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Ind. Eng. Chem. Res. 2005, 44, 5950-5954
Oxidation of SO2 and Recovery of SO3 Using Nonaqueous Solvents Napapan Wattanakasemtham, Eric Croiset, Robert R. Hudgins,* and Peter L. Silveston Department of Chemical Engineering, University of Waterloo, Waterloo, Ontario, Canada N2L 3G1
Ali Lohi Department of Chemical Engineering, Ryerson University, Toronto, Ontario, Canada M5B 2K3
An earlier study from our laboratory using nonaqueous solvents to increase the concentration of acid recovered from adsorbing and reacting SO2 to form SO3 on activated carbon is extended to additional organic solvents and the use of subcritical and supercritical CO2. However, our conclusion from the previous work that organic solvents make suitable flushing agents is contradicted because of the degradation of some solvents or the discoloration of the acid product. Desorption of SO3 from activated carbon using supercritical CO2, however, is satisfactory. A statistically designed study showed that the temperature, pressure, and superficial velocity of CO2 fed to the carbon bed control the recovery. Introduction Sulfur dioxide, SO2, is a significant atmospheric pollutant. The main emitters of SO2 are fuel combustion processes and metal smelting processes. In fuel combustion, SO2 concentrations in the so-called lean flue gas range from 300 to 20 000 ppm. Smelting processes, on the other hand, produce rich flue gas containing more than 20 000 ppm SO2. Conventional processes for the removal of SO2 from flue gas are lime scrubbers.1 These have been proven efficient and relatively economical for capturing SO2. However, they require large volumes of water, and their effluents require further treatment before being returned to the environment. In this process, SO2 is converted to calcium sulfate (gypsum), to be discarded as solid to landfill or utilized in construction. In either case, a surfeit of gypsum is produced. Alternate technologies for capturing SO2 are regenerative processes that use a catalytic reactor to recover SO2 as sulfuric acid. Activated carbon has shown excellent potential as an oxidation catalyst.2 Even at room temperature, it provides H2SO4 productivities comparable to those with vanadium pentoxide or platinum catalysts at temperatures above 400 °C. Typically, with carbon, water is used as a flushing liquid, but the H2SO4 produced is too dilute to be marketable. Therefore, studies were initiated to find a flushing agent that might yield a concentrated H2SO4 that could be easily separated from the agent. Early experiments3 employed moderately polar organic solvents such as acetone. Such organic solvents exhibit much lower hydrogen bonding in the liquid phase than water, so that the energy required for their recovery would be much less than that needed to concentrate dilute sulfuric acid. These experiments demonstrated that low-molecular-weight ketones could be used to strip SO2 from activated carbon. In this study, we explore the recovery of sulfuric acid from these solvents and extend our study to further * To whom correspondence should be addressed. E-mail:
[email protected].
solvents such as alcohols and hydrocarbons. We also examine the use of CO2 in sub- and supercritical regions but in the vicinity of the critical point of CO2, which we will define as “near-critical CO2”. The use of CO2 as a solvent has been investigated since 1950.4 Recently, it has been used for carbon regeneration.5 As a flushing agent, an ideal organic solvent should behave differently from water by not reacting with either SO2 or SO3. The effluent from stripping SO3 from activated carbon is a mixture of the organic flushing fluid and SO3. If there is any water vapor in the solvent or gas streams, some H2SO4 will also form in the mixture. For an economical process, the ideal solvent must be recoverable, probably using distillation. In previous studies,3 acetone, methyl ethyl ketone (MEK), and methyl isobutyl ketone (MIBK) were used as the flushing agents. The solubility of both SO2 and oxygen in acetone at 20 °C is higher than that in water at the same temperature.6 The current work deals mainly with the evaluation of the acid-solvent separation and extends the work of Panthaky et al.3 to additional solvents. Panthaky6 used a structured packing loaded with activated carbon, while in this study, pelletized activated carbon was used. In the present study, near-critical CO2 is included as a solvent for several reasons: it has been used for regeneration of activated carbon,7 and it should be easily separated from the effluent mixture. Furthermore, it cannot be oxidized, and it is easily removed from the product acid by sparging. The important criteria for choosing a solvent for recovering SO3 from an activated carbon catalyst are summarized in Table 1. The solvents considered in the study were the lowmolecular-weight ketones (i.e., acetone, MEK, and MIBK) studied by Panthaky,6 alcohols (i.e., methanol and ethanol), and saturated hydrocarbons such as n-heptane and n-hexane. Experimental Details Experiments were undertaken using two pieces of equipment: (1) a packed-bed reactor filled with acti-
10.1021/ie0401861 CCC: $30.25 © 2005 American Chemical Society Published on Web 02/26/2005
Ind. Eng. Chem. Res., Vol. 44, No. 16, 2005 5951 Table 1. Properties of a Desirable Flushing Agent criterion
importance
absence of a chemical reaction low boiling point low viscosity at room temperature high solubility of SO3 noncorrosive low toxicity inexpensive
complete recovery of solvent and absence of degradation products in the acid energy savings and avoidance of acid-solvent reactions improved transport properties efficient desorption of SO3 from the solid catalyst reduced equipment costs safety and low environmental impact economical process
Table 2. BPL 6 × 16 Characteristics8 properties
value
Brunauer-Emmett-Teller surface area, m2/g average pore radius, nm pore volume, mL/g apparent density, g/mL average diameter, mm
1200 1.184 0.710 0.43 2.18
vated carbon and equipped with metered SO2 and air feeds for analysis of the off-gas and (2) a batch distillation unit. The activated carbon used in this study is BPL 6 × 16, supplied by Calgon Carbon Corp.; the characteristics of BPL carbon are shown in Table 2. The schematic flow diagram of SO2 oxidation and adsorption is shown in Figure 1. The packed bed consisted of a 500-mm, 45-mm-i.d. glass column. This reactor was packed with activated carbon to a depth of 310 mm. It was fed by flue gas mixtures with volumetric compositions of 0.35% SO2 (3500 ppm) and 5% O2, with the balance being N2, simulated to be close to the flue gas composition from coal-fired power plants. Mass flow controllers (Units Instruments Inc., Orange, CA) connected to the data acquisition system (Scienmetric Instruments, Kanata, Ontario, Canada) were used to regulate the mass flow rate of each gas. Compressed air supplied by Praxair, Inc. (Kitchener, Ontario, Canada), was used as the source of oxygen. The blended flue gas mixture was fed to the reactor in a downward flow; likewise, the organic solvent was fed from the top of the reactor by means of a stainless steel spray nozzle
Figure 1. Schematic of the reactor.
(supplied by Turbotak, Inc., Waterloo, Ontario, Canada). The acid-resistant thermocouples (Omega, Laval, Quebec, Canada) were used for temperature measurements. The exiting gas passed through a gas-liquid disengager below the reactor, followed by an ice trap to remove any moisture and acid mist carried over by the gas stream. From there, the gas passed to an SO2 analyzer (model 721 AT; Western Research, Calgary, Alberta, Canada). Once the bed was saturated with SO3, the solvent was fed from the top of the reactor. For liquid flushing agents, the effluent was collected for eventual separation. The amount of SO2 converted to H2SO4 was determined by a standard titration with a NaOH solution using phenolphthalein as an indicator. The concentration of SO2 was determined by iodometry. The sample was diluted with water to ensure that all of the dissolved SO2 was measured. Organic solvents of laboratory purity were used. The breakthrough behavior was determined for SO2 oxidation on activated carbon by measuring the SO2 concentration at the exit by the online SO2 gas analyzer. A plateau in the SO2 concentration after about 3 h established breakthrough. Consequently, the column was assumed to be saturated with SO3 after 3 h online. Instead of separating SO3 from the solvent, we assumed that water addition would convert this gas to sulfuric acid. Batch distillation was used to test the feasibility of separating the H2SO4-organic solvent mixtures. Distillations were carried out at both atmospheric and reduced pressures. A mixture of 5.0 wt % of 100% H2SO4 in the solvent was freshly prepared before each experiment. A 500-mL flask containing the acid-solvent mixture was heated by means of a heating mantel. The temperature of the mixture was measured by an acid-resistant thermocouple. Once the desired temperature (i.e., close to the boiling temperature of the mixture) was reached, the supplied power was reduced to and maintained at about 30% of full power in order to continue boiling. Vapor leaving the flask passed through a condenser, and the condensate was collected in a flask. A vacuum pump (model 470-5942; Barnant Co., Barrington, IL) was connected to the outlet of the collecting flask. The system pressure was measured
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Figure 2. Schematic of the equipment for desorption of SO3/H2SO4 using high-pressure CO2.
using a mercury manometer. The distillation pressure was at 28-30 cmHg under atmospheric pressure. For oxidation of SO2 over activated carbon and desorption in pressurized CO2, a 9.7-mm-i.d. stainless steel tube was used in place of glass; the tube was rated to a pressure of 41 MPa at room temperature. Otherwise, the rest of the equipment was as described for the organic solvents. This reactor was packed with about 8.45 g of catalyst to give a 23.5-cm bed depth of activated carbon. A low-voltage heating tape was wrapped around the reactor as a heating source during operation. The temperature was measured and controlled at the center of the bed. The output signal was sent to a temperature controller. The oxidation in both the glass and steel reactors occurred at room temperature and atmospheric pressure. The setup for the experiment of SO3/H2SO4 desorption using CO2 in the vicinity of its critical point is shown in Figure 2. The critical properties for CO2 are 304 K and 7.4 MPa. Initially, the temperature of the packed bed was brought to a desired value. Next, pressurized CO2 was fed to the reactor through a high-pressure regulator. The pressure in the bed was fine-tuned, and the flow rates of CO2 were controlled manually by means of the back-pressure regulator (model BP-66; GO Regulator, Mississauga, Ontario, Canada). Downstream of the back-pressure regulator, the gaseous CO2/SO3 mixture at slightly above atmospheric pressure passed through a fine-fritted disk and into the bottom of a long, glass, H2SO4 absorption column. The intent of the disk was to subdivide the flow, thus enhancing the gas-liquid interfacial area for mass transfer and increasing the retention time of the gas in the absorbent. A volume of exactly 250 mL of water (having a height of 38 cm) was placed in the column to act as an acid absorbent. The temperature of this unit was not controlled. Samples of water were then titrated to determine the total acidity produced, from which the SO3 recovered from the activated carbon was calculated. The duration of each run was different, but the total volume of CO2 fed into the system was kept constant at each pressure and temperature. For instance, at a superficial velocity of CO2 of 0.2 cm/min, P ) 8.0 MPa
gauge, and T ) 307 K, the total volume of CO2 that flowed in 488 min was 3.34 times the size of the reactor volume including tubing and connection from the gas entrance to the gas exit at the back-pressure regulator. To flow the same volume of CO2 at a velocity of 0.5 cm/ min and at the same temperature and pressure should have taken 195 min. At the end of each run, water was used to flush the carbon bed. The flushing water was collected for calculation of the remaining SO3 on the activated carbon surface. Then the bed was blown dry using hot nitrogen gas for about 2 h at 200 °C before the next cycle of SO2 oxidation adsorption was started. Trickle-bed experiments on SO2 oxidation with water as a solvent were made in order to prove that the solvent procedure was satisfactory. It was found that a 1.5-h flush with water removed all of the adsorbed SO3 from the activated carbon. In three trials, the recovery of SO3 obtained by using water at the end of a run with nearcritical CO2 was shown to be reproducible within a standard error of 1-2%. This number was used to calculate the total sulfur recovery.
Results and Discussion Two different sets of experiments were carried out: (1) solvent-sulfuric acid separations using batch distillation and (2) SO2/SO3 desorptions using near-critical CO2. In the previous work of Panthaky,6 SO3 was successfully desorbed from activated carbon using acetone, MEK, and MIBK in place of water. Additional solvents used in this study were methanol, n-heptane, n-hexane, and ethanol. Their physical properties are given in standard texts. The two sections of Table 3 represent results from the distillation experiments done at atmospheric and subatmospheric pressures. It can be seen from the atmospheric portion of this table that for MEK a reaction occurs between the ketone and acid. This is probably a condensation reaction. Preliminary experiments, not shown in the table, indicated that acetone and MIBK
Ind. Eng. Chem. Res., Vol. 44, No. 16, 2005 5953 Table 3. Distillation Results
a
solvent
distillation temperature (°C)
[H2SO4] (wt %)
MEK methanol n-heptane n-hexane
81.0 65.0 100.0 70.0
N/Aa 30.0 32.0 59.0
acetone methanol ethanol n-heptane n-hexane
50.0 62.0 55.0 73.0 50.0
solvent loss (%)
appearance of the remaining bottoms product
Atmospheric Pressure N/Aa 0.95 4.0 1.0
thick, dark liquid cloudy, pale yellow liquid two immiscible layers: pale yellow lower layer two immiscible layers: pale yellow lower layer
Reduced Pressure (28-31 cmHg) N/Aa N/Aa 80.0 0.6-1.5 63.0 0.5-2.0 60.0 8.0-12.0 98.0 1.0-3.0
thick, dark liquid cloudy, pale yellow liquid cloudy, pale yellow liquid two immiscible layers: pale yellow lower layer two immiscible layers: pale yellow lower layer
N/A ) not able to determine.
also underwent a reaction in the presence of sulfuric acid. Reducing the partial pressure to lower the distillation temperature was not effective. Coloration of the solution was less, indicating a slower rate of reaction. Results in the table show that acetone still undergoes degradation at the lower temperature. Results for alcohols indicated that the reaction takes place because phase separation is suggested by the turbidity reported in the table under both atmospheric and subatmospheric conditions. The reaction does not seem to occur when paraffins are employed. We believe the color seen in the experiments resulted from impurities present in the reagents. Initially, the mixture of sulfuric acid and n-paraffin formed a single phase. However, during distillation, phase separation occurred, indicating that the solubility limit of acid in the nparaffin had been exceeded. The bottom products obtained from distilling n-hexane and n-heptane solvents contained two immiscible layers. The top layer was a clear and transparent liquid, while the bottom layer contained discolored solvent (pale yellow liquid) and H2SO4. This layer had an acid concentration of about 98 wt %. The solvent loss during distillation was only about 1-3%. Nevertheless, the discoloration of the acid and the value of solvent consumption suggest that the paraffins are not suitable. Because distillation experiments showed relatively low solvent loss and high recovery of acid for the paraffins, additional experiments on the stripping were performed.9 We observed that the paraffins adsorbed strongly on activated carbon so that a stripping step would have to be added to the process to restore the capacity of the carbon for SO2 oxidation. Although experiments with other paraffinic solvents were not carried out, we expect that they too adsorb on activated carbon. Consequently, organic solvents do not seem to be suitable for stripping SO3 from activated carbon to produce concentrated sulfuric acid. As an alternative to organic solvents, we used an inorganic one: supercritical or near-critical CO2. After the reactor bed was saturated with SO3, carbon dioxide at the desired pressure and temperature was introduced into the reactor at a specified superficial velocity. The pressures and temperatures were selected to fall in sub- and supercritical (i.e., near-critical) regions. CO2 exhibits its greatest solvency close to its critical point.5 The two selected temperatures were 298 and 307 K (reduced temperatures, Tr ) 0.99 and 1.02, respectively), and the two pressure values were 8.0 and 9.2 MPa (reduced pressures, Pr ) 1.10 and 1.27, respectively). We observed that superficial velocities
Table 4. Summary of the Factors and Levels Used in the Experiment levels factors temperature (T), °C pressure (P), MPa superficial velocity (v), cm/min
high
low
34 9.2 0.5
25 8.0 0.2
Table 5. Summary of the Results for Sulfuric Acid Desorption Using Sub- and Supercritical Carbon Dioxide
run
pressure (MPa)
temperature (K)
velocity of supercritical CO2 (cm/min)
% sulfur recovery
1 2 3 4 5 6 7 8
8.0 8.0 8.0 8.0 9.2 9.2 9.2 9.2
307a 307a 298 298 307a 307a 298 298
0.2 0.5 0.2 0.5 0.2 0.5 0.2 0.5
90.65 82.53 97.32 94.41 89.99 81.65 87.92 82.04
a
Supercritical CO2 conditions.
within the range of 0.2-0.5 cm/min gave a fairly small bubble size in the absorption column. This was important to provide adequate contact between the gas stream and water. After each experiment, water was used to flush the activated carbon bed to determine the remaining amount of the acids SO2/SO3 on the activated carbon as well as to initialize the activated carbon to zero acid content. A flushing time of 1.5 h was used. Three experiments on SO2 adsorption with water flushing for different time lengths were made in order to demonstrate that the water-washing procedure was reproducible. SO2 recovery as sulfuric acid varied from 99.87% to 99.93%. From these results, it was established that a 1.5-h flushing with water removes all of the adsorbed SO3 from the activated carbon. A factorial experiment was designed to collect and interpret data for the near-critical CO2 study. Factors (i.e., independent variables) considered were temperature (T), pressure (P), and superficial velocity (v). The performance indicator (i.e., the dependent variable) in this study was the percentage recovery of sulfur. The factors and values used are summarized in Table 4. Table 5 shows the experimental design grid. From the factors T, P, and v, we generate standard treatment combinations for obtaining the percent of sulfur recovery as a function of T, P, v, TP, Tv, Pv, and TPv, as shown in eq 1, where R1-R8 are coefficients to be determined from the factorial experimental results, and T, P, and v
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% sulfur recovery ) R1 + R2T + R3P + R4v + R5TP + R6Tv + R7Pv + R8TPv (1) are in units of K, MPa, and cm/min, respectively. From these eight combinations, it was possible to estimate the main effects of T, P, and v; the two-factor interactions TP, Tv, and Pv; and the three-factor interaction TPv. The factorial experimental results are shown in Table 5. From them, it was found that the effects of pressure (P) and superficial velocity (v) are greater than that of temperature (T).9 An analysis of variance (not shown) was used to assess the significance of the two- and threefactor interactions, all of which were found to be insignificant compared to the main effects and were therefore ignored. The resulting model for the operating ranges provided in Table 4 is given in eq 2. From this
% sulfur recovery ) 270.6863 - 0.4100T (K) 6.759P (MPa) - 26.6246v (cm/min) (2) model, it is clear that the three manipulated variables of temperature, pressure, and superficial velocity decrease the desorption efficiency of SO3 from activated carbon in the packed bed, with the most dominant being the pressure and superficial velocity. We may explain these influences in terms of mass transfer and solubility. Assuming that diffusion in the pores of the catalyst controls the rate of dissolution, lower pressure provides a lower density, which provides a higher diffusivity and leads to an increase in the mass-transfer rates. Lower superficial velocity contributes to higher sulfur recovery, likely because, at low superficial velocity, the retention time (or contact time) between the activated carbon and the fluid is extended, allowing for increased mass transfer. This new system of desorption of SO3 from activated carbon by means of near-critical CO2 provides several advantages and a few disadvantages, as we now indicate. Gaslike low viscosity and little to no surface tension of near-critical CO2 imply improved penetration into the pores of activated carbon, resulting in efficient flushing. The excellent heat transport capacity of nearcritical CO2 also enables effective heat removal and control of the exothermicity related to sulfuric acid production. Furthermore, the use of CO2, a more environmentally benign solvent, can circumvent the difficult remediation required by most conventional organic solvents to remove contaminants. CO2 can be easily purged by sparging with a less soluble gas or by reducing the pressure. Thus, the unique physical and transport properties of near-critical and supercritical carbon dioxide offer a number of advantages that are interesting and beneficial for our problem, even though relatively high pressures and a more meticulous equipment design are required.
Conclusions This extension of our study on using nonaqueous solvents to increase the concentration of acid recovered from adsorbing and reacting SO2 to form SO3 on activated carbon contradicts our conclusion in the previous work3 that organic solvents are suitable flushing agents. We observed in this study that distillation of the flushing solution to recover acid and recycle solvent results in the degradation of some solvents or discoloration or the acid product. Either of these difficulties renders the use of organic solvents unacceptable. Although organic solvents are good flushing agents, as has been shown, the separation of the acid produced from the solvent is impracticable. On the other hand, we find that desorption of SO3 from activated carbon using near-critical CO2 is satisfactory despite problems with bypassing in the acid-CO2 separator. The factorial design indicated that the pressure and superficial velocity of CO2 passing through the packed bed of carbon have the greatest influence on sulfur recovery. The experimental data are summarized in a multifactorial equation. The advantages and drawbacks of this new system of desorption are pointed out. Literature Cited (1) De Nevers, N. Air Pollution Control Engineering; McGrawHill: New York, 1995. (2) Haure, P. M. Periodic Operation of a Trickle-Bed Reactor. Ph.D. Thesis, Department of Chemical Engineering, University of Waterloo, Waterloo, Ontario, Canada, 1989. (3) Panthaky, M. A.; Vladea, R. V.; Lohi, A.; Hudgins, R. R.; Silveston, P. L. Trickle-Bed Removal of Flue Gas SO2 using NonAqueous Solvents. Can. J. Chem. Eng. 2001, 79, 765. (4) Mukhopadhyay, M. Natural Extracts Using Supercritical Carbon Dioxide; CRC Press: Boca Raton, FL, 2000. (5) McHugh, M. A.; Krukonis, V. J. Supercritical Fluid Extraction; Butterworth-Heinemann: Boston, 1994. (6) Panthaky, M. D. Carbon-based process for SO2 conversion into concentrated H2SO4 at room temperature. M.A.Sc. Thesis, Department of Chemical Engineering, University of Waterloo, Waterloo, Ontario, Canada, 1998. (7) Karanfil, T.; Kilduff, J. E. The Role of GAC Surface Chemistry on the Adsorption of Organic Compounds: 1. Priority Pollutants. Environ. Sci. Technol. 1999, 33, 3217. (8) Calgon Carbon Corp. http://www.calgoncarbon.com, Apr 2003. (9) Wattanakasemtham, N. M.A.Sc. Thesis, Department of Chemical Engineering, University of Waterloo, Waterloo, Ontario, Canada, 2003.
Received for review June 23, 2004 Revised manuscript received December 20, 2004 Accepted December 29, 2004 IE0401861
