ARTICLE pubs.acs.org/IECR
A New Simple Microchannel Device To Test Process Intensification E. Santacesaria,* R. Tesser, M. Di Serio, V. Russo, and R. Turco University of Naples “Federico II” Department of Chemistry, Complesso Universitario Monte S. Angelo, Via Cintia 4, IT 80126 Naples, Italy ABSTRACT: In the present work, a new laboratory device specifically developed to obtain microchannels to test the process intensification effects on a suitable test reaction will be described. Ideally, this device represents a connection between a traditional tubular packed-bed reactor and the recently appeared microreactors that are very efficient, as it is well-known, in mass- and energytransfer operations. To test the performance of this microreactor, the decomposition of hydrogen peroxide (H2O2) has been chosen as a test reaction. This reaction is of great industrial interest, because many processes use H2O2 as an oxidizing agent and the decomposition of the excess is normally performed under batch conditions by creating an alkaline medium. We observed that the H2O2 decomposition can also be promoted by stainless steel acting as a catalyst. Therefore, the decomposition of H2O2 is particularly enhanced in our device that is characterized by a relatively high surface area of stainless steel per unit of packing volume. Consequently, its result is particularly efficient in the chosen reaction. Experimental runs for H2O2 decomposition have been performed preliminarily under batch conditions, both in the homogeneous phase and in the presence of a known amount of metallic surface area as a catalyst, in the temperature range of 65-85 °C, to collect kinetic data. Successively, continuous runs at different temperatures (50-65 °C) and pH have been performed in the microchannel device, obtaining good performances and maintaining safety conditions. The obtained results have been interpreted and successfully simulated in a simplified way.
1. INTRODUCTION Many reactive systems of significant industrial interest are frequently characterized by operational difficulties that sometimes invalidates their practical application. As example, in the case of heterogeneously catalyzed liquid-solid reactions or reactions involving the mass transfer between two immiscible liquids, gradients are generated that could strongly affects the overall system performance. In other cases, the mixing is rather efficient but the thermal behavior of the reaction (either highly exothermic or endothermic) could produce some hot spots that lead to thermal degradation of catalyst and/or of the reactants decreasing the selectivity or giving place to runaway phenomena. In many industrial applications, these problems are solved by means of an intensive mechanical stirring that could not be sufficient or too expensive. From another point of view, it is well-known that, apart from the particular physical property to be controlled (temperature, pressure, or concentration), the scale of the system strongly affects the gradients and the driving forces for mass, heat, and momentum transfer. As the size is decreased, the surface/volume ratio increases and, consequently, the efficiency of each transport phenomenon also increases. This concept has been intensively deepened in recent studies and in dedicated books, and it is well-known that the microsystems in which channels of ∼300 μm are realized, allows one to strongly improve mass, heat, and momentum transfer.1-4 For the explained reasons, nowadays, the use of microdevices is more diffused for strongly exothermic reactions (nitration of toluene5), in reactions involving highly toxic substances (syntheses with ethylene oxide6), and in two-phase reactions occurring at the liquid/ liquid interface (epoxidation of 1-pentene7). The devices used are relatively reduced in size and are of complex mechanical realization. The advantages and the drawbacks of similar systems are summarized in Table 1. Moreover, the development of microreactors or r 2010 American Chemical Society
systems based on microfluidic concepts seems to be very useful in process intensification (PI), which means to realize a process characterized by a very high reaction rate, a high selectivity, and a high level of safety in a reduced volume. Considering the described advantages of PI, we have decided to investigate the decomposition of hydrogen peroxide (H2O2), as a test reaction in an opportune microdevice; in fact, this reaction is of great industrial interest being H2O2 a widely used oxidizing agent and, in the effluents from these processes, significant concentration of H2O2 are often present. It is well-known that hydrogen peroxide easily decompose to give oxygen and water, in agreement with the following reaction: H2 O2 f H2 O þ 0:5O2
ð1Þ 8-13
This is a very exothermic reaction, and many authors have confirmed that the reaction can be enhanced by three effects: increasing temperature, increasing pH, and using a catalyst. The actual technology for decomposing H2O2 in wastewater consists of adding a basic compound, such as, for example, NaOH. In a basic environment, the decomposition of H2O2 is strongly promoted and the risk of runaway is serious under these conditions. Another aspect that must be considered is that stainless steel promotes the catalytic decomposition of H2O2. Therefore, using stainless steel microreactors, which have a high surface/volume ratio, the decomposition of H2O2 can be adequately promoted. The main purpose of this work is to study all of the three previously mentioned effects to determine a kinetic law for the decomposition of H2O2 under different pH and Special Issue: IMCCRE 2010 Received: March 15, 2010 Accepted: June 1, 2010 Revised: May 26, 2010 Published: June 24, 2010 2569
dx.doi.org/10.1021/ie1006307 | Ind. Eng. Chem. Res. 2011, 50, 2569–2575
Industrial & Engineering Chemistry Research
ARTICLE
Table 1. Advantages and Drawbacks of Using Microreactors advantages
drawbacks
lowering process cost
poor accommodation for
runaway reactions prevention
high realization costs and necessity
improved mass, heat,
necessity of specialists
solid-phase systems to employ special materials and momentum transfer scale-out/numbering up
analytical methods
temperature conditions and in the presence of stainless steel as a catalyst. Another scope is to realize a simple device, characterized by the presence of microchannels such as the microreactors but having a conventional structure, such as that of a packed-bed tubular reactor. This approach could be useful for testing the effect of microchannels on the reaction behavior. Figure 1. Schematic of the apparatus employed to perform the batch runs.
2. EXPERIMENTAL SECTION 2.1. Materials. All the reactants used were obtained at the maximum purity level available. The hydrogen peroxide solution was provided by Mythen S.p.A.; the stainless steel spheres and the springs used were provided by Tridella S.r.l. Company. 2.2. Analytical Method. The residual concentration of H2O2 was estimated by an iodometric analytical method,14,15 by reacting H2O2 and KI in a 2 M H2SO4 solution, according to the following reaction: H2 O2 þ 2I - þ 2Hþ f I2 þ 2H2 O
After the iodine is formed, it is possible to titrate it with 0.1 N Na2S2O3, using starch as an indicator in the following reaction: 2S2 O3 2 - þ I2 f S4 O6 2 - þ 2I ð2Þ The result can be expressed in terms of grams of hydrogen peroxide, using the following relation: NVMH2 O2 ð3Þ g H2 O2 ¼ 2 where N is the titrating solution normality, V the volume (in liters), and MH2O2 the molecular weight of hydrogen peroxide. 2.3. Batch Reactor Setup. The batch experimental apparatus used, in which the hydrogen peroxide decomposition kinetics has been studied, consists of a four-neck glass reactor charged with the desired quantity of H2O2 (20% by weight). The temperature of the system is controlled by a recirculation thermostat, while the top of the reactor is connected, through a condenser, to a gasvolumetric measurement system (see Figure 1). This last component consists of a graduate cylinder filled with water, turned upside down and immersed in a 5-L beaker. The oxygen evolved by the reaction is collected at the top of the cylinder and water is correspondingly displaced, allowing the measurement of gas developed as a function of reaction time. The obtained results have been corrected for the water vapor pressure. An example of kinetic run is reported in Figure 2. The described system resulted useful in the investigation of hydrogen peroxide decomposition reaction by changing the experimental conditions, such as temperature, pH, and catalyst concentration (in the case of catalytic decomposition). In particular, the pH of the reactive medium was adjusted by means of controlled additions of a 0.5 M NaOH solution and was monitored using a glass electrode inserted in the reactor. With regard to the catalyst, the reaction
Figure 2. Example of the trend of a H2O2 decomposition run.
Table 2. Characteristics of the Stainless Steel (AISI 316) Springs Used as a Catalyst in the H2O2 Decomposition dimension
value
length of the wire
23.4 mm
length of the spring diameter
2.4 mm 0.25 mm
mass
0.01 g
specific surface area
18.62 cm2/g
has been promoted by stainless steel, introduced in the form of small springs with a relatively high specific surface area (see Table 2 for details). 2.4. Continuous Reactor Setup. In a second step of the present work, the H2O2 decomposition has been studied in a continuous tubular reactor with the scope of comparing the results with those obtained under batch conditions. A simplified scheme of the continuous experimental apparatus is reported in Figure 3. The microplant consists of a HPLC pump able to ensure the suitable reactant feed rate. Through this pump, H2O2 solution was fed to the tubular reactor, whose total length is 20 cm and whose external diameter is 12.7 mm (1/2 in.). This reactor is electrically heated and 2570
dx.doi.org/10.1021/ie1006307 |Ind. Eng. Chem. Res. 2011, 50, 2569–2575
Industrial & Engineering Chemistry Research
ARTICLE
Figure 3. Schematic of the employed microdevice. Legend: (1) microtubular reactor, (2) refrigerator, and (3) product collecting and sampling.
Figure 4. Ideal packing geometry of the spheres in the reactor.
the temperature is regulated by a PID controller; the reactor outlet stream is quenched in a water bath with the purpose to instantaneously stop the decomposition reaction. The residual hydrogen peroxide is then evaluated by titration. In the interior of the reactor, microchannels are realized by means of a particular packing structure made of two different-sized stainless steel spheres.16 The radii of the two types of spheres have been chosen by adopting a ratio of 0.156 between small and big spheres (r = 0.195 mm, R = 1.25 mm), which corresponds to microchannels with an average size of 300 μm. The ideal packing geometry is reported schematically in Figure 4. As it can be appreciated in the bidimensional representation of Figure 4, the interstitial cavities through which the fluid flow is ensured are uniform in size (on the order of magnitude of 300 μm).
In this reaction, molecular oxygen is formed through the interaction between undissociated hydrogen peroxide and perhydroxyl ions and also leads to the formation of hydroxyl groups. This mechanism7 suggests a reaction rate that depends on both the concentration of H2O2 and HO2-; moreover, it has been demonstrated that the decomposition is favored also by the temperature,8 and, in our investigation, we have explored the temperature range of 55-75 °C; the experimental conditions related to the batch runs are summarized in Table 3. It is wellknown that pH has a strong influence on this reaction, which is strongly promoted by the alkaline environment,7 as a consequence of the following reaction: H2 O2 þ OH- f HO2 - þ H2 O
ð5Þ
3. RESULTS AND DISCUSSION 3.1. Homogeneous Decomposition. The decomposition of
hydrogen peroxide under basic conditions occurs spontaneously, according to the following scheme: H2 O2 þ HO2 - f H2 O þ O2 þ OH-
ð4Þ
A preliminary set of runs have been performed with the scope to evaluate the effect of pH on the reaction. For this purpose, in Figure 5, the results of kinetic runs, performed at different pH, are reported for comparison. As can be seen, a small increase in pH corresponds to a large increase in the decomposition rate and below a pH threshold limit value of ∼7, the reaction does not 2571
dx.doi.org/10.1021/ie1006307 |Ind. Eng. Chem. Res. 2011, 50, 2569–2575
Industrial & Engineering Chemistry Research
ARTICLE
Table 3. List of the Experimental Runs Performed in the Homogeneous Phase without a Catalyst run
temperature, T (°C)
1
65
6
2 3
65 65
7 7.5
4
55
8
5
65
8
6
75
8
pH
Figure 6. Simulation of the 20 wt % H2O2 decomposition by assuming an apparent order of R = 2.5. These runs have been performed at pH 8.
Table 4. Kinetic Constants and Corresponding Root-MeanSquare Error (rms) Obtained by Assuming the Apparent Order of R = 2.5 run
k (min-1)
temperature, T (°C)
root-mean-square error, rms
4
0.000112
55
0.021
5
0.000172
65
0.016
6
0.000245
75
0.069
Figure 5. Graph showing the pH dependence of hydrogen peroxide decomposition rate under homogeneous conditions.
occur. The observed kinetics are in agreement with the observations reported in the literature,8 according to which a sequence of plateau, characterized by slow decomposition, can be observed, because of the progressive decomposition of the stabilizers usually present in the commercial samples of H2O2, such as tannic acid, uric acid, and benzammide. This effect was accurately studied by Rice et al.,9 comparing different decomposition runs performed with pure hydrogen peroxide solution and hydrogen peroxide containing uric acid. The presence of a plateau is justified by the presence of the mentioned stabilizer. A further investigated aspect involves the effect of temperature on the H2O2 decomposition. Experimental runs have been performed at different temperatures;55, 65 and 75 °C, at pH 8, in order to make the effect of the stabilizer negligible. The collected experimental data have been interpreted by means of a power-law expression of the type -
d½H2 O2 ¼ k½H2 O2 R dt
ð6Þ
In the above expression, the reaction pseudo-order (R) is reported by the literature to be in the range of 1-2.5.8,10,11 The best fitting of our experimental data has been obtained using a value of R = 2.5. The obtained agreement can be appreciated in Figure 6. The corresponding kinetic constants are reported in Table 4. The Arrhenius plot (Figure 7) allowed the evaluation of both the activation energy and the pre-exponential factor (see Table 5). 3.2. Kinetic Runs Performed in Batch Conditions in the Presence of Stainless Steel as the Catalyst. It is well-known that AISI 316 stainless steel is a very common material used for manufacturing engineering plants, reactors, heat exchangers, and microreactors. We observed that this material promotes the
Figure 7. Arrhenius plot for the homogeneous H2O2 decomposition at pH 8.
Table 5. Activation Energy and Pre-Exponential Factor for the Homogeneous Reaction Performed at pH 8 parameter Ea ln A
value 37.00 ( 1.4 (kJ 3 mol-1)
4.476 ( 0.511 (min-1)
decomposition of hydrogen peroxide. For this reason, we have studied the catalytic effect of the stainless steel on the decomposition of hydrogen peroxide. First of all, we have performed this study, under batch conditions, using the previously described 2572
dx.doi.org/10.1021/ie1006307 |Ind. Eng. Chem. Res. 2011, 50, 2569–2575
Industrial & Engineering Chemistry Research
ARTICLE
Table 6. List of the Runs Performed in the Presence of Stainless Steel as the Catalyst run
temperature, T (°C)
pH
Wcat(g)
1 2 3
75
7
2.0271
75 65
7 3.5
4.0222 4.0073
4
75
3.5
4.0777
5
85
3.5
4.0499
Figure 9. Decomposition runs, under batch conditions catalyzed with, respectively, (2) 2 g (run 1) and (9) 4 g (run 2) of stainless steel springs at pH 7 and T = 75 °C, and (b) a run at pH 3.5 and T = 75 °C with 4 g of steel springs (run 4).
Figure 8. Comparison between (1) a catalyzed run and ([) an uncatalyzed run at pH 7 and T = 75 °C.
apparatus. Although in the literature, no specific studies have been published about the catalytic effect of stainless steel on the hydrogen peroxide decomposition reaction, some authors12 have discussed the generic catalytic effect of metallic surfaces on this reaction. In that study,12 the main role of H2O2 consisted of the oxidation of the metallic surface, which led to the production of hydroxyl radicals through a mechanism very similar to the traditional Haber-Weiss reaction:12 H2 O2 þ Fe2þ f Fe3þ þ OH- þ OH•
ð7Þ
To investigate the catalytic activity of stainless steel, we have performed the experimental runs summarized in Table 6, by changing the temperature, the pH, and the mass of the catalyst (Wcat), with the latter being added in the form of small springs. The results shown in Figure 8 indicate that the reaction, in the presence of a metallic surface, occurs faster than the corresponding reaction in the homogeneous phase, under the same conditions of temperature (T = 75 °C) and pH (pH 7). As it can be seen, in the presence of the mentioned catalyst, the reaction also significantly occurs at pH 7. Moreover, the decomposition in the presence of different amounts of steel, when used as a catalyst, shows an almost-linear trend (see Figure 9). The absence of a plateau suggests that, in the presence of a stainless steel catalyst, at pH 7, the effect of the stabilizer is very small. A similar behavior has also been obtained at pH 3.5 and 75 °C. Therefore, these experimental data have been interpreted by adopting, as a first approximation, the following kinetic law: ! d½H2 O2 S0 Wcat ¼ k ð8Þ ½H2 O2 dt Vliq
Figure 10. Decomposition runs performed with 4 g of steel at pH 3.5 at different temperatures, under batch conditions: (2) run 3, (9) run 4, and (1) run 5.
where k is the actual kinetic constant, S0 the metallic specific surface area (18.62 cm2 g-1), Wcat the catalyst mass, and Vliq the reaction liquid volume (expressed in terms of cm3). Different runs have been performed with the intention to investigate also the effect of the temperature on the decomposition reaction in the presence of a stainless steel catalyst (see Figure 10). As expected, the decomposition rate is faster when the pH of the reaction is increased, as shown in Figure 9. Figure 10 shows that the decomposition rate strongly increases with the temperature. The assumption of a pseudo-first-order reaction for interpreting the runs reported both in Figures 9 (runs 1, 2, 4) and 10 (runs 3, 4, 5) is clearly an oversimplified kinetic approach but is acceptable for the scope of this work. The corresponding k-parameters are reported in Table 7. An approximated value for the activation energy can be estimated corresponding to 103 kJ/mol. 3.3. Runs Performed in a Tubular Packed-Bed Reactor Containing AISI 316 Small Balls. Two further experimental runs were then performed in the previously described tubular packed-bed reactor.16 From the void volume of the microreactor 2573
dx.doi.org/10.1021/ie1006307 |Ind. Eng. Chem. Res. 2011, 50, 2569–2575
Industrial & Engineering Chemistry Research
ARTICLE
Table 7. Kinetic Constant Values Obtained by Regression on the Experimental Results of the Catalyzed Reaction
Table 8. Comparison between the Specific Surfaces for Different Types of Catalyst
Wcat(g)
Vliq(cm3)
k (cm min-1)
1
2.0271
52.6
2.28 10-3
springs
18.62
2 3
4.0222 4.0073
52.6 52.6
2.09 10-3 3.43 10-4
small spheres big spheres
23.89 3.07
4
4.0777
52.6
4.30 10-4
packing
5
4.0499
52.6
2.72 10-3
run
Si (cm2 g-1)
where
continuous runs (XA = 75%; T = 65 °C)
4.63 10-4
(XA = 40%; T = 50 °C)
1.70 10-4
3.98
! S0 Wcat ½H2 O2 ra ¼ k Vliq
Here, k is the actual kinetic constant, S0 the metallic specific surface (estimated as a weighted mean between the two differentsized spheres), Wcat the catalyst mass, A the reactor’s cross section (considering the experimental void degree of 0.25), and Vliq the liquid volume in the tubular reactor (expressed in units of cm3). The kinetic constant estimated from the experimental run performed at 65 °C is k = 4.63 10-4cm min-1, which is a value comparable to the kinetic constant obtained from the batch run performed under the same conditions of pH and temperature (3.43 10-4cm min-1; see Table 7). In contrast, a value of k = 1.70 10-4cm min-1 has been obtained for the run performed, at 50 °C, under the same pH conditions. The kinetic constants obtained from the packed-bed tubular reactor are clearly less reliable, considering the formation of oxygen during the reaction and the difficulty of controlling the temperature along the reactor. Nevertheless, the objectives of intensifying the process and developing a simple device to simulate the behavior of a microchannel system has been fully reached. Figure 11. Conversion of H2O2 as a function of time for the runs performed in the microreactor at (9) 50 °C and (2) 65 °C.
(V = 4 cm3) previously measured, and with a volumetric feed flow rate Q = 0.2 cm3/min, it was possible to estimate a residence time of ∼20 min. Furthermore, we have adopted mild conditions for what concerns the reactant concentration and reactor temperature (T = 50 and 65 °C, 20 wt % H2O2) for safety reasons, because the decomposition reactions proceed very fast in the presence of steel. The results of these runs, performed at pH 3.5, are shown in Figure 11, which shows the outlet conversion of hydrogen peroxide as a function of time on stream (in this way, we have verified the attainment of steady-state conditions). It is interesting to observe that, after ∼2 h of operation, the system can be considered to be under steady-state conditions at both 50 and 65 °C, obtaining an average conversion of 40% and 75%, respectively. These performances are satisfactory by considering the relatively low contact time imposed. We have attributed this interesting result to the relatively high value of the specific surface area, whose value was estimated to be 3.98 cm2 g-1. This value has been calculated considering that the packing is composed of ∼2000 big spheres and ∼30 000 small ones (comparisons shown in Table 8). To interpret the experimental data, it has been necessary to write the mass balance on the hydrogen peroxide that flows in the reactor. This mass balance is written as follows: d½H2 O2 A ¼ - ra ð9Þ dt Q
4. CONCLUSION In this work, a process intensification has been carried out regarding the decomposition reaction of hydrogen peroxide. A preliminary study, performed under batch conditions, allowed the estimation of the kinetic parameters that proved to be useful in the design of a continuous tubular reactor with microchannels that can simulate the behavior of a microreactor. The possibility to obtain a very interesting conversion in a continuous operation has been verified, and this represent an optimal starting point for future developments of the H2O2 decomposition process. Moreover, the proposed intensified unit operation can also be used for other reactions of industrial interest. ’ AUTHOR INFORMATION Corresponding Author
*Tel.: 0039 081 674027. Fax: 0039 081 674026. E-mail:
[email protected].
’ ACKNOWLEDGMENT This study was funded by EC VII Framework Programme CPIP 228853-2 COPIRIDE. Thanks are also due to MYTHEN SpA for the initial financial support. ’ LIST OF SYMBOLS [H2O2] = hydrogen peroxide concentration (mol cm-3) R = reaction order of the decomposition reaction k = real kinetic constant (cm min-1) 2574
dx.doi.org/10.1021/ie1006307 |Ind. Eng. Chem. Res. 2011, 50, 2569–2575
Industrial & Engineering Chemistry Research
ARTICLE
S0 = metallic specific surface area (cm2 g-1) Wcat = catalyst mass (g) Vliq = liquid volume (cm3) A = empty cross section of the tubular reactor (cm2) Q = volumetric flow rate (cm3 min-1) ra = reaction rate (mol cm-3 min-1) XA = conversion degree of hydrogen peroxide
’ REFERENCES (1) Lowe, H.; Ehrfeld, W. State-of-the-art in microreaction technology: Concepts, manufacturing and applications. Electrochim. Acta 1999, 44, 3679–3689. (2) Hessel, V.; Hardt, S.; L€owe, H. Chemical Micro Process Engineering. Fundamentals, Modelling and Reactions; Wiley-VCH: Weinheim, Germany, 2004. (3) Hessel, V.; Renken, A.; Shouten, J. T.; Yoshida, J. I. Micro-Process Engineering;A Comprehensive Handbook; Wiley-VCH: Weinheim, Germany, 2009. (4) Catalogue Chemical Process Technology; Institute for Mikrotechnik (IMM), 5/2009. (5) Halder, R.; Lawal, A.; Damavarapu, R. Nitration of toluene in a microreactor. Catal. Today 2007, 125, 74–80. (6) Ehrfeld, W.; Gartner, C.; Golbig, K.; Hessel, V.; Konrad, R.; Lowe, H.; Richter, T.; Schulz, C. Proceedings of the 1st International Conference of Microreaction Technology, Frankfurt/Mainz, Germany, 1997; p 72. (7) Wan, Y. S. S.; Chau, J. L. H.; Yeung, K. L.; Gavriilidis, A. 1-Pentene epoxidation in catalytic microfabricated reactors. J. Catal. 2004, 223, 241–249. (8) Spalek, O.; Balej, J.; Paseka, I. Kinetics of the decomposition of hydrogen peroxide in alkaline solutions. J. Chem. Soc., Faraday Trans. 1 1982, 78 (8), 2349–2359. (9) Rice, F. O.; Reiff, O. M. The thermal decomposition of hydrogen peroxide. J. Phys. Chem. 1927, 31, 1352–1356. (10) McLane, C. K. Hydrogen peroxide in the thermal hydrogen oxygen reaction: 1. Thermal decomposition of hydrogen peroxide. J. Chem. Phys. 1949, 17, 379–385. (11) Livingston, R. The rate of decomposition of hydrogen peroxide in concentrated solutions. J. Phys. Chem. 1943, 47, 260–264. (12) Lin, S. S.; Gurol, M. D. Catalytic Decomposition of Hydrogen Peroxide on Iron Oxide: Kinetics, Mechanism, and Implications. Environ. Sci. Technol. 1998, 32, 1417–1423. (13) Lennon, E.; Burke, A. A.; Ocampo, M.; Besser, R. S. Microscale packed bed reactor for controlled hydrogen peroxide decomposition as a fuel cell oxidant aboard unmanned undersea vehicles. J. Power Sources 2010, 195, 299–306. (14) Kolthoff, I. M.; Mecchan, S. Treatise on Analytical Chemistry, 2nd Edition; Wiley: New York, 1978; p 888. (15) Determination of the peroxide number, NGD Method C35, 1976. (16) Santacesaria, E., Di Serio, M.; Tesser, R. (EUROCHEM srl). Ital. Patent Na 2009A000024, 2009.
2575
dx.doi.org/10.1021/ie1006307 |Ind. Eng. Chem. Res. 2011, 50, 2569–2575
