 3N6.
for achieving complete COS and CS2 conversions. Consequently, COS and CS2 end up in the tail gas contributing from 20% to 50% of the total plant sulfur emissions. To meet stringent regulations that require up to 99% plant sulfur recovery, tail gas cleanup units (TGCUs) are commonly employed, which may be as expensive as the Claus plant itself. Several of these TGCU processes are not effective or efficient in recovering/converting COS and CS2 .5,6 However, by minimization of COS and CS2 formation at the source, significant gains in the overall plant sulfur recovery and substantial reduction in sulfur emissions may be achieved. This has been our approach in addressing the sulfur emissions/recovery challenges faced by Claus plants. In fact, through a combination of experimental work and computer simulation studies, we have improved the understanding of COS reaction chemistry and kinetics7,8 as well as suggested possible solutions for minimization of COS formation in the Claus front-end units.9 Our understanding of CS2 formation under Claus plant conditions is limited, and in the following section, we provide the rationale for conducting experiments on the selected reaction systems. CS2 Formation in the Claus RF. It is widely believed that CS2 is formed from reactions between hydrocarbons and sulfur-containing species in the RF.1,3 Generally, methane is the key hydrocarbon component of the Claus plant feed and may be present at concentrations of up to a few percent in the acid gas feed stream. It is not unreasonable to expect that methane will be completely consumed via oxidation reactions in the flame region of the RF. Hence, in the past, the presence of methane in the RF/WHB exit stream had been attributed primarily to an incomplete mixing of CH4 with an air stream. However, recent laboratory data and simulation studies show that methane oxidation is much slower than the hydrogen sulfide oxidation
10.1021/ie030515+ CCC: $27.50 © 2004 American Chemical Society Published on Web 05/20/2004
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reaction10 and that the presence of methane in the RF/ WHB exit stream can be a manifestation of kinetic and not necessarily of mass-transfer limitations. That is, H2S competes successfully for O2 with CH4 under the partial oxidation condition, resulting in incomplete consumption of methane via the oxidation reaction. Thus, it may be reasonable to expect that the product gas stream leaving the flame region of the RF may contain CH4 at significantly high concentrations. Moreover, sulfur and hydrogen sulfide are also present at significant concentrations in the gas stream leaving the flame region. The sulfur in the RF is likely produced in the RF by thermal decomposition of H2S and/or by the overall reaction of SO2 with H2S; each reaction proceeds through numerous elementary reaction steps involving sulfur-, hydrogen-, and oxygen-containing radical species. It is also important to mention here that as much as 65% of the total sulfur, present as H2S in the RF feed, may be recovered from the RF product gas. In summary, the presence of sulfur and H2S at significant concentrations coupled with the high temperatures (900-1300 °C) and long residence times (0.52.0 s) in the postflame region of the RF appears to be adequate for the occurrence of methane-sulfur and methane-hydrogen sulfide reactions that may lead to the production of CS2. There is limited information on gas-phase reaction kinetics for these reactions at such conditions. The present study was initiated to gather data on the methane-sulfur and methane-hydrogen sulfide reactions from experiments conducted in quartz tubular flow reactors over the temperature range of 800-1250 °C. Previous Studies Methane and Sulfur Reaction. The overall reaction between methane and sulfur is a well-established commercial route for CS2 production. The reaction is carried out using silica gel catalysts at temperatures of 500650 °C and pressures of 400-700 kPa. The reaction for CS2 production may be written as
CH4 + 2S2 f CS2 + 2H2S
(1)
Several studies on the CH4-sulfur reaction were conducted from mid-1940 through the 1950s following a patent disclosure by de Simo.11 These investigations ranged from thermodynamic equilibrium calculations12 to catalyst evaluation and kinetic studies.13-17 An early study13 established that silica gel is an efficient catalyst for CS2 production. Subsequently, a number of studies investigated the kinetics of the catalytic reaction15-17 and that of the homogeneous gas-phase reaction14 over 500-700 °C. Interestingly, the activation energy of the homogeneous reaction14 was found to be similar to that for heterogeneous reaction15,17 and ranged from 131 to 160 kJ/mol. In a more recent study, the reaction between methane and sulfur was investigated to identify suitable catalysts for methane oxidative coupling over 500-800 °C.18 The main product of the reaction was observed to be CS2; however, no kinetic analyses was presented. Methane and Hydrogen Sulfide Reaction. In comparison to the methane-sulfur reaction, the reaction between methane and hydrogen sulfide has been less studied. One of the earliest reported studies on the reaction between methane and hydrogen sulfide was that by Waterman and van Vlodrop.19 They reported
data from a quartz tubular reactor over 1080-1280 °C and 0.4-0.7 s and inferred that the following reaction takes place:
CH4 + 2H2S f CS2 + 4H2
(2)
In a relatively recent study, Megalofonous and Pappayanankos20 reported equilibrium calculations for the CH4-H2S feed systems as well as experimental data for the homogeneous and heterogeneous reactions over 713-860 °C. Although they reported hydrogen yields, no experimental data for CS2 were presented. More recently, two studies funded by the U.S. Department of Energy (DOE) investigated two different aspects of the H2S-CH4 reaction system. In the first study,21 several catalysts were evaluated for their effectiveness in CS2 production over 600-1200 °C at a fixed residence time of 1.0 s. The second study22 reported thermodynamic equilibrium calculations for the CH4-H2S system, wherein carbon was considered as one of the products. Their calculation showed that, at a high H2S-CH4 feed ratio, the occurrence of coke formation was not thermodynamically favored. In summary, we reach the conclusion that there is very limited information on gas-phase reactions of methane with sulfur and hydrogen sulfide at temperatures encountered in the Claus front-end RF, i.e., over the temperature range of 900-1300 °C. Experimental Section Equipment Description. The details of the experimental apparatus and procedure have been described elsewhere.7,23 Briefly, the experimental setup consisted of one of four coiled quartz tubular reactors (available in lengths of 1.6, 3.2, 6.4, and 16 m), each 5.0 mm in diameter, placed in a Thermolyne 46240CM muffle furnace (internal dimensions: 47 cm × 48 cm × 71 cm) that could be operated over the temperature range of 600-1700 °C. The gas temperature profile along the reactor was measured with shielded thermocouples and has been reported elsewhere.7 The measured gas temperatures along the reactor length were found to be equal to the furnace set-point temperature within the accuracy of the thermocouple measurements ((0.5%). Thus, we consider that all experiments were conducted under isothermal reactor conditions. The reactant gas flow rates were controlled using an eight-channel Linde mass flow controller. The hot product gas stream from the reactor was first quenched to temperatures of less than 500 °C in less than 15 ms by a combination of free and forced convection in a short length of reactor extending out of the furnace. The quenched gas was further cooled to temperatures of less than 150 °C in a cocurrent double-pipe quartz heat exchanger using cold water. Any sulfur present in the product stream was removed partially in the quartz heat exchanger and completely in a downstream sulfur trap, which consisted of a 25-mm-diameter glass tube packed tightly with glass wool. The inlet and the sulfur-free outlet gas streams were analyzed online by a Varian gas chromatograph equipped with a thermal conductivity detector and a pulsed flame photometric detector. Preparation of a Sulfur-Laden Stream. A description for the preparation of a sulfur-containing stream and the measurement of sulfur concentrations have been described earlier.7,23 Briefly, it involved allowing the flow of the nitrogen gas over liquid sulfur contained
3306 Ind. Eng. Chem. Res., Vol. 43, No. 13, 2004 Table 1. Experimental Conditions for CH4-S2 and CH4-H2S Reactions feed composition (mol %) reaction system
CH4
S2
CH4-S2 CH4-H2S
2.45 2.45
0.9 ( 0.09
range of test conditions H2S
inlet pressure (kPa)
reactor temp (°C)
residence time (ms)
1.75
101-106 108-111
800-1250 800-1250
90-530 380-1400
Table 2. Results of CH4-S2 Experiments Conducted with a Feed Composition of 2.45 mol % CH4 and 0.9 mol % S2 in Nitrogena
temp (°C)
residence time (s)
element balance (1-total in product/total in feed) × 100
product composition (mol %) CH4
CS2
H2S
C2H4
C
H
S
0.1 1.0 0.9 5.1 4.6 3.7 22.8
0.4 1.0 1.2 13.32 22.8 30.2 54.7
97.9b 71.7b 36.3b -0.2 -17.8 -0.3 -10.7
800 900 1000 1100 1150 1200 1250
0.13 0.12 0.11 0.10 0.10 0.09 0.09
2.43 2.30 2.13 1.74 1.52 1.36 0.88
0.01 0.13 0.29 0.54 0.73 0.84 0.85
Data from the 1.6 m Reactor 0.01 0.00 0.25 0.00 0.57 0.01 0.72 0.02 0.66 0.04 0.53 0.08 0.30 0.08
800 850 900 950 1000 1050 1100 1150 1200 1250
0.26 0.25 0.24 0.23 0.22 0.21 0.20 0.20 0.19 0.18
2.34 2.29 2.23 2.13 2.03 1.92 1.70 1.37 1.13 0.57
0.06 0.12 0.22 0.28 0.36 0.47 0.61 0.67 0.75 0.76
Data from the 3.2 m Reactor 0.14 0.00 0.23 0.00 0.45 0.00 0.64 0.01 0.78 0.01 0.81 0.02 0.67 0.02 0.53 0.03 0.33 0.05 0.16 0.45
1.9 1.7 -0.2 1.1 1.5 1.2 4.3 14.4 19.0 9.0
1.7 1.9 -0.2 0.0 0.8 4.6 16.3 32.2 45.1 55.2
85.3b 73.5b 49.8b 33.0b 17.1b 3.3 -4.5 -3.9 -1.9 6.1
800 850 900 950 1000 1050 1100 1150 1200 1250
0.53 0.50 0.48 0.47 0.45 0.43 0.42 0.40 0.39 0.38
2.33 2.20 2.13 2.00 1.89 1.89 1.69 1.31 0.68 0.08
0.12 0.21 0.32 0.39 0.43 0.56 0.62 0.74 0.80 0.79
Data from the 6.4 m Reactor 0.25 0.00 0.45 0.00 0.65 0.00 0.81 0.00 0.92 0.02 0.78 0.02 0.50 0.05 0.29 0.09 0.14 0.38 0.05 0.54
0.1 1.8 -0.2 2.1 3.9 -1.9 1.8 9.0 8.9 20.7
-0.3 1.1 -0.4 1.6 3.3 6.0 18.8 37.0 53.8 73.8
73.3 52.3 28.4 11.6 1.0 -5.9 3.2 2.0 3.4 9.8
a A value of zero represents perfect element balance; a positive value indicates that some of the element is likely present in species unaccounted for in the product stream; a negative value theoretically should not occur and is indicative of the errors associated with the compositional analysis. b The high positive S balance is a result of the unreacted S2 in the product that was not accounted for.
in a round-bottomed flask placed in a heated mantle. The liquid sulfur was maintained at a constant temperature. The amount of sulfur carried by the nitrogen stream was a function of the nitrogen flow rate and the liquid sulfur temperature. The sulfur content of the gas stream was determined in separate experiments in which the sulfur-carrying stream was allowed to flow for a fixed amount of time (10 or 15 min) through a cooled tube containing glass wool to trap the solid sulfur. The amount of sulfur trapped on the glass wool was determined gravimetrically. From the known nitrogen gas flow rate, the time of flow, and the measured amount of sulfur trapped, the sulfur concentration was calculated and expressed as mole percent of S2. The choice of S2 as a representative species for sulfur was based on the fact that S2 is thermodynamically predicted to be the dominant sulfur species present at dilute conditions and high temperatures of our experiments. Five experimental runs were carried out to determine the sulfur concentration, and the average deviation was found to be (10%. Compositional Analysis. The concentration of methane in the reactant stream and the composition of the product gas stream was analyzed by an inline gas chromatograph. For each experimental run (i.e., at a fixed temperature and reactor length), two duplicates
of two sets of compositional analyses were performed. One set of analyses involved measurement of CH4, C2H6, C2H4, C2H2, and H2S concentrations. The separation of these species was achieved in a HayeSep Q column. The second set of analyses, which was performed soon after the first one, allowed the measurement of CS2 and H2S concentrations, with the separation achieved in a Chromosil 310 column. It is important to mention that because of complexities involved with the alteration of gas chromatography operating conditions, unfortunately, no analysis for hydrogen gas was performed. The errors in gas concentrations for each measured species were within 4%. Experimental Conditions. The details of the experimental conditions for both the CH4-S2 and CH4H2S reactions are reported in Table 1. Results and Discussion CH4-Sulfur Reaction. The results of the methanesulfur reaction experiments are provided in Table 2. In addition to the distributions of measured product compositions, element balances are given from known reactant compositions and measured product compositions. All experiments were conducted under conditions of sulfur as a limiting reactant with respect to the CS2 formation reaction (1). Because S2 is by far the most
Ind. Eng. Chem. Res., Vol. 43, No. 13, 2004 3307
Figure 1. Experimentally measured CS2 concentration in the reactor product stream as a function of the temperature.
prominent sulfur species at the temperatures of our experiments, any reference to sulfur as a reactant would correspond to S2. (a) CS2 Formation. The concentration of CS2 in the product stream exiting the three quartz reactors as a function of temperature is shown in Figure 1. Clearly, the CS2 concentration can be observed to increase with temperature and then level off at around 1200 °C. At temperatures of less than 1200 °C, CS2 formation increases with the reaction or residence time. The data at 1200 and 1250 °C indicate no particular trend with the residence time; however, at these high temperatures an almost complete conversion of the sulfur into CS2 occurs. Possible explanations for the observed CS2 concentration at levels lower than that corresponding to complete conversion of feed S2 to CS2 include the consumption of CS2 via its reaction with products of the methane pyrolysis reaction such as H2, C2H4, etc., and variation in the feed sulfur concentration (measured to be on the order of 10%). The data at 1000 °C and higher temperatures from the three reactors indicate that a majority of CS2 formation occurs in less than 100 ms and a relatively smaller amount of CS2 is produced in the following 200-300 ms. This is expected because the experiments were conducted under conditions of S2 as a limiting reactant. Thus, the rate of the CS2 formation reaction is high initially but falls rapidly with increasing residence time as the S2 concentration continues to decrease. Assuming that CS2 and H2S are the only known sulfur-containing products of the reaction, it can be seen from the sulfur balance data that almost a complete consumption of S2 occurs at 1000 °C in the 6.4 m reactor (i.e., within 400 ms) and at 1100 °C in the 1.6 m reactor (i.e., within 100 ms). Thus, the increase in the amount of CS2 formed in the two longer reactors (6.4 and 16.0 m reactors) at 1100 °C and higher temperature occurs at the expense of H2S formation. (b) H2S Formation. The concentration of H2S measured in the quenched product gas stream exiting the three quartz reactors is presented in Figure 2. The figure also shows predicted H2S concentrations, and this aspect is discussed later. Figure 2 has several notable features. Overall, it can be seen that for each reactor the H2S concentration increases with temperature initially, attains a maximum, and then falls rapidly with a further increase in temperature. The maximum H2S concentration is observed at 1000, 1050, and 1100 °C
Figure 2. Experimentally measured H2S concentration in the reactor product stream as a function of the temperature. Lines are predictions from eq 10.
for the 6.4, 3.2, and 1.6 m reactors, respectively. The maximum value of ∼0.9 mol % H2S for the 6.4 m reactor also corresponds to the complete conversion of feed sulfur to stoichiometric amounts of H2S and CS2 according to reaction (2). It may be noted that, at lower temperatures (less than 1000 °C), the amount of H2S formed at any temperature increases with increasing reaction or residence time. On the other hand, at higher temperatures, the concentration of H2S decreases with increasing residence time, indicating the occurrence of H2S consumption reaction(s). It is evident that, at higher temperatures (>1000 °C), H2S is formed in short residence times and consumed with increasing reaction time. However, it is not clear whether all of the feed sulfur is first converted to CS2 and H2S according to overall reaction (1) or whether a portion of the feed sulfur is consumed by other parallel reaction pathways that do not lead to H2S formation. For example, it is possible that some of the S2 is consumed according to the following reaction (3), which proceeds without the formation of H2S. Unfortunately, the data do not lend
CH4 + S2 f CS2 + 2H2
(3)
themselves to conclusively establishing whether and to the extent reaction (3) occurs. The initiation step for reaction (3) is likely the same elementary reaction as that for reaction (1). However, while the H2S-forming reaction is the only termination pathway worth considering at lower temperature, other reaction pathways leading to H2 formation may also become important at higher temperatures. This is true because higher temperatures favor H2 formation over H2S production. (c) Formation of C2H4. At temperatures greater than 1000 °C, pyrolysis of methane is expected to proceed at appreciable rates, yielding measurable concentrations of key reaction products that include H2 and C2 hydrocarbons such as C2H2, C2H4, and C2H6. However, in our experiments, the only product of methane pyrolysis and/or oxidative coupling reaction observed at measurable concentrations was C2H4. The concentration of C2H4 in the product gas streams from three quartz reactors has been plotted as a function of temperature in Figure 3. It can be seen from Figure 3 that ethylene formation is negligible at temperatures below 900 °C. The amount of ethylene produced increases only slightly
3308 Ind. Eng. Chem. Res., Vol. 43, No. 13, 2004
Figure 3. Experimentally measured C2H4 concentration in the reactor product stream as a function of the temperature.
as the reactor temperature is increased from 900 to 1150 °C. However, a sharp increase in the C2H4 concentration occurs at 1150 and 1200 °C for the 6.4 and 3.2 m reactors, respectively. For the shortest reactor (1.6 m) employed in this study, the residence time is likely not long enough to effect a significant amount of C2H4 production, as evidenced by only a slight increase in the C2H4 concentration in the product gas at temperatures up to 1250 °C. (d) Kinetics of the CH4-Sulfur Reaction. The discussion in this section is based on the consideration that the reaction between methane and sulfur occurs predominantly in the gas phase. Whereas it is known that silica may promote the methane-sulfur reaction, we believe that, for the catalytic reaction to make a significant contribution to the overall methane-sulfur reaction, the surface area should be orders of magnitude higher than that provided by our quartz tubular reactor. Further, from the analysis of data in previous experimental studies,17,18 we inferred that this is a reasonable assumption. For example, Anderson et al.18 reported data for experiments conducted in an empty quartz tubular reactor and in a reactor containing catalyst held on a silica frit by silica wool. The reactant conversions at 700 °C for at least four catalysts were found to be equal to those from an empty reactor. The fact that the reactant conversions from a catalyst-packed reactor were similar to those from an empty quartz tubular reactor despite the availability of a high surface area of silica frit and silica wool implies that the surface reaction had a minimal contribution to the overall reaction. Accordingly, we assume that the reactions in our quartz reactor occur primarily in the gas phase. Assuming that the formation of CS2 occurs via the gas-phase reaction of CH4 and S2 according to reaction (1), it would be expected that the measured H2S-CS2 ratio will be 2. From the plot of the H2S-CS2 ratio presented in Figure 4, it may be observed that the H2SCS2 ratio remains around 2 at temperatures below 1000 °C but decreases rapidly as the temperatures exceed 1000 °C. The decrease in the H2S-CS2 ratio is a result of an increase in the CS2 concentration with a simultaneous and corresponding decrease in the H2S concentration. The data at 1000 °C and lower temperatures also indicate that a relatively small amount of products of side reactions are produced. Therefore, these lower temperature data were used to obtain the Arrhenius parameters for the rate constant k1 of the following
Figure 4. H2S/CS2 ratio in the reactor product stream as a function of the temperature.
Figure 5. Rate constant for CS2 formation by the methane-sulfur reaction (k in m3/kmol‚s).
reaction rate expression describing CS2 formation according to overall reaction (1):
rCS2 ) k1CCH4CS2
(4)
Because reaction (1) is accompanied by no net volume change and because the reactor can be treated as an isothermal reactor, the rate of formation of CS2 can be represented as rCS2 ) dCCS2/dtres. An analytical expression for the CS2 concentration was obtained, through integration of eq 4, as a function of rate constant k1, actual residence time tres, and the initial concentrations and conversions of CH4 and S2. From measured concentrations and known residence time, the rate constant was calculated for each experimental run. The Arrhenius plot for the rate constant has been shown in Figure 5. It may be observed from Figure 5 that the calculated rate constant corresponding to the data at 800 °C from the 1.6 m reactor does not follow the trend exhibited by other data. These particular data were deemed to be an outlier point, and the deviation from the expected trend is likely the result of the inaccuracy associated with the measurement of a very low CS2 concentration. All data on the plot except the outlier was regressed to obtain the following Arrhenius expression for the rate constant k1:
Ind. Eng. Chem. Res., Vol. 43, No. 13, 2004 3309
k1 ) 5.53 × 1010 exp[-19320/T] m3/kmol‚s
(5)
The activation energy of 160 kJ/mol is higher than those reported for both the homogeneous gas-phase reaction14 and the catalytic reaction on silica gel.15,17 In what follows, we examine the most likely initial reaction between sulfur and methane and its likelihood of being the rate-controlling step for the CS2 formation reaction. It should be recalled that the reactant gas is a mixture of CH4 and sulfur vapor in inert nitrogen. At the high temperatures employed in our study, S2 is expected to be the dominant sulfur species; however, S will also be present albeit at very low concentrations. The key initial reactions for methane involving the most abundant species in the reacting mixture are expected to be as follows:
CH4 + S f CH3 + HS CH4 f CH3 + H CH4 + S2 f CH3 + HS2
∆H ) 83 kJ/mol ∆H ) 439 kJ/mol
(6) (7)
∆H ) 197.5 kJ/mol (8)
The heats of formation of all species except HS2 in the above reactions (6)-(8) and those in reactions (11)-(15) were taken from JANAF Thermochemical Tables.24 The heat of formation of HS2 was obtained from Alzueta et al.25 Using the rate kinetics for reactions (7) and (6) given in refs 26 and 27, respectively, and assuming that the sulfur vapor contains S and S2 at equilibrium concentrations, it can be calculated that methane will react at a rate 2 orders of magnitude faster via reaction (7) than via reaction (6). Considering the species abundance and endothermicities of reactions (6)-(8), it is estimated that reaction (8) is the most dominant pathway for the initial reaction of methane. Applying the Marcus relationship28 to estimate the activation energy for the highly endothermic reaction (6), one obtains an activation energy of 198 kJ/mol, equal to the heat of reaction. This activation energy is significantly higher than that obtained in the present study for the overall reaction. Therefore, it appears that, although reaction (8) may be the most dominant initial reaction for methane, it is likely not the rate-controlling step for the overall reaction. (e) Prediction of H2S Formation/Consumption. We have attempted to model the H2S formation/ consumption behavior. A simple model comprising a single H2S formation reaction and a single H2S consumption reaction was considered. H2S formation was assumed to occur via reaction (1), whereas consumption was assumed to occur via the following H2S decomposition reaction:
H2S + M f products
(9)
Now, in the absence of other reacting species, the thermal decomposition in the temperature range of our interest (i.e., 800-1300 °C) proceeds through a series of elementary reaction steps involving H- and S-containing radicals to produce stable species H2 and S2. However, in the presence of CH4, it is likely that several of these radical species produced as a result of H2S decomposition will directly interact with methane or with radical species derived from methane. Thus, in the presence of CH4, new termination pathways for radicals that do not necessarily lead to the formation of S2, also referred to hereafter as free sulfur, will likely open up.
The assumption of whether H2S decomposition leads to S2 formation directly influences the results of the simplistic modeling of H2S formation/consumption because any S2 produced via H2S decomposition is treated in the model as a reactant for H2S formation via reaction (1). Unfortunately, without the knowledge of key elementary reactions involved in this rather complex system, it is difficult to ascertain the extent to which H2S decomposition results in the formation of free S2. In the simple model, the net rate of formation of H2S was expressed by the following equation:
(rH2S)net ) 2k1CCH4CS2 - k2CtCH2S
(10)
where k1 is the rate constant obtained in this study and k2 ()1.68 × 1011 exp[-28900/T] m3/kmol‚s) is the rate constant for H2S decomposition obtained from one of our previous studies.29 The differential equations resulting from material balance on H2S, CH4, and S2 for isothermal plug-flow reactors were solved using the ordinary differential equation solver in MATLAB. Calculations were performed for two cases: (case i) H2S consumption leading to the formation of free S2 and (case ii) H2S consumption not resulting in the formation of free S2. Accordingly, for case ii, any H2S that is consumed produces S2, which, in turn, can react with CH4 to produce more CS2 and H2S. However, for case ii, it is assumed that H2S is consumed but does not produce S2 and, therefore, cannot react with CH4 for further production of CS2. The simulation results for the two cases showed that the net H2S formation at temperatures lower than 1000 °C for the two cases did not differ. This is not surprising because H2S decomposition at these lower temperatures occurs at insignificant rates and the net rate of H2S formation is dictated by the formation reaction (1). At higher temperatures, however, expectedly the predicted H2S concentrations for case i were significantly higher than those for case ii. However, the predictions for case ii were closer to the measured data than those for case i. A comparison of the measured H2S concentration with predictions for case ii is shown in Figure 2. The difference between the measured and predicted H2S concentrations is not surprising considering the simplistic kinetic model employed. Nonetheless, the two-reaction model is able to capture the H2S formation/consumption behavior reasonably well. A possible explanation for measured H2S concentrations being lower than predicted concentrations is that some of S2 reacts with CH4 to form CS2 via reaction (3) without formation of H2S. This has not been considered in the simple model. It needs to be emphasized that this model describes only H2S formation/consumption and that inclusion of other sulfur reactions needs to be incorporated for reliable prediction of CS2 formation and for closure on sulfur mass balance. CH4-Hydrogen Sulfide Reaction. The results of the methane-hydrogen sulfide reaction experiments are provided in Table 3. The table contains the product composition as well as data on element balances calculated from known reactant composition and measured product composition. The concentration of hydrogen was not measured in the experiments. (a) CS2 Formation and H2S Consumption. The experimental data on CS2 formation from CH4 and H2S mixtures are reported in Figure 6. These data were obtained from the 6.4 and 16.0 m reactors. The longest reactor used for the CH4-S2 reaction was 6.4 m. The
3310 Ind. Eng. Chem. Res., Vol. 43, No. 13, 2004 Table 3. Results of CH4-H2S Experiments Conducted with a Feed Composition of 2.45 mol % CH4 and 1.75 mol % H2S in Nitrogena element balance (1-total in product/total in feed) × 100
product composition temp (°C)
residence time (s)
CH4
CS2
800 850 900 950 1000 1050 1100 1150 1200 1250
0.54 0.52 0.49 0.47 0.45 0.44 0.42 0.41 0.39 0.38
2.46 2.45 2.44 2.42 2.41 2.27 2.13 1.83 0.97 0.32
800 850 900 950 1000 1050 1100 1150 1200 1250
1.39 1.32 1.26 1.22 1.17 1.13 1.08 1.05 1.01 0.98
2.50 2.47 2.42 2.32 2.30 2.22 1.84 1.12 0.55 0.11
H2S
C
H
Data from the 6.4 m Reactor 0.00 1.75 0.00 0.00 1.73 0.00 0.00 1.73 0.00 0.00 1.64 0.00 0.01 1.49 0.00 0.09 1.30 0.00 0.23 0.99 0.02 0.47 0.49 0.05 0.61 0.28 0.14 0.68 0.09 0.45
-0.4 0.1 0.5 1.1 1.2 3.6 2.5 1.8 24.1 22.2
0.0 0.6 1.0 2.7 5.5 12.4 20.8 36.1 62.6 75.4
0.0 0.9 1.3 6.0 13.7 14.9 17.8 18.3 13.9 17.0
Data from the 16.0 m Reactor 0.00 1.75 0.00 0.00 1.73 0.00 0.02 1.64 0.00 0.10 1.47 0.00 0.14 1.21 0.01 0.23 0.88 0.05 0.49 0.53 0.11 0.57 0.24 0.28 0.67 0.06 0.47 0.70 0.04 0.68
-2.0 -0.7 0.4 1.0 -0.5 -3.9 -4.0 7.7 12.1 11.4
-1.2 0.0 2.8 8.3 12.6 18.8 33.7 54.3 68.6 75.8
0.0 0.9 3.8 4.2 14.1 23.2 14.2 20.6 20.4 17,5
C2H4
S
a A value of zero represents perfect element balance; a positive value indicates that some of the element is likely present as chemical species unaccounted for in the product stream; a negative value theoretically should not occur and is indicative of the errors associated with the compositional analysis.
Figure 6. Experimentally measured CS2 concentration in the reactor product stream as a function of the temperature.
Figure 7. H2S conversions in two quartz reactors as a function of the temperature. (symbols correspond to data, and lines represent predictions).
CS2 concentration as a function of the temperature shows a trend similar to that observed for the CH4-S2 system. However, the rates of formation of CS2 in the CH4-H2S system were much lower than those in the CH4-S2 system. Moreover, little CS2 was observed at temperatures below 900 °C, when an insignificant amount of H2S decomposition occurs. The conversion of H2S as a function of temperature is shown in Figure 7. The solid lines are the predicted H2S conversions with the consideration that H2S thermal decomposition is the sole reaction occurring. The H2S conversions were calculated using the rate expression discussed earlier, i.e., (-rH2S) ) k2CtCH2S. At conditions of our experiments, the H2S decomposition reaction in the absence of any sulfur-consuming reaction would be equilibrium-limited because of the occurrence of a reverse reaction between H2 and S2. However, in the presence of methane, sulfur is converted into CS2, shifting the equilibrium toward
further decomposition of H2S. The close agreement between the measured and predicted H2S conversions was quite surprising. This suggests that the H2S decomposition reaction proceeds independently of any side reactions and that the products of this reaction then interact with methane and methane-derived radical species to produce CS2. From the sulfur balance data, it may be noted that a majority but not all of the H2S consumed is converted into CS2. It is possible that the balance of converted H2S was converted to either free sulfur (i.e., S or S2) or other sulfur-containing products that were not analyzed for. Because the product gas stream was not analyzed for other sulfur-containing species such as methyl mercaptan and because free sulfur could not be analyzed on the analytical system, it is not possible to comment on where the balance of the sulfur not present as CS2 ends up. It may be mentioned that any free sulfur in the product stream
Ind. Eng. Chem. Res., Vol. 43, No. 13, 2004 3311
H2S + M f H2 + S +M
∆H298 ) 297.6 kJ/mol (11)
H2S + M f H + SH +M
∆H298 ) 377.8 kJ/mol (12)
In the past, the scission of the S-H bond of H2S according to reaction (12) was considered to be the initiation reaction. However, recent experimental studies30,31 have shown that while both reactions (11) and (12) occur at the early stages, reaction (11) is the more dominant pathway for H2S decomposition. This is also favored energetically based on the endothermicity of the reaction. During thermal decomposition of H2S, the S and H radicals produced via reactions (11) and (12) react further with H2S according to the following reactions:
Figure 8. Experimentally measured C2H4 concentration in the reactor product stream as a function of the temperature. (solid symbols, CH4-H2S feed; open symbols, CH4-S2 feed).
that may have dropped out as solid yellow sulfur in the condenser and the sulfur trap could not be visually observed as a result of deposition of coke in these components of the reactor system. Deposition of coke was also observed in the reactors and is briefly discussed in a later section. (b) C2H4 Formation. As in case of CH4-S2 systems, ethylene was also observed in the product gas at the reactor exit for CH4-H2S systems. In Figure 8, C2H4 produced in the 6.4 and 16.0 m reactors from CH4-H2S feed mixtures are presented. For comparison, the ethylene data from the 6.4 m reactor for the CH4-S2 system are also shown in Figure 8. For the CH4-H2S system, ethylene formation is negligible below 1000 °C but increases steadily at higher temperatures. Upon comparison of the data from the 6.4 m reactor for CH4H2S and CH4-S2 systems, it can be seen that a lower amount of C2H4 is formed in the CH4-H2S feed system. These data may suggest that H2S is more effective than S2 in suppressing C2H4 formation. However, a closer look at the data indicates that this may not necessarily be true. To elaborate further on this point, if we consider the data of Table 1, it can be seen that, for the CH4-S2 system, the feed S2 is completely consumed in less than 200 ms (i.e., in the 3.2 m reactor) at 1000 °C and within 100 ms at temperatures of 1100 °C and higher. Therefore, in the 6.4 m reactor, a majority of the S2 is consumed within the initial short section of the reactor, leaving plenty of reaction time for methane (rich in a S2 devoid reaction mixture) to pyrolyze and produce C2H4. Now, in the CH4-S2 feed systems, the reaction mixture after S2 has apparently been consumed completely does contain H2S. However, the concentration of H2S is lower than that in a CH4-H2S system. Thus, it does appear that H2S is effective in suppressing C2H4 formation, but it is not necessarily true that it more effective than S2. (c) Possible Mechanism of the CH4-Hydrogen Sulfide Reaction. The decomposition of H2S is the critical step in formation of CS2 from the reaction between CH4 and H2S. The thermal decomposition of H2S has been studied by several researchers and is generally accepted to follow two initiation steps:
H2S + S f HS + HS
∆H ) 17 kJ/mol (13)
H2S + H f H2 + SH +M
∆H ) -62.5 kJ/mol (14)
Using the rate constants for reactions (13) and (14) reported in ref 32, it can be calculated that the two reactions proceed rapidly,23 with the net result being that, for every H2S molecule producing a S or H radical, another molecule is consumed according to reactions (13) and/or (14). Thus, the net rate of H2S consumption is observed to be twice that due to the initiation reaction. In the presence of CH4, the S and H radicals could also react with CH4 according to the following reactions:
CH4 + S f CH3 + HS
∆H ) 83 kJ/mol
CH4 + H f CH3 + H2
∆H ) 36 kJ/mol (15)
(6)
Using the reaction rate constant for reactions (6) and (15) given in ref 27 and that for reactions (13) and (14) reported in ref 32, it can be calculated that the consumption of the H radical via reaction with H2S is 3 or 4 orders of magnitude greater than that with CH4; however, the consumption of the S radical via reaction with H2S is merely 1 order of magnitude higher than that with CH4. It is recognized that there are other possible reactions for consumption of S and H radicals; however, we have focused merely on the likely initial steps of the reactions. Now, the predictions of H2S conversions shown in Figure 7 were based on the assumption that the net H2S consumption is twice that due to the initiation step. The agreement between the measured and predicted H2S conversions indicates that the radical species S and H react more readily with H2S than with CH4. This leads to an abundance of SH radicals, which would be expected to play the critical role in formation of CS2 via its reaction with CH4. Now, SH could be involved in the key initiation step for CH4 consumption in the CH4H2S reacting system according to the following reaction:
CH4 + SH f CH3 + H2S
∆H298 ) 58 kJ/mol (16)
However, this reaction produces H2S, and it would be expected that the net rate of H2S conversion in the presence of CH4 would be lower than that in the absence of CH4. The agreement of experimentally measured H2S conversion with prediction based on the assumption that H2S is converted only via thermal decomposition suggests that reaction (16) plays a less important role in
3312 Ind. Eng. Chem. Res., Vol. 43, No. 13, 2004
Figure 9. Picture of a 3.2 m reactor showing the absence of coke deposition on the reactor tube immediately downstream of the point of mixing (mixer) of the two reactant streams (CH4-S2 feed).
Figure 10. Picture of a 6.4 m reactor showing coke deposition on the reactor tube downstream of the mixer (CH4-H2S feed).
the critical step for CS2 formation, i.e., formation of the CH3 radical. There are several other possible mechanisms for CH3 formation including scission of the C-H bond via thermal decomposition. The preceding discussion only indicates the complex nature of the seemingly simple reaction system. More detailed investigation is required to establish the mechanism of CS2 formation from the reaction between CH4 and H2S. Coke Deposition. The presence of sulfur appears to affect coke deposition, which was observed in the reactors at temperatures exceeding 1050 °C for both reacting systems, i.e., CH4-S2 and CH4-H2S. Here we define deposition as the actual layout of the coke layer on the surface of the tube irrespective of whether coke formation originates in the gas phase or on the surface. The photographs of coke deposition in the reactors as observed at the end of 1200 °C experimental runs are shown in Figures 9 and 10 for methane-sulfur and methane-hydrogen sulfide systems, respectively. In the figures, the two vertical inlets act as preheat sections to bring the two reactant streams, methane and sulfur or hydrogen sulfide, close to the reaction temperature
prior to reacting with each other. The two gas streams containing the individual reactants collide head-on at the mixer, and the reactants in a well-mixed stream continue to react downstream along the reactor. From Figure 9, it can be seen that coking had already initiated in the heat-up section or inlet carrying the methane stream. An interesting feature to note is that no coke deposition is observed in the short length of the reactor downstream of the mixer. However, coke deposition can be observed in the further portion of the reactor. This indicates that either the coking reaction is inhibited or the deposition of coke is inhibited when the methane stream comes in contact with a sulfur-laden gas stream. After a majority of sulfur is consumed, coke deposition starts to occur again. In Figure 10, coke deposition in the reactors during the reaction between methane and hydrogen sulfide is shown. The coking in the reactor can be observed in the inlet to the reactor that supplied the methane-containing gas stream. Unlike the case for the methane-sulfur reaction, for the methane-hydrogen sulfide reaction system, deposition of coke was observed at the mixer as well as the section immediately and farther downstream of the mixer. These results indicate that the coking as well as coke deposition is unaffected by the presence of hydrogen sulfide. It may also be inferred that sulfur acts as a more effective coking/coke deposition inhibitor. Further investigation would be required to establish whether sulfur actually acts as an inhibitor to the coking reaction or merely prevents deposition of coke on the walls of the reactor. Conclusions An experimental study on the gas-phase reaction between methane and sulfur and between methane and hydrogen sulfide was completed over a temperature range of 800-1250 °C, which is representative of conditions in a Claus front-end RF. From the experiments, it was concluded that methane is kinetically favored to react with sulfur to produce carbon disulfide. The reaction between sulfur and methane is very rapid, and a complete conversion of the limiting reactant, sulfur, occurs in less than 100 ms at 1100 °C. At lower temperatures (
