Infrared Spectroscopic Study of Reaction of Carbon Dioxide with

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Infrared Spectroscopic Study of Reaction of Carbon Dioxide with Aqueous Monoethanolamine Solutions Chenhu Sun and Prabir K. Dutta* Department of Chemistry and Biochemistry, The Ohio State University, Columbus, Ohio 43210, United States S Supporting Information *

ABSTRACT: Increasing levels of atmospheric CO2 emissions from fossil fuel combustion is altering global climate, and several strategies are being examined for CO2 capture and sequestration. Aqueous amine solutions, in particular, monoethanolamine (MEA) is extensively used for CO2 adsorption. Our objective in this study was to follow the reaction of CO2 with 15 wt % monoethanolamine (MEA) by horizontal attenuated total reflection Fourier transform infrared spectroscopy (HATR-FTIR), with particular focus on the evolution of MEA, protonated MEA (MEAH+), MEA carbamate, and carbonate/ bicarbonate species. Since the path length in HATR-FTIR is well-defined, by use of standards and Beer’s law, the molar absorptivity was calculated. This allowed us to monitor the concentrations of MEA and MEAH+ quantitatively. Using the characteristic infrared bands of the various amine-derived species, influence of SO2 (151 ppm) on CO2 absorption by MEA was studied, including the cycling process of adsorption (40 °C) and desorption (100 °C). No effect of SO2 was noted on CO2 uptake by MEA. With blends of piperazine and MEA, the CO2 loading increased, and evolution of vibrational bands due to MEA species was slower in the presence of piperazine. CO2 + OH− ⇌ HCO3−

1. INTRODUCTION Flue gas from the combustion of fossil fuel in power plants is a significant source of CO2 emission and was 37% of total CO2 emissions in the United States for 2013.1 Because of the dominant role CO2 plays in global warming, there is considerable ongoing research in CO2 capture technologies from point source emissions.2 There are several competing technologies for CO2 capture including physical and chemical absorption, adsorption, and membrane processes.3 None of these technologies is mature enough to handle the large amounts of postcombustion flue gas from power plants. The most advanced technology for CO2 capture is chemical absorption using aqueous alkalnolamines, with monoethanolamine (MEA) being the most widely studied.4,5 Secondary and tertiary amines as well as blends of amines are also well studied.6 There are numerous investigations on the performance of MEA aqueous solutions as CO2 absorbents.7−11 Development of new species, their fate during CO2 absorption, and desorption by amines determines the efficacy of adsorption. For the reaction of MEA with CO2, the prominent mechanism involves the formation of a zwitterionic form of the carbamate followed by a slow proton exchange reaction with a base (denoted by B, e.g., MEA).3,6 CO2 + RNH 2 ⇌ RNH 2+CO2−

(1)

RNH 2+CO2− + B ⇌ RNHCO2− + BH+

(2)

Equilibrium constants and kinetics of these reactions are well studied.12 However, for speciation, spectroscopic methods such as nuclear magnetic resonance (NMR) spectroscopy, Raman spectroscopy, and infrared spectroscopy are more suitable.13−15 Identification of amine, protonated amine, carbamate, carbonate and bicarbonate species formed upon absorption of CO2 in aqueous solutions of MEA, as well as other amines, 2-amino2-methyl-1-propanol (AMP), methyldiethanolamine (MDEA), and heterocyclic amines, with IR spectroscopy have been reported.16−18 The degradation of alkalnolamines is also a critical issue. Besides thermal and oxidative degradation, chemical degradation of MEA by reaction with SO2 in flue gas stream is of concern. Typical levels of SO2 are in the range of 6−300 ppm after wet desulfurization.19 The overall reaction of the MEA− SO2−H2O system is SO2 + 2RNH 2 + H 2O → 2RNH3+ + SO32 −

(5)

This reaction results in formation of heat stable amine salts such as sulfites and sulfates, and thereby degrades the performance of the amine.20 Another area of interest is the use of amine blends. For example, the addition of piperazine (PZ) appears to be an Received: January 3, 2016 Revised: May 13, 2016 Accepted: May 14, 2016

In addition, CO2 can react with water: CO2 + H 2O ⇌ H 2CO3

(4)

(3) © XXXX American Chemical Society

A

DOI: 10.1021/acs.iecr.6b00017 Ind. Eng. Chem. Res. XXXX, XXX, XXX−XXX

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Industrial & Engineering Chemistry Research

Figure 1. Schematic of CO2 absorption apparatus.

thereby enhancing the rate of absorption. In this paper, we present an infrared spectroscopic study using the horizontal attenuated total reflectance (HATR) method for assessing CO2 absorption (15%) with 15 wt % MEA. There are three new aspects of this study. First, since the HATR technique has reproducible path lengths of propagation of IR radiation, quantitative studies were carried out using calibration standards and use of Beer’s law. Second, the effect of SO2 on the evolution of the vibrational bands of MEA species during reaction with CO2 and during adsorption/desorption cycles was examined. Third, we investigate how piperazine blended with MEA influences speciation with CO2 loading.

crystal has a trough plate configuration, with a large, recessed germanium (Ge) crystal to accommodate the liquid sample. Since ATR sampling depth depends on the wavelength of light, a software correction (provided by PerkinElmer) was used to get comparable intensities across the spectrum. If the peaks were isolated, the relative peak height was obtained by drawing a baseline tangent to both shoulders of the band of interest and then measuring the height of the maximum of the curve to the baseline. The relative peak area is integrated from the area between the curve and the linear baseline. The error bars in the figures represent standard deviations for a minimum of three measurements. For overlapping bands, curve deconvolution was used to resolve the bands. For quantitative analysis, calibration curves were set up with standard solutions of MEA and methyl carbamate (a model for MEA carbamate), and the absorptivities were obtained. For the quantification of protonated amine (MEAH+), the pH-dependent protonation equilibria with log K = 9.06 was used to obtain the absorptivity.15 These procedures are described in the Supplementary Section. The pH value of the solution during gas absorption was monitored by an Accumet AB15 pH meter (Fisher Scientific).

2. MATERIALS AND METHODS A schematic of the absorption reactor used to analyze CO2 reactivity with MEA is shown in Figure 1. Research grade O2, N2, CO2, and SO2 were obtained from Praxair (Columbus, USA). A gas stream with various compositions (the O2 is fixed at 6% in all experiments) was bubbled through an aqueous solution of MEA (ACS reagent, ≥ 99.0%, SigmaAldrich) in a three-neck flask maintained at different temperatures (40−100 °C) by a temperature-controlled water bath. The gas flow rate was maintained at 100 sccm. The gas outlet reactor was connected to a condenser for preventing solvent loss. Solution volumes of 50 mL were used for all experiments. Pure N2 was used to purge the reactor before the absorption experiments. A 0.5 mL sample of the aqueous solution was withdrawn from the reactor at appropriate intervals and was analyzed by a PerkinElmer Spectrum 400 infrared spectrometer coupled with a 20-reflection PIKE HATR accessory. Each spectrum was recorded as the average of 256 scans over the spectral range of 4000−800 cm−1 with a resolution of 8 cm−1. The HATR

3. RESULTS 3.1. Reactivity of MEA with CO2. All studies were done with 15 wt % MEA with CO2 (15% CO2/6% O2) at 40 °C using the reactor shown in Figure 1. The reaction between CO2 and MEA is controlled by several factors, including diffusion, convection, and the chemical reaction.24 Gas molecules need to overcome the resistance of the gas film and move to the gas− liquid interface, and a similar situation exists for MEA through the liquid film. The mass transport determines the equilibration times, and can be long, even if the chemical reaction between CO2 and MEA is fast. In industrial absorption processes, reactor design including packed bed, spray column, and bubble column are used to enhance gas−liquid contact.24 Another strategy to enhance mass transfer involves using a rotating packed bed.24 The reactor shown in Figure 1 is not optimized to reduce mass transport. There are reports in the literature of similar designs, especially for spectroscopic measurements, and the equilibration times are of the order of hours, comparable to our observations.11,25−28 Since similar conditions are main-

effective promoter for CO2 absorption with monoethanolamine (MEA), methyldiethanolamine (MDEA), and potassium carbonate (K2CO3) solutions.21,22 Piperazine, because of its low aqueous solubility is typically used as a “catalyst” to accelerate absorption. In the CO2−MDEA−PZ system,23 CO2 was reported to be transferred via PZ-bound CO2 to MDEA (denoted by R3N) by the reaction PZ−CO2 + R3N ⇌ PZ + R3N−CO2

(6)

B


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Figure 2. (a) Evolution of infrared spectra of 15 wt % MEA with 15% CO2/6% O2 at 40 °C. (b) Curve deconvolution of the 1280−1420 cm−1 region for the MEA sample with CO2 bubbling for 5 h into peaks of carbonate at 1386 cm−1, bicarbonate at 1360 cm−1, and carbamate at 1322 cm−1 (c) Change in IR intensity of MEA, MEAH+, and MEA carbamate of 15 wt % MEA at 40 °C upon reaction with 15% CO2/6% O2 as a function of time. (d) Concentration profiles of MEA and MEAH+ for 15 wt % MEA upon reaction with 15% CO2/6% O2 at 40 °C as a function of time.

1568, 1486, and 1322 cm−1 are observed and assigned to COO− asymmetric and symmetric stretching, and N−COO− stretching vibration of the carbamate species, respectively.16,18,31,32 In addition, there is also protonation of MEA (MEAH+), resulting in peaks shifts, in particular, C−N and C− OH stretching modes shift from 1075 to 1066 cm−1 and 1025 to 1017 cm−1, respectively.15 The 1066 cm−1 overlaps with the MEA band at 1076 cm−1, so curve deconvolution as shown in Figure S1 was used to estimate the intensity of the 1066 cm−1 band. The weak band at 1386 cm−1 is assigned to the doubly degenerate stretching mode of carbonate.33 The carbonate band at 1386 cm−1 decreases with time, and a band develops at 1362 cm−1, which is assigned to the C−O stretching of bicarbonate.33,34 Previous studies have noted the interference due to overlapping bands in this spectral region.15 The carbonate/bicarbonate split is clearly seen in Figure 2b from the curve deconvolution of the 1280−1400 cm−1 region (sample in which CO2 was bubbled for 5 h). The rate constants for CO2 (aq) reaction with OH− and MEA are 1.24 × 104 M−1 s−1,25 and 6.11 × 103 M−1 s−1.35 The comparable reactivity results in CO2 reacting with the amine preferentially because of its higher concentration. Yet the fact that we observe

tained for all reactions in the present study, the reactivity of MEA (always 15 wt %) for reaction with CO2, SO2 can be compared, as well as the influence of PZ. This information regarding reactivity should be transferable to any reactor design. Figure 2a shows the evolution of the infrared spectra. The peak assignments are listed in Table 1. MEA prior to CO2 exposure exhibited vibrational bands at 955 cm−1 (C−N−H out-of-plane wagging and C−NH2 twisting) and 1362 cm−1 (NH2 twisting).29,30 In the presence of CO2, new bands at Table 1. Infrared Peak Assignments for CO2-Loaded MEA Solution.16,18,23,32 wavenumber (cm−1) 1568 1517 1486 1386 1362 1322 1066 955

assignment −

COO (carbamate) NH3+ (protonated MEA) COO− (carbamate) C−O (carbonate) NH2 (MEA) N−COO− (carbamate) C−N (protonated MEA, carbamate) C−NH2 (MEA) C


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Industrial & Engineering Chemistry Research the carbonate band indicates that the CO2 is reacting with the OH−, due to the higher pH with initially high amine concentrations. We discuss later the origin of the bicarbonate band. The relative peak height of the characteristic MEA, MEAH+, and carbamate versus time of CO2 reaction is shown in Figure 2c. The intensity of the MEA peak (1362 cm−1) dropped down to almost zero intensity after 2 h, while the band due to protonated MEA at 1066 cm−1 keeps increasing. The band at 1322 cm−1 due to the carbamate reaches a maximum at 3 h and starts to decrease and growth of protonated MEA also shows a smaller increase after this time. The focus of the quantitation was on three species, MEA, MEAH+, and MEA-carbamate. We did not consider the presence of carbamic acid, since previous studies have known absence of this species at measurable concentrations.36 The absorptivity was calculated using calibration standards for MEA and pH titration for MEAH+ (details in Table S1 and Figure S2), and found to be 0.0018 and 0.012 (units of M−1 distance−1, distance is not known, though it is constant in all experiments since the path length of the HATR is fixed), respectively. The calculated concentrations of MEA and MEAH+ are plotted in Figure 2d. The concentrations of MEA tends toward zero (not detectable in the IR) in 3 h, and MEAH+ tends to 1.7 M after 5 h. Previous studies have noted that the absorptivity for MEAcarbamate is lacking in the literature.14 MEA carbamate is not commercially available, nor can it be readily isolated in water from MEA−CO2 reaction products. Thus, for estimating the amount of MEA carbamate, we used methyl carbamate as the standard to calculate the absorptivity. The formulas of methyl carbamate and MEA carbamate are H 2 NCOOCH 3 , HOCH2CH2NHCOO−, with N-COO− stretching vibrations at 1375 and 1322 cm−1, respectively. Even though a good calibration curve was obtained for the 1375 cm−1 band of methyl carbamate with absorptivity of 0.0468 (Figure S2), the quantitation results for MEA−carbamate were not meaningful. For example using the absorptivity of 1375 cm−1 band of methyl carbamate and the intensity of the 1322 cm−1 band of MEA carbamate in Figure 2a, the calculated concentration of MEA carbamate at 1 h was found to be 0.3 M, which does not meet the mass balance of total MEA concentration (2.5 M) with both MEAH+ and free MEA calculated at 0.6 M. Methyl carbamate is therefore not an appropriate standard for MEA carbamate. Figure 3 shows the change in pH as a function of time as CO2 (15% CO2/6% O2) is bubbled through the MEA solution. The pH value for 15 wt % MEA drops from 11.7 to 8.5 over 6 h, consistent with previous observations.11 3.2. Influence of SO2. With 15 wt % MEA at 40 °C, a band at 936 cm−1 appears after 5 days with continuous bubbling of 151 ppm of SO2/6% O2, and a slight redshift of the peak at 1076 cm−1 indicates the formation of MEAH+, as shown in Figure S3. The band at 936 cm−1 is assigned to the sulfite ion, SO32−.37 The effect of SO2 on CO2 absorption by MEA over a shorter 6 h time scale by MEA was examined. A gas stream of 151 ppm of SO2 with 15% CO2 and 6% O2 was bubbled into the MEA solutions. The evolution of the relative peak height of MEA (1362 cm−1), MEAH+ (1066 cm−1), and carbamate (1322 cm−1) as a function of time is shown in Figure 4 for 15 wt % MEA (5 h) with and without SO2. There is no change in the infrared intensity with time indicating that the influence of 151

Figure 3. pH change of 15 wt % MEA in different gases at 40 °C.

ppm of SO2 on CO2 absorption by MEA is negligible within the 6 h period. Figure 3 also shows the effect of bubbling 151 ppm of SO2 along with the CO2 into the 15 wt % MEA solutions. The pH change in the presence of 151 ppm of SO2 is insignificant as compared with 15% CO2. CO2 capture technologies with MEA involves absorption of CO2 by the amine, followed by heating to release CO2 in a concentrated stream.5,38 This process was simulated for the 15 wt % MEA by bubbling 15% CO2 and 6% O2 at 40 °C until saturated (the concentration of MEAH+ stops increasing, 5 h), followed by desorption with N2 at 100 °C for 4 h. Figure 5 plots the intensity of the MEA (1362 cm−1), MEAH+ (1066 cm−1) and carbamate (1322 cm−1) peaks during the cyclic process, which was carried out over three cycles. There are minimal changes between each cycle. The process was repeated with 151 ppm of SO2. In the presence of 15% CO2/6% O2, the amount of MEA regenerated after three cycles is 39% of initial MEA. In the presence of 151 ppm of SO2/15% CO2/6% O2, there is 40% MEA regenerated. Therefore, the presence of 151 ppm of SO2 does not affect the MEA regeneration over the 25 h process. 3.3. Blended Amines: MEA and Piperazine (PZ). The CO2 absorption characteristics of a 15 wt % MEA aqueous solution containing 5 wt % piperazine in the presence of 15% CO2/6% O2 was investigated at 40 °C, and the spectra are shown in Figure 6. Besides the MEA (1362 cm−1), MEAH+ (1066 cm−1) and carbamate (1322 cm−1) bands, bands at 1520 and 1432 cm−1 assigned to asymmetric and symmetric vibrations of the COO− of PZ monocarbamate, respectively, and the 1289 and 1276 cm−1 bands assigned to N−C stretching vibration of the PZ carbamate are observed.17 When CO2 is added to aqueous PZ solutions, the main equilibria involved are17 CO2 + PZ + H 2O ⇌ PZCOO− + H3O+ −

−

(7) −

+

CO2 + PZCOO + H 2O ⇌ OOCPZCOO + H3O

(8)

PZ is a diamine and forms a hydrolysis-resistant carbamate (PZCOO−), as well as a dicarbamate (−OOCPZCOO−).39 With increasing CO2 absorption levels (>3 h) the PZ monocarbamate band at 1432 cm−1 shifts to 1426 cm−1, (inset D


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Figure 5. Cyclic experiments for 15 wt % MEA in the presence of 15% CO2/6% O2. Change of IR intensities of MEA, MEA carbamate, and MEAH+ of 15 wt % MEA during three cycles of adsorption and desorption of CO2 at 40 and 100 °C, respectively.

Figure 6. Evolution of IR spectra of 15 wt % MEA/5 wt % PZ blend during reaction with 15% CO2/6% O2 at 40 °C. (Inset: Curve deconvolution in the 1410−1450 cm−1 region for the 2 h sample with PZ carbamate at 1432 cm−1 and PZ dicarbamate at 1426 cm−1).

MEA carbamate. Though PZ is the minor constituent, the intensity of the PZ derived species is quite prominent. The relative peak intensity of free MEA (1362 cm−1), protonated MEA (1066 cm−1), and MEA carbamate (1322 cm−1) is plotted over time with and without PZ in Figure 8. The results demonstrate that the MEA consumption, and MEAH+ and carbamate formation become slower when PZ is included, with the MEA carbamate reaching saturation at 4 h instead of 3 h in the absence of PZ.

Figure 4. Comparison of IR peak intensities of reactants and products formed upon reaction of 15 wt % MEA with 15% CO2/6% O2 at 40 °C with or without 151 ppm of SO2 (a) MEA at 1362 cm−1, (b) carbamate at 1322 cm−1, (c) MEAH+ at 1066 cm−1.

4. DISCUSSION Speciation by infrared spectroscopy of the MEA−CO2 system has been reported extensively in the literature,15−18,40 and our spectral observations are in general agreement with these studies. The MEA−CO 2 reaction involves the formation of carbamate with a 2:1 stoichiometry.41

of figure shows the curve deconvolution), and assigned to the transformation of monocarbamate (PZ-COO− ) to the dicarbamate (−OOC-PZ-COO−).17 Figure 7a shows that the amine peak (955 cm−1) essentially disappears in 3 h. Figure 7b shows the evolution of the MEA and PZ species over a 6 h period. The PZ-dicarbamate is more stable as compared to the

CO2 + 2RNH 2 ⇌ RNHCO2− + RNH3+ E

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Figure 7. IR intensity profiles of reaction products of 15 wt % MEA/ 5% PZ with 15% CO2/6% O2 at 40 °C as a function of time: (a) amine, (b) MEA and PZ products.

This 2:1 MEA to CO2 reaction stoichiometry restricts CO2 absorption capacity of primary amines to a theoretical upper limit of 0.5 mol CO2/mol amine. Once this stage is reached, there will be no more absorption of CO2. With the 15 wt % MEA sample, the infrared band due to MEA disappears after 2 h (drops below 0.5 M, the detection limit of the IR). The carbamate band decreases after 3 h (Figure 2c), along with growth of the bicarbonate band (Figure 2a). Once the MEA has all reacted (3 h), further CO2 absorption leads to a drop in pH to 8.5 (Figure 3), and the carbamate begins to hydrolyze RNHCOO− + H 2O ⇌ RNH 2 + HCO3−

(10)

The hydrolysis is facilitated by drop in pH, since the carbmate is converted to carbamic acid, which undergoes rapid hydrolysis.11,15 The product of the carbamate hydrolysis, MEA gets protonated at the lower pH, and at the conclusion of the experiment at 5 h approaches 1.7 M. On the basis of the MEAH+ concentration, the CO2 loading in the 15 wt % MEA sample approaches 68%, higher than the 50% predicted by the 2:1 stoichiometry, due to the hydrolysis of the carbamate and the reaction of MEA with CO2. Observation of bicarbonate is a good indication of further uptake of CO2 beyond the stoichiometric limit. Effect of SO2. With 151 ppm of SO2 in the gas stream, the evolution of the vibrational bands of MEA, MEAH+, and the carbamate are unaltered (Figure 4). Also, the presence of SO2 does not appear to influence the thermal recycling process (Figure 5). The solubilities of SO2 and CO2 at 25 °C are 1.2 and 0.034 mol/L atm, respectively, but the effect of the much

Figure 8. Comparison of IR peak intensities of 15 wt % MEA with or without 5% PZ in 15% CO2/6% O2 at 40 °C as a function of time for (a) amine at 955 cm−1, (b) carbamate at 1486 cm−1, (c) MEAH+ at 1066 cm−1.

higher concentration of CO2 results in reducing the amount of SO2. Thus, we conclude that the IR is not showing any new species within the time frame (6 h) studied. It is reported in the literature that amine capacity for CO2 capture is reduced by SO2 as a result of irreversible reaction of amine with sulfur species. However, the experimental time of most such studies is much longer than that of this work.20,42,43 No reaction was observed by 1H NMR spectroscopic studies F



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between MEA and SO 2 over an year. 44,45 Also, low concentrations of sulfur species are detected, typically by ion chromatography,44 which is more sensitive than IR. Effects of Blending in PZ with MEA. In the presence of PZ (Figure 8), the vibrational bands from the products of the MEA + CO2 reaction (MEAH+ and carbamate) are slower to develop. More CO2 is being absorbed in the blended system since it took 4 h of bubbling CO2 for the MEA carbamate to reach saturation, as compared to 3 h in the absence of PZ (Figure 8b). We conclude that PZ and MEA are independently reacting with the CO2, and since the rate constant of PZ reaction with CO2 is an order of magnitude higher than that of MEA,46 the reaction of CO2 with MEA is retarded.

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS

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REFERENCES

This work is funded by The U.S. Department of Energy under Award Number DE-FE0007632. However, any opinions, findings, conclusions, or recommendations expressed herein are those of the author(s) and do not necessarily reflect the views of the DOE.

5. CONCLUSIONS Horizontal attenuated total reflection Fourier transform infrared (HATR-FTIR) spectroscopy is a suitable method for monitoring reactivity of aqueous MEA solution and for qualitative and quantitative analysis of the speciation during CO2 absorption. Infrared spectra of CO2 absorption with 15 wt % MEA were recorded at 40 °C. Specific bands were assigned to characteristic vibrational modes of MEA, protonated MEA (MEAH+), MEA carbamate, and carbonate/bicarbonate species. The evolution of each species was evaluated by plotting the relative absorbance as a function of time, which clearly demonstrated that the MEA−CO2 reaction involves the formation of carbamate with a 2:1 stoichiometry and subsequent carbamate hydrolysis. The change in pH was another indicator of this trend. Using standard solutions and pH titration, the absorptivity of MEA and MEAH+ were obtained, and actual concentrations during the reaction of CO2 with MEA were estimated. However, the quantitative result of MEA carbamate evolution was not acquired because methyl carbamate turned out to be an inappropriate infrared standard for MEA carbamate. Although sulfite was identified in the infrared spectrum after 5 days with continuous exposure of MEA to 151 ppm of SO2 in the presence of 15% CO2, the effect of 151 ppm of SO2 on MEA reaction with 15% CO2 was not significant over a short 6- h time scale. SO2 also had no influence with recycling experiments over 24 h. Qualitative identification of amine and carbamate species during the absorption of CO2 in a 15 wt % MEA aqueous solution containing 5 wt % piperazine was obtained by HATR. All PZ derived species could be differentiated in the infrared spectrum from MEA derived species. The speciation of MEA/PZ blends demonstrate that the CO2 absorption occurs independently for both amines, and leads to slower reaction with MEA and higher CO2 loading than MEA solution by itself.

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ASSOCIATED CONTENT

S Supporting Information *

. The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.iecr.6b00017. Calculation of absorptivity for MEAH+; pH-dependent amine protonation equilibria (Table S1); peak fitting for 4.91 M MEA in 15% CO2/6% O2 absorption at 40 °C for 5 h (Figure S1); calculation of absorptivity for MEA and methyl carbamate; Beer’s Law graphs of MEA and methyl carbamate (Figure S2); evolution of FTIR absorbance spectra of 15 wt % MEA in 151 ppm of SO2/6% O2 at 40 °C over 6 days (Figure S3)(PDF) G


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Infrared Spectroscopic Study of Reaction of Carbon Dioxide with

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