Carbon Dioxide Postcombustion Capture - American Chemical Society

Jul 17, 2009 - ROBERT BURNS, §. MARCEL MAEDER, §. AND. MOETAZ ATTALLA †. CSIRO Energy Technology, PO Box 330, Newcastle NSW 2300,...
0 downloads 0 Views 983KB Size
Download PDF
Environ. Sci. Technol. 2009, 43, 6427–6433

Carbon Dioxide Postcombustion Capture: A Novel Screening Study of the Carbon Dioxide Absorption Performance of 76 Amines G R A E M E P U X T Y , * ,† R O B E R T R O W L A N D , † ANDREW ALLPORT,† QI YANG,‡ MARK BOWN,‡ ROBERT BURNS,§ MARCEL MAEDER,§ AND MOETAZ ATTALLA† CSIRO Energy Technology, PO Box 330, Newcastle NSW 2300, Australia, CSIRO Molecular and Health Technologies, Private Bag 10, Clayton VIC 3169, Australia, and Discipline of Chemistry, The University of Newcastle, University Drive, Callaghan NSW 2308, Australia

Received May 7, 2009. Accepted June 24, 2009.

The significant and rapid reduction of greenhouse gas emissions is recognized as necessary to mitigate the potential climate effects from global warming. The postcombustion capture (PCC) and storage of carbon dioxide (CO2) produced from the use of fossil fuels for electricity generation is a key technology needed to achieve these reductions. The most mature technology for CO2 capture is reversible chemical absorption into an aqueous amine solution. In this study the results from measurements of the CO2 absorption capacity of aqueous amine solutions for 76 different amines are presented. Measurements were made using both a novel isothermal gravimetric analysis (IGA) method and a traditional absorption apparatus. Seven amines, consisting of one primary, three secondary, and three tertiary amines, were identified as exhibiting outstanding absorption capacities. Most have a number of structural features in common including steric hindrance and hydroxyl functionality 2 or 3 carbons from the nitrogen. Initial CO2 absorption rate data from the IGA measurements was also used to indicate relative absorption rates. Most of the outstanding performers in terms of capacity also showed initial absorption rates comparable to the industry standard monoethanolamine (MEA). This indicates, in terms of both absorption capacity and kinetics, that they are promising candidates for further investigation.

Introduction The capture, reversible release, and storage of carbon dioxide (CO2) from combustion flue gases (postcombustion capture, PCC) is now recognized by government and industry as a viable near-term option for greenhouse gas abatement (1, 2). It is particularly relevant to electricity generation from fossil fuels (coal, oil, and gas) which accounts for approximately 25% of global CO2 emissions (3) with this figure set to increase drastically in the next 25 years (4). PCC has a number of practical advantages over other methods such as oxy-firing * Corresponding author e-mail: [email protected]. † CSIRO, Division of Energy Technology. ‡ CSIRO, Division of Molecular and Health Technologies. § The University of Newcastle. 10.1021/es901376a CCC: $40.75

Published on Web 07/17/2009

 2009 American Chemical Society

to produce a pure CO2 stream and IGCC with precombustion capture and is of similar economic cost (5). In particular, PCC can be retro-fitted to existing power stations and integrated into new ones. Additionally, the parasitic energy demand of a PCC plant on a power station can be reduced (at the cost of CO2 removal efficiency) according to electricity demand if additional electricity output is required from a power station during times of peak load or optimal electricity pricing. Mature technology to separate CO2 from H2 or CH4 and release it as a pure gas stream already exists in the gas processing and ammonia production industries (6). Traditionally, the CO2 is vented or used for enhanced oil recovery, food manufacture, and chemical production (6). The most mature and applied technology for the capture and release of CO2 is cyclic chemical absorption/desorption using an aqueous amine solution of monoethanolamine (MEA) 30% w/w. Aqueous solutions of other amines that have better performance characteristics than MEA are also available as proprietary formulations from a number of commercial suppliers (e.g., Fluor, BASF and MHI). The standard chemical process for PCC by chemical absorption/desorption is shown in Figure 1. The gas (entering from the bottom) and amine solution (entering from the top) are contacted at atmospheric pressure in a counter-current fashion inside a packed absorber column at low temperature (∼40 °C). The CO2 rich amine solution is then circulated to the top of a desorption column which is at high temperature (∼120 °C). At this temperature the chemical equilibria and CO2 solubility are shifted such that the absorbed CO2 is released and exits at the top of the column. The regenerated amine solution is then circulated back to the absorber column. The application of PCC to combustion flue gases from electricity generation poses a number of technical challenges. The main issue is the energy requirement of the process, which using current industry standard technology is expected to reduce the efficiency of a coal fired electricity plant by ∼21% (5). The main energy cost is the heat and steam required for the desorption column and to pump the absorbing solution around the system. By increasing the capacity of a chemical absorbent in terms of the amount of CO2 that can be absorbed and desorbed per unit mass of solution, this energy requirement may be more than halved (7). In this work 76 different amine moeties have been screened for their ability to absorb CO2. These include primary, secondary, and tertiary amines; alkanolamines; polyamines of a mixed or single type; cyclic and aromatic amines; amino acids; and sterically free and hindered amines. The screening process involved an initial screening method using isothermal gravimetric analysis (IGA) to measure CO2 absorption as a mass increase. This method required only milligrams of material for screening, allowing testing of amines that were prohibitively expensive or difficult to obtain in large quantities. The initial rate of CO2 absorption was also used as an indicator of the kinetic behavior. Many of the amines that were identified as performing well using this method were also tested with a more traditional larger scale CO2 absorption apparatus to validate the initial screening results and provide a more reliable estimate of the CO2 absorption capacity. The measured CO2 absorption capacity was compared to the absorption capacity from model calculations. This allowed identification of a number of amines showing absorption capacities greater than current understanding of the CO2 absorption mechanism would predict. The results of the screening study and modeling VOL. 43, NO. 16, 2009 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

6427

FIGURE 1. The chemical absorption and desorption process for cyclic CO2 capture and release. were then used to identify the chemical properties of the amines most relevant to their performance as chemical absorbents.

Experimental Section Micro-Scale Isothermal Gravimetric Analysis (IGA). A Setaram TG-DTA/DSC thermal gravimetric analyzer was used in isothermal mode at 40 °C to monitor the increase in mass of an aqueous amine solution when exposed to an atmosphere of 15% CO2 and 85% N2 by volume at ambient pressure. This CO2 concentration was chosen as it closely resembles that of a coal fired power station flue gas. The instrument was setup with two gas stream inputs, one providing a mix of 15% CO2 and 85% N2 (99.9% purity, BOC Australia) and the other N2 (>99.99% purity, BOC Australia). Both flows were controlled using mass flow controllers (Bronkhurst HighTech El-Flow) with a total gas flow of 30 mL.min-1 used for all experiments. Two separate IGA experiments were performed in order to determine the total CO2 uptake of the amine test solution. The first experimental run determined the mass loss due to evaporation and the second determined the mass increase of the test solution when exposed to CO2 over the same length of time. Each experiment was performed on a fresh 100 µL aliquot of the test solution in a 100 µL alumina crucible (Setaram). The samples were allowed to reach thermal equilibrium under N2 for 20 min. Following this, the experiment was continued under N2 to measure evaporative loss or the gas flow changed to the CO2-N2 mix to measure absorption for a period of six hours. The test solutions were made to an amine concentration of 30% w/w in deionized water, unless solubility constraints only allowed lower concentration. All chemicals were purchased from SigmaAldrich and used without further purification. The purity of all the amines used was in the range 95-99.5% with the highest purity available always chosen. For calculations of absorption capacities the mass at time t from the evaporation run is subtracted from the mass at time t of the absorption run. This is illustrated in the Supporting Information (SI). Large figures showing representative experimental curves are also given in the SI. The CO2 absorption capacity is then determined from the 6428

9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 43, NO. 16, 2009

maximum point on the absorption curve. Total mass loss due to water evaporation at the point of maximum CO2 absorption was generally 20%, with the rate of evaporation unaffected by CO2 absorption. Macro-Scale CO2 Absorption. Many of the amines that showed promise from the IGA experiments were also tested in a larger scale absorption apparatus based on that used by Ma’Mun et al. (8). In this apparatus a mixture of 13% CO2 (99.5% purity, BOC Australia) and 87% N2 (>99.99% purity, BOC Australia) was delivered using mass-flow controllers (Bronkhurst High-Tech El-Flow). The gas was passed through a mixing chamber, humidified to maintain the water balance and then sparged via a sintered glass frit through a 30% w/w amine solution in a glass reactor vessel at ambient pressure. Gas flow rates and solution volumes of 1.7 mL.min-1 and 300 mL or 1.0 mL.min-1 and 20 mL, respectively, were used depending on the quantity of amine available. The CO2 content of the gas outflow was measured using a Horiba VS-3001 general purpose gas sampling unit and Horiba VA3000 NDIR multigas analyzer. The difference between the inflow and outflow CO2 concentration was used to determine the amount of CO2 absorbed. The humidifier and amine solution were thermostatted to 40 °C by immersion in a temperature controlled water bath (Techne). A schematic of this apparatus is shown in the SI. The experiments were run by initially allowing the humidifier and reactor containing the amine solution to thermally equilibrate. During this period the system was flushed with N2 and calibration of the CO2 analyzer was completed. The appropriate gas mix was first established with the gas flow passing through the saturator but bypassing the reactor and passing directly to the CO2 analyzer. The gas flow was then switched to pass through the reactor to begin the experiment. Each experiment was run until the measured CO2 concentration in the outflow returned to the original value (typical runs lasted between 30 min and 3 h).

Results and Discussion The definition of CO2 absorption capacity often used is the moles of CO2 absorbed per mole of amine in solution (molar capacity, nCO2/namine). As many of the amines tested were

FIGURE 2. Plot of CO2 absorption capacity from the micro-scale and macro-scale methods for 10 amines (two fall on the same point) showing the level of agreement between the two methods. The correlation coefficient between the data sets (G) is 0.89. polyamines, and each amine functional group can contribute, the molar absorption capacity can be further normalized to the number of amine functional groups, or N atoms, present in each molecule (normalized molar capacity, nCO2/nN). In this manuscript nCO2/nN is used when discussing absorption capacity to allow a fair comparison of absorption capacity between poly and mono amines. Of the 76 amines tested using the microscale IGA method, 24 were then tested using the larger macro-scale apparatus. These included amines that gave unreliable results using the IGA method due to excessive evaporation or precipitation, and validation of the values for amines found to have high absorption capacities. The results from the IGA method were validated against literature data and results from the macroscale apparatus. Figure 2 shows a plot of the CO2 absorption capacity determined reliably via both methods for 10 amines. If results from both methods were available the macro-scale results were taken as being more reliable as this method is less influenced by evaporation or precipitation. From five replicate runs using MEA, both methods achieved an error of approximately 5% in terms of reproducibility. Previous modeling results and experimental work (7-14) has indicated that the capacity of an aqueous amine solution to chemically absorb CO2 is a function of the pathway via which CO2 reacts and the basicity of the amine, which can be expressed using its pKa. Generally, primary and secondary amines (represented as R1R2NH) can react with dissolved CO2 to form a carbamic acid (R1R2NCOOH). Depending upon its acidity, it may then give up a proton to a second amine molecule forming a carbamate (R1R2NCOO-) according to an overall stoichiometry of 2 as shown in eq 1. Via this pathway two moles of amine are consumed per mole of CO2 if the carbamic acid is acidic, which is generally assumed to be the case. Kinetically and thermodynamically this reaction pathway is generally favored for primary and secondary amines (14). CO2 + R1R2NH h R1R2NCOOH R1R2NCOO- + H+ h R1R2NCOOH R1R2NH + H+ h R1R2NH+ 2

(1)

A second reaction pathway that also contributes to CO2 absorption is CO2 hydration to form bicarbonate. In this pathway an amine molecule (represented as R1R2R3N) simply

acts as a proton accepting base, and possibly a catalyst, for the hydration of CO2 (11, 15). The overall stoichiometry for this second pathway is given by eq 2. Via this pathway one mole of amine is consumed per mole of CO2, so in terms of capacity it is more efficient. For tertiary and some sterically hindered primary and secondary amines this is the only pathway contributing to absorption. However, this pathway is generally less favorable kinetically than carbamate formation (14). + CO2 + H2O h HCO3 + H

(2)

R1R2R3N + H+ h R1R2R3NH+

Assuming that either or both pathways can contribute to absorption for a given amine, the molar absorption capacity should vary between 0.5 and 1.0 if the reaction is driven to the products (which is the case if CO2(g) is constantly dosed). Both pathways depend on the availability of unprotonated base to accept a proton. Clearly the basic strength or pKa of a given amine will influence how far both of these reaction pathways can proceed to the products. Amines that are weak bases may not achieve the 0.5-1.0 capacity range. Figure 3 is a plot of the CO2 absorption capacity versus the amine group pKa. For polyamines it is the pKa of the most basic amine group. The pKa values were either taken from literature if available (16) or calculated using the ACD/pKa 8.3 software package (17). The amines have been divided into four classes: primary amines; secondary amines; tertiary amines; and polyamines of mixed type (e.g., primary and secondary amine functionality in the one molecule). Figure 3 also shows model calculations of the capacity as a function of pKa. This modeling was done by simulating all of the relevant elementary steps as equilibria according to a previously published method (7). The dotted line represents the predicted capacity if only the pathway of eq 1 occurs with a large stability constant for carbamic acid formation (K ) 0.1 L mol-1) and a small carbamic acid pKa (pKa ) 4). The dashed line is the predicted capacity if only the pathway of eq 2 occurs. The area between the two lines represents where amines in which both pathways contribute would fall. The primary and secondary amines in general do not appear to show a strong correlation with pKa. This is not surprising as the sensitivity of carbamate formation to pH, and thus amine pKa, is dependent on the carbamate stability constant which varies from amine to amine. The tertiary amines do VOL. 43, NO. 16, 2009 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

6429

FIGURE 3. Plot of the absorption capacity versus amine group pKa (most basic amine for polyamines) at 40 °C. MEA is labeled as a black dot. The seven amines showing outstanding performance are also labeled. The dashed line represents model predictions of absorption capacity as a function of pKa if only the pathway of eq 1 occurs (with a large stability constant and small pKa for carbamaic acid) and the dotted line if only the pathway of eq 2 occurs. show a strong dependence on pKa consistent with bicarbonate formation being the dominant reaction pathway for CO2 absorption. This is because the CO2 hydration reaction is independent of the amine but is strongly pH dependent due to the small stability constant for bicarbonate formation. The mixed amines all contain primary or secondary functionality and also show little correlation with pKa. Capacity, Poor Performers. A significant number of the amines fall below the area bounded by the lines of predicted capacity. From comparing the structures of these amines two structural features are apparent. Either the amines contain a carboxylic acid group or are a polyamine. The effect of both these features on absorption capacity can be easily rationalized. An acidic functional group such as a carboxylic acid will act to protonate a basic functional group such as an amine on dissolution, reducing the ability of the amine to participate in reactions with CO2. For the case of polyamines, those that fall below the lines have at least one amine functional group with a pKa less than 6. These amine functionalities with low pKa are weak bases and are effectively spectators that do not contribute to the overall absorption capacity. Capacity, Outstanding Performers. Also apparent in Figure 3 are seven amines that are significantly above (more than one absorption capacity unit) the area enclosed by the lines. They are achieving capacities greater than would be predicted based on their pKa. These seven amines are made up of one primary, three secondary and three tertiary monoalkanolamines and their structures, absorption capacities, and pKas are given in Figure 4. Three of the four primary and secondary amines have some form of steric hindrance around the amine functionality according to the definition of Satori et al. (18) (2-amino-2-methyl-1,2-propanediol, 2-piperidineethanol and 2-piperidinemethanol). This is a feature recognized as increasing capacity while maintaining a fast reaction rate. Satori et al. (18) propose that this large capacity and fast rate occurs through formation of an unstable carbamate intermediate that hydrolyses to bicarbonate and protonated amine (overall the same pathway as eq 2). Two of the secondary and two of the tertiary amines are heterocycles. Hetercycles such as piperidine and piperazine are known to have fast reaction rates toward carbamate formation (19-23). Being secondary amines, they may also 6430

9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 43, NO. 16, 2009

be good base catalysts of the CO2 hydration reaction (15), which may help improve their capacity. The most interesting result is that all of these amines share a common structural feature, a hydroxyl group within 2 or 3 carbons of the amine functionality. While it is unclear what the role of this structural feature is, the distance of the hydroxyl functionality from the amine and the structural features around it appears crucial. The structures of three structurally similar series and their capacities highlighting this are given in Figure 5. For example, 2-piperidineethanol and 2-piperidinemethanol achieved capacities of ∼1, whereas 3-piperidinemethanol only achieved a capacity of 0.8. This indicates that the proximity of the hydroxyl group and its freedom to move are important. One possibility is that a hydroxyl group the appropriate distance from the amine functionality, and with the appropriate structural features surrounding it, is able to form a stable intramolecular hydrogen bond with the nitrogen to form a five or six member ring structure. Intramolecular hydrogen bond formation between amine and hydroxyl groups has been predicted by simulation (24, 25) and shown experimentally (24, 26, 27) and was previously suggested by Sharma (23). While this may decrease the amine pKa, for primary and secondary amines it may also destabilize carbamate formation and push the absorption toward the more stoichiometrically efficient hydration pathway of eq 2. The type of intramolecular hydrogen bond proposed is illustrated in Figure 6. A second possible explanation for primary and secondary amines is that the carbamic acid formed is a weak acid. It is generally assumed that the pKa of the MEA carbamic acid is meso-2,3-diaminosuccinic acid > piperazine > 1,4,7triazacyclononane > spermine > trans-N-hydroxyethyl1,4-diaminocyclohexane). Such amines would be expected to exhibit faster absorption rates due to the simple fact that statistically the per mole likelihood of CO2 encoun-

tering amine functionality is increased. It is worth noting that all of the primary and secondary amines identified as having an outstanding capacity also had initial absorption rates no more than 25% slower than MEA. This means kinetic performance is not likely to exclude them from potential CO2 capture applications. Also, it has been shown that the rate of CO2 absorption can be enhanced and high absorption capacity maintained by adding at low concentration an amine that exhibits a fast absorption rate to an amine solution that exhibits high absorption capacity (11). Of the 76 amines tested seven were found to have an outstanding CO2 absorption capacity compared to modeling predictions. Of the four primary and secondary amines showing outstanding absorption capacity, all showed initial absorption rates similar to MEA. To further evaluate these amines as new candidate molecules for large scale CO2 capture significant further testing is required. The particular information required is detailed information on CO2 absorption rate as a function of temperature, amine concentration and CO2 loading; the energy requirement of and their capacity to cyclically capture and release CO2; their resistance to oxidative and thermal degradation; their corrosiveness; their resistance to degradation by flue gas impurities (SOx, NOx, and trace elements); and their toxicity and the toxicity of degradation products. The results from this work indicate that there is still significant scope for improvement in the use of aqueous

FIGURE 7. The IGA microscale curve showing capacity as a function of time for MEA 30% w/w. The black dashed line shows the linear region used to calculate the initial rate of CO2 absorption.

FIGURE 8. Plot showing the initial absorption rate versus amine group pKa calculated from the IGA microscale data. 6432

9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 43, NO. 16, 2009

amine solutions for CO2 capture by chemical absorption. However, what is also clear is that there is a lack of understanding about the chemistry involved. Understanding how the amines that show outstanding CO2 absorption capacities do so while maintaining good rates of CO2 absorption is fundamental to achieving an optimal formulation that maximizes efficiency and minimizes cost and sustainability for PCC on an industrial scale. Existing understanding of the reaction pathways are unable to account for these characteristics.

Supporting Information Available Figures of the experimental apparatus, figures showing representative results from the microscale method, a table containing all amines tested and their associated results and figures of absorption capacity and initial rate versus pKa with numbering according to the tables. This material is available free of charge via the Internet at http:// pubs.acs.org.

(14)

(15) (16) (17) (18) (19) (20) (21)

Literature Cited (1) IEA. Capturing CO2; IEA Greenhouse Gas R&D Program: Cheltenham, UK, 2007; p 23. (2) IPCC. IPCC Special Report on Carbon Dioxide Capture and Storage. Prepared by Working Group III of the Intergovernmental Panel on Climate Change; Cambridge University Press: Cambridge, 2005; p 440. (3) WRI Climate Analysis Indicators Tool. http://www.wri.org/. (4) CSIRO. The Heat Is On: The Future of Energy in Australia; Commonwealth Scientific and Industrial Research Organisation: Canberra, 2006. (5) Davison, J. Performance and costs of power plants with capture and storage of CO2. Energy 2007, 32, 1163–1176. (6) IEA. Prospects for CO2 Capture and Storage; OECD/IEA: Paris, 2004. (7) McCann, N.; Maeder, M.; Attalla, M. Simulation of enthalpy and capacity of CO2 absorption by aqueous amine systems. Ind. Eng. Chem. Res. 2008, 47 (6), 2002–2009. (8) Ma’mun, S.; Svendsen, H. F.; Hoff, K. A.; Juliussen, O. Selection of new absorbents for carbon dioxide capture. Energy Convers. Manage. 2007, 48, 251–258. (9) Blauwhoff, P. M. M.; Versteeg, G. F.; Swaaij, W. P. M. V. A study on the reactions between CO2 and alkanolamines in aqueous solution. Chem. Eng. Sci. 1983, 38 (9), 1411–1429. (10) Freguia, S.; Rochelle, G. Modelling of CO2 capture by aqueous ethanolamine. Seperations 200349 (7), 1676–1686. (11) Savage, D. W.; Satori, G.; Astarita, G. Amines as rate promoters for carbon dioxide hydrolysis. Faraday Discuss. Chem. Soc. 1984, 77, 17–31. (12) Tobiesen, F. A.; Svendsen, H. F. Experimental validation of a rigorous absorber model for CO2 postcombustion capture. AIChE J. 2007, 53 (4), 846–865. (13) Tomizaki, K.-y.; Shimizu, S.; Onoda, M.; Fujioka, Y. An acid dissociation constant (pKa)-based screening of chemical absorbents that preferably capture and release pressurized carbon

(22) (23)

(24) (25)

(26)

(27)

(28) (29)

(30)

dioxide for greenhouse gas control. Chem. Lett. 2008, 37 (5), 516–517. Versteeg, G. F.; Dijck, L. A. J. V.; Swaaij, W. P. M. V. On the kinetics between CO2 and alkanolamines both in aqueous and non-aqueous solutions. An overview. Chem. Eng. Commun. 1996, 144, 113–158. Donaldson, T. L.; Nguyen, Y. N. Carbon dioxide reaction kinetics and transport in aqueous amine membranes. Ind. Eng. Chem. Fundam. 1980, 19, 260–266. NIST. NIST Critically Selected Stability Constants of Metal Complexes; The National Institute of Standards and Technology: Gaithersburg, MD, 2001. ACD/I-Lab ACD/pKa 8.3. http://ilab.acdlabs.com/. Satori, G.; Savage, D. W. Sterically hindered amines for CO2 removal from gases. Ind. Eng. Chem. Fundam. 1983, 22, 239– 249. Bishnoi, S.; Rochelle, G. T. Absorption of carbon dioxide into aqueous piperazine: Reaction kinetics, mass transfer and solubility. Chem. Eng. Sci. 2000, 55, 5531–5543. Cullinane, J. T.; Rochelle, G. Kinetics of carbon dioxide absorption into aqueous potassium carbonate and piperazine. Ind. Eng. Chem. Res. 2006, 45, 2531–2545. Derks, P. W. J.; Kleingeld, T.; Aken, C. v.; Hogendoorn, J. A.; Versteeg, G. F. Kinetics of absorption of carbon dioxide in aqueous piperazine solutions. Chem. Eng. Sci. 2006, 61, 6837– 6854. Samanta, A.; Bandyopadhyay, S. S. Kinetics and modeling of carbon dioxide absorption into aqueous solutions of piperazine. Chem. Eng. Sci. 2007, 62, 7312–7319. Sharma, M. M. Kinetics of reactions of carbonyl sulphide and carbon dioxide with amines and catalysis by bronsted bases of hydrolysis of COS. Trans. Faraday Soc. 1965, 61 (508P), 681– 688. Cacela, C.; Fausto, R.; Duarte, M. L. A combined matrix-isolation infrared spectroscopy and MO study of 1-amino-2-propanol. Vib. Spectrosc. 2001, 26 (1), 113–131. Goldblum, A.; Deeb, O.; Leow, G. H. Semiempirical Mndo/H calculations of opiates part 1. Building blocks: Conformations of piperidine derivatives and the effect of hydrogen bonding. J. Mol. Struct. (Theochem) 1990, 207, 1–14. Przeslawska, M.; Melikowa, S. M.; Lipkowski, P.; Koll, A. Gas phase FT-IR spectra and structure of aminoalcohols with intramolecular hydrogen bonds. Vib. Spectrosc. 1999, 20 (1), 69–83. Ribeiro da Silva, M. A. V.; Cabral, J. I. T. A. Standard molar enthalpies of formation of 1-methyl-2-piperidinemethanol, 1-piperidineethanol, and 2-piperidineethanol. J. Chem. Thermodyn. 2006, 38, 1461–1466. Danckwerts, P. V. The reaction of CO2 with ethanolamine. Chem. Eng. Sci. 1979, 34, 443–446. Christensson, F.; Koefoed, H. C. S.; Peterson, A. C.; Rasmussen, K. Equilibrium constants in the ammonium carbonate-carbamate system. The acid dissociation constant of carbamic acid. Acta Chem. Scand. A 1978, 32, 15–17. McCann, N.; Phan, D.; Wang, X.; Conway, W.; Burns, R.; Attalla, M.; Puxty, G.; Maeder, M. Kinetics and mechanism of carbamate formation from CO2(aq), carbonate species, and monoethanolamine in aqueous solution. J. Phys. Chem. A 2009, 113, 5022–5029.

ES901376A

VOL. 43, NO. 16, 2009 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

6433