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Energy & Fuels 2008, 22, 1714–1719
Combined Process of CO2 Capture by Potassium Carbonate and Production of Basic Zinc(II) Carbonates: CO2 Release from Bicarbonate Solutions at Room Temperature and Pressure Fabrizio Mani,*,† Maurizio Peruzzini,‡ and Piero Stoppioni† Department of Chemistry, UniVersity of Florence, Via della Lastruccia, 3-50019 Sesto Fiorentino, Firenze, Italy, and ICCOM CNR, Via Madonna del Piano, 10-50019 Sesto Fiorentino, Firenze, Italy ReceiVed NoVember 19, 2007. ReVised Manuscript ReceiVed February 11, 2008
The absorption of CO2 in aqueous solutions of K2CO3 has been investigated at six different temperatures in the range 10–60 °C and at atmospheric pressure. The CO2 absorption is fast and occurs with high efficiency without any added catalytic activator. The CO2 capture ranges between 83 and 99% expressed as the percentage of absorbed CO2 with respect to the neat CO2 flowing through the absorbent solution. The efficiency of CO2 removal increases with the temperature, which contrasts with the exoenthalpic nature of the reaction between CO2 and CO32- and the diminished solubility of CO2 at high temperature. Our finding indicates that the increase of the CO2/CO32- reaction rate with the temperature prevails over the opposite effect on the equilibrium due to the temperature increase. The stripping of pure CO2 from the HCO3- solutions has been achieved by adding a zinc(II) salt to the absorbed solutions which releases the captured CO2 rapidly and completely at room temperature and pressure with the simultaneous separation of solid compounds identified as mixtures of commercially valuable basic zinc carbonates. Although the exact evaluation of the overall merits of the proposed CO2 capturing technology would require a complete analysis of the energy profile of the process, it is obvious that any improvement of the CO2 capture process resulting in the associated synthesis of a valuable product without additional energy burdening reduces the cost of CO2 capture and maximizes the neat balance CO2(captured) - CO2(emitted) where CO2(emitted) represents the total energy bill (electrical, thermal, mechanical) required to sustain the entire cycle of CO2 abatement.
Introduction Zinc(II) carbonates with formulas ranging from ZnCO3 · 2H2O to 2ZnCO3 · 3Zn(OH)2 are generally prepared by reacting hydrated Zn(II) salts with alkali metal carbonates and bicarbonates or ammonium bicarbonate, typically in the presence of NaOH or urea.1–4 Depending on several factors including the pH, the temperature, and the reagents’ ratio, pure neutral or basic zinc carbonate or their mixtures with different compositions may be obtained. Basic zinc carbonates are valuable industrial products with applications in a variety of fields ranging from the manufacturing of modified polyesters and cellulosic fibers5,6 and fire retardant resins7 to the production of cosmetic and pharmaceutical * To whom correspondence should be addressed. E-mail:
[email protected]. † University of Florence. ‡ ICCOM CNR. (1) Castellano, M.; Matijevicˇ, E. Chem. Mater. 1989, 1, 78. (2) Hashimoto, T.; Ogawa, T.; Kasai, T. Sumimoto Search 1988, 37, 75. (3) Schindler, P.; Reinert, M.; Gamsjaeger, H. HelV. Chim. Acta 1969, 52, 2327. (4) Mansmann, M.; Schneider, K. Bayer A. G., DE Patent 2404049, 1975. (5) Qu, M.; Zhang, J.; Liu, A.; Du, Z.; Lu, S. Chemical Fiber Group Co., Ltd., CN Patent 1353146, 2002. (6) Sakai Y.; Miura H.; Tsujimoto H.; Honda K.; Fujitani T. New Japan Chemical Co. Ltd., JPN Patent 2002004172, 2002. (7) Iwaisawa, O.; Okano, O.; Ishida, K. Sakai Chemical Industry Co, Ltd., JPN Patent 2002275472, 2002. (8) Yokoyama, H.; Tomomasa, S.; Sakuma, K.; Yoshikawa, N.; Kaway, E.; Ogawa, S. Shiseido Co, Ltd., JPN Patent 200296254, 2002.
products.8,9 Their use as catalysts for hydrotreating of vegetable oils for the production of biodiesel,10,11 and as starting material for the manufacturing of zinc oxide, which is extensively employed in rubber, paint, and ceramic industry, are emerging as new and potentially important applications of basic zinc carbonate.12 Remarkably, all these applications may find additional benefit from the physical nature of the basic zinc carbonate particles with particular interest being addressed to the formation of homogeneous finely dispersed particles with micrometric dimensions. Carbon dioxide is a greenhouse gas and it is nowadays well accepted that the introduction of huge amounts of anthropogenic CO2 in the atmosphere is contributing to the global warming and related climate changes.13 In order to control and, possibly, remove part of the CO2 brought in the atmosphere by costacceptable processes, it is mandatory to develop new and innovative CO2 absorption technologies. The use of alkali (sodium and, mainly, potassium) metal carbonates, either as solids or in solution, for CO2 capture has (9) Schwartz, J. R.; Johnson, E. S.; King, B. T.; Akred, R. J.; Polson, G.; Turley, P. A. The Procter & Gamble Company, US Patent 2004213751, 2004. (10) Teruyama, K.; Kobashi, A.; Nobori, H. J. Chem. Soc. Jpn., Ind. Chem. Sect. 1951, 54, 401. (11) Suppes, G. J. US Patent 2002035282, 2002. (12) Walde, G.; Rudy, A. Brueggerman, L. K-G. DE Patent 3900243, 1990. (13) The Science of Climate Change; Houghton, J. T., Meira Filho, L. G., Callender, B. A., Harris, N., Kattenberg, A., Maskell, K., Eds.; Cambridge University Press: Cambridge, UK, 1995; pp 572.
10.1021/ef7006936 CCC: $40.75 2008 American Chemical Society Published on Web 03/28/2008
Absorption of CO2 in Aqueous Solutions of K2CO3
been studied for a long time.14–22 Portable devices featuring highly efficient and regenerable CO2 absorbers based on solid potassium carbonate have been investigated by NASA for human space travels.23 Na2CO3 would be a CO2 absorbent more useful than K2CO3 in view of its higher natural abundance and reduced cost which makes it suitable for large scale applications. However, CO2 capture processes based on aqueous alkali metal carbonates suffer two main drawbacks due to the low rate of CO2 absorption at room temperature and pressure and the energy cost of the regeneration step of the absorbent solution. A perusal of the data available about the CO2 capture by aqueous K2CO3 solutions shows that high CO2 capture rate generally requires large amounts of catalytic activators such as hindered amines (5–30% w/w)24–29 or arsenic(III) derivatives, 30 and corrosion and foam inhibitors. Increasing the concentration of the absorbent (to 20–40% w/w) and/or raising the temperature close to the boiling point of the solution (>100 °C) represent alternative ways to increment the absorption ability of the carbonate solution.31 The regeneration of carbonate and the desorption of CO2 by heating large volumes of aqueous solutions is a still highly energy consuming step raising serious concerns to the use of this technology to CO2 removal. In this paper, we report our recent results in this area detailing a study which highlights the fast and efficient absorption of CO2 by aqueous potassium carbonate and the preparation of high quality basic zinc carbonates. A patent claiming the technology illustrated here has been recently filed.32 In the process herein presented, the captured CO2 is entirely recovered in pure form from the aqueous absorbent solution at room temperature and pressure without any energy demand. The effects of temperature on CO2 capture and desorption have been also investigated and are here presented. As a general consideration, it must be pointed out that the method here discussed is not regenerative and consequently it is unsuitable for large scale applications such as those dealing with CO2 removal from power plants emissions, due to the relatively small (about 105 metric tons) world market of basic zinc carbonate. However, the process, nicely integrating CO2 capture with basic zinc carbonate production, could take (14) Pérez-Salado Kamps, A.; Meyer, E.; Rumpf, B.; Maurer, G. J. Chem. Eng. Data 2007, 52, 817. (15) Green, D. A.; Turk, B. S.; Gupta, B. S.; Raghubir, P.; Portzer, J. W.; McMichael, W. J.; William, J.; Harrison, D. P. Int. J. EnViron. Manage. 2004, 4, 53. (16) Corti, A. Energy 2004, 29, 415. (17) Hayashi, H.; Taniuchi, J.; Furuyashiki, J.; Sugiyama, S.; Hirano, S.; Shigemoto, N.; Nonaka, T. Ind. Eng. Chem. Res. 1998, 37, 185. (18) Hirano, S.; Shigemoto, N.; Yamada, S.; Hayashi, H. Bull. Chem. Soc. Jpn. 1995, 68, 1030. (19) Roper, G. H. Chem. Eng. Sci. 1995, 4, 255. (20) Williamson, R. V.; Mathews, J. H. J. Ind. Eng. Chem. 1924, 16, 1157. (21) Sartori, G.; Savage, D. W. Ind. Eng. Chem. Fundam. 1983, 239. (22) Benson, H. E.; Field, J. H.; Haynes, W. P. Chem. Eng. Progr. 1956, 52, 433. (23) Onischak, J.; Baker, B. J. Eng. Ind. 1978, 100, 383. (24) Culliname, J. T.; Rochelle, G. T. Ind. Eng. Chem. Res. 2006, 45, 2531. (25) Culliname, J. T.; Rochelle, G. T. Fluid Phase Equilib. 2005, 227, 197. (26) Culliname, J. T.; Rochelle, G. T. Chem. Eng. Sci. 2004, 59, 3619. (27) Yih, S.; Sun, C. Chem. Eng. J. 1987, 34, 65. (28) Say, G. R.; Heinzelmann, F. J.; Iyengar, J. N.; Savage, D. W.; Bisio, A.; Sartori, G. Chem. Eng. Prog. 1984, 80, 72. (29) Giavarini, C.; Rinaldi, G. Chim. Ind. 1977, 59, 621. (30) Giammarco, G.; Giammarco, P. DE Patent 2946193, 1980. (31) Savage, D. W.; Astarita, G.; Joshi, S. Chem. Eng. Sci. 1980, 35, 1513. (32) Mani, F.; Peruzzini, M.; Stoppioni, P. IT Patent FI2007A000199, 2007.
Energy & Fuels, Vol. 22, No. 3, 2008 1715
Figure 1. Simplified scheme of the absorber used in this work. Fittings for both thermometer and pH electrode are not shown.
economic advantage for applications such as gas purification from CO2 in H2 or NH3 production, in coal gasification processes, and in CO2 scrubbing of waste gas streams from, for example, thermal and power plants of chemical factories. Remarkably, the present integrated process of CO2 capture and production of a valuable chemical commodity has been carried out at room temperature and atmospheric pressure, using the same reagents and without any energy penalty with respect to the conventional technology, but avoiding the energy intensive process of the regeneration step. Experimental Section General Information. The absorber device was a home-built glass cylinder with a diameter of 56 mm and height 300 mm fitted with three polyethylene disks threaded on a 2-mm glass rod (Figure 1) and equipped with a thermometer and a combined pH electrode (Crison GLP 22.02 model pH meter calibrated with standard buffer solutions at pH 7.0 and 9.0). The absorber is charged with 0.300 dm3 of the absorbent solution. CO2 absorption experiments were carried out with 0.30 M K2CO3 (4.0% w/w) solutions at six temperatures comprised between 10 and 60 °C. For comparison purposes, room temperature absorption reactions with 0.30 M Na2CO3 (3.1% w/w) solutions were also performed. The temperature of the absorber was kept constant by means of a thermostatted water bath (Julabo model F33-MC refrigerated bath) regulated at the required absorption temperatures. To mimic flue gas, a gas mixture containing 12% (v/v) CO2 in N2 (pure gas Rivoira), was continuously fed into the absorber through a sintered glass diffuser (16–40 µm pores) at the bottom of the absorbent solution with a flow rate of 14 dm3/h. Flow rates of N2 and CO2 were measured with gas mass flow meters (Cole Parmer) equipped with a gas controller (Aalborg). The inlet and outlet CO2 concentrations in the flue gas mixture were measured with a Varian CP-4900 gas chromatograph calibrated with a 10% v/v CO2/N2 reference gas (Rivoira). The vent gas exited from the top of the absorber. The inlet gas mixture was humidified by bubbling through water at the operating temperature before it entered the absorbent reactor. The outlet gas was dried by flowing in turn through a condenser cooled at -5 °C, a concentrated H2SO4 solution, and a gas purification tower filled with P2O5, before being analyzed by the gas chromatograph. The 13C NMR spectra of the absorbed solutions were obtained with a Varian Gemini g300bb spectrometer operating at 75.46 MHz. Chemical shifts are to high frequency relative to tetramethylsilane as external standard at 0.00 ppm. CD3CN was used as internal reference. The H2O solutions were added with ca. 15% (v/v) D2O (Aldrich) contained in a sealed capillary to provide enough signal for deuterium lock system without changes in the concentration of the solutions. NMR spectra were recorded in three successive
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Table 1. Removal Efficiency at Increasing Absorption Steps and at the End of the Absorption Experiment at Variable Temperature CO2 removal efficiency (%) at increasing time (min)a
av CO2 removal efficiencyb
entry
temp (°c)
CO2 flow (mol/h)
15
30
45
60
71.75
%
time (min)c
final removal (%)d
1 2 3 4 5 6
10 20 30 40 50 60
0.0710 0.0688 0.0669 0.0680 0.0706 0.0697
96.8 97.5 98.5 98.9 99.1 99.3
94.7 96.4 97.8 98.2 98.4 98.4
92.4 95.1 96.8 97.1 96.7 96.0
90.5 93.2 95.2 94.9 93.2 91.3
88.4 90.7 92.4 91.3 88.5 86.6
83.1 85.6 88.1 88.1 87.7 86.6
86.58 85 82 78 73.75 71.75
49.1 48.2 50.1 48.2 48.2 48.8
a Average values measured at increasing CO absorption steps and at the last variable step before stopping the experiment. b Average values measured 2 for the entire duration of the absorption experiment. c Duration of each experiment (min). d Final values measured at the end of the absorption experiment.
Table 2. Residual CO32- (mol %) at Increasing CO2 Absorption Steps and at the End of the Experiments as a Function of the Temperaturea av CO32- percentage (mol) at increasing time steps (min)b
final CO32and duration time
NMR exptl CO32- percentage (mol) at different times (min)c
entry
temp (°C)
15
30
45
60
71.75
% mol
time (min)d
30
60
final step
1 2 3 4 5 6
10 20 30 40 50 60
67.9 68.6 67.9 68.5 67.4 67.8
45.6 46.1 46.7 45.9 44.3 44.8
29.1 29.4 27.9 29.0 27.5 28.4
16.7 16.8 17.1 16.6 15.5 17.2
9.1 9.3 9.8 9.6 9.2 11.0
2.7 3.5 5.5 7.1 8.6 11.0
86.58 85 82 78 73.75 71.75
45.9 46.6 46.5 45.7 45.0 44.7
17.3 17.0 17.2 16.3 15.3 16.3
2.2 2.9 4.4 6.1 8.0 11.0
a Experimental values obtained by 13C NMR spectra recorded at each of the three absorption steps are reported for comparison purposes. b Percentage of unreacted CO32- with respect to the sum of CO32- and HCO3- species. c Values determined from 13C NMR analysis of the chemical shift for the fast exchanging CO32-/HCO3- species (see ref 35). d Duration of each experiment (min).
absorption steps, namely after 30 min and 60 min, and at the end of every absorption experiment fixed at about 50% of the amount of CO2 absorption. All reagents were reagent grade. Na2CO3, K2CO3, ZnCl2, and ZnSO4 · 7H2O (Sigma-Aldrich) were used as received. Formation of Zinc Carbonates from CO2 Absorption Solutions. Zinc carbonate was produced at the same time of CO2 release by adding a solution of the appropriate zinc salt to different solutions recovered after each final absorption step. In a typical experiment, different portions of CO2 absorbed solutions (75 mL) were reacted under vigorous stirring with the amount of a zinc salt (ZnSO4 · 7H2O 5.37–6.78 g; ZnCl2 2.71–3.30 g) strictly necessary to release the same quantity of CO2 captured during the absorption step. In a different series of experiments, an excess of zinc salt was used (Zn2+/CO2(absorbed) ) 2/1 as molar ratio). The experiments were carried at constant temperatures between 20 and 70 °C. The release of pure CO2 was followed using a gas-tight apparatus which comprises a 100 mL conical flask, containing the HCO3-/CO32solution obtained during the absorption step. The system was equipped with a pressure-equalizing dropping funnel containing the zinc(II) salt solution and connected to two 100 mL gas burets (Volac) equipped with pressure-equalizing devices. Both burets and pressure-equalizing devices are filled with CO2-saturated water. By means of three-way valves, one buret is filled with CO2 while the other is emptied, thus allowing continuous collection of the gas. The buret internal gas pressure was continuously equilibrated with the external one. The accuracy in the overall volume measurements was about (5 mL. Precipitation of zinc carbonate and release of CO2 were immediate and most of the CO2 (>80%) was recovered within a few minutes. The complete release of CO2 did not take more than 15 min at room temperature. The rate of CO2 desorption greatly increased as the temperature was raised. The voluminous colloidal precipitate was filtered, washed with water, ethanol and diethyl ether in turn before being dried at 60 °C to constant mass. The dry basic zinc carbonate consists of finely dispersed homogeneous powder with a great volume with respect to the mass. Mother liquids of the precipitation reactions were evaporated to dryness at 50 °C and dried in a oven at 60 °C to give solid mixtures which were redissolved in water (35 mL) leaving an insoluble fraction which was separated by filtration, dried, and weighed. All these solutions were checked by 13C NMR spectroscopy and pH measurement.
In order to determine the (molar mass)/CO32- ratio of the samples of zinc carbonates obtained as described above, accurately weighed portions of the precipitates (1–2 g) were completely decomposed with excess of concentrated H2SO4 and the CO2 released was collected and measured using the apparatus previously described. The reliability of the method was checked by analyzing samples of pure K2CO3 and KHCO3, which confirmed a reproducibility within 5%. Each experiment was repeated, at least, in triplicate showing a high reproducibility of the results.
Results and Discussion Absorption of CO2 by K2CO3 Aqueous Solutions. The absorption experiments (Tables 1 and 2) were carried out at variable temperatures using the thermostatted glass absorber described in the Experimental Section. During each absorption experiment the K2CO3 solution was not circulated while the CO2/N2 gas mixture simulating the flue gas was continuously flowing at the bottom of the absorbent through a sintered glass diffuser. The outlet gas was dried before being analyzed by gas chromatography. Samples of the solution (0.5 mL) were withdrawn from the absorbent solution at increasing steps of CO2 absorption, i.e., after 30 and 60 min and at the end of each experiment, and checked by 13C NMR spectroscopy. Changes of pH were also measured throughout the CO2 absorption process showing values decreasing from 11.5 to 8.4. The main reaction of CO2 (Ka(293) ) 4.08 × 10-7)33 (the reference standard of each species is 1 mol dm-3) with the CO32- (Kb(293) ) 2.1 × 10-4)33 solution is CO2 + CO32- + H2O a 2HCO3- Keq(293) ) 8.6 × 104 (1) Reaction 1 shows a favorable equilibrium which is substantially shifted toward the formation of bicarbonate when comparable amounts of CO2 and CO32- are reacted. However, the rate of CO2 absorption by carbonate is slow compared with a similar reaction involving NH3 or amines, making the process inefficient for CO2capture. (33) Edwards, T. J.; Newman, J.; Prausnitz, J. M. Ind. Eng. Chem. Fundam. 1978, 17.
Absorption of CO2 in Aqueous Solutions of K2CO3
Figure 2. Average CO2 absorption (mol %) at increasing steps (time in minutes), and at different temperatures (°C).
In our experiments carried out at ambient temperature using relatively low K2CO3 concentration (4.0% w/w), the rate of CO2 capture by carbonate solution was greatly enhanced by increasing the liquid–gas contact surface and extending the contact time between the gas and the liquid absorber. To this purpose, the gas mixture is introduced at the bottom of the absorbent through a sintered glass diffuser (16–40 µm pore) favoring the formation of a great amount of microbubbles. Additionally, three disks spaced within the solution (see Figure 1) increase the liquid turbulence and then the contact time between the flowing gas and the absorber solution. Each experiment was stopped when no more than about 50% of CO2 is absorbed by residual carbonate and repeated, at least, in triplicate showing high reproducibility of the results.34 The experimental conditions and the results of typical absorption process are reported in Tables 1 and 2. Inspection of Table 1 shows that, at each investigated temperature, the total efficiency of CO2 removal was greater than 83% with residual unreacted CO32- in solution in the range 3–11% (Table 2). The CO32- and HCO3- percentages, calculated from the amounts of absorbed CO2 (see reaction 1), well agree with the experimental values determined by recording 13C NMR spectra at each of the three absorptions steps, using the procedure previously reported by us.35 As long as an excess of carbonate is present in the absorbent solution (i.e., during the first 30 min of the absorption process) the average removal efficiency of CO2 is higher than 94%, with a maximum value of 99.3% for the reaction at 60 °C, after 15 min of CO2 absorption. Within 60 min the average removal efficiency was always greater than 90% (Table 1) with residual CO32- being 15–17% (Table 2). In Figure 2 the absorption efficiency at each investigated temperature is plotted versus the reaction time, for increasing amount of absorbed CO2. As the maximum CO2 removal efficiency increases with the temperature, the CO2 absorption at 60 °C requires the least time to reach the end of experiment, fixed at about 50% of total CO2 absorption. These results clearly confirm that the kinetic constraints of the CO2/CO32-/H2O reaction prevail over the (34) During the different absorption experiments, percent changes of each solution species do not exceed more than 2 units %. (35) Mani, F.; Peruzzini, M.; Stoppioni, P. Green. Chem. 2006, 8, 995.
Energy & Fuels, Vol. 22, No. 3, 2008 1717
Figure 3. Average CO2 absorption (mol %) at different temperatures (°C) and steps (time in min). The curve labeled as “final” refers to the end of each experiment.
thermodynamic ones. Even if the reaction equilibrium (1) is disfavored and the CO2 solubility in water decreases with the temperature, the increasing rate of reaction 1 at higher temperature allows more CO2 to react, provided sufficient amount of CO32- is present in solution. When the solution contains a large excess of HCO3- with respect to CO32-, as occurs close to the end of the absorption reaction, increasing the temperature over 30 °C results in a overall decreasing of the absorption efficiency. This is shown in Figure 3, where the absorption efficiency is plotted against the temperature for six increasing CO2 absorption steps. The data shown in Figure 3 suggest that the best compromise between absorption efficiency, time, and temperature is achieved by carrying out the reaction between 30 and 40 °C, with a final overall absorption efficiency of 88.1%. Replacing Na2CO3 for K2CO3 under the same experimental conditions showed an averaged 15–20% reduction of CO2 absorption. In this respect, we cannot provide any explanation accounting for the different behavior of the two alkaline metal ions in the Lewis acid–base reaction (1). In view of the low cost of the naturally abundant Na2CO3 with respect to K2CO3, it would be of importance to develop efficient carbon dioxide absorption systems based on sodium carbonate. In this regard, we are currently investigating CO2 absorption by dilute Na2CO3 solutions with removal efficiency comparable with K2CO3. Release of CO2 at Room Temperature and Pressure and the Production of Basic Zinc Carbonates. Portions (75 mL) of the solutions recovered after each final step of the CO2 absorption process contain 0.036–0.042 mol of HCO3- (about 89–97% on molar basis) and unreacted CO32- with a pH between 8.40 and 8.96. Treatment of these solutions at room temperature with aqueous Zn2+ solution (as sulfate or chloride), as described in the Experimental Section, causes the immediate formation of a precipitate followed in a few seconds by vigorous CO2 evolution. The released gas was collected and measured with the gas-tight apparatus described in the Experimental Section. Several experiments (from a minimum of six for each solution) were necessary to determine the minimum amount of zinc salt required to desorb the same amount of CO2 captured during the absorption process. A perusal of a series of precipitation/desorption experiments carried out at room temperature using 25 different CO2 absorbed solutions and diverse
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Table 3. Experimental Conditions for the Precipitation of Zinc Carbonate in the Presence of Zinc(II) Salts entry
reaction temp (°C)
solution pH
% HCO3(mol)
CO2(abs) (mol)
Zn2+/ CO2 (abs)
CO2(des)/ CO2 (abs)
solid carbonate (mol)
molar mass M/CO32-
1 3 3 4 6 6 6 7c 7c
20 20 20 40 20 40 20 20 20
8.40 8.56 8.56 8.67 8.96 8.96 8.96 9.70 9.70
97.3 94.5 94.5 93.0 89.0 89.0 89.0 54.4 54.4
0.0213 0.0201 0.0201 0.0196 0.0181 0.0181 0.0181 0.0085 0.0085
1.13a 1.12b 2.01b 1.07b 1.12a 1.05a 0.99b 1.92a 1.86b
0.99 1.00 1.54 1.40 0.99 1.26 1.00 1.06 1.03
0.0120 0.0110 0.0103 0.0088 0.0109 0.0075 0.0096 0.0068 0.0072
219 ( 4 228 ( 2 250 ( 2 256 ( 2 218 ( 4 259 ( 1 218 ( 6 273 ( 5 272 ( 1
a
Zinc(II) chloride. b Zinc(Ii) sulfate. abs ) absorbed; des ) desorbed. c Absorption experiment stopped after 30 min.
zinc salts (ZnSO4 · 7H2O or ZnCl2) indicates that the molar ratio between desorbed CO2 and Zn2+ approaches the value of 1/1. Table 3 collects the results of a few selected experiments. Increasing the reaction temperature in the range 40–70 °C, increases the CO2(desorbed)/Zn2+ ratio, and a maximum value of 1.6/1 is reached at the highest investigated temperature. On the other hand, increasing the concentration of Zn2+ at room temperature causes an increment (in most experiments ca. 10%) of the CO2(desorbed)/CO2(absorbed) molar ratio. The reaction of Zn2+ solutions with the CO2 absorbed solutions at the final step of the absorption process may be described by a complicated picture of simultaneous homo- and heterogeneous equilibria which involve, apart from Zn(II) species, H2O, HCO3-, and CO32- as predominant species and dissolved CO2, H2CO3, and OH- as less abundant ingredients. However, such complex thermodynamic behavior can be much simplified by neglecting unimportant equilibria featuring species with comparatively smaller concentration or equilibrium constants. The main reactions to be taken into account to rationalize the decomposition of HCO3- and the formation of CO2 are 2HCO3- a CO2 + CO32- + H2O Keq(293) ) 1.2 × 10-5 (2) expressed by reverse reaction (1), and HCO3- a CO2 + OH- Keq(293) ) 2.4 × 10-8
(3)
Both reactions are substantially left-shifted, and negligible amounts of CO2 are in equilibrium with HCO3- in the homogeneous equilibria at room temperature. However, the precipitation of insoluble carbonates and hydroxides at room temperature and/or the increase of the temperature may shift to right both reaction 2 and, to a less extent, reaction 3 and then produce appreciable amounts of gaseous CO2 due to its poor water solubility (298 K, 0.145 g CO2/100 g H2O; 333 K, 0.058 g CO2/100 g H2O).36 A number of insoluble zinc carbonates with formulas ranging from ZnCO3 · 2H2O37 (1) and 2ZnCO3 · 3Zn(OH)238 (2) (i.e., Zn5(CO3)2(OH)6) have been reported depending on the experimental conditions, namely the concentration of either HCO3or CO32-, their molar ratio, pH, temperature, and, to a lesser extent, the nature of the zinc salt used. In the presence of Zn2+, the reactions 2 and 3 can be rewritten as Zn2+ + 2HCO3- a CO2 + ZnCO3 + H2O
(4)
Zn2+ + 2HCO3- a 2CO2 + Zn(OH)2
(5)
(36) CRC Handbook of Chemistry and Physics; Weast, R. C., Ed.; CRC Press, Inc., Boca Raton, FL, 1986; p, B82. (37) Smith, H. J. J. Am. Chem. Soc. 1918, 40, 883. (38) Ghose, S. Acta Crystallogr. 1964, 17, 1051.
Both zinc(II) carbonate and hydroxide are slightly soluble in water, even if widely different solubility values have been reported for Zn(OH)2 (1-10 × 10-5 g/100 g H2O) and ZnCO3 (2 × 10-2 to 6 × 10-4 g/100 g H2O).39 The overall reaction yielding the basic zinc carbonate (2) can be written as 5Zn2+ + 10HCO3- a Zn5(CO3)2(OH)6 + 8CO2 + 2H2O(6) Moreover, the CO32- and the OH- species in solution may produce basic zinc(II) carbonates without any CO2 evolution. Quantitative experiments were carried out to rationalize the reaction between the HCO3-/CO32- solutions and Zn(II) salts. The results of a selected number of experiments are reported in Table 3. In a typical experiment, 0.0226 mol of hydrated zinc sulfate in 10 mL of H2O was slowly added at room temperature to a 75 mL aliquot of the absorbed solution containing 0.0403 mol of HCO3- and 0.00236 mol of CO32- at pH ) 8.56 (Table 3, entry 3). The precipitate of zinc carbonate which separated out from the solution containing K2SO4, was filtered, thoroughly washed with water and dried at 60 °C to constant weight. The mass of the filtered solid was 2.58 g, while a volume of 0.483 dm3 (293 K, 1.01 bar) of CO2 (0.0201 mol) was measured, corresponding to 100% of the CO2 previously captured in the absorption process. In order to estimate the yield of the reactions 4 and 5, the mother liquid was evaporated to dryness at 50 °C and the residue dried at 60 °C in the oven. The mass of the solid mixture was 4.03 g and contained all of the K2SO4 which formed as a secondary product of the process. Addition of 35 mL of water completely redissolved the K2SO4 contained in this solid mixture and left 0.110 g of an insoluble residue which was confirmed by chemicophysical measurements to be zinc(II) carbonate. On this basis, it is confirmed that both reactions 4 and 5 are practically complete at room temperature leaving less than 5% of zinc carbonate in solution. In a second experiment, 0.0210 mol of zinc sulfate solution was reacted with a 75 mL aliquot of the absorbed solution thermostatted at 40 °C (Table 3, entry 4). The recovered mass of basic zinc(II) carbonate was 2.24 g while the released CO2 was 0.666 dm3 (295 K, 1.01 bar; 0.0275 mol) providing a CO2(desorbed)/CO2(absorbed) molar ratio of 1.40/1 and no insoluble residue after complete evaporation of the solution. In order to estimate the composition of the precipitated zinc(II) carbonate, 25 samples of the solids obtained from different CO2 desorption reactions were completely decomposed with concentrated sulfuric acid and the evolved CO2 was collected and measured with the apparatus described in the experimental section. Mass (g) of the decomposed sample and amount (mol) of the collected CO2 were used to calculate the (39) Aylett B. J. In ComprehensiVe Inorganic Chemistry; Bailar J. C., Jr.;. Emeléus, H. T., Nyholm, R., Trotman-Dickenson, A. F., Eds.; Pergamon Press: Oxford, UK, 1973; Vol. 3, pp 224–234.
Absorption of CO2 in Aqueous Solutions of K2CO3
(molar mass (M))/CO32- of each sample. For most of the samples, M/CO32- values in the range 220 ( 10 (ca. 50% of the samples) and 260 ( 10 (ca. 40%) were obtained. These results point out that neither 1 nor 2 was obtained as pure compound, but a mixture of them in different proportions was always recovered. Thus, M/CO32- values of 220 are in agreement with 1/2 mixtures with a molar ratio value around 2 (calculated value ) 218). M/CO32- values of 260 ( 10 are in agreement with 1/2 mixtures with a molar ratio value of about 0.5 (calculated value ) 252). Apparently, there is not any clear relationship between the Zn2+/HCO3- molar ratio, the pH of the solution, the nature of zinc(II) salt (either sulfate or chloride), and the calculated M/CO32- values. However, increasing the reaction temperature and/or the Zn2+/HCO3- molar ratio causes more CO2 release and gives mixtures with a greater content of 2, as indicated by the greater values of M/CO32-. Pure compound 2 seems to form only when comparable amount of HCO3- and CO32- are present in solution as occurs when the CO2 absorption reaction is stopped after 30 min (Table 3, entries 7). Conclusions The efficient absorption of CO2 (g83%) by a dilute water solution of K2CO3 at room temperature is the focus of this work. The absorption efficiency decreases with increasing absorbed CO2, i.e., HCO3- concentration, but, providing that comparable amounts of CO32- and HCO3- are present in the absorbent solution, the absorption efficiency is higher than 94% at any temperature investigated (10–60 °C). The rate of absorption increases with the temperature, and the efficiency of the process reaches the maximum value of 99.3% at 60 °C, the highest temperature investigated, without the need of adding any activator.
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More interestingly, we have devised a simple and straightforward procedure to overcome the highly expensive heating step for CO2 stripping from the HCO3- solutions. The addition of a zinc salt to the solution resulting from the investigated CO2 absorption process rapidly releases at room temperature and pressure practically all of the captured CO2, with the simultaneous separation of solid products identified as a mixture of different basic zinc carbonates in the form of homogeneous and very finely dispersed white powder. The recovery from the absorbed solutions of valuable solid products, such as basic zinc carbonates, and the release of highly pure CO2 at room temperature and pressure, might lower the costs of the CO2 uptake process based on K2CO3 solutions, which could become a more and more attractive method for CO2 capture. In order to successfully scale up the process of CO2 absorption/ desorption, from the laboratory scale to applied dimensions, it will be necessary to replace K2CO3 with more inexpensive chemicals, such as native Na2CO3 and utilize secondary zinc salts recovered, for example, from steelmaking flue dust or iron galvanizing processes. Although, the problem of CO2 capture in flue gas is certainly far to being resolved, our study pinpoints the significant advantages of a process which combines high absorption efficiency and production of useful solid materials minimizing energy requirements. Acknowledgment. Financial support from MIUR (Rome, Italy) and Ente Cassa di Risparmio di Firenze (Florence, Italy), through Florence Hydrolab Project, is gratefully acknowledged. EF7006936
