Amino Acids as Carbon Capture Solvents: Chemical Kinetics and

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Amino Acids as Carbon Capture Solvents: Chemical Kinetics and Mechanism of the Glycine + CO2 Reaction Dongfang Guo, Hendy Thee, Chun Y Tan, Jian Chen, Weiyang Fei, Sandra E. Kentish, Geoffrey W Stevens, and Gabriel da Silva Energy Fuels, Just Accepted Manuscript • DOI: 10.1021/ef400413r • Publication Date (Web): 04 Jun 2013 Downloaded from http://pubs.acs.org on June 11, 2013

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Amino Acids as Carbon Capture Solvents: Chemical Kinetics and Mechanism of the Glycine + CO2 Reaction Dongfang Guo,1,2 ‡ Hendy Thee,1 Chun Y. Tan,1 Jian Chen,2 Weiyang Fei,2 Sandra Kentish,1 Geoffrey W. Stevens,1 Gabriel da Silva1 * 1

Department of Chemical and Biomolecular Engineering, The University of Melbourne Victoria 3010, Australia 2

Department of Chemical Engineering, Tsinghua University Beijing 100084, China

Abstract Amino acids are potential solvents for carbon dioxide separation processes, but the kinetics and mechanism of amino acid – CO2 reactions are not well described. In this paper, we present a study of the reaction of glycine with CO2 in aqueous media using stopped-flow UV/visible spectrophotometry as well as by gas/liquid absorption into a wetted wall column. Combining these two techniques we have observed the direct reaction of dissolved CO2 with glycine under dilute, idealized conditions, as well as the reactive absorption of gaseous CO2 into alkaline glycinate solvents under industrially relevant temperatures and concentrations. From stopped-flow experiments between 25 and 40 °C we find that the glycine anion NH2CHCO2- reacts with CO2(aq) with k [M-1 s-1] = 1.24×1012 exp(-5459/T [K]), with an activation energy of 45.4±2.2 kJ mol-1. Rate constants derived from wetted wall column measurements between 50 and 60 °C are in good agreement with an extrapolation of this Arrhenius expression. Stopped-flow studies at low pH also identify a much slower reaction between neutral glycine and CO2, with k [M-1 s-1] = 8.18×1012exp(-8624/T [K]) and activation energy of 71.7±9.6 kJ mol-1. Similar results are observed for the related amino acid alanine, where rate constants for the respective neutral and base forms are 1.02±0.40 and 6250±540

* To whom all correspondence should be addressed. E-mail: [email protected] ‡ Current address: China Huaneng Group Clean Energy Research Institute, 48 Zhichun Rd, Haidian District, Beijing 100098, P. R. China. E-mail: [email protected].

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M-1 s-1 at 25 °C (versus 2.08±0.18 and 13900±750 M-1 s-1 for glycine). This work has implications for the operation of carbon capture systems with amino acid solvents, and also provides insight into how functional groups affect amine reactivity towards CO2.

Keywords: Carbon dioxide; kinetics; amino acid; carbon capture; absorption

Introduction

The removal of carbon dioxide from the flue gas of power stations for sequestration is being actively developed as a technology to help mitigate anthropogenic global warming [1, 2]. The most widely used technology involves reactive absorption into aqueous solutions of amines such as monoethanolamine (MEA), which has been relatively well developed for use in natural gas processing and other areas. However, the scale-up of conventional CO2 separation technology to the level required to combat climate change through carbon capture and storage (CCS) presents significant challenges, due to the energy requirements for CO2 regeneration and solvent loss from amine degradation and vaporization [2].

Amino acids present potential benefits over amines in CO2 separation systems. Amino acids typically have significantly lower vapor pressures than amines, resulting in reduced solvent loss to the atmosphere. Furthermore, amino acids tend to demonstrate greater resistance to oxidative degradation and lower toxicity than typical alkanolamine solvents [3, 4]. Amino acids such as glycine, alanine, proline, and taurine have been proposed as alternatives to amines in recent years [3-12], and have been used commercially in acidic gas treating processes, including in carbonate solutions (the GV process), in the BASF Alkazid solvent, and in membrane gas absorption units [3, 13]. However, there remain significant gaps in our knowledge of how amines and amino acids react with CO2, and how structure and functional groups control reactivity.

It is generally accepted that amines react with CO2 via a zwitterionic mechanism [2]. In this mechanism, CO2 binds with a primary or secondary amino group to form a zwitterion (Eq. 1), which rapidly deprotonates, exchanging H+ with water (Eq. 2) or any other base present in 2


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solution. This final step is exothermic, and produces a carbamate that can be thermally driven back to the amine + CO2. RR′NH + CO2 → RR′N+HCO2-

(1)

RR′N+HCO2- ⇌ RR′NCO2- + H+

(2)

Amino acids can exist in aqueous solution as the protonated, neutral, or deprotonated base form (e.g., for glycine see Eq. 3). The distribution between these three forms at 25 °C for glycine is depicted in Figure 1 [14] (numerical values are listed in the Supporting Information); we observe that above around pH 5, which is relevant to solvent absorption systems, the total amino acid concentration can be described as the sum of the neutral (AA0) and base (AA-) forms. Kinetic studies of amino acid – CO2 reactions have generally assumed that the base form is the active species, since this anion is expected to dominate under the alkaline solvent conditions. +

NH3CH2CO2H ⇌ +NH3CH2CO2- + H+ ⇌ NH2CH2CO2- + 2H+

(3)

A number of prior studies have considered the reaction between glycine and CO2. However, no kinetic data is available at temperatures representative of industrial absorption processes and the only studies that examine temperatures significantly above room temperature [4,6] report rate constants that are over an order of magnitude smaller than those obtained at around 300 K [3,5]. We aim to resolve this by investigating the kinetics and mechanism of CO2 reacting with glycine (Gly) using the complimentary techniques of stopped-flow UV/visible spectrophotometry and reactive CO2(g) absorption into a wetted-wall column. Combining these techniques allows us to span wide ranges of temperature and glycine concentration. We consider the reaction of CO2(aq) with both the neutral (k1, Eq. 4) and base (k2, Eq. 5) forms of glycine (Gly0 and Gly-, respectively), and provide further comparison to the amino acid alanine (at 25 °C). Our results demonstrate that the neutral forms of both amino acids do react with CO2, although they are far less reactive than their deprotonated counterparts. This work has implications for the operation of amino acid solvent systems for CCS, and more generally provides insight into the fundamental kinetics and mechanism of substituted amine reactions with CO2.

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k

1 AA0 + CO2 → products

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(4)

k

AA- + CO2 →2 products

(5)

Experimental Section

Materials All reagents employed here were of analytical reagent grade. Two indicators were purchased and used without further purification: 4-nitrophenol (≥99.5 %, Sigma) and thymol blue (Sigma). Sodium hydroxide (≥97.0 %, Chem-Supply) and sulfuric acid (Scharlau Chemie S.A.) were used to adjust pH values in the stopped-flow experiments, whereas potassium hydroxide (≥90 %, Sigma) was used in the wetted-wall column studies. Glycine (≥98.5 %) and L-alanine (≥99.0 %) were purchased from Chem-Supply. Gas mixtures of CO2/N2 were obtained from BOC Gases Australia Limited; Mixtures of 10.1 % and 14.8 % CO2 in N2 were used for the preparation of aqueous CO2 solutions, with mixtures of 10.2 % and 89.8 % CO2 in N2 used for the wetted-wall column experiments and for the calibration of a MGA3000C CO2 gas analyser (ANRI Instrument and Control Pty. Ltd.).

Stopped-Flow Kinetics The stopped-flow pH indicator technique [15-17] was adopted to study reactions of amino acids with CO2. A more detailed description of the apparatus and experimental procedure can be found in our prior work [17]. Reaction was initiated by mixing a solution containing dissolved CO2 with an aqueous solution containing the pH indicator and other reagents. CO2 solutions were freshly prepared for each batch of experiments by bubbling CO2/N2 gas mixture through a gas absorption bottle at atmospheric pressure containing deionized, distilled water. Bubbling was continued for at least one hour before the experiment began, with the flow of the mixture gas continued as long as the solution was in use. A gas-tight syringe with a needle was used to withdraw and contain the CO2 solution, which was then immediately connected to the sample tube of the stopped-flow apparatus, so as to avoid the solution contacting with air. The dissolved CO2 concentration was determined from published 4


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solubility data [18], and was typically in the range of 1.5 to 2.3 mM at the temperatures used here. For the solutions, pH was measured using a Metrohm Titrando 809 autotitrator (Switzerland), and adjusted to the desired value by adding sulfuric acid or sodium hydroxide. The corresponding concentrations [AA0] and [AA-] were determined by reference to Figure 1 [14].

Kinetic measurements were performed on an Applied Photophysics SX.18MV-R StoppedFlow Reaction Analyzer using the 10 mm optical path length configuration (Applied Photophysics Ltd., United Kingdom). The stopped-flow cell and reagent reservoirs were temperature-regulated to within ±0.1 °C using a Grant water bath. Prior to kinetic experiments, spectrophotometric properties for all reagents were measured between 200 nm and 800 nm in the stopped-flow cell. Extinction coefficients (ε) for the acid and base forms of the pH indicator thymol blue at 598 nm are 230 M-1 cm-1 and 3.57×104 M-1 cm-1, respectively, whereas those for the indicator 4-nitrophenol at 400 nm are 100 M-1 cm-1 and 1.81×104 M-1 cm-1, respectively.

The typical kinetic experiment was initiated by mixing the aqueous reagent solutions in a 1:1 ratio and monitoring the indicator absorbance versus time. Experiments were performed in which both pH and temperature were varied. Seven repeat runs were conducted at each set of conditions, and values reported here represent averages. Initial reaction rates (corresponding to around 10 % conversion) were obtained from the recorded absorbance traces by an exponential regression based on the Marquardt algorithm. Standard deviations from repeat experiments were used to determine uncertainty intervals (twice the standard error) for observed rate constants, which were propagated through the kinetic analysis to arrive at uncertainties in the final second order rate constants and activation energies. Observed rate constants determined using the stopped-flow device are provided as Supporting Information.

Wetted-Wall Column The kinetics of CO2 absorption were studied using a wetted-wall column (WWC), a device that allows contact between a gas and a liquid phase with controlled and measureable surface area for mass transfer, and therefore, accurate measurement of the flux of CO2 into potassium glycinate solutions. A detailed description of this apparatus and the characterization of its performance as well as a process flow diagram of the full experimental setup can be found in

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Thee et al. [19]. Briefly, the WWC consists of a stainless steel tube with a total contact area of 4840 mm2. The chamber is housed inside a second thick-walled steel chamber for temperature control. Glycinate solvent flows up through the middle of the tube and is evenly distributed on the outer surface of the column. The liquid is collected at the bottom of the column and recirculated. The gas supply which has been pre-saturated with water flows counter-currently past the liquid film before downstream concentration analysis.

Aqueous solutions of potassium glycinate were prepared by adding the amino acid glycine and an equimolar amount of potassium hydroxide to deionized water. Experiments were performed in which both glycine concentration and temperature were varied. During each experiment, data points were collected at a steady-state flux with a bulk CO2 partial pressure of 90 kPa. The physical properties of the glycinate solutions including density, viscosity, the Henry’s Law constant and the diffusivity constant of CO2 were obtained from the literature [5, 20]. The corresponding [AA0] and [AA-] concentrations were determined by reference to Figure 1 [14]. Observed rate constants obtained on the wetted wall column are listed as Supporting Information.

Results and Discussion Observed first order rate constants (kobs, s-1) for dissolved CO2 reacting with aqueous glycine and alanine solutions at 25 °C in the stopped-flow device are illustrated in Figure 2, as a function of [AA-] (i.e., pH). We observe that both amino acids exhibit similar reactivity, with the observed reaction rate increasing with increasing concentration of the base form (decreasing neutral form concentration). Furthermore, in both cases the intercept at [AA-] = 0 is small, indicating that the neutral amino acid is relatively unreactive towards CO2. These results are consistent with the base form of the amino acid being the dominant reactive species, although in order to extract accurate rate constants for these reactions we need to first characterize the reactivity of the neutral amino acid and also correct for CO2 hydration by H2O and OH-. In the following section we detail a series of stopped-flow kinetics experiments performed to 6


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determine rate constants for the neutral forms of glycine and alanine reacting with CO2. Following this, rate constant measurements for the deprotonated base forms are described. Finally, wetted-wall column experiments on CO2 absorption into alkaline glycine solutions at higher concentrations and temperatures are introduced and comparison is made to the stopped-flow results and to prior work.

CO2 Reaction with the Neutral Amino Acids Forms It is apparent from Figure 2 that the neutral amino acid + CO2 reactions are slow (relative to the base forms), making their rate constants difficult to accurately observe using the pH indicator technique in buffered solutions. However, the slow reaction rate permits us to study these reactions in essentially unbuffered media, as [H+] changes relatively little over the course of the reaction. In an aqueous amino acid solution the total reaction rate (r, M s-1) can be defined as the rate of CO2 consumption:

r = kobs [CO 2 ] = −

 dA d [H + ]  d [CO 2 ] d [H + ] = =  dt t = 0 dt t =0  dt dA  t =0,A 0

(6)

Here, kobs (s-1) is the observed pseudo-first-order rate constant and A is the spectrophotometric absorbance. The observed rate constant will include contributions from the reaction of CO2 with the neutral (k1) and base (k2) forms of the amino acid, as well as contributions from the CO2 + H2O (k3 = 4.07×106exp(-5584/T[K])) and CO2 + OH- (k4 = 9.88×1013exp(-6956/T[K])) reactions measured under similar conditions in our previous study [12]. The observed rate constant can thus be expressed as: kobs = k1[AA0] + k2[AA-] + k3 + k4[OH-]

(7)

In order to exclude reactions of the acid and base forms of alanine and glycine, experimental conditions have been chosen at around pH 6.0 (cf. Figure 1). Note also that similar experiments were performed at pH 6.9 (not shown) with near-identical results. The item d[H+]/dA in Eqn 6 can be obtained by mixing the aqueous glycine solutions containing 4nitrophenol as the indicator with known amounts of sulfuric acid solutions in a 1:1 ratio using

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the stopped-flow apparatus. The relation between added [H+] and absorbance displays a good linear fit, which is shown in Figure 3 for glycine.

Observed rate constants obtained from the absorbance versus time traces were used to determine k1 for both amino acid reactions according to Eq. 7, which corrects for CO2 hydration by water and hydroxide. Under our experimental conditions the hydroxide mechanism is negligible, although water hydration is somewhat important. Calculated rate constants, k1, are provided in Table 1 from 25 °C to 40 °C for glycine. Fitting the glycine rate constants to the Arrhenius equation (Figure 4) provides an activation energy of 71.7±9.6 kJ mol-1 and the rate constant expression k1 [M-1 s-1] = 8.18×1012 exp(-8624/T [K]). Reliable temperature-dependent pKa data was not located for alanine, so experiments were only performed at 25 °C. The rate constant determined for alanine at this temperature (1.02±0.40 M-1 s-1) is around half that determined for glycine + CO2 (cf. Table 1) indicating similar reactivity for the two neutral amino acids. Note that the relatively large uncertainties assigned to these neutral-form rate constants and activation energies are a consequence of the unbuffered technique developed to observe the slow reaction kinetics at around neutral pH.

CO2 Reaction with the Base Amino Acid Forms Alkaline conditions are required to investigate the reaction between CO2 and the base form of glycine, where equilibrium between the neutral and basic species acts as the buffer. Assuming that the ionization equilibrium of the buffer and indicator (thymol blue) are instantaneous, the initial reaction rate can be defined as follows:

kobs = −

1 d [CO 2 ] Q dA =− [CO2 ]0 dt t =0 [CO 2 ]0 dt t =0

(8)

Where Q is the buffer factor [10], defined by the following relation:

dA ∂A d [H + ] 1 d [H + ] 1 d [CO 2 ] = ( )t ( )= = dt ∂x dt Q dt Q dt

(9)

The buffer factor is obtained via Eqn 10. In this equation CB, CIn are the total buffer and indicator concentrations, respectively; αB, αIn are mole fractions of the base forms of the buffer and indicator, respectively; b is the optical path length; ∆ε is the difference in 8


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extinction coefficients between the acid and base forms of indicator. Values for the buffer factor calculated according to Eqn 10 were compared to values from spectrophotometric titration of the buffer-indicator system, and were found to agree well.

Q=

α B (1 − α B )CB α In (1 − α In )bCIn ∆ε

(10)

Figure 5 shows a plot of the observed reaction rate k′obs for glycine that is attributable to the amino acid only (i.e., k′obs = kobs - k3 - k4[OH-]), plotted against the concentration of the base form, [AA-]. We observe a good linear relationship between k′obs and [AA-], indicating pseudo-first-order kinetics. The slope of the least squares linear regression is (k2 - k1), and using the value of k1 determined above we have extracted values for k2 as a function of temperature, as listed in Table 2 (note that they change little when corrected for k1). Figure 6 reveals that the rate constants conform to a relatively linear Arrhenius plot, with rate constant expression k2 [M-1 s-1] = 1.24×1012exp(-5459/T [K]) and activation energy of 45.4±2.2 kJ mol-1. The rate constant for the base form of alanine (6250±540 M-1 s-1 at 25 °C) is again a factor of two smaller than that for glycine (cf. Table 2).

Wetted-Wall Column Experiments Absorption of CO2(g) into alkaline (pH = 8.9 to 10.0) solutions containing up to 2 M glycine was studied at 53 and 62 °C. Absorption rates for CO2(g) are plotted in Figure 7 as a function of glycine concentration and temperature. It is apparent that the presence of glycine dramatically increases the flux of CO2 into the solvent, which is consistent with large enhancement factors (on the order of 50 to 200) determined across the range of amino acid concentrations and temperatures considered here. Observed pseudo-first-order rate constants uncorrected (kobs, s-1) and corrected (k'obs, s-1) for reaction of CO2 with OH- are depicted in Figure 8 as a function of glycine anion concentration (note that correction is not made here for the slow reactions of CO2 with water and neutral glycine). We observe a first order dependence on the glycine anion concentration, consistent with the stopped-flow results and the postulated reaction mechanism. From the slopes of these plots second order rate constants k2 for the glycine anion + CO2 reaction are determined and presented in Table 2. Figure 9 compares these WWC rate constants to the corresponding stopped-flow values, along with an extrapolation of the fitted Arrhenius expression. We find that the WWC rate constants at higher temperatures conform to the Arrhenius expression determined using stopped-flow 9


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spectrophotometry at lower temperatures (and lower glycine concentrations).

Discussion We have performed a comprehensive kinetic and mechanistic study for the reactions of the simple amino acid glycine with carbon dioxide, demonstrating that the deprotonated base form of the amino acid is several orders of magnitude more reactive than the neutral form. This is an important result, as it helps guide operating conditions (specifically pH) for carbon capture systems employing amino acid solvents, as well as aiding in the design of novel amino acid (and amine) solvents. Carboxylic acid groups in amino acids typically deprotonate at alkaline conditions (cf. Figure 1), with the amino acid expected to be almost exclusively in the base form at around pH 12 and above. Concentrated aqueous solutions of amines and amino acid salts tend to have pH values several units below pH 12, where an amino acid such as glycine would be distributed between the active base and relatively inactive neutral forms. Reactive absorption of CO2, which releases H+ following deprotonation of the zwitterion intermediate to yield a carbamate, will further decrease the solvent pH. With glycine, for example, we find that operating at around pH 9 vs pH 12 would result in an approximate order of magnitude decrease in the apparent rate constant. Reaction rate improvements could therefore be achieved by maintaining a high pH, or indeed by utilizing amino acids with additional functional groups that make them reactive in their neutral forms, or with amine groups that deprotonate at lower pH. As noted in the introduction, an aim of this study was to provide kinetic data for the Gly- + CO2 reaction under industrially-relevant conditions, addressing the discrepancy between rate constants previously reported for this reaction. To this end, Figure 10 compares the values of k2(T) obtained in this study with the available literature data [3-7,10,21-23]. We find that our results are in good agreement with the low-temperature measurements of several groups [10,21-23], which are all somewhat below the values obtained by Kumar et al. [3] and Portugal et al. [5]. Note, however, that our results are orders of magnitude greater than those of Park and co-workers [4,6], who provided the only prior results at temperatures significantly above 300 K. These data appear to be in error and the Arrhenius expression provided here should yield a more accurate description of glycine + CO2 kinetics across the broadest range of conditions yet considered.

It is also of interest to compare the rate constant for the glycine anion + CO2 reaction to that 10


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of other amines. For the widely used primary alkanolamine, MEA, Thee et al. [24] reported a rate constant of 11 000 M-1 s-1 at 25 °C, whereas we obtained a slightly larger value of 13 900 M-1 s-1 at the same temperature for glycine. Moreover, the activation energy for the MEA reaction (31.8 kJ mol-1) is smaller than that of the glycine reaction reported here (45.4 kJ mol1

), which results in rate constants for glycine that are even higher than those of MEA at the

higher temperatures representative of typical industrial CO2 absorption processes (e.g., 57 000 vs. 30 000 M-1 s-1 at 50 °C).

Finally, the nature of the slow reaction between neutral amino acids and CO2 identified here cannot be assigned definitively from this work, although it is likely to be the reaction of the dominant zwitterion species. We note that there is some prior support for a cationic mechanism of CO2 hydration, and in this case the cationic –NH3+ functional group in neutral glycine and alanine could be facilitating such a reaction [25,26]. Interestingly, this mechanism does not proceed via a carbamate intermediate, and instead the ammonium cation site is acting as a catalyst for hydration of CO2 to H2CO3. Alternatively, we may be observing a relatively rapid reaction between the minor neutral form of glycine, NH2CH2CO2H, via a conventional carbamate mechanism. In this case the slow reaction kinetics that we observe would be predominantly attributable to the low concentration of this species versus the dominant zwitterion isomer.

Supporting Information Available: Glycine speciation data as a function of pH and temperature; observed rate constants from SF and WWC experiments.

Acknowledgments The authors gratefully acknowledge financial support from the China Scholarship Council, the Cooperative Research Centre for Greenhouse Gas Technologies (CO2CRC), the Particulate Fluids Processing Centre (PFPC) and the Australian Government through its CRC Program.

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Lackner, K. S. Science 2003, 300, 1677-1678. Rochelle, G. T. Science 2009, 325, 1652-1654. Kumar, P. S.; Hogendoorn, J. A.; Versteeg, G. F.; Feron, P. H. M. AIChE J. 2003, 49, 203-213. Park, S.-W., Son, Y.-S.; Park, D.-W.; Oh, K.-J. Sep. Sci. Technol. 2008, 43, 3003-3019. Portugal, A. F.; Derks, P. W. J.; Versteeg, G. F.; Magalhaes, F. D.; Mendes, A. Chem. Eng. Sci. 2007, 62, 6534-6547. Lee, S.; Song, H.-J.; Maken, S.; Park, J.-W. Ind. Eng. Chem. Res. 2007, 46, 1578-1583. Portugal, A. F.; Magalhaes, F. D.; Mendes, A. Chem. Eng. Sci. 2008, 63, 3493-3503. van Holst, J.; Kersten, S. R. A.; Hogendoorn, K. J. A. J. Chem. Eng. Data 2008, 53, 1286-1291. van Holst, J.; Versteeg, G. F.; Brilman, D. W. F.; Hogendoorn, J. A. Chem. Eng. Sci. 2009, 64, 59-68. Vaidya, P. D.; Konduru, P.; Vaidyanathan, M. Ind. Eng. Chem. Res. 2010, 49, 11067-11072. Hwang, K.-S.; Park, D.-W.; Oh, K.-J.; Kim, S.-S.; Park, S.-W. Sep. Sci. Technol. 2010, 45, 497-507. Paul, S.; Thomsen, K. Int. J. Greenhouse Gas Contr. 2012, 8, 169-179. Lu, J.-G.; Zhang, H.; Ji, Y.; Fan, F.; Liu, C. Energy Fuels 2010, 24, 4617-4626. Goldberg, R. N.; Kishore, N.; Lennen, R. M. J. Phys. Chem. Ref. Data 2002, 31, 231-370. Khalifah, R. G. J. Bio. Chem. 1971, 246, 2561-2573. Pocker, Y.; Bjorkquist, D. W. J. Amer. Chem. Soc. 1977, 99, 6537-6543. Guo, D.; Thee, H.; da Silva, G.; Chen, J.; Fei, W.; Kentish, S.; Stevens, G. W. Environ. Sci. Technol. 2011, 45, 4802-4807. Lide, D. R.; Bruno, T. J. CRC Handbook of Chemistry and Physics, CRC Press (2012). Thee, H.; Smith, K. H.; da Silva, G.; Kentish, S. E.; Stevens, G. W. Chem. Eng. J. 2012, 181–182, 694701. Kumar, P. S.; Hogendoorn, J. A.; Feron, P. H. M.; Versteeg, G. F. J. Chem. Eng. Data 2001, 46, 13571361. Penny, D. E.; Ritter, T. J. J. Chem. Soc., Faraday Trans. I 1983, 79, 2103-2109. Jensen, A.; Jensen, J. B.; Faurholt, C. Acta Chemica Scandinavica 1952, 395-397. Caplow, M. J. Amer. Chem. Soc. 1968, 90, 6795-6803. Thee, H.; Suryaputradinata,Y. A.; Mumford, K. A.; Smith, K. H.; da Silva, G.; Kentish, S. E.; Stevens, G. W. Chem. Eng. J. 2012, 210, 271-279. Jean, Y.; Lehn, J.-M. Chem. Phys. 1979, 39, 111-122. Jean, Y.; Volatron, F. Chem. Phys. 1982, 65, 107-111.

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TABLES

Table 1: Rate constants for CO2 reacting with neutral amino acids T (°C)

25.0

30.0

35.0

38.0

-1 -1

Glycine k1 (M s ) 2.08±0.18 4.22±0.53 5.31±0.74 7.51±1.08 Alanine k1 (M-1 s-1) 1.02±0.40

-

-

-

Table 2: Rate constants for CO2 reacting with basic amino acids T (°C) a -1 -1

25.0 b

30.0

35.2

40.0

53

62

Glycine

k2 (M s )

13900±750

19100±870

26800±1140

31900±2100

71800

91301

Alanine

k2 (M-1 s-1)

6250±540

-

-

-

-

-

a

Rate constants at 25 to 40 °C from stopped-flow experiments and at 53 to 62 °C from wetted-wall column

experiments. b 25.5 °C for glycine.

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Energy & Fuels

FIGURES

100 90 80 70 Distribution %

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 14 of 23

60

acidic neutral base

50 40 30 20 10 0

0

2

4

6

8

10

12

14

pH

Figure 1: Distribution (%) of glycine between the acid, neutral, and base forms in the pH range 0 – 14 at 25 °C [14].

14


Page 15 of 23

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Energy & Fuels

Figure 2: Plot of kobs versus [AA-] at 25 °C for glycine and alanine.

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Energy & Fuels

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 16 of 23

Figure 3: Typical absorbance trace of CO2 reacting with the neutral form of glycine at pH 6.9 and 40 °C (Insert: calibration plot of added known amounts of sulfuric acid versus absorbance).

16


Page 17 of 23

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Energy & Fuels

Figure 4: Arrhenius plot of lnk1 versus 1000/T for glycine + CO2.

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Energy & Fuels

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 18 of 23

Figure 5: Plot of k′obs (kobs - k3 - k4[OH-]) versus [Gly-] at 25 °C – 40 °C.

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Page 19 of 23

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Energy & Fuels

Figure 6: Arrhenius plot of lnk2 versus 1000/T for glycine + CO2.

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Energy & Fuels

500

CO2 Absorption Rate / mL/min

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 20 of 23

53°C 62°C 400

300

200

100

0 0.0

0.5

1.0

1.5

2.0

[Gly]Total / M

Figure 7: Wetted-wall column CO2 absorption rates as a function of total glycine concentration (0.5 to 2.0 M) and temperature (53 and 62 °C).

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Page 21 of 23

200000

(a)

53°C 62°C

kobs / s-1

150000

100000

50000

0 0.0

0.5

1.0

1.5

2.0

−

[Gly ] / M

200000

(b)

53°C 62°C 150000

k'obs / s-1

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Energy & Fuels

100000

50000

0 0.0

0.5

1.0

1.5

2.0

−

[Gly ] / M

Figure 8: Observed first-order rate constants for CO2 hydration as a function of glycine concentration, [Gly-]. Rate constants are (a) uncorrected (kobs) and (b) corrected (k'obs) for hydration via the CO2 + OH- reaction.

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Energy & Fuels

11.6

WWC data Stopped-flow data 11.2

ln [k2 / M-1 s-1]

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 22 of 23

10.8

10.4

10.0

9.6 3.0

3.1

3.2

3.3

3.4

1000/T [K]

Figure 9: Plot of lnk2 versus 1000/T for glycine + CO2 from wetted-wall column (WWC) experiments compared to stopped-flow results and extrapolated Arrhenius fit (dashed line).

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Page 23 of 23

14

Present work [SF] Present work [WWC] Penny et al. [21] Kumar et al. [3] Portugal et al. [5] Jensen et al. [22] Caplow et al. [23] Vaidya et al. [10] Park et al. [4] Lee et al. [6]

12

ln [k2-Gly/ M-1 s-1]

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Energy & Fuels

10

8

6

4 3.0

3.2

3.4

3.6

3.8

4.0

4.2

1000/T [K]

Figure 10: Comparison of k2(T) measurements for Gly- + CO2 made in this work (SF: stopped-flow; WWC: wetted-wall column) and in the prior literature.

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Amino Acids as Carbon Capture Solvents: Chemical Kinetics and

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