In the Laboratory
Catalytic Oxidation of Sulfur Dioxide by Activated Carbon A Physical Chemistry Experiment E. Raymundo-Piñero and D. Cazorla-Amorós* Departamento de Química Inorgánica, Universidad de Alicante, Alicante, Spain; *
[email protected] E. Morallón Departamento de Química-Física, Universidad de Alicante, Alicante, Spain
The chemistry and chemical engineering curriculum includes the concept of catalysis, which is relevant from both fundamental and applied points of view. It is important to explain homogeneous and heterogeneous catalysis and their applications to students, employing examples that are as simple as possible. In this contribution, we present an experiment that is indeed simple to perform and presents a good example of heterogeneous catalysis. Moreover, with the increased interest in environmental chemistry, more student experiments related to this area are desirable. The experiment described here is of particular interest as relates to well-known environmental problems, the formation of acid rain and the removal of SO2. In recent years, one of the applications of heterogeneous catalysis receiving considerable attention is the abatement of pollutants (1). Among these pollutants, SO2 and NOx emissions produced by anthropogenic sources, mainly coal-based power generation and internal combustion engines, have significant environmental impact (2–4 ). For this reason, researchers are devoting considerable resources to developing processes to remove these gases. The atmospheric oxidation of SO2, one of the agents responsible for acid rain, occurs not only in the gas phase but also in the aqueous phase (in clouds, fog, etc.). Owing to its high solubility in the aqueous phase, the oxidation of SO2 in this medium is often more important than gas phase oxidation (5). Conventional desulfurization technology uses calcium carbonate to produce calcium sulfite and calcium sulfate after reaction with SO 2. This compound is a by-product with a low commercial value. Newer technologies apply recyclable SO2 adsorbents such as activated carbon (6 ). One possibility being actively explored is the use of carbonaceous adsorbents as catalysts for low-temperature SO2 oxidation to sulfuric acid in the presence of oxygen and water (6–12). In this process, sulfuric acid solutions are obtained as the end product. The laboratory experiment described here is concerned with the use of activated carbon as a catalyst for the oxidation of aqueous solutions of SO2 to sulfuric acid. The experiment, which is quite simple to perform, is a good example of a heterogeneous catalytic process. It provides a better understanding of acid rain formation and an example of SO2 removal.
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Experimental Procedure CAUTION: The experimental procedure involves the following steps that must be carried out in a fume cupboard. SO2 is handled during the experiment. Note that is both an oxidant and a good reducing agent. SO2 is also toxic and an irritant by inhalation (13). 1. The preparation of a saturated aqueous solution of SO2 from the reaction between Na2 SO3 and H2SO 4 (14). 2. The preparation of diluted solutions of SO 2 of concentration 2 × 10{3 M or lower. 3. The catalytic oxidation of the diluted solutions of SO2 to H2SO 4 and the use of a conductometer to follow the reaction.
The first step of the experiment can either be done by the students during the experimental session or the solution can be prepared by the teacher before the lab session. The latter reduces the duration of the experiment and avoids problems related to the preparation and manipulation of gaseous SO2 and its solutions by the students. The teacher should choose between these two possibilities depending on the level of the course and the time dedicated to the experiment.
Preparation of Aqueous Solutions of SO2 A saturated solution of this gas can be prepared using a conventional gas generator (Fig. 1) (15). The 100-mL dropping funnel (A) should contain concentrated H2SO4 (about 10 mL is enough). The 100-mL round-bottom flask (B) should have Na2SO3 (about 6 g to prepare 30 mL of a solution of SO2 with a concentration close to or higher than 1 M). Thirty milliliters of distilled water should be placed in bottle D (bottle C is used for safety reasons), and beaker E should contain a solution of NaOH to retain the excess of SO2 as sulfite anions. Because gaseous SO2 is readily soluble in water (3927 cm3 in 100 g of H2O at 20 °C [16 ]), a solution of concentration near to 1 M can be easily obtained in bottle D, through bubbling of the SO2 through the bottle D for several minutes. Following this, the accurate concentration of the solution should be determined by titration with iodine using starch as indicator (10). Diluted solutions with concentrations of
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2 × 10{3 M or lower can also be prepared from this saturated solution and titrated with iodine to determine their concentration (concentrations of 10{5 M can be easily detected).
Catalytic Oxidation of Aqueous Solutions of SO2 In this test, the reaction is followed by a high precision conductometer (sensitivity of 10-3 mS cm{1), which allows the performance of the reaction with concentrations of SO2 lower than 2 × 10{3 M to be followed with a high degree of reproducibility. The measurement of conductance instead of pH is superior because the complete conversion of SO2 produces a drop of pH of only about 0.3 units. The experiment is already in use in the investigations of carbon as a catalyst in oxidation reactions (7, 17 ) and is simple to perform as can be seen from the following. The experimental setup is shown in Figure 2. A volume of 250 mL of diluted SO2 solution of known concentration (ca. 2 × 10{3 M) is introduced into a three-necked round-bottom flask. In one neck of the flask, the conductivity cell is fitted, and in the second, a thermometer is inserted to record the temperature of the reaction. In the third neck, a flow of air is introduced to obtain water saturated with oxygen. The flow of air is maintained during the catalytic oxidation. By using a hot-plate with magnetic stirring, it is possible to control the temperature of the reaction from room temperature to 50 °C. If available, a thermostatic bath would be preferable for this purpose. It should be noted that even at temperatures higher than 50 °C, there is no evolution of SO2 gas because the concentration used is much lower than the saturation concentration (the solubility of SO2 at 90 °C is about 0.09 M [18]). In addition, the concentration of oxygen at 50 °C in saturated water is about 10{3 M (18). Once the temperature has stabilized, the conductivity of the solution is measured. At 25 °C, the specific conductivity of a solution of SO2 of concentration 1.65 × 10{3 M is 0.57 mS cm{1 and the concentration of oxygen in O2-saturated water is higher than 10{3 M (18). After measuring the conductivity of the SO2 solution, ca. 0.3 g of activated carbon is placed into the reactor vessel. In principle, any activated carbon may be used for this purpose, although it is more convenient to use granular activated carbon with a particle size of 2–3 mm. Because the activated carbon may be hydrophobic, it is necessary to verify its wettability by placing a sample in water. If the activated carbon is found to be hydrophobic, then it will be necessary to wet it for a designated period in a known volume of water (3–4 mL) before adding it into the reactor. In this case, a blank measurement of the conductivity of the SO2 solution should be done after diluting with the volume used for wetting the carbon. After loading the activated carbon, the conductivity is measured as a function of time and curves like those shown in Figure 3 are obtained. The figure also includes conductivity measured in absence of catalyst. Initially, a small decrease or constancy in the conductivity may occur due to the adsorption of SO2 on the activated carbon. Once the reaction reaches the steady state, conductivity increases linearly with time although, after a long time, a downward deviation will occur and the highest conductivity will be reached when most of the SO2 reacts. This is what occurs with sample ACF40: for an initial concentration of SO2 of 1.65 × 10 {3 M (conductivity = 0.57 mS cm{1), the total oxidation of SO2 results in a conductivity value of the solution of 1.05 mS cm{1. It must be noted
Figure 1. Experimental system used for the generation of SO2. A: dropping funnel containing H2SO4; B: flask containing Na2SO3; C: safety bottle; D: collection bottle; E: trap for excess SO2.
Figure 2. Experimental system used for the catalytic oxidation of SO2.
that without the activated carbon catalyst, SO2 oxidation does not take place during the same period of time (Fig. 3). At the end of the reaction, students can detect the formation of SO42{ anions and the consumption of SO2 using conventional analytical reactions. SO2 can be easily detected using iodine and starch as indicator and SO42{ by adding a Ba2+ solution. BaSO4 is insoluble and is not dissolved by HCl, whereas BaSO3 is dissolved by HCl. Before proceeding to discuss the information obtained from these experiments, we will present an overview of the kinetics of this reaction obtained from the literature (11, 12). Kinetics of the SO2 Oxidation Reaction A very detailed study of the reaction mechanism can be found in refs 11 and 12. The process is thought to follow four steps: (i) sulfur dioxide adsorption to a carbon active site, (ii) oxidation of the adsorbed sulfur dioxide with oxygen to form sulfur trioxide, (iii) hydration of the adsorbed sulfur trioxide to form bound sulfuric acid, and (iv) desorption of the acid. This mechanism is described as: σv + SO 2(aq) σSO2 +
1/ O (aq) 2 2
σSO3 + H2O
σSO2 σSO3 σH2SO4
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(1) (2) (3)
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In the Laboratory σH2SO4
H 2SO4(aq) + σv
(4)
where σv, σSO , σSO , and σ H SO are the free carbon sites, adsorbed sulfur dioxide, adsorbed sulfur trioxide, and adsorbed sulfuric acid, respectively. Reactions 1–3 are assumed to be reversible and at equilibrium. If reaction 4 controls the rate of the overall process, the apparent rate of sulfur dioxide conversion is: 2
3
2
4
d[H2SO4(aq)]/dt = { d[SO2(aq)]/dt = k4 σH SO 2
4
(5)
This equation assumes no reversible adsorption of sulfuric acid on the sites. Using the reversible assumption for eqs 1–3, the concentrations of the occupied surface sites are: σSO2 = K 1 σv [SO 2(aq)]
(6)
σSO3 = K 2 σSO2 [O 2(aq)]1/2
(7)
σH2SO4 = K 3 σSO3 [H2O]
(8)
where Ki represents the ratio ki/k{i of the forward and reverse rate constants of eqs 1–3. The free carbon sites can be related with the total carbon sites (σtotal), which is a constant, using eqs 6–8 and eq 9: σtotal = σv + σSO2 + σSO3 + σH 2SO4
(9)
σtotal = σv(1 + K1[SO2(aq)] + K1K 2[SO2(aq)][O2(aq)]1/2 +
K 1K2K3 [SO2(aq)][O2(aq)]1/2[H2O])
(10)
Using these equations, the reaction rate (eq 5) can be written as: 1/2
d[SO2(aq)]/dt = k′[SO2(aq)][O2(aq)] [H2O]/
whereas the second one has a magnitude similar to that of the first for H2SO3. Thus, the H2SO4 formed can be considered to be completely dissociated, giving two H+ and one SO42{ anion. In this sense, we can consider that the solution behaves as a strong electrolyte and that the theory of conductivity for these systems can be applied (19). Hence, under the experimental conditions employed in these experiments, the overall reaction can be written as: HSO3{ + 1/2 O2 → SO42{ + H+ C0 – x
1/2
(1 + K1 [SO2(aq)] + K1K2 [SO 2(aq)][O2(aq)] + (11) K 1K2K3 [SO 2(aq)][O2(aq)]1/2[H2O]) where k′ = k4K1K2K3σtotal . Equation 11 states that the rate of sulfur dioxide conversion is directly proportional to the number of active carbon sites, the water concentration, the oxygen concentration, and the concentration of sulfur dioxide (11, 12). It can also be deduced from eq 11 that the reaction order for SO2 may vary between 0 and 1. An alternative approach to the use of the complex eq 11 is the use of a power-law model to fit the rate constant and the exponents of each concentration term. Equation 12 demonstrates this model (11, 12): { d[SO 2(aq)]/dt = kW([O2(aq)])a([SO 2(aq)])b ([H2O])c (12) where k is an apparent rate constant and W is the mass of the catalyst in the reactor. Equation 12 thus assumes a firstorder dependence of the rate on the mass of catalyst in the reactor. The literature values (12) for the apparent reaction orders of a, b, and c are 0.25, 0.123, and 1.01, respectively. Analysis of the SO2 Oxidation Results The increase in conductivity observed in the experiments, as shown in Figure 3, is due to the oxidation of SO2. Because the concentration of SO2 is low (≤ 2 × 10{3 M), most of the aqueous SO 2 exists as the HSO3{ anion (pK1 = 1.9 [18]— corresponding to 90% dissociation, for a concentration 1.65 × 10{3 M)—and the concentration of SO32{ anion is negligible (pK2 = 7.2 [18]). H 2SO4 is formed as a product of the reaction. Its first dissociation constant is very large,
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Figure 3. Catalytic oxidation of SO2 using an activated carbon fiber (ACF40) and a granular activated carbon (KUA1) as catalysts. Reaction temperature 25 °C. Initial concentration: 1.65 × 10{3 M for ACF40 and 1.76 × 10{3 M for KUA1.
x
C0 + x
where C0 is the initial concentration of SO2 and x the amount reacted at a given time. Since the number of ions formed from the reaction increases with time, then so does the conductivity. The change in concentration of H+ and the concentration of SO42{ can both lead to a change in conductivity. Assuming that the concentration is sufficiently low, the increase in conductivity with respect to the initial SO2 concentration should be k – k0 = x(Λ °(H+) + Λ °(SO42{) – Λ °(HSO3{)) where k – k0 is the increase in conductivity and Λ is the limiting molar conductivity of the ions in S cm 2 mol {1 (19) (Λ °(H+) = 349.6, Λ°(SO42{) = 160.0, Λ°(HSO3{) ≈ 35). The major contribution to the increase in conductivity arises from the formation of one H+ for each HSO3{ anion oxidized. From this approach, this change in conductivity is directly related to the conversion of SO2 (x). As shown in Figure 3, and depending on the activated carbon, there is a time lag before reaching steady state (see Fig. 3) where, at a low conversion of SO2, the conductivity should change linearly with time (see eq 12). Once the steady state is reached, the initial reaction rate (i.e., the rate at time t = 0) and the activity, as µ mol of SO2 reacted per minute and per gram of carbon, can be calculated as follows from the slope of the straight line of the conductivity versus time plot: a = (∆k/∆t) F /m where a is the activity in units of µ mol min{1 g{1; ∆k/∆t is the slope of the straight line in mS cm{1 min {1; F (in µ mol cm mS{1) is a factor evaluated by dividing the number of moles of the gas dissolved in the volume of solution by the conductivity
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In the Laboratory
of the initial SO2 solution; and m is the mass of carbon in g. (At 25 °C and a SO2 concentration of 1.65 × 10{3 M, F = 723.7 µmol cm mS{1 for a volume of 250 mL). This approach is strictly valid for a low SO 2 conversion (i.e., for the rate at time t = 0). It is interesting to note that the straight line extends over a wide change of conductivity (i.e., up to a high conversion of SO2). According to the kinetic equation (eq 12), this means that, under these experimental conditions, the apparent reaction order for SO2 is nearly 0. It must be remembered that the oxygen concentration remains constant throughout the experiment. This observation agrees with the reaction order found in the literature (b in eq 12 is 0.123 [12]). The measurements shown in Figure 3 correspond to the catalytic oxidation of SO2 at 25 °C using two activated carbons as catalysts: (i) an activated carbon prepared by chemical activation with a Dubinin–Radushkevich (DR) surface area of 1900 m2 g{1 and an activity of 9.6 µmol min {1 g{1 and (ii) a commercial activated carbon fiber with a DR surface area of 1240 m2 g{1 and an activity of 16.8 µmol min{1 g{1. The time lag before reaching steady state is larger for the activated carbon fibers. This is a consequence of the differences that exist between activated carbons of different origin, porosity, and surface chemistry. It is also a consequence of the greater difficulty of doing the experiment with fibers than with granular activated carbon. A general conclusion from the work can be that the number of active carbon sites is not only related to the porosity of the activated carbon but also to its surface chemistry. Conclusions This experiment is a good example of heterogeneous catalysis. It can be used to understand the formation of acid rain and as a possibile way to eliminate pollutants. It is quite simple to perform and is suitable for physical chemistry, chemical engineering, and environmental chemistry courses. The experiment described can be used as a basis for developing a wider study, planned by the teacher according to the amount of time available for heterogeneous catalysis and environmental chemistry. A more complete kinetic study may be accomplished by doing the experiment at different temperatures (to estimate the activation energy) or using different concentrations of SO2 in order to check the reaction order with respect to this reactant. Additionally, the surface chemistry and porosity of a given activated carbon could be modified, for example by oxidation with HNO3 (20), and the effects of this on the catalyst’s properties could be analyzed. This experiment is a good opportunity to introduce students to the complex but important field of carbon surface chemistry.
Finally, it must be noted that we have compared the catalytic activities obtained with this experimental system with results from other setups (viz., flow reactors and a different analytical system), and the results are very concordant. Acknowledgments We thank DGICYT and CICYT (projects AMB96-0799 and QUI97-2051-CE) for financial support. We wish to thank L. E. A. Berlouis for his help in the revision of the English. E. Raymundo-Piñero thanks the MEC for a Ph.D. Thesis grant. Literature Cited 1. See issues of the journal Applied Catalysis B: Environmental. 2. Harter, P. Sulphates in the Atmosphere; IEA Coal Research: London, 1985. 3. Bosch, H.; Janssen, F. Catalysis Today 1988, 2, 4. 4. Trogler, W. C. J. Chem. Educ. 1995, 72, 973. 5. Baird, C. Environmental Chemistry; Freeman: New York, 1995. 6. Stencel, J. M.; Rubel, A. M. Proc. 8th Int. Conf. on Coal Science; Oviedo, Spain, Sep 10–15, 1995; Vol. II, p 1791. 7. Kurth, R.; Tereczki, B.; Boehm, H. P. Proc. 15th Carbon Conference; Philadelphia, PA, 1981; p 244. 8. Knoblauch, K.; Ritcher, E.; Jüntgen, H. Fuel 1981, 60, 832. 9. Zawadzki, J. Carbon 1987, 25, 495. 10. Zhao, X. S.; Cai, G. Y.; Wang, Z. Z.; Wang, Q. X.; Yang, Y. H.; Luo, J. S. Appl. Catal., B 1994, 3, 229. 11. Mochida, I.; Kuroda, K.; Kawano, Sh.; Matsumura, Y.; Yoshikawa, M. Fuel 1997, 76, 533. 12. Mochida, I.; Kuroda, K.; Kawano, Sh.; Matsumura, Y.; Yoshikawas, M.; Grulke, E.; Andrews, R. Fuel 1997, 76, 537. 13. Hawley’s Condensed Chemical Dictionary; Revised by Sax, N. I.; Lewis, R. J. Sr.; Van Nostrand Reinhold: New York, 1987. 14. Grubitsch, H. Anorganisch-Präparative Chemie; Springer: Berlin, 1949. 15. Cazorla-Amorós, D.; Martínez-Mira, I.; Román-Martínez, M. C. Experimentació en Síntesi Química: Química Inorgánica; Universidad de Alicante, Secretariat de Normalització Lingüística: Alicante, Spain, 1996. 16. Greenwood, N. N.; Earnshaw, A. Chemistry of the Elements: Pergamon: Elmsford, NY, 1984. 17. Tereczki, B. Ph.D. Thesis, University of Ludwig-Maximilians, Munich, Germany, 1979. 18. Handbook of Chemistry and Physics, 69th ed.; Lide, D. R., Ed.; CRC: Boca Raton, FL, 1992-93. 19. Atkins, P. W. Physical Chemistry, 5th ed.; Oxford University Press: New York, 1994. 20. Bansal, R. C.; Donnet, J. B.; Stoeckli, F. Active Carbon; Dekker: New York, 1988.
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