Initiated Sulfonation of Methane to Methanesulfonic - American

MSA was achieved at 1000 psi of CH4, 30 wt % of SO3, and 350 µmol of K2S2O8 at 333 K ... At both 333 and 363 K, the SO3 conversion to MSA passes thro...
6 downloads 0 Views 110KB Size
Download PDF
736

Ind. Eng. Chem. Res. 2001, 40, 736-742

K2S2O8-Initiated Sulfonation of Methane to Methanesulfonic Acid Lisa J. Lobree and Alexis T. Bell* Department of Chemical Engineering, University of California, Berkeley, California 94720

The direct synthesis of methanesulfonic acid (MSA) from CH4 and SO3 has been successfully achieved in an H2SO4 solvent using K2S2O8 as an initiator. The maximum SO3 conversion to MSA was achieved at 1000 psi of CH4, 30 wt % of SO3, and 350 µmol of K2S2O8 at 333 K for a 2-h reaction time. SO3 conversion to MSA is approximately independent of reaction time above 2 h, at which time the SO3 conversion is observed to plateau at conversions much less than 100%. This plateau has been attributed to a loss of solution SO3 as a result of the formation of CH3(SO3)nH and (CH3SO2)2O. The rate of MSA formation exhibits a positive-order dependence on the CH4 pressure, and this pressure dependence is nonlinear, rising more rapidly for pressures equal to or greater than 800 psi. Addition of SO3 to the solvent medium at concentrations greater than 10% results in a rapid decrease in SO3 conversion as a result of the formation of significant byproducts and CH3(SO3)nH. At both 333 and 363 K, the SO3 conversion to MSA passes through a maximum with increasing K2S2O8 concentration, and then declines to a plateau for higher K2S2O8 concentrations. The position of the maximum is temperature-dependent and occurs at 100 µmol of K2S2O8 at 363 K and 350 µmol at 333 K. A maximum in SO3 conversion to MSA is also observed at 333 K when the temperature is increased from 313 to 363 K. A free radical mechanism has been proposed to account for the observed phenomena. Introduction Methanesulfonic acid (MSA) is a clear, colorless, strong organic acid that is used in catalysis, electroplating, biological, and agricultural chemical preparation and as an electrolyte in electrochemical applications.1 Currently, MSA is produced via the chlorine oxidation of methylmercaptan.2

(1) CH3SH + 3Cl2 + 2H2O f CH3SO2Cl + 5HCl (2)

CH3SO2Cl + H2O f CH3SO3H + HCl CH3SH + 3Cl2 + 3H2O f CH3SO3H + 6HCl

In the first step, methylmercaptan reacts with chlorine and water to form methanesulfonyl chloride, which is then hydrolyzed in the second step to form MSA and hydrochloric acid. For every 1 mol of MSA, 6 mol of hydrochloric acid are formed. The formation of large quantities of HCl is a potential problem, as there is no guarantee that a buyer for HCl will always exist. If the demand for the byproduct HCl falls relative to the demand for MSA, then the remaining HCl must be treated as waste. Consequently, there are strong incentives to develop new methods for the synthesis of MSA that eliminate the coproduction of HCl. There is a precedent in the literature for synthesizing MSA directly from CH4 and SO3. Sen and co-workers investigated the radical-initiated functionalization of CH4.3 These experiments were carried out in fuming sulfuric acid (30 wt % SO3) to which was added 0.010.1 mmol of initiator. The CH4 pressure was set to 1000 psi, and the reaction temperatures were 363-433 K. Several initiators were used, including K2S2O8, metal * To whom correspondence should be addressed. E-mail: [email protected]. Phone: (510) 642-1536. Fax: (510) 642-4778.

salts, and hydrogen peroxide. A limited amount of MSA formation was observed when HgSO4, Ce(SO4)2, and H2O2 were used as the initiators. However, significantly higher rates of MSA formation were observed when K2S2O8 was used as the free radical initiator. The goal of the present work was to investigate the effect of reaction conditions on the synthesis of MSA from CH4 and SO3 using K2S2O8 as an initiator to activate CH4. The effects of CH4 pressure, SO3 concentration, K2S2O8 concentration, temperature, and time were probed. The observed dependence of SO3 conversion to MSA on these parameters is discussed in light of a proposed reaction mechanism. Experimental Section Experiments were performed using the apparatus shown in Figure 1. Reactions were carried out in a 100cm3 high-pressure batch reactor (Parr Instrument Company, 3000 psi maximum) constructed of Hastelloy B. A gas sampling chamber, a gas expansion vessel, and a series of gas scrubbers are located downstream of the reactor. The gas sampling chamber is used to withdraw gaseous samples for analysis by gas chromatography. The gas expansion vessel allows for the pressure in the reactor to drop to approximately 20% of its original value prior to entering the gas scrubbers. Both the NaOH scrubber and the gas-washing bottle containing Carusorb are used to remove sulfur(II) compounds. The mineral oil bubbler allows for visual indication of gas flow under all conditions. Unless otherwise stated, the following procedure was used: (1) K2S2O8 (Mallinckrodt, 99.1%) was added to a glass liner containing a Teflonencased stirring bar, followed by the addition of 3 cm3 of fuming sulfuric acid (Aldrich Chemical Co., ∼27-33 wt % SO3); (2) the glass liner was then transferred to the autoclave, which was sealed and attached to a gashandling system; and (3) the reactor and gas lines were purged with CH4, and then the autoclave was pressur-

10.1021/ie000725b CCC: $20.00 © 2001 American Chemical Society Published on Web 01/06/2001

Ind. Eng. Chem. Res., Vol. 40, No. 3, 2001 737

Figure 1. Schematic of reactor/scrubber experimental apparatus used for MSA synthesis.

ized with CH4 to the desired pressure. The reactor was heated to the desired reaction temperature (heating time of 20-45 min, depending on the final temperature), and this temperature was maintained for 2 h, unless otherwise stated. Following reaction, the vessel was cooled under ambient conditions to room temperature (cooling time of 2.5 h), the gases were vented, and the system was purged with N2. The product solution was then removed from the glass liner and slowly added to 0.5-1 cm3 of H2O to convert the remaining SO3 to H2SO4. Reaction products were characterized by 1H and 13C NMR spectrometry. All spectra were acquired using a Bruker AMX-400 MHz FT-NMR spectrometer. A capillary containing D2O and CH3OH, immersed within the capillary containing the sample, was used as a lock, reference, and integration standard. For calibration, pure MSA (Aldrich Chemical Co., 99.5%) was mixed in varying concentrations with H2SO4 (Fisher Chemical, 95.5 wt % in H2O). Additional SO3 used for the varying SO3 experiments was obtained from Aldrich Chemical Co. (99%). Ultra-high-purity CH4 was obtained from Matheson, and N2 was obtained onsite. The N2 was passed through an oxysorb trap, an ascarite trap, and a molecular sieve trap, in that order, for additional purification. Results To identify and quantify MSA in the reaction products, calibration mixtures containing varying concentrations of MSA in H2SO4 were prepared, and 1H NMR spectra were acquired for each mixture. Figure 2 shows the spectrum for a sample containing 20 wt % MSA with a capillary insert containing D2O and CH3OH. In the region from 2 to 5 ppm, three strong peaks are observed at 4.76, 3.22, and 2.98 ppm. The peak at 4.76 ppm is due to HDO (proton-deuterium exchange between D2O

Figure 2. 1H NMR spectrum of a standard sample containing 20 wt % MSA in H2SO4 (95.5 wt % in H2O).

and CH3OH), while the peak at 3.22 ppm is due to the methyl group in CH3OH. The peak for MSA is located at 2.98 ppm.4 Consistent with the literature, the MSA peak location varies with concentration and shifts downfield with decreasing MSA concentration.5 Figure 3 shows that the reaction of 3 mL of H2SO4(SO3), 100 µmol of K2S2O8, and 1000 psi of CH4 at 363 K for 2 h results in the formation of MSA, as indicated by the peak at 2.94 ppm observed in the 1H NMR spectrum of the reaction products. Integration of this peak reveals that 7.5 mmol of MSA is formed, which agrees well with the quantity observed by Sen and coworkers under identical conditions (7.6 mmol).3 The SO3 conversion to MSA is 38%. Only 3% of the CH4 is converted under these conditions, as CH4 is present in a 14-fold excess relative to SO3. A 35-fold expansion of the region between 3.50 and 4.60 ppm shows that there are three small features at 4.56, 3.72, and 3.69 ppm;

738

Ind. Eng. Chem. Res., Vol. 40, No. 3, 2001

Figure 3. 1H NMR spectrum acquired following the reaction of 3 mL of H2SO4(SO3), 100 µmol of K2S2O8, and 1000 psi of CH4 at 363 K for 2 h.

identification of these peaks is discussed in more detail below. Although not shown, experiments were conducted in the absence of either SO3 or K2S2O8, and in both cases, neither MSA nor any other product observed was detected. A series of experiments was carried out to determine the effects of reaction conditions on the formation of MSA. The effects of reaction time on the conversion of SO3 to MSA are shown in Figure 4A. It is evident that 30% of the SO3 is converted to MSA within the time necessary to heat the reaction from 298 to 363 K (0.75 h) and to cool it back down to 298 K (2.5 h). Increasing the time at 363 K slowly increases the level of SO3 conversion to 40%, most of this gain being achieved within 120 min. The data at zero time in Figure 4A show that significant reaction occurs during reactor heat-up and cooldown; hence, it was of interest to determine the SO3 conversion at reaction temperature while limiting the amount of MSA formed during heat-up and cooldown. This was achieved in a second set of experiments conducted with the following modifications: (1) CH4 was added to the autoclave once the reaction temperature was achieved instead of at room temperature, and (2) the reactor was rapidly quenched, using an ice bath, following reaction (0.5 h). These experiments were conducted for varying times at 333 and 363 K. The results are shown in Figure 4B. At 363 K, the SO3 conversion rises to 15% by 0.5 h, and this conversion level is approximately constant for reaction times up to 6 h. At 333 K, the conversion is observed to rise initially to a value of 23% at 0.5 h, and it then continues to slowly rise to 37% after 4 h of reaction. To determine whether the plateau with time was due to a depletion of K2S2O8, an experiment was conducted in which the reaction was stopped (for the experiment shown in Figure 4B at T ) 363 K and t ) 2 h), and an additional 100 µmol of K2S2O8 was added to the solution. The experiment was repeated for an additional 2 h. No additional MSA formation was observed, suggesting that this plateau is not due to a depletion of K2S2O8. A second experiment of a similar type was carried out, only this time an amount of SO3 equivalent to that present at the onset of the experiment was added after 2 h of reaction. Here, too, no additional MSA was formed, suggesting that the added SO3 is consumed in a manner that makes it unavailable to form additional MSA.

Figure 4. (A) SO3 and CH4 conversion to MSA as a function of time for the reaction of 3 mL of H2SO4(SO3), 100 µmol of K2S2O8, and 1000 psi of CH4 at 363 K for 0-6 h. (B) SO3 conversion to MSA as a function of time for the reaction of 3 mL of H2SO4(SO3), 100 µmol of K2S2O8, and 1000 psi of CH4 (added at reaction temperature) at 333 and 363 K for 0-6 h. Reactor quenched to room temperature, using ice water over 0.5 h.

Figure 5. SO3 and CH4 conversion to MSA as a function of CH4 pressure for the reaction of 3 mL of H2SO4(SO3), 100 µmol of K2S2O8, and 200-1200 psi of CH4 at 363 K for 2 h.

The effect of CH4 pressure is shown in Figure 5. The conversion of SO3 to MSA rises more rapidly than linearly and attains a value of 56% at a CH4 pressure of 1200 psi. The concentration of SO3 in the liquid was varied between 0 and 75 wt % by either diluting fuming sulfuric acid with water (SO3,l + H2Ol a H2SO4,l) or mixing pure SO3 with fuming sulfuric acid. Figure 6

Ind. Eng. Chem. Res., Vol. 40, No. 3, 2001 739

Figure 6. SO3 and CH4 conversion to MSA as a function of SO3 concentration for the reaction of 3 mL of H2SO4 (SO3 ≈ 0-75 wt %), 100 µmol of K2S2O8, and 1000 psi of CH4 at 363 K for 2 h.

1H

Figure 7. NMR of samples with varying SO3 concentrations for the reaction of 3 mL of H2SO4 (SO3 ≈ 30-75 wt %), 100 µmol of K2S2O8, and 1000 psi of CH4 at 363 K for 2 h. Table 1. Identification of Byproducts 1H

chemical shift (ppm)

chemical -CHx shift (ppm) identity assignment

13C

4.40 (4.12)a

65.8 (66.3)a

-CH2

3.57 (3.67)b 3.49

57.7 (59.8)b 53.5

-CH3 -CH3

name

CH2(SO3H)2 methanedisulfonic acid CH3OSO3H methyl bisulfate unidentified -

a For 50 wt % CH (SO H) standard in H O. b For 5 wt % 2 3 2 2 CH3OSO3H standard in H2SO4.

shows that the conversion of SO3 passes through a maximum of 67% at a SO3 concentration of 10 wt %, as the concentration of SO3 increases. Figure 7 shows that the decrease in the formation of MSA with increasing SO3 concentration is accompanied by an increase in the intensity of three 1H NMR peaks occurring at 4.40-4.57 ppm, 3.57-3.72 ppm, and 3.493.69 ppm. A detailed analysis of these features was carried out on the basis of 1H, 13C, and 2-D NMR (HMQC with BIRD Filter and GARP decoupling)6 techniques for the run in which the initial loading of SO3 was 74.8 wt %. These results, shown in Table 1, suggest that the peak observed in the range of 4.404.57 ppm is due to methanedisulfonic acid [CH2(SO3H)2] and the peak observed in the range of 3.57-3.72 ppm is due to methyl bisulfate [CH3OSO3H]. It is not clear at this time what species is associated with the peak observed between 3.49 and 3.69 ppm; however, it is known that this molecule contains a CH3 group. The

Figure 8. (A) SO3 and CH4 conversion to MSA as a function of the initial quantity of K2S2O8 for the reaction of 3 mL of H2SO4(SO3), 0-1000 µmol of K2S2O8, and 1000 psi of CH4 at 363 K for 2 h. (B) SO3 and CH4 conversion to MSA as a function of the initial quantity of K2S2O8 for the reaction of 3 mL of H2SO4(SO3), 0-1000 µmol of K2S2O8, and 1000 psi of CH4 at 333 K for 2 h.

possibility that any of the observed peaks is due to methanesulfonic acid anhydride [(CH3SO2)2O] or the mixed anhydride [CH3SO2OSO3H] can be excluded, as these products would be converted quantitatively to MSA during the addition of the reaction products to water at the completion of each experiment. It has been previously reported that the rate of MSA formation is independent of the initial concentration of K2S2O8.3 To confirm this, experiments were conducted with varying initial quantities of K2S2O8 at 333 and 363 K. Figure 8A and B shows that the conversion of SO3 does depend on the initial quantity of K2S2O8 present in the reaction mixture. At 363 K (Figure 8A), the conversion of SO3 rises to a maximum of 40% for 100 µmol K2S2O8. Above this point, the SO3 conversion decreases and becomes constant at 10% as the initial quantity of K2S2O8 approaches 1000 µmol. Similarly, the conversion of SO3 to MSA reaches a maximum when the same experiments are conducted at 333 K (Figure 8B); however, the maximum SO3 conversion occurs at 350 µmol of K2S2O8 and is more than two times larger than that observed at 363 K. Above 350 µmol of K2S2O8, the SO3 conversion to MSA decreases and plateaus just above 40%. The effect of temperature on the conversion of SO3 to MSA for a constant initial concentration of K2S2O8 is illustrated in Figure 9. With increasing temperature, the conversion of SO3 passes through a maximum at 333 K.

740

Ind. Eng. Chem. Res., Vol. 40, No. 3, 2001

Figure 9. SO3 and CH4 conversion to MSA as a function of temperature for the reaction of 3 mL of H2SO4(SO3), 100 µmol of K2S2O8, and 600 psi of CH4 at 313-363 K for 2 h.

Figure 10. Proposed reaction mechanism for K2S2O8-initiated MSA synthesis from CH4 and SO3 in fuming sulfuric acid.

Discussion It is apparent that, with the exception of the influence of CH4 pressure, the effects of all other reaction conditions are complex. To interpret the effects of reaction time, SO3 concentration, K2S2O8 concentration, and temperature, it is useful to first consider the consequences of using fuming sulfuric acid as a reaction solvent. Because the equilibrium constant for the reaction SO3,l + H2Ol a H2SO4,l is very high, the solvent can be treated as completely anhydrous, and any ionic species present will be solvated by H2SO4. Moreover, it is known that, as the concentration of SO3 in H2SO4 increases, SO3 will react with H2SO4 to form H2S2O7, H2S3O10, H2S4O13, and so on.7-9 As a consequence, the concentration of SO3 in solution is dictated by the amount of SO3 dissolved and the equilibria between SO3 and H2(SO3)nO (n ) 2, 3, 4, ...). Figure 10 illustrates a series of elementary steps that can be proposed to describe the reaction pathway through which K2S2O8 promotes the synthesis of MSA from CH4 and SO3 in oleum. This mechanism is an extension of that originally proposed by Sen and coworkers,3 which contained steps 1, 4, 5, and 6. Step 1 represents the dissociation of S2O82- anions to form two SO4-‚ radical anions. In H2O, SO4-‚ is known to react with H2O to form OH‚ and HSO4-.10 SO4-‚ is also known to attack organic reagents of the form RH to produce R‚ and HSO4-.11 It is, therefore, reasonable to hypothesize that, in the absence of H2O and high concentration of organics, SO4-‚ will react preferentially with H2SO4, in accordance with step 2. The reaction of SO4-‚ with

S2O82- is also possible (step 3) and is analogous to what has been reported for SO4-‚ reaction with C2O42-.10 Molecular SO4, produced as a product of this step, can decompose to form SO3 and O2.10 The formation of MSA is initiated by the reaction of SO4-‚ with CH4 dissolved in the oleum. CH3‚ radicals produced in step 4 can then react with SO3 to form CH3SO3‚ (step 5), which, in turn, forms MSA via abstraction of a hydrogen atom from CH4 (step 6). CH3SO3‚ can also react with SO3 via step 7 to form CH3S2O6, which then abstracts a hydrogen atom from CH4 to form CH3S2O6H (step 8), which is the methanesulfonic acid mixed anhydride. It is also possible that higher-order polyacid species [e.g., CH3(SO3)nH] might form as a consequence of CH3S2O6‚ reacting with additional SO3. Such products would be analogous to those formed in oleum at high SO3 concentrations [i.e., HO(SO3)nH], as noted above. In addition to steps 1-8, it is possible for MSA to form methanesulfonic acid anhydride (MSAA) via step 9. The H2O produced in step 9 would react immediately with SO3 to form H2SO4. The approach to a plateau in the plot of SO3 conversion to MSA versus time seen in Figure 4A and B suggests that the reaction has come to equilibrium or that the initiator has been fully consumed. Both of these interpretations can be easily ruled out. The standard Gibbs free energy for the reaction CH4,g + SO3,l f CH3SO3Hl is estimated to be -50.3 kcal/mol, and the enthalpy of reaction at 298 K is estimated to be -54.9 kcal/mol.12 For the reaction conditions of this study, the equilibrium conversion of SO3 to MSA is predicted to be 100% in all cases. Because the plateau in SO3 conversion shown in Figure 4B is well below this level, achievement of equilibrium can be ruled out as the cause for the plateau. The second possibility, complete consumption of K2S2O8, can also be excluded, because, as describe above, addition of K2S2O8 after reaction at 363 K for 2 h produced no additional MSA. A more plausible explanation is that significant amounts of CH3(SO3)nH and (CH3SO2)2O (MSAA) were formed during reaction. Although both of these products will revert to MSA upon addition of water to the reaction products, each 1 mol of CH3(SO3)nH represents the consumption of n mol of SO3, and each 1 mol of (CH3SO2)2O represents the consumption of 3 mol of SO3. The net effect of forming CH3(SO3)nH and (CH3SO2)2O is to increase the conversion of SO3 above that determined on the basis of MSA formed. Consequently, it is suggested that the plateaus observed in Figure 4A and B are the result of complete consumption of SO3, only a part of which ends up in MSA. The nonlinear increase in SO3 conversion to MSA as the CH4 pressure rises (Figure 5) is not likely to be due to an increase in the solubility of CH4 in H2SO4 with increasing pressure. Literature data reporting the solubility of CH4 in water as a function of pressure and temperature show that the solubility of CH4 in H2O rises linearly and then begins to plateau at pressures greater than or equal to 1100 psi.13 The maximum in the conversion of SO3 to MSA with increasing initial concentration of SO3 seen in Figure 6 can be attributed to two phenomena. The first is the formation of the byproducts methanedisulfonic acid and methyl bisulfate, and the second is the formation of CH3(SO3)nH. At low initial SO3 concentrations, the conversion of SO3 to MSA rises because of the increasing driving force to form MSA (step 5 in Figure 10). At some

Ind. Eng. Chem. Res., Vol. 40, No. 3, 2001 741

point, however, the formation of byproducts becomes increasingly important, and this diverts SO3 from conversion to MSA. Although the presence of CH3(SO3)nH could not be detected by 1H or 13C NMR spectroscopy, the formation of this product with increasing values of n as the initial SO3 concentration increaseswould rapidly deplete the reaction solution of SO3. Figure 8A and B shows that, at both 363 and 333 K, conversion of SO3 to MSA rises to a maximum value with increasing initial concentration of K2S2O8 and then decreases and approaches a plateau at yet higher initial concentrations. The maximum in SO3 conversion occurs at a lower initial concentration of K2S2O8 at 363 K than at 333 K, and the plateau in SO3 conversion to MSA is lower at the higher temperature. These trends can be explained, at least in part, by considering the reaction scheme proposed above. Doing so requires that we first assess the dependence of the rate of MSA formation on the initial concentration of K2S2O8. If steps 4-8 are taken to be irreversible, then the rate of MSA formation can be written as

rMSA ) k4[SO4-‚][CH4]

(1)

where [X] is the concentration of species X in oleum and ki is the rate coefficient for the ith reaction. Assuming that the concentration of SO4-‚ rapidly reaches a value representative of steady state, [SO4-‚] can be written as

[S2O82-] [SO4-‚] ) 2k1 k2[H2SO4] + k4[CH4]

(2)

and eq 1 can be rewritten as

[S2O82-][CH4] rMSA ) 2k1k4 k2[H2SO4] + k4[CH4]

(3)

The time dependence of the concentration of S2O82([S2O82-]) is given by

d[S2O82-] ) -k1[S2O82-] - k3[SO4-‚][S2O82-] (4) dt Using eq 2, eq 4 can be rewritten as

d[S2O82-] [S2O82-]2 2- k1[S2O8 ] - 2k1k3 dt k2[H2SO4] + k4[CH4] (5) If we define τ ) k1t and [S2O82-]* ) [S2O82-]/[S2O82-]o, where [S2O82-]o is the initial concentration of S2O82-, then eq 5 can be rewritten in dimensionless form as 2-

d[S2O8 ]* ) -[S2O82-]* - C([S2O82-]*)2 dτ

(6)

where

C)

2k3[S2O82-]o k2[H2SO4] + k4[CH4]

(7)

Because the concentrations of H2SO4 and CH4 are essentially constant, C depends only on the temperature and the initial concentration of K2S2O8. The solution to eq 6, assuming [S2O82-]* ) 1 for τ ) 0, is

(1 + C)[S2O82-]* (1 + C[S2O82-]*)

) e-τ

(8)

It is apparent from eq 8 that, when C ) 0, the solution [S2O82-]* approaches a pure exponential and, when C > 0, [S2O82-]* decreases more rapidly with τ. Solving eq 8 for [S2O82-] and inserting this result into eq 3 results in

rMSA )

d[MSA] k4 e-τ ) [CH4] dτ k3 (1 + 1/C - e-τ)

(9)

Integrating eq 9 assuming [CH4] is constant and [MSA] ) 0 for τ ) 0 yields

[MSA] )

k4 [CH4] ln[1 + C(1 - e-τ)] k3

(10)

Evaluation of eq 10 for a fixed reaction time (τ ) k1t) and variable C (∝ [S2O82-]o) yields a rapid monotonic increase in MSA concentration (∝ SO3 conversion) to a slowly rising plateau. This solution suggests that the reaction mechanism presented in Figure 10 is plausible and helps to account for the experimental observations in Figure 8A and B. However, this mechanism does not account for the maxima observed in Figure 8A and B at 100 and 350 µmol for 363 and 333 K, respectively, suggesting that the model in Figure 10 does not capture all of the elementary processes involved in the formation of MSA. Figure 9 shows that SO3 conversion to MSA increases as the temperature rises from 313 to 333 K and then decreases above 333 K to approximately 15% at 363 K. It was thought that this might be an effect of vaporliquid equilibrium, namely, that, as the temperature rises above 333 K, the concentrations of CH4 and SO3 in the liquid decrease and hence the rate of MSA formation decreases. However, calculations of liquid concentrations using vapor-liquid equilibrium data14,15 for SO3 and CH4 show that there is an insignificant change in the quantity of these reactants in the liquid over the temperature range investigated. Additionally, increased temperature does not result in the formation of additional undesired byproducts. A possible explanation is that, at lower temperatures, increasing the temperature increases the rate of MSA formation, but at higher temperatures, increasing the temperature increases the rate at which the initiator is destroyed, e.g., steps 2 and 3 in Figure 10. Conclusions A study has been conducted of the effects of reaction conditions on the synthesis of methanesulfonic acid from CH4 and SO3 using K2S2O8 as an initiator. The reaction time, CH4 pressure, SO3 concentration, K2S2O8 concentration, and temperature have been varied, and the highest SO3 conversion to MSA was observed for PCH4 ) 1000 psi, CSO3 ) 30 wt %, nK2S2O8 ) 350 µmol, and T ) 333 K for a reaction time of 2 h. The effect of varying these reaction parameters resulted in a set of interesting observations. It was found that SO3 conversion to MSA reaches a plateau after an initial period of time. Additionally, SO3 conversion to MSA exhibits a maximum with temperature at 333 K. The SO3 conversion to MSA was observed to have a positive-order nonlinear depen-

742

Ind. Eng. Chem. Res., Vol. 40, No. 3, 2001

dence on CH4 pressure. Similar to the effect of temperature, SO3 conversion to MSA also passes through a maximum with SO3 concentration. Finally, it was also observed that the SO3 conversion to MSA reaches a maximum when the initial concentration of K2S2O8 in the solution is varied and that this maximum depends on the reaction temperature. A sequence of elementary steps for the reaction of CH4 and SO3 to MSA in the presence of K2S2O8 has been proposed and has been successfully applied to capture some of these observed phenomena. Acknowledgment This work was supported by a grant from Elf Atochem North America Inc. Literature Cited (1) Kroschwitz, J. I.; Howe-Grant, M. Kirk Othmer Encyclopedia of Chemical Technology; Wiley: New York, 1991. (2) Guertin, R. U.S. Patent 3,626,004, 1971. (3) Basickes, N.; Hogan, T. E.; Sen, A. Radical-Initiated Functionalization of Methane and Ethane in Fuming Sulfuric Acid. J. Am. Chem. Soc. 1996, 118, 13111. (4) Pouchert, C. Aldrich Library of NMR Data; Aldrich Chemical Inc.: Milwaukee, WI, 1993. (5) Koeberg-Telder, A.; Cerfontain, H. Solutes in Sulfuric Acid. Part VI. A Nuclear Magnetic Resonance Study of Organic Sulfonic Acids and 1H Nuclear Magnetic Resonance Standards. pKBH Determination of Sulfonic Acids. J. Chem. Soc., Perkin Trans. 2 1975, 226.

(6) Bax, A.; Subramanian, S. Sensitivity-enhanced 2-D Heteronuclear Shift Correlation NMR Spectroscopy. J. Magn. Reson. 1986, 67 (3), 565. (7) Douglas, B.; McDaniel, D.; Alexander, J. Concepts and Models of Inorganic Chemistry, 3rd ed.; Wiley: New York, 1994. (8) Gillespie, R. J.; Robinson, E. A. Non-Aqueous Solvent Systems; Waddington, T. C., Ed.; Academic Press: London, 1965; p 162. (9) Gillespie, R. J. Cryoscopic Measurements in Sulphuric Acid. Part IV. Reactions of Ionised Sulphates in Sulphuric Oleum. Selfionisation Equilibria in Sulphuric Acid, and Ionic Equilibria in Oleum. The Polysulphuric Acids. J. Chem. Soc. 1950, 2516. (10) House, D. A. Kinetics and Mechanism of Oxidations by Peroxydisulfate. Chem. Rev. 1962, 62 (3), 185. (11) Lin, M.; Sen, A. Oxidation and Oxidative Carbonylation of Methane and Ethane by Hexaoxo-µ-peroxodisulfate(2-) Ion in Aqueous Medium. A Model for Alkane Oxidation through the Hydrogen-atom Abstraction Pathway. J. Chem. Soc., Chem. Commun. 1992, 892. (12) Guthrie, J. P.; Stein, A. R.; Huntington, A. P. Thermodynamics of Methanesulfonic Acid, Methanesulfonyl Chloride, and Methyl Methanesulfonate. Can. J. Chem. 1998, 76 (6), 929. (13) Berecz, E.; Balla-Achs, M. Studies in Inorganic Chemistry (4)sGas Hydrates; Elsevier: New York, 1983; p 75. (14) Sulfur Trioxide and Oleum: Storage and Handling. Dupont Corporation, Wilmington, DE. (15) Wankat, P. C. Separations in Chemical Engineering, Equilibrium Staged Separations, 1st ed.; Elsevier Science Publishing Co., Inc.: New York, 1988.

Received for review August 7, 2000 Revised manuscript received November 6, 2000 Accepted November 9, 2000 IE000725B