Energy Fuels 2009, 23, 4768–4773 Published on Web 06/23/2009
: DOI:10.1021/ef9001082
Study on the Interaction between NOx and K2CO3 during CO2 Absorption† Xinglei Zhao,‡ Michael A. Simioni,§ Kathryn H. Smith,§ Sandra E. Kentish,§ Weiyang Fei,‡ and Geoffrey W. Stevens*,§ The State Key Laboratory of Chemical Engineering, Department of Chemical Engineering, Tsinghua University, Beijing 100084, China, and Cooperative Research Centre for Greenhouse Gas Technologies, Department of Chemical and Biomolecular Engineering, The University of Melbourne, Melbourne, Victoria 3010, Australia. ‡ Tsinghua University. § The University of Melbourne. Received February 7, 2009 . Revised Manuscript Received May 8, 2009
The trend toward global warming has become more and more apparent in recent years. It has been widely accepted that CO2 is the main contributor to this process and that CO2 emissions urgently need to be reduced. However, the capital cost for CO2 capture from a coal-fired power station using traditional solvent absorption technology is relatively high. A possible solution to reduce the cost is to achieve absorption of SO2, NOx, and CO2 in a one-step process. Potassium carbonate (K2CO3) is a potential solvent for absorbing all of these components simultaneously. To apply this absorption process in an industrial situation, it is very important to understand the interaction between the solvent and minor components contained in the flue gas, such as SO2 and NOx. The effect of SO2 on CO2 absorption using K2CO3 solutions has been previously discussed. This paper studies the influence of NOx upon CO2 absorption using 30 wt % K2CO3 solvent. It was found that the reaction between NOx and K2CO3 was irreversible, which would result in the accumulation of NO2- and NO3- in the solvent using an industrial flue gas process, with a simultaneous reduction in the CO32- concentration.
from the reactions between nitrogen and oxygen at very high temperatures. This chemical reaction produces more than 90% NO and 5-10% NO2.4 NO is a colorless, nonflammable, and odorless gas, which is not a real pollutant by itself. However, it is easily converted to NO2, which is one of the most dangerous gases known.5 NO2 is visible as a brown-red gas and has a very characteristic annoying smell. It is also very reactive and a strong oxidizing agent, acting as a deep irritant in the lungs. There is toxicological evidence of pulmonary inflammation, deterioration of respiratory defense mechanisms, and increased susceptibility to respiratory pathogens after NO2 exposure.6-9 NO2 also plays a key role in the atmospheric reactions that produce ozone and smog and is also assumed responsible for acid rain.10,11 Emissions of NOx are projected to increase from 12 million tons in 1995 to around 28 million tons by 2020.2 NOx emissions from coalfired power stations currently account for 20.8% of the total emissions.12
1. Introduction The Intergovernmental Panel on Climate Change has concluded that warming of the climate system is unequivocal. Further, most of the observed increase in global average temperatures since the mid-20th century is very likely due to the observed increase in anthropogenic greenhouse gas concentrations, such as carbon dioxide, methane, and nitrous oxide.1 Carbon capture and storage has the potential to play a major role in reducing greenhouse gas emissions from stationary emitters of CO2, such as coal-fired power stations. However, the high costs associated with implementing CO2capture technologies, such as solvent absorption, on a large scale is a significant limitation. The presence of minor components in the power station flue gas, such as SO2 and NOx, also make the CO2-capture process using solvent absorption more challenging because these gases can degrade the solvent, resulting in inefficient operation. To reduce acid rain, desulphurization equipment has been used for many years by power plants to remove the SO2 from the flue gas. However, there are very few pollution controls or other emission reduction measures in place for NOx, even though it is also a major greenhouse gas1 as well as a contributor to acid rain and ground-level ozone.2 The mixture of nitric oxide (NO) and nitrogen dioxide (NO2) is generally referred to as NOx.3 These gases result
(4) Bi, T. C. Environ. Prot. Petrochem. Ind. 2005, 28 (3), 55–60. (5) Paliatsos, A. G.; Kaldellis, J. K.; Koronakis, P. S. Environ. Bull. 2002, 11 (12b), 1119–1126. (6) World Health Organization (WHO). WHO Regional Publications. 1987; pp 297-314. (7) U.S. Environmental Protection Agency. Air quality criteria for oxides of nitrogen. Office of Research and Development, Washington, D.C., 1993. (8) Berglund, M.; Bostroem, C. E.; Bylin, G. Scand. J. Work, Environ. Health 1993, 19 (Supplement 2), 14–20. (9) WHO regional Office for Europe, N. L. World Health Organization Regional Office for Europe, Copenhagen Report. 1995. (10) Cape, J. N.; Fowler, D.; Davison, A. Environ. Int. 2003, 29 (2-3), 201–211. (11) Erisman, J. W.; Draaijers, G. Environ. Pollut. 2003, 124 (2), 379– 388. (12) Schnelle, K. B.; Brown, C. A. Air Pollution Control Technology Handbook; CRC Press: Boca Raton, FL, 2002.
† Progress in Coal-Based Energy and Fuel Production. *To whom correspondence should be addressed. Telephone: þ61-38344-6621. Fax: þ61-3-8344-8824. E-mail:
[email protected]. (1) Solomon, S.; Qin, D.; Manning, M. Contribution of Working Group I to the Fourth Assessment Report of the Intergovernmental Panel on Climate Change. Summary for Policymakers. 2007. (2) Streets, D. G.; Waldhoff, S. T. Atmos. Environ. 2000, 34, 363–374. (3) L opez, A. B.; Garcia, A. G. Fuel Process. Technol. 2005, 86, 1745– 1759.
r 2009 American Chemical Society
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: DOI:10.1021/ef9001082
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Values of HNO2, the Henry’s law coefficient, and k8, the second-order aqueous-phase rate constant, were determined to be (7.0 ( 0.5) 10-5 M atm-1 and (1.0 ( 0.1) 108 M-1 s-1 at 22 °C, respectively. Zhang et al.19 introduced the idea of NOx absorption using the alkaline solution Na2CO3 according to the following steps: Step 1. NOx diffuses to the gas-liquid interface and dissolves. Step 2. Dissolved NOx reacts with H2O and forms HNO2 and HNO3 as described in eqs 9 and 10. N2 O3 ðor NO þ NO2 Þ þ H2 O f 2HNO2 ð9Þ
To reduce greenhouse gas emissions and also reduce the capital cost of CO2 capture, a one-step solvent absorption process may be possible for removing CO2 as well as NOx and SO2. Potassium carbonate solution, K2CO3, is a potential solvent for this application. The effect of SO2 on CO2 capture using K2CO3 solutions has been previously considered;13 however, the effect of NOx on CO2 absorption using K2CO3 solvent is not widely available. Therefore, the main aim of this study is to study the interaction between NOx and K2CO3 during CO2 absorption. To do this, K2CO3 solutions were loaded with NOx and, subsequently, with CO2 with the resulting solutions analyzed for nitrite (NO2-) and nitrate (NO3-) concentrations in the liquid phase and nitric oxide (NO) concentrations in the gas phase. Ion chromatography was used to analyze the liquid samples, while gas chromatography was used to determine the gas-phase compositions.
2NO2 ðor N2 O4 Þ þ H2 O f HNO2 þ HNO3
Step 3. A neutralizing reaction between HNO2, HNO3, and Na2CO3 as follows: Na2 CO3 þ 2HNO2 f 2NaNO2 þ H2 O þ CO2 ð11Þ
2. Theory
Na2 CO3 þ 2HNO3 f 2NaNO3 þ H2 O þ CO2
The combustion of fossil fuels predominantly generates NO. However, this gas is partially converted to NO2 at low temperatures when vented to the atmosphere, as shown in eq 1.14 Sathiamoorthy et al.15 reported that NO is generally converted to NO2 through a variety of reactions that may include direct reaction with dissociated oxygen or reactions with ozone in the presence of oxygen. NO2 may be removed through subsequent reactions with hydroxyl radicals created from organic species or water vapor. ð1Þ 2NO þ O2 ¼ 2NO2
N2 O3 þ H2 O f 2HNO2
ð13Þ Combining eqs 2, 4, 10, 11, and 12 and replacing Na by K gives 2NO2 ðor N2 O4 Þ þ K2 CO3 ¼ KNO2 þ KNO3 þ CO2
ð3Þ
ðk4 ¼ 554 s -1 Þ
ð4Þ
ðk5 ¼ 8:68 106 s -1 Þ
ð5Þ
N2 O4 þ H2 O f HNO2 þ HNO3
ð14Þ Furthermore, CO2 formed from eqs 13 and 14 or CO2 present in the flue gas could react with K2CO3 as follows: K2 CO3 þ CO2 þ H2 O ¼ 2KHCO3 ð15Þ
The equilibrium constants for eqs 2 and 3 are as follows: K2 ¼ ½N2 O4 =½NO2 2 ¼ 4858 m3 mol -1
ð6Þ
K3 ¼ ½N2 O3 =½NO2 ½NO ¼ 21:56 m3 mol -1
ð7Þ
ð12Þ
Na and K belong to the same group in the element periodic table. Thus, similar performance would be expected between K2CO3 and Na2CO3 when reacting with NOx. In this study, K2CO3 has been used as the absorption solution because of its higher solubility in water when compared to Na2CO3. Higher water solubility allows higher solvent concentrations to be used and, hence, better absorption capacity without precipitation problems. Combining eqs 1, 3, 9, and 11 and replacing Na by K gives the following: 2NO þ 1=2O2 ðor NO þ NO2 Þ þ K2 CO3 ¼ 2KNO2 þ CO2
Kormyama and Inoue16 studied the absorption of nitrogen oxides into water and presented the reaction model under atmospheric pressure and temperature as follows: ð2Þ 2NO2 ¼ N2 O4 NO þ NO2 ¼ N2 O3
ð10Þ
This study aims to verify the occurrence of eqs 13 and 14 during CO2 capture using K2CO3 solvent from model flue gases containing NOx.
Javeda et al.17 concluded that the removal of NO2 from a gas stream is relatively easy because it reacts with water and air to form nitric acid and, hence, may be removed by aqueous scrubbing. Lee and Schwartz18 reported that the following reaction occurred, which is a combination of eqs 2 and 4 above: ð8Þ 2NO2 ðgÞ þ H2 O ðlÞ f 2H þ þ NO3 - þ NO2 -
3. Experimental Section 3.1. Materials. Potassium carbonate (99%) was obtained from Thasco Chemical Co., Ltd. (Amphur Muang, Rayong, Thailand). Ferrous sulfate (88%), sodium nitrite (97%), and potassium nitrate (99%) were purchased from Chem-Supply (Gillman, Adelaide, Australia). Sulfuric acid (98%) and sodium bicarbonate (99.7%) were from Nuplex Industries (Aust) Pty. Ltd. (Torrensville, Adelaide, Australia). Sodium carbonate (99.9%) was supplied by Merck (Kilsyth, Victoria, Australia). Carbon dioxide (99.5 vol %), nitrogen (99.99 vol %), O2 (99.9 vol %), and NO (2430 ppmv in nitrogen) were obtained from BOC Gases Australia Limited North Ryde, New South Wales, Australia. 3.2. NOx Generation and Absorption in K2CO3 Solution. The NOx generation and absorption experimental rig can be found
(13) Wappel, D. The effect of SO2 on the absorption of CO2 using potassium carbonate solution. Diploma Thesis, The University of Melbourne, Melbourne, Victoria, Australia, 2006. (14) Francois, G. Appl. Catal., A 2001, 222 (1-2), 183–219. (15) Sathiamoorthy, G.; Kalyana, S.; Finney, W. C.; Clark, R. J.; Locke, B. R. Ind. Eng. Chem. Res. 1999, 38 (5), 1844–1855. (16) Komiyama, H.; Inoue, H. Chem. Eng. Sci. 1980, 35 (1-2), 154– 161. (17) Javeda, M. T.; Irfana, N.; Gibbs, B. M. J. Environ. Manage. 2006, 78 (3), 1. (18) Lee, Y. N.; Schwartz, S. E. J. Phys. Chem. 1981, 85 (7), 840–848.
(19) Zhang, Y.; Chen, J. S. Anhui Chem. Ind. 2006, 32 (2), 54–55.
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Figure 1. Experimental rig for NOx generation and gas absorption into potassium carbonate solution.
was again manipulated to provide a molar ratio of NO/O2 consistent with that seen in power plant flue gas (Table 1). In this case, a series of absorption experiments were conducted at temperatures of 15, 40, 70, and 90 °C. In each case, the absorption time was 12 h, and the temperature was controlled to within (2 °C via a hot plate. In some cases after absorption of NO, carbon dioxide (from a gas cylinder) was also absorbed into the potassium carbonate solution at a flow rate of 300 mL/min for about 3 h. This led to CO2 loadings between 0.08 and 0.12 mol of CO2/mol of K2CO3. The CO2 loading of the solution was then determined using the following equation: moles of CO2 absorbed ½HCO3 - CO2 loading ¼ ¼ moles of K2 CO3 2½CO3 2 - þ ½HCO3 - ð17Þ
Table 1. NO Generation and Absorption Experimental Conditions NO source gas flow rate (mL/min) gas composition (mol %) molar ratio of NO/O2 absorption time (h)
NO N2 O2 NO N2 O2
NO produced by chemical reaction
NO sourced from cylinder
0.045
0.19 90 7.8 0.19 91.9 8.0 0.025 12
1.9 2.3 97.7 0.024 5
in Figure 1. The rig consisted of three gas cylinders (O2, NO/N2, and CO2), three rotameters for the different gases, four Erlenmeyer flasks, a long-stem funnel, magnetic stirrers, and valves. A leakage test was completed before every run using water at gas flow rates higher than those used in the experiments. The NOx used in these experiments was obtained via two methods. In the first method, NO was generated via a chemical reaction. Inside the first flask, nitric oxide could be generated by the reaction of sodium nitrite (NaNO2) with ferrous sulfate (FeSO4) and sulfuric acid according to the following reaction:21,22
3.3. Vapor-Liquid Equilibrium (VLE) Rig. After NOx (and in some cases, CO2) was absorbed into the solution, the reversibility of the reactions was studied at 100 °C, which is typical of regeneration conditions used for stripping CO2 from a loaded solution, in a VLE rig (Figure 2). This apparatus can be used to determine the VLE data between a solvent and different gas mixtures and was constructed in-house based on the papers of Nasir and Mather22 and Isaacs et al.23 It consists of two equilibrium vessels, each with a volume of 200 mL, a condenser, a bubble flow meter, a rotameter, and a gas sampling port. The equilibrium vessels were placed in a temperature-controlled oil bath. In these experiments, 250 mL of NOx-loaded solution, prepared as above, was divided equally between the two equilibrium vessels, which were then sealed. Pure nitrogen was passed via a rotameter and then through the two vessels at a low flow rate (5-10 mL/min) until an equilibrium gas/liquid mixture was established in the second vessel. Water evaporated from the first vessel in saturating the nitrogen stream was continuously returned by condensation from the exhaust gases at ambient temperature. The time taken to reach equilibrium was studied by taking regular gas samples until a constant composition of CO2, NO2, and N2 was obtained. Generally, equilibrium conditions were reached after approximately 1 h, but all experiments were completed for 2 h to ensure that samples were taken at
2NaNO2 þ 2FeSO4 þ 3H2 SO4 f Fe2 ðSO4 Þ3 þ 2NaHSO4 þ 2H2 O þ 2NOv
ð16Þ
The resulting NO supply was then mixed with oxygen before absorption into a 2.77 mol/L (or 30 wt %) K2CO3 solution, which is the typical solvent concentration used for CO2 absorption. The oxygen flow rate was manipulated to provide a molar ratio of NO/O2 consistent with that seen in power plant flue gas (Table 1). These absorption experiments were all conducted at ambient temperature (15-16 °C) for a period of 5 h. The outlet gas pressure was controlled to slightly less than 101.3 kPa using a vacuum pump to obtain a steady gas flow rate. In other experiments, a cylinder supply of NO in N2 (2430 ppmv) was used in place of the reaction. The oxygen flow rate (20) Wei, Z. S.; Lin, Z. H.; Qiu, R. L.; He, H. M. Acta Sci., Nat. Univ. Sunyatseni 2006, 45, 103. (21) Xu, Y. Q.; Wang, A. Q. J. Henan Urban Constr. Junior Coll. 2001, 10 (2), 48–50. (22) Nasir, P.; Mather, A. E. Can. J. Chem. Eng. 1977, 55 (6), 715–717.
(23) Isaacs, E. E.; Otto, F. D.; Mather, A. E. J. Chem. Eng. Data 1980, 25 (2), 118–120.
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Table 2. Ion Concentrations Following NO Absorption (without Any CO2 Absorption) Using 30 wt % K2CO3 Solution at Ambient Temperatures (15-16 °C) anion concentration in the solvent after absorption NO2- (mmol/L) NO3- (mmol/L) CO32- (mmol/L) HCO3- (mmol/L) ratio NO2-/NO3total physical absorption of NO and NO2 predicted from Henry’s law constantsa (mmol/L)
5 h of absorption using 2.3 mol % NO 200 9.1 2560 40 22 0.25
12 h of absorption using 0.19 mol % NO 4.4 0.2 2730 20 0.021
a Henry’s law constant for NO2 = 7.0 10-3 M/atm 32 with ΔH/R = 1800 K.33 Henry’s law constant for NO = 1.9 10-3 M/atm with ΔH/R = 1500 K.34
Figure 2. Schematic of the VLE rig.
equilibrium conditions. The exit gas stream was then passed through a 5% triethanolamine (TEA) solution to capture the NO2 content24 prior to exhaust. After equilibration with nitrogen, some experiments were also conducted with oxygen as the circulating gas to determine the likelihood of NO2- oxidation to NO3-. 3.4. Analytical Methods. There are several anions that could exist in the solvent after absorption of NOx and CO2, including NO2-, NO3-, CO32-, and HCO3-. Acid-base titration was used to determine the CO32- and HCO3- concentrations. The titration was performed using a Metrohm-Titrando 809 autotitrator (Switzerland). Any potential effect from NO2- and NO3- absorption on CO32- and HCO3- titration accuracy was tested by adding different amounts of NO2- and NO3- (up to 20 g/L) into a standard solution. The results showed that there was no noticeable difference in the concentration of CO32- and HCO3- after adding NO2- and NO3- to the solution. There are several methods reported in the literature for NO2and NO3- analysis, including Fourier transform infrared spectroscopy (FTIR),25,26 UV-vis spectrophotometry (UV),27 and ion chromatography (IC).28-30 Our analysis showed that the peak height for NO3- and NO2- was too small when measured by FTIR, even when the NO2- concentration reached 2000 ppm. UV also had associated errors because of the high CO32concentration and some possible side reactions. In comparison to FTIR and UV, IC provided better performance at high CO32- concentrations. IC was carried out using a Dionex 4500i ion chromatograph (from RMIT University) equipped with a Dionex GP40 programmable gradient pump and a Dionex Ion Pac AG4A-SC guard column in series with a Dionex Ion Pac AS4A-SC 4 mm column. This was followed by either one or two Dionex anion micromembrane suppressors (AMMSII, P/N 043074) and then a Dionex conductance detector. A Shimadzu C-R3A computing integrator was used to acquire peak heights and peak areas. Samples were filtered and injected through 0.45 μm Nylon 66 membrane filters (Alltech Assoc. Aust. Pty. Ltd.) into a 50 μL injection loop.
At least three replicate injections of each sample were made, and the results were averaged. The eluent was 1.7 mmol/L Na2CO3-1.82 mmol/L NaHCO3 at a flow rate of 1.0 mL/min. The suppressor regenerant was 12.5 mmol/L sulfuric acid. The NO concentration in the gas phase from the first flask (Figure 1) and the equilibrium vessels (Figure 2) was determined by gas chromatography (GC-8A Shimadzu) with a Hayesep DB column 100/120 (Altech) and thermal conductivity detector (TCD). The NO2 concentration was determined by analysis of the downstream TEA solution using ion chromatography. Makoto et al.24 reported that NO2 could be completely absorbed into a 5% TEA solution.
4. Results and Discussion The change in NO2- and NO3- concentrations after absorption of NO into 30 wt % K2CO3 solution is shown in Table 2 and Figure 3. The presence of substantial quantities of NO2- and NO3- in the solution proves that NOx can be absorbed chemically through eq 14. The total absorption does not appear to be strongly temperature-dependent, although trends in the concentration with temperature may be masked by experimental error given the relatively small quantities of both ions present after absorption from 0.19 mol % NO. Results are comparable for potassium carbonate solutions that are also loaded with carbon dioxide. The molar ratio of NO2-/NO3- is greater than 2.5 in all cases (see Table 2 and Figure 4), which leads to the conclusion that eqs 13 and 14 occurred simultaneously. The ratio of NO2-/NO3- would be approximately 1 if only eq 14 occurred and would be infinite if only eq 13 occurred. Yan et al.31 also reported that weak acid salts containing sodium and potassium can react with N2O3 (NO þ NO2) to form NO2- and with NO2 to form NO2- and NO3-, as in eqs 13 and 14. As shown in Figure 4, the ratio of NO2-/NO3- decreases with an increasing temperature, indicating a shift toward eq 14 as the temperature increases. Use of the published Henry’s law constants for NO2 and NO illustrates that chemical sorption of nitrous oxides clearly dominates over physical absorption (see Table 2).
(24) Makoto, N.; Toshiyuki, H.; Eigo, K.; Takeshi, M.; Makoto, S. J. Chromatogr., A 1996, 739 (1-2), 301–306. (25) Tuazon, E. C.; Winer, A. M.; Graham, R. A.; Pitts, J. J. N. Adv. Environ. Sci. Technol. 1980, 10, 259–299. (26) Ramis, G. Appl. Catal. 1990, 64, 243. (27) Khorassani, H. E.; Theraulaz, F.; Thomas, O. Acta Hydrochim. Hydrobiol. 1998, 26 (5), 296–299. (28) Simon, P. K.; Dasgupta, P. K. Anal. Chem. 1993, 65 (9), 1134– 1139. (29) Oms, M. T.; Jongejan, P. A. C.; Veltkamp, A. C.; Wyers, G. P. Int. J. Environ. Anal. Chem. 1996, 62, 207–218. (30) Acker, K.; Moller, D.; Wieprecht, W.; Auel, R.; Kalax, D.; Tscherwenka, W. Air Soil Pollut. 2001, 130, 331–336.
(31) Yan, Y.; Wei, E. Low Temp. Specialty Gases 2000, 18 (4), 24–30. (32) Lee, Y. N.; Schwartz, S. E. J. Phys. Chem. 1981, 85, 840–848. (33) Berdnikov, V. M.; Bazhin, N. M. J. Phys. Chem., Engl. Transl. 1970, 44, 395–398. (34) Schwartz, S. E.; White, W. H. Advances in Environmental Science and Engineering; Gordon and Breach Science Publishers: New York, 1981; Vol. 4, pp 1-45.
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Figure 3. Total nitrogen (NOx) absorbed into 30 wt % potassium carbonate solution.
Figure 5. Change in the absorbed NOx concentration during the equilibrium process after absorption of 0.19 mol % NO for 12 h at 15 °C.
Figure 6. Impact of O2 on NO2- and NO3- concentrations in solution. In these experiments, pure oxygen was passed through a solution of 250 mL at 14 mL/min for 4 h.
NO2-/NO3-
Figure 4. Molar ratio of after NOx absorption into 30 wt % potassium carbonate solution without any CO2 absorption.
equilibration for 2 h at 100 °C when N2 is passed through the two equilibrium vessels. This is also true when CO2 is loaded into the solution. GC analysis of the gas phase captured from the rig after equilibration also showed no evidence of NO in this phase. The very small quantity desorbed is consistent with physical desorption alone. Use of Henry’s law constants available for NO2 and NO, as presented in Table 2, suggests that desorption of the gas physically absorbed at 16 °C would release a total of 0.25 mmol/L. Thus, the chemical reactions between NOx and K2CO3 are essentially irreversible. When O2 is passed through the solution for 4 h at different temperatures, the NO2- and NO3- concentrations remain almost the same (Figure 6). Therefore, NO2- is not readily oxidized to NO3- in the system, implying that both species would persist in solution even under the oxidizing conditions typical of flue gas capture.
For the experiments conducted at 0.19 mol % NO, the corresponding reduction in CO32- concentration from the initial value of 2770 mmol/L (or 30 wt % K2CO3) is too small to be reliably determined. However, the experiments that used pure NO generated by the reaction (2.3 mol % NO) show a substantial reduction in CO32-, and a measurable concentration of HCO3- is generated (Table 2). This is the direct result of the CO2 generated in eqs 13 and 14 being consumed by eq 15. The implication of this for carbon dioxide capture is a reduction in the absorption capacity for CO2. When the NOx-loaded solution prepared by absorption with 2.3 mol % NO at 16 °C is heated to 100 °C in the VLE rig to regenerate the solvent, a small amount of NO2 is desorbed. This is evidenced by the presence of NO2- and NO3- in the 5% TEA solution downstream of the VLE rig. However, analysis of this TEA solution by IC suggests that the loss of NOx from the solution is small. A total of 0.09 mmol is recovered from 70 mL of TEA solution, which corresponds to the total concentration of NOx in the equilibrated solution, falling by only 0.35 mmol/L or 0.2 mol %. This is consistent with the observed change in concentration in the aqueous phase. Figure 5 shows that NO2- and NO3- absorbed at 15 °C are almost totally retained after
5. Conclusions This work has shown that NOx will readily dissolve into K2CO3 solution via both chemical and physical absorptions. However, chemical absorption is dominant, representing at least 90% of the total absorption. The chemical reactions 4772
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between NOx in the flue gas and K2CO3 are well-described by eqs 13 and 14. These reactions have also been shown to be irreversible under the conditions used in this study. This is an important finding because it implies that NO2- and NO3would accumulate in the K2CO3 solvent during flue gas absorption processes used for CO2 capture. The NO2- anion is not readily oxidized to the NO3- species, implying that both species would persist in solution. Furthermore, CO2 is generated by the chemical absorption of NOx, and this leads to a
reduction in the available dioxide from flue gas.
CO32-
for the reaction with carbon
Acknowledgment. The authors acknowledge the Cooperative Research Centre for Greenhouse Gas Technologies (CO2CRC) and the Particulate Fluid Processing ARC Special Research Centre (PFPC) for financial support for this work and Peter Carpenter (RMIT University) for providing use of the ion chromatograph.
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