In Situ Nuclear Magnetic Resonance Mechanistic Studies of Carbon

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IN SITU NMR MECHANISTIC STUDIES OF CARBON DIOXIDE REACTIONS WITH LIQUID AMINES IN AQUEOUS SYSTEMS: NEW INSIGHTS ON CARBON CAPTURE REACTION PATHWAYS Pavel V. Kortunov*, Michael Siskin, Lisa Saunders Baugh, David C. Calabro Corporate Strategic Research Laboratory ExxonMobil Research and Engineering Co. 1545 Rt. 22 East, Annandale, NJ, USA 08801

Supplementary Material Available Abstract A series of closely related primary, secondary and tertiary alkanolamine model compounds were monitored in real time in aqueous solution via in-situ nuclear magnetic resonance (NMR) spectroscopy while purging CO2– rich gas through the solution over a range of temperatures. The real-time in-situ spectroscopic monitoring of this reaction chemistry provides new insight about reaction pathways through identification of primary products and their transformations into secondary products. New mechanistic pathways were observed and elucidated. The effects of CO2 loadings, relative absorption and desorption kinetics, pH, temperature, and other critical features of the amine/CO2 reaction system are discussed in detail. The effect of amine basicity and structure on these parameters was further elucidated by studying complementary electron-rich and -poor amines (pKa ~4.5-11) and guanidines (pKa ~ 14-15). While tertiary amines act only as simple proton acceptors, primary and secondary amines function as both bases and nucleophiles to form carbamates and (bi)carbonates whose product ratio is a function of both reaction conditions and amine steric and electronic properties. Water is also acting as a Lewis base by hydrolysis of carbamate species into bicarbonate which results in a more beneficial 1:1 CO2:amine ratio. Primary and secondary amines tend to react with CO2 similarly at different CO2 partial pressures, showing weak pressure dependence on CO2 loading; in contrast, reaction efficiencies of tertiary amines which generally form less stable carbonate and bicarbonate products are a strong function of CO2 pressure. Primary and secondary amines capture significantly less CO2 per mole of amine than tertiary amines (lower CO2 loading capacities) due to the formation of carbamate species. Their faster reaction rates with CO2 and high capture efficiencies at low CO2 partial pressures are advantageous. In contrast, tertiary amines more effectively react with CO2 at lower temperatures, capturing up to 1 CO2 per amine; initially, and unexpectedly, carbonate and bicarbonate species are initially formed simultaneously. Even at high pH carbonates evolve into a final bicarbonate product. The secondary benefit of forming bicarbonates are their lower thermal stability (greater ease of desorption). Unexpectedly guanidines do not form bicarbonates directly; reaction proceeds via exclusive initial formation of the guanidinium carbonate. In summary, varying amine basicity leads to significant changes in the carbamate/(bi)carbonate equilibrium and stability of reaction products.

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Introduction and Background (MEA, HOCH2CH2NH2, bp 170 ºC), in about 30 wt% concentration in water, is a widely used alkanolamine absorbent in this chemical capture because of its low cost and low molecular weight (more moles per pound). Regeneration of the CO2 saturated aqueous MEA solution is very energy intensive since thermal decomposition of its carbamate salt requires ~120 ºC and supplying the heat of vaporization of large amounts of water. Secondary amines form less stable carbamates. Tertiary amines lack a transferable proton and therefore cannot form carbamates; they only form bicarbonates. Generally, secondary and tertiary amines are more costly. Chemical reaction/absorption is preferred in cases with low concentrations or amounts of CO2 in the combusted gases. Regeneration is accomplished by applying heat to dissociate the carbon-nitrogen chemical bond formed by amine reaction with the CO2.

Carbon dioxide is a ubiquitous component of emissions from the combustion of fossil fuels for power generation, oil refining, transportation, and during natural gas production and most chemical manufacturing. Methane-rich natural gas from underground wells is usually contaminated with acid gases, mainly carbon dioxide (and to a lesser degree, sulfur-containing components such as hydrogen sulfide, mercaptans and carbonyl sulfide). The chemical separation/capture of carbon dioxide from such gaseous mixtures using aqueous amine solutions is the main focus of this paper. Various commercial processes are used for separating CO2 from gas mixtures. These technologies are classified into processes where: (a) a chemical reaction between an amine and CO2 occurs; (b) where the CO2 is only physically dissolved into a solvent e.g., methanol, glycols); or (c) a combination of both. The quantity of absorbent required is a function of the quantity of CO2 to be removed, the partial pressure of CO2 and absorbent Vapor Liquid Equilibrium (VLE). The main chemical and physical solvents used in acid gas separation have been summarized by Nirula and Ashraf.1 Detailed kinetics, and postulated reaction pathways have been comprehensively reviewed and have provided fundamental insights into CO2 reaction mechanisms for chemical separations.2-4

Physical absorption is a bulk solubility phenomenon where inorganic or organic liquids may be used to preferentially absorb CO2 from the gaseous mixture. Physical absorption by solvents is favored by high pressures of CO2 and low concentrations of inert gases. The solvent regeneration step is relatively simple for physical absorption because it is carried out by increasing the temperature or by reducing the pressure. The absorption process depends on the operating temperature and pressure, as well as the nature of the gases and the absorption liquid. The solubility of a gas in an absorbent solution is proportional to the solubility of each gas and the partial pressures of that gas in the feed (Henry's Law). The magnitude of the heat of absorption is lower than that in chemical absorption. Solvents such as methanol, used in the Rectisol® Process, sulfolane used in the Sulfinol®-M and Sulfinol®-D processes and dimethyl ethers of polyethylene glycols used in the Selexol™ Process are preferred1. Higher boiling solvents are preferred to minimize solvent losses and to prevent contamination of the released gas with solvent vapors. This type of process could be a very efficient approach for processing such high-pressure CO2-rich streams as those encountered in advanced power-generation systems, such as Integrated Gasification Combined Cycle (IGCC).

The most common chemical process is an acidbase reaction of CO2 with an amine in aqueous solution. For example, the process for selective absorption of CO2 from a normally acid gas mixture emitted in the flue gas from power plants involves: (a) counter-currently contacting the gaseous mixture with a basic solution of an amine compound, typically an aqueous solution, under conditions such that CO2 is selectively absorbed from the mixture; (b) regenerating, at least partially, the absorbent solution containing CO2 to give the original amine and CO2 gas; and (c) recycling the regenerated solution for the selective absorption of CO2 back to step (a). The regeneration step is usually carried out by heating and stripping with steam. Chemical absorption involves one or more reversible chemical reactions between acidic CO2 and a basic aqueous solution of an absorbent, such as an alkanolamine (vide infra). This is an elementary rapid acid-base reaction in which the “acidic” CO2 is effectively neutralized by reaction with two moles of the amine to form a carbamate salt (R2NH2+.R2NCO2-). Monoethanolamine

Selection of a suitable process for acid gas removal is affected by both the recovered product gas purity specification and a rather large number of interacting parameters, such as the partial pressure


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of the acid gas impurities, total pressure of the feed gas, inlet temperatures of the feed gas, degree of removal required, trace impurities, utilities available, economic considerations (energy and capital costs), required plant life, size and weight, location, and environmental constraints. The possibility of integrating the acid gas removal unit within the overall plant, utilizing for example, lowgrade heat for regeneration, can also influence the choice of CO2 capture technology.

to nucleophilic attack by various N- and O-donors. The traditional reaction chemistry of aqueous amine/CO2 reactions is summarized in Scheme 1. For primary and secondary amines in aqueous solution, two types of reaction products are formed. The amine can act as a Brønsted base to neutralize the carbonic acid formed by reaction of the gaseous CO2 dissolved in the water (i.e., nucleophilic attack by water); this results in alkylammonium bicarbonate/ carbonate products, Scheme 1(a). Note that for an ammonium bicarbonate, the amine to CO2 ratio is 1:1, meaning that each amine has “captured” one mole of CO2, whereas for an ammonium carbonate, the amine to CO2 ratio is 2:1, or only 0.5 mole of CO2 captured per amine. The amine can also function as a nucleophile (i.e., Lewis base), adding to the carbonyl group of the CO2 molecule to form a zwitterion15, which is the ionic form of the analogous carbamic acid (1:1 amine to CO2 ratio), Scheme 1(b). The zwitterion/carbamic acid species is then rapidly deprotonated by a second mole of free amine to give a more stable ammonium carbamate (2:1 amine to CO2 ratio). Water can subsequently hydrate the carbamate species, releasing free amine, to produce the bicarbonate product shown in (a). Lacking a proton, tertiary amines are not able to rearrange from zwitterion to carbamic acid via intramolecular proton transfer; thus, they do not generate stable carbamate products with CO2 and only form carbonate and bicarbonate products in aqueous solution.

This paper focuses specifically on fundamental understanding of CO2 reaction mechanisms in aqueous systems, using various amines having a wide range of basicities (pKa ~4.5-15.5; anilines guanidines). These studies provide insights for the design of more economically efficient, selective CO2 scrubbing from commercially important source gases, e.g., power plant flue gases. In-situ 13C and 1H NMR spectroscopy using a built-in micro reactor was used to provide real time insights on reaction mechanisms and product speciation under various conditions. These learnings allow us to select and fine-tune absorbent structure, reaction mechanisms, and products that could provide substantially increased overall process efficiency. NMR has been used by others to identify the reaction products of amine bases with carbon dioxide5-9 but not for in-situ real time mechanistic and relative kinetic studies as presented here; as a function of time, temperature, CO2 partial pressure, and pH, for both the absorption and desorption portions of the process. The use of titration10 and Fourier Transform Infrared Spectroscopy (FTIR) methods have been reported11-13 however, quantification and real time / flow unit identification of unexpected species is cumbersome. We report here the first in a series of in-situ NMR studies concerning CO2 reaction chemistry of amine, and guanidine bases. This paper describes aqueous amine-CO2 chemistry; future papers in this series will focus on non-aqueous, mixed base, ionic liquid, and hindered base systems.14

Reaction Chemistry Sorption

of

An important practical consequence of these different reaction pathways, and a motivation for the model amine work described in this and subsequent publications, is that the theoretical maximum CO2 sorption capacity for an amine system varies based on which products are formed. When carbamates are the preferred reaction product (aqueous or non-aqueous conditions), the maximum possible CO2 sorption capacity is only 50 mole % per amine. However, if bicarbonates are preferred, the theoretical maximum CO2 sorption capacity doubles to 100 mole % per amine. Since carbamate and (bi)carbonate species have different relative stabilities and energies of formation, the balance between these two species may be tuned by tailoring amine structure, nucleophilicity and pKa, as well as reaction parameters such as temperature and amine concentration, providing more efficient amine-based CO2 scrubbing processes.

Amine/CO2

The chemistry of CO2 derives from the high electron deficiency of its carbon atom bonded to two highly electronegative oxygens. The electrophilicity of this carbon makes it susceptible


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Scheme 1. Pathways for CO2 reactions with amines in aqueous solution: (a) direct formation of the ammonium bicarbonate/carbonate; (b) formation of a carbamate via zwitterion/carbamic acid solubility of the amine and its reaction products with CO2. Lower-pKa aromatic and nitrile-amines (7-9) and higher-pKa cyclic amines and guanidines (10-12) were also studied to probe larger effects of basicity. Additional amines studied (but not discussed in detail in the text) are tabulated in Supplementary Material. These model amines were chosen for their commercial availability, high solubility in water and relatively high boiling points. The diamines 4-6 were also studied to determine if the presence of the second amine center (and elimination of the potentially reactive OH functionality) would alter the preferred reaction stoichiometry. Other than a pronounced effect on solution viscosity in some cases, the presence of the second amine group in the molecules did not appear to affect behavior as compared to model compounds bearing only one amine and an OH group per molecule.

Experimental Design Using in-situ NMR monitoring, we will first describe the temporal evolution of product formation for the reaction of CO2 with a series of primary (1°), secondary (2°), and tertiary (3°) amines (Table 1) in aqueous solution as a function of reaction temperature. The quantitative results illustrate stepwise product formation and decomposition under absorption and desorption conditions, respectively. Subsequently, we will discuss the behavior of amine and guanidine sorbents covering a wide range of pKa values. These studies employ two homologous Nmethylated series of monoalkanolamines and their ether dimers to study the effect of amine structural type on reactivity (Table 1, alkanolamines 1-3 and etheramines 4-6). Alkanolamines and etheramines are preferred over the use of amines lacking polar functional groups, in order to facilitate high water


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Table 1. Physical properties of amines used in aqueous CO2 sorption experiments.

Compound No.

Mol. Wt (g/mol), density (g/mL), boiling pt. (0C)

Structure

Methylaminoethanol (MAE)

61.08 1.01 170 75.11 0.935 158

Dimethylaminoethanol (DMAE)

89.14 0.886 134

1 Monoethanolamine (MEA) 2

3

4 1,5-Diamino-3-oxapentane (DAOP) 5 1,5-Bis(methylamino)-3-oxapentane (DMAOP) 6

104.15 0.98 64/4 mm 132.21 0.872 76-78/20 mm 160.26 0.841 189

Bis[2-(N,N-dimethylamino)ethyl]ether (DMAEE)

ACD predicted pKa of conj. acida 9.16

9.40

8.88

9.07

9.87

9.12

Aminoacetonitrile (AAN)

70.09 0.958 185 56.04 0.956 58-66/8-15 mm

4.61

Aniline

93.13 1.022 184

10.45

Piperidine (PP)

85.15 0.862 106

Piperazine (PZ)

86.14 1.1 (solid) 145-146

9.55 5.41

115.18 0.918 52-54/11 mm

15.20

7 Aminopropionitrile (APN) 8

9

10

11

7.14

5.43

NH 12

N

N

1,1,3,3-tetramethylguanidine (TMG) a

Predicted using ACD/Labs (Advanced Chemistry Development) software V8.14 for Solaris© as reported by Chemical Abstracts Service (SciFinder), 25 °C. pKas of other species relevant for aqueous CO2 uptake: H3O+, -1.74; H2O, 15.7; H2CO3 (carbonic acid), 6.35; HCO3- (bicarbonate anion), 10.33. [March, J. Advanced Organic Chemistry, 3rd Ed.; John Wiley & Sons: New York, 1985, p. 221.]


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are 25 °C predicted values obtained using ACD/Labs (Advanced Chemistry Development) software V8.14 for Solaris© as reported by Chemical Abstracts Service (SciFinder).16,17

Experimental Section General NMR Procedure for CO2 Uptake and Desorption The experimental setup for in-situ monitoring of CO2 uptake by the amine solutions was built inside a wide bore 400 MHz Bruker Avance™ NMR spectrometer equipped with variable temperature capability, a 10 mm Broad-Band liquid NMR probe, and Bruker TopSpinTM 1.3 software. A plasticcapped 10 mm NMR tube containing the solution to be tested was placed inside the probe (Figure 1). A Sigma-Aldrich micro pH glass combination electrode (3.5 mm diameter), connected to an external pH meter, was positioned inside the solution but above the NMR monitoring region in order not to interfere with NMR measurements. In some experiments, a sealed glass capillary tube containing ethylene glycol was added to the solution for accurate temperature monitoring of the solution during reactions. The tube was sealed with a plastic cap fitted with two thin plastic tubes for CO2 gas flow in and out of the solution. The CO2 inlet tube was positioned below the solution surface. The gas outlet tube was connected to laboratory ventilation that set an ambient, e.g. 0 psig (or 1.0 bar absolute) pressure in the NMR tube. The gas flow rate was controlled by calibrated Brooks 5896TM electronic flow regulators in the range of 1.0 – 50.0 cc/min, depending on the CO2 partial pressure in the feed gas. Although the majority of experiments reported here were performed by purging pure CO2 gas (PCO2 = 1.0 bar), special mixtures of gases were also used to study effects of CO2 partial pressure. For example, 10 mol% / 90 mol% CO2/N2 mixture, purchased from Matheson Tri-Gas, was used to monitor reaction at CO2 partial pressure 0.1 bar. Amine solutions (typically 15-35 wt% amine) were prepared directly in 10 mm diameter, 8-inch NMR tubes using distilled, deionized water or D2O, with or without a drop of d6-DMSO (taken from a recently open ampule and/or stored in a glove box) to provide a reference 13C signal. Molarity and per cent by weight of amine of the specific solutions discussed are noted in the text. Exact amounts and concentrations were sometimes adjusted slightly in order to maintain a sufficient height of solution to cover the pH meter probe tip. A typical solution used ca. 700 mg of amine and ca. 4 g of solvent. All amines and solvents were purchased from commercial sources and used as received. In order to utilize the most consistently obtained pKa values (and since literature values were not located for some amines), all amine pKa values reported herein

After bringing the solution to the desired temperature, CO2 flow was initiated. The solution temperature was controlled by a pre-heated N2 purge (either house N2 or liquid N2 vapor) flowing at 1200 L/h through the NMR probe with the solution tube. A thermocouple was mounted 10 mm below the sample. The temperature range for the experiments was limited by the amine/solvent physical properties (boiling and freezing points), and was narrower than the NMR instrument capabilities (-150 to +120 °C). Desorption experiments were performed by changing the feed gas to N2, using the same flow rate, and increasing the solution temperature if needed. When applicable, additional 1D and 2D NMR analysis of the solutions was carried out in 5 mm NMR tubes using a Bruker Avance IIITM narrow bore 400 MHz spectrometer equipped with a 5 mm QNP probe and Bruker TopSpinTM 2.1 software.

Spectral acquisition For 13C NMR quantitative analysis of the starting solution and final product(s) of CO2 absorption, a standard single-pulse sequence with proton decoupling (zgig pulse sequence) with repetition delay equal or longer than 60 seconds was used. At least 64 scans were typically taken to generate the 13 C spectrum. In order to observe intermediate reaction products qualitatively on a short time scale, NOE signal enhancement (zgpg or zgpg30) was used with a shorter repetition delay between 2-5 seconds. Further calibration of 13C peak intensities was performed after every reaction on the final reaction products by comparing NMR spectra taken with and without NOE enhancement. For 1H NMR quantitative analysis of the starting solution, intermediate products, and final products of CO2 sorption, a single-pulse zg sequence was used with a repetition delay between 10 and 60 seconds. At least 8 scans were typically taken to generate a 1H spectrum. Manual tuning and matching procedures for the NMR probe were performed between experiments in order to correct impedance changes of 13C and 1H circuits during the reaction caused by the formation of new chemical compounds. Spectral analysis It has to be noted that only the 13C isotope of carbon is detectable by NMR due to a spin quantum number of ½ (12C has zero spin quantum number and, therefore, shows no magnetic activity). The natural abundance of 13C carbons is approximately


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~1.11 %. To eliminate a potential effect of 13C natural abundance variability on result analysis and interpretation, we performed carbon isotope characterization of MEA (1), DMAE (3), TMG (12), and CO2. The corresponding δ13C values of -27.0‰, -37.3‰, -41.3‰, and -11.0‰ (with respect to the Pee Dee Belemnite standard, 1.111% 13C) indicate that 13C abundance in selected compounds varies over the range 1.066% (for TMG) to- 1.099% (for CO2). Based on the low variability of 13C abundance in amines and gaseous CO2, 13C NMR spectroscopy can be used to quantitatively analyze amount of reacted CO2 versus the amount of amines present in the solution.

discussed in subsequent publications10, 1H NMR provides key information about CO2-amine reaction products, their evolution, and equilibrium6.

Percent carbamate loading was calculated by integration of the 13C NMR carbamate carbonyl resonance (1 carbon, usually at ca. 163.0-164.5 ppm) versus resonances representing the total amine methylene and (if present) methyl groups. After normalizing for the number of carbons represented in the amine aliphatic region, the carbonyl integral was divided into the amine aliphatic integral to obtain the mole percent CO2 present per molecule and per amine group (correcting when necessary for 2 amines per molecule). Percent carbonate/bicarbonate (not resolved due to rapid exchange7) loading was similarly calculated using the 13C NMR carbonate/bicarbonate carbonyl resonance (1 carbon, usually found at 160.0-168.0 ppm).

Tertiary amines The tertiary amine dimethylaminoethanol (DMAE) (3) illustrates the simplest chemical behavior in the presence of CO2. Unlike primary and secondary amines, it can only form the two types of stable products shown in Scheme 1(a) – ammonium bicarbonate and/or carbonate (dependent on the pH of the solution). Initially, carbonic acid is the expected product after CO2 gas is dissolved and hydrated in water (Scheme 2). At ambient pressure and temperature, the concentration of carbonic acid is low18 which limits the rate of bicarbonate and/or carbonate formation. The C=O resonance for free CO2 (~124.5 ppm) and aqueous carbonic acid (predicted at 156.5 ppm) was not observed at any stage in the experiment, reinforcing the very low concentration of these species. After pure CO2 was introduced to an aqueous solution of DMAE (3 M/28 wt % in protio-H2O) at 1.0 bar and 30 °C, one 13C NMR peak appeared in the carbonyl region at ~164.5 ppm, as shown in Figure 2.2 (top). Since the resonances for pure carbonate and bicarbonate species are predicted at approximately 168.5 ppm and 160.5 ppm∗, respectively, this initial peak was assigned as a bicarbonate/carbonate mixture in rapid equilibrium (Scheme 2)6. As more CO2 is introduced into the solution, the resulting 13C resonance peak increases in intensity and gradually shifts upfield confirming preferable formation of bicarbonate over carbonate species when a CO2:amine stoichiometry close to 1:1 is reached (Figures 3 and 2.1, bottom).

It is important to note that under the conditions used for these measurements, the gas-liquid mass transfer is rate-limiting and the time-dependent data is not reflective of the intrinsic reaction kinetics. Rather than a detriment, this attenuation of the real kinetics has enabled careful monitoring of the sequence of reactions and relative kinetics occurring as a function of concentration and temperature.

Results and Discussion

1

H NMR was primarily used to verify amine concentration in water solution. The signal of amine protons were integrated versus the water signal. The 1H structural resonances of the amine (– CH2–, –CH–, and/or –CH3, usually, at 1.0-4.0 ppm) did not require any normalization and could be detected directly. However, the –OH resonance of water (at 4.78 ppm) interferes with the -OH, –NH, and –NH+ resonances of the amine through fast proton exchange, resulting in a single resonance shifted downfield (usually at 4.7-5.5 ppm). In order to extract the integral value of the proton signal coming from water, the proton signals coming from the amino- and hydroxy- protons were calculated based on previously measured structural protons of the amine (methylene, methyl, etc.) and subtracted from the gross integral. In principle, 1H NMR may give quantitative information about reaction products by accounting for –CH– peak splitting. In this study, such a method was not used for quantification of reaction since it gives only an indirect evaluation of the reaction whereas 13C NMR can directly detect reacted CO2 and all reaction products. However, in water-free non-aqueous solutions, which will be

∗

C=O resonance of pure NaHCO3 in D2O = 161.44 (relative to d8-THF); shift of pure Na2CO3 = 169.95 ppm; shift of a 1:1 molar mixture = 165.62


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Scheme 2. Pathways for CO2 reactions with tertiary amines in aqueous solution: direct formation of the ammonium bicarbonate/carbonate via deprotonation of carbonic acid as illustrated for DMAE bicarbonate/carbonate 13C NMR resonance at 160165 ppm versus the resonance representing the amine HOCH2- groups. The total absorbed CO2 concentration in the solution (bicarbonate + carbonate) gradually grows, reaching a maximum amount of reacted CO2 after approximately 3 hours of bubbling at 10 sccm and 30 0C. The equilibrium CO2 loading reached 0.94 moles of CO2 per amine. This corresponds to an approximately 45.5 amine weight % uptake. The fact that the amount of reacted CO2 that reacted is near the theoretical limit of 1.00 also implies that the product is present mainly as bicarbonate, rather than carbonate (the theoretical uptake limit for carbonate is 0.5 mole of CO2 per amine group). The apparent 94% loading either implies that 6% of the amines remain unreacted (unresolvable from ammonium species), or that 12% of the amines are ammonium groups associated with divalent carbonates rather than with monovalent bicarbonate species. Figure 3 further shows pH monitoring of the DMAE solution simultaneously with CO2 uptake. The initial solution pH of 12.05 dropped with the formation of (bi)carbonate species and reached 8.17 at maximum uptake.

The amine backbone carbons of DMAE showed sensitivity to amine protonation due to formation of the (bi)carbonate (Figure 2.2, bottom right). Since free and protonated amines exist in a fast exchange mode through rapid proton transfer, the slight shift upfield can be explained by increased concentration of protonated amines in the solution (in conjunction with increasing concentration of carbonate/bicarbonate). We used this approach to calculate the concentration of DMAE existing as the free base and protonated DMAE based on the 13C peak position. The shifts of both DMAE carbons, NCH2CH2OH, are very sensitive to protonation. The most sensitive backbone resonance (as gauged by the degree of ppm shift upon reaction with CO2) appears to be the HOCH2– carbon, rather than the – CH2N– or –NMe2 carbon. Therefore, we fitted the time evolution of the HOCH2– peak (Figure 2.2 bottom, red curve) and calculated the mole fraction of free amine and protonated amine in the solution during the reaction with CO2. We used values 59.11 ppm and 55.64 ppm (see Figure 2.1) as 13C resonances of the HOCH2- carbon of the free base and fully protonated amine, respectively. Approximately half of the amine was protonated after 30 minutes of CO2 flow as shown in Figure 3 by the solid red curve, with equilibrium (near-full conversion to ammonium bicarbonate) reached at approximately 3 hours. It is worth comparing the 3 hour reaction time for protonation of DMAE (pKa ~8.88) to that for the significantly more basic 1,1,3,3tetramethylguanidine (TMG (12), pKa ~15.2, vide infra) which was fully protonated in approximately 60 m under similar reaction conditions.

In speciating the 13C NMR C=O region of the amine/CO2-products, the assignment of the peaks to specific products such as bicarbonate, carbonate, carbamate, carbamic acid, zwitterionic structures, and other potential byproducts such as ureas is complicated by the overlapping range of ppm shifts (typically 170-155 ppm) seen for these various species. For the DMAE/CO2 product shown in Figure 2.1 (bottom), the single observed C=O peak was assigned to equilibrating bicarbonate/carbonate based on the reaction chemistry described above. This single peak represents both bicarbonate and

Figure 3 also shows the ratio of total captured CO2 per amine group over time (solid blue curve) as quantified by the integration of the


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carbonate species present in a fast equilibrium The relatively upfield 160.85 ppm shift of the CO2/DMAE reaction peak likely indicates that the equilibrium reaction product is predominantly in the bicarbonate form, consistent with the pH of the final solution and the total reacted CO2/amine ratio in Figure 3. We have thereby used the 13C resonance value of the carbonyl peak during this reaction to quantify the equilibrium of bicarbonate and carbonate species in solution as a function of reaction time. By fitting the time evolution of the carbonate/bicarbonate 13C NMR peak (blue curve on Figure 2.2, bottom left), and basing our analysis on the endpoint 13C NMR resonances of pure tetramethylguanidinium carbonate and bicarbonate (168.35 and 160.40 ppm, respectively (vide infra), we obtained the ratio of carbonate and bicarbonate species in the solution. The total CO2 uptake curve (Figure 3, solid blue curve) was used to calculate the mole fraction of CO2 per amine in carbonate and bicarbonate species (dash and dash-dot curves, respectively).

As CO2 is introduced into the amine solution, carbonic acid is formed at low concentration. It is not detected by NMR because it is rapidly deprotonated by DMAE to form the carbonate/bicarbonate mixture (preferentially bicarbonate as opposed to pure carbonate expected at high pH) shown in Scheme 2. With increasing protonation of DMAE, the solution pH drops favoring the formation of bicarbonate from carbonic acid even more. A second route to bicarbonate formation becomes possible at low pH. After a majority of amines are protonated and form carbonate and bicarbonate species with CO2, carbonate anions (conjugate base) can deprotonate carbonic acid to form two bicarbonate molecules (Scheme 3):

Figure 3 shows that CO2 reacts with the less basic aqueous solution of DMAE (pKa 8.88) by unexpectedly forming carbonate and bicarbonate products simultaneously starting at an early stage of the reaction. Bicarbonate appears to be the favored reaction product at all stages of the reaction, even at high solution pH. The carbonate concentration initially increases at high pH reaching approximately 0.16 CO2 per amine (32 mol% of all amines in solution) and then declines as pH decreases to 0.05 CO2 per amine (10 mol% of all amines in solution). As will be discussed later in this paper, this is contrary to the much stronger guanidine bases (pKa ~15) which initially form only carbonate at high pH.


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Scheme 3. Alternative pathway for bicarbonate formation via carbonic acid deprotonation by the ammonium carbonate with tertiary amines in aqueous solution as illustrated for DMAE This mechanism provides another pathway for the decline of carbonate as bicarbonate increases with progressing levels of DMAE protonation (Figure 3). Scheme 3 may also explain the observed difference (vide infra) in the carbonate/bicarbonate product distribution for the weak (DMAE; pKa 8.88) and strong (TMG; pKa 15.2) bases. Clearly stronger Brønsted basicity (higher pKa) affords lower equilibrium concentrations of the free carbonic acid at all CO2 absorption levels, thereby diminishing the importance of Scheme 3 as a pathway for bicarbonate formation.

studied under similar conditions to DMAE (30 °C, 15 wt% solution) and produced similar product mixtures. At equilibrium, the molar loading of CO2 per amine site was 0.91, and the bicarbonate/carbonate resonance was seen at 160.57 ppm (see Figures S1.1, S1.2 in Supplementary Material). These values, given experimental error limits for peak integration, are very close to those for DMAE.

Primary and Secondary Amines in Aqueous Solution In addition to the potential for protonation by carbonic acid to form (bi)carbonates, as discussed above for tertiary amines, primary and secondary amines are known to directly attack free CO2 to form a zwitterion as shown in Scheme 1. This species rapidly rearrange to the carbamic acid via intramolecular proton transfer, as shown below for the primary amine MEA.

It is worth noting that in the described reaction schemes, the rate limiting step is the nucleophilic attack of water (a weak Lewis base) on acidic CO2 to form carbonic acid. The subsequent steps of amine protonation and carbonate-bicarbonate interchange are considerably faster. The tertiary dimethylamino)ethyl]

diamine bis[2-(N,Nether (DMAEE, 6) was

:

Scheme 4. Rearrangement of the zwitterion to the carbamic acid with primary amines in aqueous solution as illustrated for MEA (pKa 9.16)

In the presence of another basic amine, the carbamic acid may be converted into an ammonium carbamate via intermolecular proton transfer:


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Scheme 5. Deprotonation of the carbamic acid by a free amine to form carbamate with primary amines in aqueous solution as illustrated for MEA (pKa 9.16) known stoichiometry and CO2 reaction chemistry of carbamate formation and conversion to bicarbonate for amines in aqueous solution.

To confirm this expected reaction mechanism, we studied CO2 uptake by 3M (18-23 wt%) aqueous solutions of the primary/secondary alkanolamine pair MEA / methylaminoethanol (MAE) (1-2), as well as 15 wt% (2.4-3.1 M) aqueous solutions of the etheramine pair 1,5-diamino-3-oxapentane (DAOP) / 1,5-bis(methylamino-)3-oxapentane (DMAOP) (45).

Figure 5 shows a graph of the ratio of captured CO2 per MEA or DAOP primary amine group (as quantified by the integral of the carbamate and bicarbonate 13C NMR resonances) over time as CO2 was bubbled through their aqueous solutions at either 30 or 10 °C. After approximately 2 h of bubbling, the maximum amount of reacted CO2 was reached and was 0.70 CO2 per MEA amine group (about 40 mole % carbonate/bicarbonate and about 30 mole % carbamate on a CO2 basis) and approximately 0.85 CO2 (about 70 mol % carbamate about and about 15 mole % carbonate/bicarbonate) per each amino group of DAOP. The large difference in CO2 loading and product equilibrium is influenced by the temperature and will be discussed in the following section. In each case a quantitative conversion of the amine groups into products (either ammonium counterions or carbamate species) was observed; i.e., for MEA, the 30 mole % carbamate value implies that another 30 mole % of amines are present as ammonium counter-ions to the carbamate species (= 60 mole % of amines reacted with CO2 to form ammonium carbamate + 40 mol% reacted with CO2 to form ammonium (bi)carbonates = 100%). The pH of the solution was monitored simultaneously with CO2 uptake (shown for DAOP). The initial solution pH of 12.4 dropped with the formation of carbamate species to 8.2 and reached 7.5 at maximum uptake.

As pure CO2 at 1.0 bar was introduced into the MEA/H2O solution at 30 °C, one product was initially seen in the C=O region of the 13C NMR spectrum at 164.9 ppm (Figure 4, top). The intensity of this peak increased with time and reached a maximum at a loading of approximately 0.5 CO2 groups per amine, at which point it began to decline as a second singlet resonance appeared at 161.5 ppm and grew larger, shifting upfield. At full uptake, the second peak showed higher relative intensity (0.39 CO2 per amine) and was detected at 160.7 ppm, whereas the first (now 164.7 ppm) species represented a smaller fraction (0.30 CO2 per amine). We assign these two peak clusters as bicarbonate/carbonate (160.7 ppm, similar to the ppm shift for tertiary DMAE-H+ bicarbonate/carbonate) and carbamate (164.7 ppm). The backbone carbons of the MEA molecule show related shifting and splitting in the 13C NMR spectrum, as shown in Figure 4, bottom. The HOCH2- carbon shifts upfield, as was observed for the tertiary amine DMAE; however, each resonance also splits into two peak clusters, the first representing a MEA-carbamate anion structure HOCH2CH2NH(COO-) which correlates with the carbamate C=O peak; the second an ammoniumMEA counter-ion structure (HOCH2CH2NH3+).

Based on monitoring of the MEA reaction intermediates, we conclude that at an early reaction stage the nucleophilic primary amine forms exclusively the ammonium carbamate with relatively fast reaction rates (Scheme 1). Only after reaching the theoretical maximum of carbamate formation (CO2/N=0.5), does the less nucleophilic water hydrolyze the carbamate primary product to bicarbonate (see Scheme 6 below), ultimately generating a carbamate/bicarbonate equilibrium mixture at long reaction times (>140 min in Figure 5).

A structurally similar non-hydroxylated primary amine, DAOP (4), reacted similarly to MEA to give the spectra shown in the Supplementary Material Figures S2.2-S2.4 (splitting was initially observed in the 164.0-164.6 ppm carbamate region due to the difunctionality). The decline of this peak after a maximum of ~0.5 CO2 per amine group, and the associated appearance of a bicarbonate-region peak near 160.6 ppm, are completely consistent with the


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Scheme 6. Hydrolysis of the carbamate to form the ammonium bicarbonate with primary amines in aqueous solution as illustrated for MEA (pKa 9.16). Carbamate hydrolysis liberates a free amine which, with continued introduction of fresh CO2 forms either carbamate, which quickly decomposes into bicarbonate, or directly forms the bicarbonate salt via aqueous carbonic acid (in a similar fashion to reaction with tertiary amines) at this reduced pH, as described in Scheme 1. The resulting carbamatebicarbonate equilibrium depends on many parameters, such as amine basicity, amine nucleophilicity, amine concentration, solution temperature, and CO2 partial pressure, and will be discussed in the following sections.

The resulting lower nucleophilicity (at similar basicity) of the secondary amines more readily accepts the small proton from carbonic acid, with the predominant formation of bicarbonate species rather than formation of the N-C carbamate bond. Thus, under similar conditions, primary amines tend to preferentially form carbamate over bicarbonate, resulting in lower CO2 capture per solution volume. Additionally, carbamate products tend to be more thermally stable than bicarbonates, and significantly influence the conditions required to desorb CO2 and regenerate free amine, such as desorption temperature and/or partial pressure of CO2.2,18

The secondary amines MAE (2) and DMAOP (5) show similar reaction mechanisms with CO2 in aqueous solution – initial carbamate formation followed by partial carbamate hydrolysis into bicarbonate and quantitative overall conversion of the amine groups with CO2 into mixed carbamate and bicarbonate ammonium salts. The results are included in the Supplementary Material (Figures S.3.1-S3.5). It’s worth noting that the carbamate/bicarbonate equilibrium for primary and secondary amines at identical conditions is different (discussed in detail in the following sections).

In summary, we have analyzed CO2-amine reaction pathways and equilibria between carbamate and (bi)carbonate species for primary, secondary, and tertiary amines of similar basicity, pKa ~9-10. Behavior-wise, secondary amines can be considered as moderately sterically hindered derivatives of primary amines. However, the relatively lower affinity of secondary amine nitrogens for the carbon of CO2 (lower nucleophilicity) at similar basicity favors proton accepting and shifts the CO2 / 20-amine reaction equilibrium to bicarbonate species. The least nucleophilic tertiary amines do not form a stable NC bond, but can accept a proton (Brønsted basicity) to form only carbonate/bicarbonate species with CO2 in water.

When the CO2-amine equilibrium point is reached at 30 °C, the primary amine MEA (pKa 9.14) tends to form more carbamate species than its secondary analogue MAE (pKa 9.40): 60 mol% total amines in the ammonium carbamate form for MEA versus 30% for MAE. At 10 °C, primary amine DAOP also forms more carbamate than the analogous secondary amine DMAOP (30 and 4 mol% total amines as ammonium carbamates, respectively). At similar conditions of CO2 pressure, temperature, and amine concentration, the carbamate/bicarbonate equilibrium for the primary amines was shifted towards carbamate relative to the analogous secondary amines. This observation may reflect a lower affinity of secondary amine towards the carbon of CO2 due to the slight steric hindrance created by the methyl group on nitrogen.

The efficiency of CO2 uptake on a weight basis is an important criterion for liquid CO2 sorption processes. Polyfunctional amines, due to their high density of reactive sites per weight, can provide higher CO2 capacity at lower amine concentrations than monofunctional amines. Table 2 summarizes the reaction of CO2 with multifunctional amines 4-6 (Table 1), as well as some additional amines having similar pKas, at 1.0 bar of CO2 and 30 0C. Corresponding 13C NMR spectra are available in Supplementary Materials (Figures S4.1—S4.8).


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Table 2. Summary of the aqueous reactions of CO2 with multi-functional amines at 30 °C.

Mol. Wt, density,

Structure

0 a

b.p., C

CO2/Amine group mole ratio

Number of amine groups

Concentration, wt%

2

15

1.42

0.71

2

15

1.80

0.90

2

15

1.82

0.91

2

15

0.92

0.46

2

10

1.09

0.54

2

30

1.18

0.59

2

44

1.10

0.55

3

32

1.62

0.54

4

33

2.13

0.53

3

30

1.33

0.44

2

30

0.85

0.42

CO2/Amine mole ratio

104.1 0.98 1,5-diamino-3-oxapentane

64/4 mm 132.21 0.872

1,5-dimethylamino-3-oxapentane

76/20 mm 160.26 0.841

Bis[2-(N,N-dimethylamino)ethyl] ether

189 112.17 powder

1,4-diazabicyclo[2.2.2]octane

174 60.1 0.899

ethylenediamine

118 60.1 0.899

ethylenediamine

118 60.1 0.899

ethylenediamine

118 103.17 0.955

diethylenetriamine

204 129.1 0.982

triethylenetetraamine

267 129.2 0.984

aminoethylpiperazine

222 130.19 1.061

hydroxyethylpiperazine

246


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A 15 wt% APN solution in deionized water was treated with 1.0 bar of CO2 at 10 °C. Figure 6 (top) represents the progress of the CO2-amine reaction and reaction product distribution as calculated from 13 C NMR taken during the reaction. The APN reaction with CO2 is shown below.

WEAKLY BASIC AMINES IN AQUEOUS SOLUTION (pKa < 9) 3-Aminopropionitrile (pKa 7.14) In the following section, the chemistry of CO2 with weakly basic amines will be analyzed. To eliminate the effects of steric hindrance, a series of primary amines were chosen as representative model compounds of less basic amines: aminopropionitrile (APN (7), pKa 7.14), aminoacetonitrile (AAN (8), pKa 5.43), and aniline (AN (9) (pKa 4.61) (Table 1). When considering effects of basicity, it is worth remembering that this “basicity” not only reflects the tendency of an amine to accept a proton (donate electrons into the proton’s empty 1s orbital), but also the amine’s ability to donate electrons to and bond with the relatively electron-poor carbon atom of CO2. All Brønsted bases are also Lewis bases, but the opposite is not always true19. Thus, for amines with lower pKas, one would expect a lower stability of NH bonds in carbonate and bicarbonate species, as well as less stable N-C bonds in carbamates or carbamic acids.

The starting pH of the APN solution is lower than for the more basic primary amines MEA (11.4 vs. ~12.4). However, the reaction mechanism between CO2 and APN is similar – initial formation of carbamate with significant drop of solution pH followed by conversion of carbamate to bicarbonate. After APN saturates with CO2 at the carbamate stage, subsequent hydrolysis gave a final product distribution of 40 mol% carbamate and 18 mol% (bi)carbonate (total CO2 loading = 58 mol%), and a final solution pH of 7.4 (compared to 8.2-8.4 for MEA and DAOP under similar conditions), which might be too low for efficient bicarbonate stabilization. Due to the low Brønsted basicity of APN, bicarbonate species are less favorable (compare Fig. 6, top with DAOP Fig. 5, bottom) and the final equilibrium mixture favors carbamate.


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Scheme 7. Proposed mechanism for CO2 reactions with the mildly basic primary amines in aqueous solution: carbamate formation (via zwitterion/carbamic acid) followed by its hydrolysis into the ammonium bicarbonate as illustrated for APN (pKa 7.14). Aminoacetonitrile (pKa 5.43) Figure 6 (bottom) shows the CO2 reaction evolution of even less basic aminoacetonitrile (AAN (8)) as a 15 wt% aqueous solution at identical conditions. The starting solution pH is very low (9.0 vs. 11.4 for APN and 12.4 for MEA). It is possible that the reaction equilibrium between CO2 and AAN might be different, as shown in the reaction scheme below: due to its low basicity, the initial carbamic acid/zwitterion species may not be deprotonated by a second molecule of AAN, in which case the carbamic acid may remain in the product mixture.


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Scheme 8. Proposed mechanism for CO2 reactions with the weakly basic primary amines in aqueous solution: equilibrium of the carbamate with the carbamic acid and free amine as illustrated for AAN (pKa 5.43). The 13C NMR of the CO2-AAN reaction products shows a resonance at 163.1 ppm while the carbamate peak of the other products studied here usually appears in the 164-165 ppm range. At equilibrium, the solution pH has plateaued at ~6.0. This low pH can be considered as additional evidence of solution acidity, supporting the presence of a carbamic acid product rather than a carbamate with AAN.

Stronger Bases in Aqueous Solution (Including Amidines and Guanidines) This section will describe the reaction chemistry of strong bases (pKa > ~10) with CO2, including piperidine (PP (10), pKa 10.45) and piperazine (PZ (11), pKa 9.55). Although the pKa of piperazine is similar to that of DMAOP (5), it is included in this section since it exhibits unusual behavior due to its two closely located basic secondary amine groups and is known to function as a promoter for CO2 uptake with other amines.21-24 Selected amidines and guanidines were also studied (see Supplementary Material); data for 1,1,3,3tetramethylguanidine (TMG (12), pKa 15.20) is discussed here.25

Due to the very low basicity of AAN, bicarbonate species are very unstable at ambient temperature, but were detected at 10 0C (13C NMR peak at 160.4 ppm, not shown). The low stability of the reaction product is also reflected by the low reaction yield: the AAN did not react quantitatively with 1.0 bar CO2 at 10 ºC. At equilibrium, its total CO2 loading per amine was 37 mol% (35 mol % carbamate or carbamic acid and 2 mol% (bi)carbonate).

Piperidine (pKa 10.45) Figure 7 shows the reaction evolution of a 15 wt% aqueous solution of piperidine with CO2 at 10 0C. As CO2 was introduced into the amine solution, one product was initially seen in the C=O region of the 13 C NMR spectrum at 163.2 ppm. This peak reached a maximum at a loading of approximately 0.5 CO2 groups per amine, at which point it began to decline as a second singlet resonance appeared at 162.0 ppm and grew larger shifting upfield. At full uptake, the second resonance (now at 160.5 ppm) reached loading of approximately 1 CO2 group per amine, whereas the first 163.2 ppm species had completely disappeared. We assign these two peak clusters as bicarbonate/carbonate (160.5 ppm, similar ppm shift as bicarbonate/carbonate in previously discussed tertiary, secondary and primary amine spectra) and carbamate (163.2 ppm, based on chemical shift and splitting of structural carbons).

Aniline (pKa 4.61) Aniline (AN) (9), with a pKa of 4.61, does not react with CO2 in aqueous solution under similar conditions. Thus, a pKa of about 5.0-5.5 can be considered a minimum required value for amines to react with CO2 at ambient pressure and temperature in aqueous solution without promoters. However, bases with pKa < 5.0, such as aniline, may still react with CO2 under conditions of either significantly lower temperatures, higher CO2 pressures, or both20. In summary, we have found that weakly basic amines with pKa < ~9 do not form stable bicarbonates at ambient conditions. The major aqueous CO2 reaction product for weakly basic amines is the carbamate, which is significantly less thermally stable than carbamates formed with stronger bases, e.g., MEA. The lowered thermal stability of the carbamate can be explained by lower affinity of the less basic amine nitrogen to the CO2 carbon of a carbonyl group (weaker N-C bond).

Figure 7 (bottom) shows a graph of the ratio of captured CO2 per amine group (as quantified by the integral of the bicarbonate/carbonate 13C NMR resonance) over time. After approximately 30 min of bubbling, the majority of the PP molecules in solution reacted with CO2 to form carbamate


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species. Complete carbamate formation caused a solution pH drop from 10.6 to approximately 8.5. At this point, similar to MEA and the other previously discussed amines, the carbamate product began to be converted into bicarbonate. After approximately 2 h, the maximum amount of reacted CO2 was reached at a quantitative ratio of 1 CO2 per PP (now completely in bicarbonate form). The pH of the solution reached 7.10 at maximum uptake. Based on the reaction mechanism and products shown below, piperidine behaves similarly to the primary amines MEA and APN (see Fig. 5, bottom, and 6, top). However, the more strongly basic properties of PP favor protonation of the secondary amine nitrogen over nucleophilic addition to the CO2 carbon, hence the initial carbamate product is fully hydrolyzed to bicarbonate, as indicated in Figure 7 (bottom).


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Scheme 9. Proposed mechanism for CO2 reactions with piperidine (pKa 10.45) in aqueous solution: carbamate formation (via zwitterion/carbamic acid) followed by its hydrolysis into the ammonium bicarbonate Piperazine (pKa 9.8)

162.2 ppm) species represented the major product fraction (0.65 CO2 per PZ) with a total combined loading of about 1 CO2 per PZ molecule.12-14 This quantitative loading of CO2 per PZ molecule requires that the piperazinium carbamate species present must either be intramolecular (as shown below) or involve doubly protonated/doubly carboxylated PZ species that maintain an overall ratio of 1 CO2 per molecule. Alternatively, some carbamates bearing unreacted NH groups (having a 2:1 amine:CO2 molar ratio) could be compensated for by the presence of doubly protonated diammonium carbonate/bicarbonate species (1:2 amine:CO2 molar ratio).

Figure 8 shows the evolution of the CO2 reaction with an aqueous solution of 15 wt% piperazine at 30 0 C. As CO2 was introduced into the amine solution, one product was initially seen in the C=O region of the 13C NMR spectrum at 163.2 ppm (Fig. 8, top), assigned as carbamate. While this peak grew and shifted upfield (presumably due to protonation of the 2nd nitrogen on the PZ molecule), a second resonance appeared at 160.5 (bicarbonate), as shown in the reaction scheme below. At equilibrium, the second bicarbonate resonance (now at 160.4 ppm) was the minor species (0.35 CO2 per PZ molecule) whereas the first carbamate (now


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Scheme 10. Proposed pathways for CO2 reactions with piperazine (pKa 9.8) in aqueous solution: (a) formation of the ammonium bicarbonate via zwitterion hydrolysis and (b) carbamate formation (via zwitterion/carbamic acid) followed by its partial hydrolysis into the ammonium bicarbonate

In contrast to the more basic piperidine, piperazine preferentially forms a carbamate species over a bicarbonate and can be considered as more nucleophilic towards reaction with CO2. This might be caused by the difunctional nature of piperazine. After protonation of the first nitrogen, the second nitrogen is less basic (second pKa value of 5.41, Table 1), but may remain nucleophilic. In this case, one piperazine molecule can potentially participate in the formation of intermolecular reaction species in combination with CO2 bicarbonate and carbamate units on neighboring molecules:

the PZ reacted with CO2 to form carbamate species with a ratio exceeding 0.5 CO2 per PZ molecule. Carbamate formation caused the solution pH to drop to approximately 9.5. At this point, bicarbonate species appeared while the amount of carbamate remained unchanged. After approximately an additional 10 min, the reaction reached equilibrium with a final pH of ~7.1. Thus, in contrast to previously described CO2 reactions with primary and secondary amines (where carbamate formation was followed by conversion to bicarbonate after pH dropped to ~8.5), PZ forms bicarbonate on the second nitrogen with no change in carbamate loading at low solution pH. 1,1,3,3,-Tetramethylguanidine (pKa 15.20)5 Aliphatic amidines and guanidines typically have pKas in the range of 10 - 16 and can be considered as the most basic nitrogen-containing structures. They reversibly react with CO2 at conditions close to ambient (unlike sodium and potassium carbonates, which also react with CO2 but regenerate at ~175 0 C). A series of amidines and guanidines was studied for aqueous CO2 uptake, including 1,1,3,3tetramethylguanidine (TMG (12), pKa 15.2, shown in the reaction scheme below). 2-Methyl-2imidazoline (2-MI, pKa 10.98), 1,4,5,6tetrahydropyrimidine (THP, pKa 12.21), triazabicyclodecene (TBD, pKa 14.47), methyltriazabicyclodecene (MTBD, pKa 14.37), 1,1dimethylguanidine (pKa 14.54, hydroxyethylguanidine (predicted pKa not

The formation of intra-molecular carbamates with ionic interactions between the carboxylate and ammonium sides of one isolated molecule only is not anticipated because it requires a diaxial boat conformation. The richness of potential intermolecular reaction products and configurations is supported by the complex 13C NMR of reacted piperazine backbone methylenes in the range 40-45 ppm (Figs. 8, bottom and S6.1). Figure 9 shows a graph of the ratio of captured CO2 per PZ molecule (as quantified by the integral of the bicarbonate/carbonate 13C NMR resonance) over time as CO2 was bubbled through a 1 M aqueous solution of PZ at 30 0C. After approximately 10 min of bubbling, the majority of


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reported), hydroxypropylguanidine (predicted pKa not reported), and 2-amino-1-ethyl-imidazoline (predicted pKa not reported) were also studied15; in aqueous solution, they show similar CO2 reaction properties as those described below for TMG (see Supplementary Material).

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fractions of free base and protonated TMG (red curves) and carbonate, bicarbonate, and total CO2 absorbed in the solution (blue curves) as CO2:amine mole ratio. It is important to note that TMG as the free base was completely protonated to TMG∙H+ at a very early reaction stage (C=N peak reaching 161.5 ppm after approximately 40 min) in the highly basic aqueous solution. At this time, the TMG reaction with CO2 was still in progress and approximately 80% of reacted TMG molecules had formed carbonate species with CO2 (two TMG∙H+ species per product), with only 20% of the reacted TMG forming the expected bicarbonate species (Figure 11, top). With excess CO2 continuing to enter the solution and TMG already saturated as the protonated form (TMG∙H+), the product equilibrium started to shift from carbonate to bicarbonate (Figure 11, top), resulting in a higher CO2 loading in the solution. 13C NMR can detect this transformation via the higher intensity and upfield shifting of the (bi)carbonate carbonyl peak (Fig. 10, top). At equilibrium, the carbonate species was completely decomposed and the reaction products consisted of bicarbonate species only (13C NMR peak at 160.5 ppm, Figure S6.2) with a corresponding total loading of 1.0 CO2 per TMG molecule. It is worth noting that the solution remained relatively basic (pH = 9.7). Based on the pathway described above, we propose the following reaction mechanism for CO2 reaction with the very strong base TMG in aqueous solution:

As CO2 was introduced into a 3 M (35 wt %), highly basic (pH ~14) aqueous solution of TMG at 30 °C, the TMG C=N (168.1 ppm, Figure 10) shifts upfield indicating protonation of the TMG structure to a guanidinium cation. The first reaction product was initially seen in the C=O region of the 13C NMR spectrum at 168.2 ppm (Fig. 10, top), which we assign to an equilibrating carbonate-bicarbonate species present predominantly as carbonate (for a discussion of this equilibrium, see previous discussion for the tertiary amine DMAE and Figure 2.2). With further introduction of CO2 to the solution, this peak grows and shifts upfield to 160.5 ppm, which indicates a predominantly bicarbonate concentration with trace amount of carbonate. A complex analysis was performed to determine the concentration of protonated TMG species and the carbonate/bicarbonate equilibrium as a function of reaction time based on the trajectory of two 13C peaks initially detected at 168.1 and 168.2 ppm, respectively (Figure 10, bottom). These calculations assume the following positions for the 13C resonances of the species involved: free TMG base, 168.1 ppm; fully protonated TMG, 161.5 ppm; TMG∙H+ carbonate, 168.2 ppm; and TMG∙H+ bicarbonate, 160.5 ppm. Figure 11 shows the


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Scheme 11. Proposed mechanism for CO2 reactions with stronger base TMG (pKa 15.20) in aqueous solution: deprotonation of carbonic acid to form the ammonium carbonate followed by deprotonation of a second molecule of carbonic acid to form the guanidinium bicarbonate molecule). This lower CO2 partial pressure and higher reaction temperature led to a longer reaction time, but only moderately decreased the equilibrium CO2 loading to 92 mol%. This can be explained by the very high pH of the starting solution (~14 vs. 12.5 for amines with pKa ~9-10). From the high pH of the CO2-saturated solution at ambient temperature, we expect a high stability of protonated TMG∙H+ and a high reaction yield at elevated temperatures.

As CO2 is introduced into the aqueous TMG solution, it is expected to form very low concentrations of carbonic acid. The carbonic acid is most likely not detected by 13C NMR, because it is rapidly deprotonated by two molecules of the very basic TMG to form exclusively carbonate species. This reaction pathway continues until the majority of the amine (~90 mol%) is protonated and has formed carbonates with CO2. At this time, only a trace amount of bicarbonate species are detected. As more TMG species are protonated, the solution pH drops, favoring bicarbonate formation. However, bicarbonate is not formed as was previously described in Scheme 1 for amines because all of the TMG species are already protonated. Thus, in contrast to standard ammonium bicarbonate formation through the direct deprotonation of carbonic acid by a free amine base (Scheme 1), very strong bases such as TMG can form bicarbonate through reaction of their carbonate species with carbonic acid, giving two bicarbonate species, as shown in Scheme 3 above.

Effect of Amine Temperature on Amine-CO2 Reaction Equilibrium CO2-amine thermodynamic equilibrium is significantly affected by the solution temperature. Our study showed that the CO2-amine reaction equilibrium is shifted towards carbamate and carbonate, and further to free CO2 and free amine, at elevated temperature. The temperature range at which the CO2-amine reaction is favorable depends on amine type/structure, basicity, and reaction product stabilities.

Figure 11, bottom, shows the extent of the CO2/TMG reaction in 30 wt% TMG solution with 0.1 bar of CO2 (90:10 N2/CO2 gas mixture) at 45 0C as quantified by the integral of the bicarbonate/carbonate 13C NMR resonance over time. After approximately 5.0 h of slow CO2/N2 mixture bubbling, reaction equilibrium was reached by formation of bicarbonate (1 CO2 per 1 TMG

To study temperature effects, seven aqueous solutions of 15 wt% DMAEE (6), DAOP (4), DMAOP (5), AAN (7), APN (8), PP (10), and TMG (12) were saturated with 1.0 bar of CO2 at 10 0C and then heated to 30 °C, 50 °C, 70 °C, and finally 90 °C under a continuous 1.0 bar pressure of CO2. The equilibrium CO2 loadings for these temperatures


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bicarbonate equilibrium to changes of solution pH at elevated temperatures. The pH of the starting (CO2-free) DAOP solution decreases as a function of temperature (Table 3). The lower pH value of the starting solution at elevated temperature leads to a lower solution pH after carbamate formation. At conditions when solution pH is approaching a neutral value of 7.0 after carbamate formation, carbamate decomposition into the even less thermally stable bicarbonate is less favorable.

are shown in Figure 12 and Table 3. The comparison was performed with all amines at similar concentration (~15 wt%) and fixed CO2 pressure to avoid secondary concentration effects; small deviations in molar concentration caused by differences in molecular weight will have a minor effect on the solvent/amine ratio and can be neglected. As shown in Figures 12 and S1.2, the CO2-amine equilibrium for a 15 wt% aqueous solution of the difunctional tertiary amine DMAEE (pKa 9.12) is a function of solution temperature. The highest CO2 loading capacity of 0.98 CO2 per each amine of DMAEE (6) was achieved at the lowest temperature studied, 10 ºC. At higher solution temperatures, CO2 loading drops to 0.91 at 30 ºC, 0.80 at 50 ºC, 0.56 at 70 ºC and 0.16 at 90 0C. It is worth mentioning that at 50-70 ºC all amines are still protonated and the loading drop is associated with bicarbonate decomposition into carbonate. However, at 90 0C, free unreacted amines can be found. Thus, at 90 ºC, the solution pH is too low to maintain a higher reaction yield. This property of tertiary amines can be very beneficial for certain applications, including low-temperature CO2 capture. If a current desorption process relies on partial CO2 desorption at 120 0C, which drives the high energy cost of desorption, the use of tertiary amines may offer a significantly lower desorption temperature and savings of regeneration energy. One can design a CO2 capture system based on variation of temperature (TSA, temperature swing adsorption). The CO2 desorption will be especially effective if the partial pressure of CO2 is dropped as well, as described in the following section.

In many cases, the carbamate-bicarbonate equilibrium in aqueous amine solution can be explained by solution pH as a direct measure of amine basicity. However, amine nucleophilicity, steric hindrance, amine concentration, and CO2 partial pressure also affect the CO2-amine equilibrium, particularly the carbamate-bicarbonate equilibrium. For example, the difunctional secondary amine DMAOP (pKa~9.87) shows a trend for carbamate-bicarbonate equilibrium similar to primary amine DAOP (Table 3). However, the higher basicity of DMAOP and its slightly higher steric hindrance (due to the presence of the methyl groups on the amines) shift the carbamatebicarbonate equilibrium to bicarbonate, which favors higher CO2 loading capacity. The thermal stability of the reaction products formed with the less basic amines such as APN (pKa~7.14) and AAN (pKa~5.43) was lower than that for products from DAOP, DMAOP, DMAEE, and similar amines. This latter finding might be critically important for the development of cyclic CO2 capture systems, since lower regeneration energy might be even more beneficial than high CO2 capacity.

The highest CO2 loading capacity seen with the 15 wt% aqueous solution of the difunctional primary amine DAOP (pKa~9.06) was also achieved at the lowest temperature (10 0C). At higher solution temperature, CO2 loading drops to 0.71 at 30 0C, 0.61 at 50 0C, 0.51 at 70 °C, and 0.36 CO2 at 90 0C (Figures 12, S2.4 and Table 3). In contrast to the tertiary amine DMAEE, for which bicarbonate decomposes into carbonate at higher temperatures, the bicarbonate of the nucleophilic primary amine DAOP decomposes back into carbamate. Heating the aqueous DAOP solution from 10 °C to 30 °C leads to lower CO2 concentration in the solution due to a higher carbamate concentration, which increases from 0.15 at 10 °C to 0.25 CO2 at 30 °C (e.g., from 30 to 50 total mol% of amines in ammonium carbamate species) while (bi)carbonate concentration drops from 0.70 to 0.46 CO2 (from 70 to 46 mol% of amines in ammonium (bi)carbonate species). We attribute this shift of the carbamate-

Figures 12, S5.2 and Table 3 confirm that both reaction products (carbamate and bicarbonate) for APN have low temperature stability. Bicarbonate concentration drops at higher temperatures as a consequence of lower thermal stability at lower solution pH at elevated temperature. Moreover, carbamate species are also unstable at temperatures above ambient. At 50 0C, 64 mol % of amines react with CO2 to form ammonium carbamates (0.32 CO2 per amine); 42 mol % at 70 0C (0.21 CO2 per amine); and only 24% (0.12 CO2 per amine) at 90 0C. The low stability of its reaction products makes weakly basic APN a desirable system for CO2 capture at ambient and sub ambient temperatures, providing for a less energy-intensive regeneration of amine and CO2 at 70-90 0C (versus 120 0C for MEA). However, costs associated with refrigeration would also need to be factored into an evaluation of such a system.


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0C (compared to 0.61 at 50 °C and 0.51 at 70 °C for DAOP).

The CO2 reaction products formed with weakly basic AAN (pKa~ 5.43) are very unstable, and the product yield at 30 0C and 50 0C is low (Figs. 12, S5.3 and Table 3). As described above, the bicarbonate exists only as a trace product. By forming such a very weak N-C bond with CO2, AAN can be considered as a candidate for CO2 capture at sub ambient temperatures (below 0 0C), with effective amine and CO2 regeneration at 50-70 0C (versus 120 0C for MEA).

Piperazine (PZ) forms both carbamate and bicarbonate with CO2 in aqueous solution. Both of these species are stable over a broad temperature range. The CO2 capacity for PZ at 30 0C and 50 0C is 1.0 CO2/PZ, and drops to 0.84 at 70 0C and 0.69 at 90 0C (Fig. S6.2 and Table 3). Nucleophilicity of PZ and the relatively strong N-C bond to CO2 that is formed dominate the reaction equilibrium, providing a majority of carbamates at all temperatures. Similarly to PP, PZ can be considered as a candidate molecule for CO2 capture at elevated temperatures.

In contrast to primary and secondary amines with pKa ~ 9-10, which tend to form carbamates over bicarbonates with CO2 at temperatures between 5090 0C, more basic PP favors bicarbonate species over carbamates over a broad range of temperatures, due to its slightly higher basicity. There are two major benefits of this property: higher CO2 capacity (1.0 versus 0.5 CO2/amine) and the effective capture of CO2 at higher temperatures. For example, CO2 loading at 50 0C for PP remains near 1.0 CO2/amine and drops slightly to 0.78 at 70

The very strong base TMG (pKa ~ 15.2) forms bicarbonate species that are stable over a broad temperature range (Figure 12, Table 3). Its maximum CO2 uptake capacity of 1.0 CO2/TMG was achieved at 30 0C and 50 0C, and drops only to 0.95 at 70 0C and 0.86 at 90 0C. TMG can also be considered a candidate molecule for CO2 capture at elevated temperatures.


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Table 3. Summary of aqueous CO2 reaction with amines at various temperatures. Structure

DAOP

DMAOP

DMAEE

APN

AAN

PP

PZ

CO2/Aminemole ratio

a

total

(bi)carbonate

carbamate

pH of solution

10

0.85

0.70

0.15

12.20

30

0.71

0.46

0.25

12.01

50

0.61

0.28

0.33

11.45

70

0.51

0.18

0.33

10.98

90

0.36

0.06

0.29

10.28

10

0.97

0.95

0.02

12.60

0

T, C

30

0.90

0.82

0.08

12.35

50

0.72

0.60

0.12

11.81

70

0.54

0.32

0.22

11.29

90

0.38

0.12

0.26

10.67

10

0.97

0.97

0

12.00

30

0.91

0.91

0

–

50

0.79

0.79

0

–

70

0.57

0.57

0

–

90

0.16

0.16

0

–

10

0.58

0.18

0.40

11.40

30

0.51

0.10

0.41

11.03

50

0.32

0.03

0.29

10.59

70

0.21

0

0.21

10.12

90

0.12

0

0.12

9.63

10

0.37

0.02

0.35

9.05

30

0.20

0.01

0.19

–

50

0.06

0

0.06

–

70

0

0

0

–

90

0

0

0

–

10

1.01

1.01

0

10.6

30

1.00

1.00

0

–

50

0.98

0.98

0

–

70

0.78

0.78

0

–

90

0.64

0.64

0

–

10

-

-

–

–

30

0.99

0.35

0.64

11.8

50

1.00

0.35

0.65

–

70

0.84

0.31

0.53

–

90

0.69

0.19

0.50

–


starting

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Structure

NH N

TMG a

N

0

T, C

CO2/Aminemole ratio total

a

(bi)carbonate

carbamate

pH of solution

10

1.00

1.00

0

14.0

30

0.99

0.99

0

–

50

0.99

0.99

0

–

70

0.95

0.95

0

–

90

0.86

0.86

0

–

starting

Loading is per amine site (not per molecule) for difunctional amines. capacity due to the formation of less bicarbonate and more carbonate species (0.45 CO2/DMAE at 1.0 bar of CO2, 0.15 at 0.1 bar, and 0.0 at 0.01 bar (Supplementary Material). The different behavior of the 7 M solution can be explained by exceeding the critical amine concentration (amine/solvent ratio) above which the equilibrium formation of divalent carbonate species is favored over monovalent bicarbonate. Similar effects were detected with primary and secondary amines (discussed below), which tend to form carbamate species (2:1 amine:carbon dioxide) over monovalent bicarbonates at high amine concentration. Shifting the CO2/amine equilibrium to carbonate at higher amine concentration may give potential advantages for CO2 sorption process development, e.g., faster reaction kinetics, the option to run a process at higher pH, and lower CO2 desorption energy. However, this shift presents disadvantages of lower CO2 absorption capacity (on a CO2/amine basis) and higher solution viscosity.

Effect of Amine Concentration and CO2 Partial Pressure on Amine-CO2 Reaction The CO2-amine thermodynamic equilibrium varies significantly with changing CO2 partial pressure and amine concentration in the solution. At low partial pressure of CO2, the reaction driving force is lower, which may significantly affect the reaction yield. This effect is especially important for less stable reaction products such as bicarbonate. The effects of CO2 and amine concentration variations were studied for DMAE at 45 °C. As shown in Figure 13, at a given amine concentration (3 or 5 molar), the CO2/amine equilibrium molar loading changes over a wide range from about 0.08 up to 0.92 depending on partial pressure of CO2 in the gas stream. The study of varying CO2 partial pressure was performed by first treating the DMAE solution with 0.1 bar CO2 in N2 (10 mol% CO2/90 mol% N2), then switching the reaction gas to pure (1 bar) CO2, then finally switching to the 0.01 bar mixture (1 mol% CO2/99 mol% N2). At a 3 M amine concentration, the highest CO2 loading of 0.92 per amine was achieved at highest partial pressure of CO2 (1.0 bar) indicating that the majority of amines form bicarbonate species. At 0.1 bar CO2 partial pressure, the observed loading of 0.57 indicates that the bicarbonate-carbonate equilibrium is shifted towards carbonate, even at lower pH. At 0.01 bar CO2, the CO2/amine loading is only 0.18, indicating that the majority of previously chemisorbed CO2 was released into the gas phase and removed from solution as free amine was regenerated by the N2 component. One can build a CO2 capture system based on variation of partial pressure (PSA, Pressure Swing Adsorption), with or without a temperature swing (TSA) as discussed below.

The CO2-amine thermodynamic equilibrium for primary and secondary amines also significantly varies with changing CO2 partial pressure and amine concentration for the reasons discussed above. The effects of low CO2 partial pressure are important for the less stable (bi)carbonate reaction products, but less pronounced for more thermally stable carbamate species formed with primary amines (MEA, DAOP) and secondary amines (MAE, DMAOP). As shown in Figure 13, at a given amine concentration and temperature (3M, 45 °C), the CO2/amine equilibrium loading for the secondary amine MAE changes over the range of 0.45 to 0.74 depending on partial pressure of CO2 in the gas stream. At low CO2 pressure, carbamate is the predominant reaction product while (bi)carbonate formation is less favorable. At low CO2 pressure, the rate of hydration of CO2 to (bi)carbonate anion is very slow. At higher partial pressure of CO2, the carbamate/(bi)carbonate equilibrium is able to shift toward (bi)carbonate to increase CO2 capacity in the amine solution. Thus, at a CO2 partial pressure

It is less intuitive to expect amine concentration to influence reaction yield. While the reaction equilibria for 3 and 5M solutions are very similar (Figure 13), DMAE at 7 M shows a lower sorption


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of 0.1 bar, approximately 76 mol% of the amines form ammonium carbamate species (38 mol% CO2 per amine) while 22 mol% are in the (bi)carbonate form with a resulting total CO2 loading of about 0.60 per amine. At 1.0 bar of CO2, only 32 mol% of the amines are in the carbamate form (16 mol % CO2 per amine x 2) while 42% mol% are in the bicarbonate and carbonate form with a total loading 0.74 CO2 per amine (Figure 13). The maximum molar CO2 capacity of 1:1 CO2/amine for the secondary amine MAE at 45 0C can be theoretically achieved at pressures at or above 1.0 bar.

Page 26 of 46

parameter to measure since it reflects electron donating ability, which depends on both electronic and steric factors. Lacking an H atom and an affinity for the CO2 carbon, tertiary amines act as proton acceptors only and form bicarbonate and carbonate species with CO2 in water. At lower 10/20 amine concentrations, the bicarbonate/carbamate equilibrium may be shifted towards bicarbonate. However, low amine concentrations are not practical for commercial utilization of CO2 capture technologies due to low molarity of captured CO2 and therefore low process capacity. At more practical amine concentrations of above 3 molar, carbamate species are the predominant reaction product between CO2 and primary and secondary alkanolamines and etheramines. Therefore, the performance of primary and secondary amines for CO2 capture is primarily determined by the stability of their carbamate species, while tertiary amine performance is characterized by the stability of their carbonate/bicarbonate species, which is generally lower. We have evaluated the impact of different reaction products on amine performance for CO2 capture under different conditions. It is worth mentioning that the overall efficiency of an amine to capture CO2 also depends on reaction rates. Although we do not discuss intrinsic reaction kinetics here, it is well known that formation of carbamate species is generally faster than formation of bicarbonates.

Figure 13 also compares CO2/MAE and CO2/MEA reaction products at a fixed CO2 partial pressure of 1.0 bar and various amine concentrations at 3 and 5M. At low amine concentration (3M; H2O/amine ratio of about 10:1), reaction equilibrium is shifted towards bicarbonate. As shown in the previous section, higher amine concentrations favor carbamate formation because higher solution pH is maintained, disfavoring bicarbonate formation. This may be a reflection of exceeding the critical amine/solvent ratio in solution, after which formation of carbamate is more favorable than bicarbonates. In more dilute amine solution, the amine is titrated with CO2 more rapidly, causing a more rapid rate of pH decrease and shifting the reaction equilibrium towards bicarbonate. At higher amine concentrations in solution, the transformation of carbamate into (bi)carbonate is slower. Note that the 3 M CO2 loading for the secondary amine MAE is higher than that for MEA (Figure 13, top), which is primarily a result of the higher concentration of bicarbonate species present with secondary amines versus primary amines.

Figure 13 compares CO2/amine reaction equilibria for the analogous series of primary (MEA), secondary (MAE), and tertiary (DMAE) alkanolamines as a function of CO2 partial pressure in the gas phase at a fixed temperature (45 0C) and at two molar concentrations (3 and 5M). Whereas the primary and secondary amines behave similarly over the whole range of CO2 partial pressures and show a weak pressure dependence on CO2 loading, reaction efficiency with the tertiary amine is a strong function of CO2 pressure. At a low CO2 partial pressure of 0.01 bar, only 10-20 mol% of the tertiary amine DMAE reacts with CO2 forming (bi)carbonate species. This finding makes the tertiary amine undesirable for deep purification of CO2-containing gas streams. However, at 0.1 bar of CO2, DMAE captures CO2 in an amount comparable to the 10 amine MEA and the 20 amine MAE, and increases CO2 equilibrium loading to its maximum by 1.0 bar. Since Figure 13 indicates that maximum capacity can theoretically be achieved at 1.0-2.0 bar of CO2, further increasing CO2 pressure beyond 2.0 bar is not likely to lead to higher sorption capacity for tertiary amines. This behavior suggests that the

Comparison of 10 / 20 / 30 amines Detailed in situ NMR studies of amine/CO2 reaction pathways with primary, secondary and tertiary amines in aqueous solution showed complex reaction pathways. Primary and secondary amines may act as bases and/or nucleophiles and form several reaction products with CO2, including carbamates, bicarbonates, and carbonates. The carbamate/bicarbonate equilibrium depends on the conditions of reaction, such as amine concentration in solution, solution temperature, and CO2 partial pressure in the gas phase, as well as on intrinsic properties of the amines, such as Brønsted basicity (the affinity of amine to accept a proton) and nucleophilicity, or Lewis basicity (the affinity of the amine to attack the electrophilic carbon of CO2). Brønsted basicity can be characterized by pKa value, while nucleophilicity is a more challenging


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0.5-0.6 CO2 per amine can be desorbed at 90 0C. Dropping CO2 partial pressure is beneficial for the CO2 desorption process, as can be seen from comparing the 1.0 to 0.1 bar data in Figure 14. Ignoring their generally lower reaction rates, tertiary amines such as DMAE can be effectively utilized for low temperature CO2 capture and offer significant benefits in lower regeneration energy.

optimal CO2 partial pressure range for the utilization of tertiary amines for CO2 capture at 45 0 C lies between 0.1 bar and 2.0 bar. Since equilibrium loadings vary with temperature, the optimum partial pressure range depends on the temperature at which CO2 has to be captured from the gas stream. For example, effective utilization of tertiary amines (e.g., achieving maximum loading capacity) for CO2 capture from hot natural gas may require 5-10 bar of CO2 in the stream.

Variation of basicity

Primary and secondary amines capture significantly less CO2 per mole of amine than tertiary amines due to formation of carbamate species. However, faster reaction rates with CO2 and high capture efficiencies at low CO2 partial pressures, which are beneficial for deep gas purification, are two prime advantages for primary and secondary amines. The high-pressure performance of primary and secondary amines has to be studied in more detail to evaluate their potential for natural gas purification.

Variation of amine basicity leads to significant changes in carbamate/bicarbonate equilibrium and stability of reaction products. Generally, amines with lower pKa tend to form less stable bicarbonate species. However, the affinity of the amino nitrogen for the CO2 carbon of less basic amines is also lower. In previous sections, we have discussed the tendency of aminopropionitrile (pKa 7.14) to preferentially form a carbamate, which is less stable than the carbamate formed from monoethanolamine (MEA) or methylaminoethanol (MAE) (pKa ~9.1-9.4). The very weakly basic aminoacetonitrile (AAN) (pKa 5.43) is still able to attack CO2 but forms a carbamic acid rather than a carbamate. The N-C bond of the resulting carbamic acid is very weak and CO2 can be regenerated by slight heating of the solution.

Figure 14 compares the CO2 capacity of the primary, secondary and tertiary amines MEA, MAE, and DMAE as a function of temperature at fixed CO2 partial pressures (1.0 bar and 0.1 bar) and fixed amine concentration (3 and 5M). The primary and secondary amines show lower CO2 loading capacities at ambient temperature than the tertiary amine, due to predominant formation of carbamates with a 1:2 CO2:amine molar ratio. However, the consistent loading performance of these amines in the range of 45-90 0C makes them potential candidates for CO2 capture at elevated temperatures. Due to formation of the thermally stable carbamate species, the loading capacity at 90 0 C for the primary and secondary amines approaches 0.5 CO2/amine at 1.0 bar of CO2 and between 0.3-0.4 CO2/amine at 0.1 bar of CO2. However, conditions for carbamate decomposition and CO2 desorption at an appreciable rate require heating the amine solution to temperatures above 120 0C with simultaneous decreasing of the CO2 partial pressure (with steam or a purge gas that is inert to the amine) to below 0.1 bar, which is generally very energy intensive.

In order to analyze the effect of basicity on the CO2/amine reaction equilibrium, we compared several amines and guanidines having a basicity range from ~5.4 (AAN) to ~15.2 (tetramethylguanidine (TMG)) (Figure 12). At temperatures above 70 0C, weakly basic AAN (pKa 5.43) does not react with CO2. At lower temperatures, reaction yield increases and AAN By captures ~0.4 CO2 per amine at 10 0C. extrapolation of the AAN/CO2 absorption isobar, the maximum loading capacity of AAN is predicted to occur at -40 0C, which is significantly lower than the freezing point of the absorbent and therefore impossible to achieve. Slightly more basic aminopropionitrile (APN, pKa 7.14) shows absorption capacities higher than those of AAN in the 10 - 90 0C temperature range. Higher amine basicity and nucleophilicity lead to greater thermal stability of reaction products; i.e., decomposition of reaction product beginning at higher temperatures. The effects of basicity can therefore be interpreted by the shift of the absorption isobar to a higher temperature range (for amines with higher pKa) or to lower temperatures (for amines with lower pKas). Thus, the ~1.7 pKa units difference between APN and AAN leads to a

In contrast to primary and secondary amines, tertiary amines more effectively react with CO2 at lower temperatures, capturing up to 1 CO2 per amine through the formation of bicarbonate species. The secondary benefit of bicarbonate formation is its lower thermal stability. Thus, as shown in Figure 14 for DMAE at 1 bar at both concentrations, approximately 0.2 CO2 per amine can be desorbed by heating the solution from 45 to 65 0C (at a constant pressure of CO2) and another


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shift of the absorption isobar of approximately 20-30 0 C.

Page 28 of 46

Conclusions Detailed in situ NMR studies of amine/CO2 reaction pathways with primary, secondary and tertiary amines in aqueous solution were performed for amines (plus amidines and guanidines) having a wide range of basicities. New mechanistic pathways were observed and elucidated. Varying amine basicity leads to significant changes in the carbamate/bicarbonate equilibrium and stability of reaction products. Generally, amines with lower pKa values are less nucleophilic and tend to form less stable carbamate and bicarbonate species.

The primary amine 1,5-diamino-3-oxapentane (DAOP) (pKa 9.07) has a higher CO2 absorption capacity than APN and AAN due to formation of bicarbonate species at low temperatures and stable carbamates formed at higher temperatures. The additional ~2.0 pKa units of basicity for DAOP versus APN shifted the absorption isobar by approximately 30-40 0C and significantly improved CO2 capture performance for DAOP at elevated temperatures. Thus, DAOP, MEA, MAE and other primary and secondary amines with pKas of ~9-10 can effectively capture CO2 at temperatures between 10 and 90 0C.

While tertiary amines act as simple proton acceptors to form bicarbonates and carbonates, primary and secondary amines function as both bases and nucleophiles to form carbamates and (bi)carbonates whose product ratio is a function of reaction conditions, basicity, and amine steric and electronic properties.

Although the tertiary amine bis[2-(N,Ndimethylamino)ethylether (DMAEE) has a similar pKa to DAOP (9.12), it cannot be directly compared with primary and secondary amines due to its exclusive formation of (bi)carbonate. The absorption isobar of DMAEE, as well as those of similar tertiary amines such as dimethylaminoethanol (DMAE) shows a relatively narrow temperature range between 0 and 40 0C, where tertiary amines effectively react with CO2 with high CO2/amine loading. As discussed previously, the benefit of using a tertiary amine is the low thermal stability of its reaction products. In contrast to primary and secondary amines of similar basicity, tertiary amines decompose a significant portion of their bicarbonate products (e.g., desorb CO2) by 90 0C even without reducing CO2 partial pressure.

The performance of primary and secondary amines for CO2 capture is primarily determined by the stability of their carbamate products. Tertiary amine performance is characterized by the stability of their carbonate and bicarbonate products, which is generally lower. Primary and secondary amines tend to react with CO2 similarly at different CO2 partial pressures, showing weak pressure dependence on CO2 loading. In contrast, reaction efficiencies of tertiary amines are a strong function of CO2 pressure, and they unexpectedly form both carbonate and bicarbonate products simultaneously, even at high pH. Primary and secondary amines capture significantly less CO2 per mole of amine than tertiary amines (lower CO2 loading capacities) due to the formation of carbamate species that require two moles of amine per mole of CO2 captured. Nonetheless, their faster reaction rates with CO2 and high capture efficiencies at low CO2 partial pressures are advantageous. In contrast, tertiary amines more effectively react with CO2 at lower temperatures, capturing up to 1 CO2 per amine as bicarbonate species. The secondary benefit of the bicarbonate is its lower thermal stability (lower energy requirement for regeneration). Stronger bases, e.g., guanidines, were found to initially form carbonates which then, even in the strongly basic aqueous medium, decompose to the bicarbonate product that is traditionally reported.

Due to its higher basicity, piperidine (PP) (pKa 10.45) primarily forms a bicarbonate with CO2 in aqueous solution, with a 1:1 CO2/amine ratio. The uptake capacity of PP is significantly higher than that for less basic primary and secondary amines having pKas of ~9-10 at all temperatures studied, and reaches its maximum below 50 0C. The disadvantage of this high efficiency is the formation of thermally stable bicarbonate reaction products (in contrast to the unstable bicarbonates formed by the tertiary amines just discussed), which need to be regenerated at temperatures significantly higher than required for regeneration of less basic primary and secondary amines. The very strong base 1,1,3,3-tetramethylguanidine (TMG) (pKa 15.20) forms the most thermally stable bicarbonate detected so far in our experiments. Its reaction rate and efficiency with CO2 remain very high up to 90 0C, with a loading capacity exceeding 0.8 CO2/TMG at 1.0 bar of CO2


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SUPPLEMENTARY MATERIAL

AUTHOR INFORMATION

1

Corresponding Author

H and 13C NMR spectra of unreacted and reacted amines discussed in the paper as well as results for other primary, secondary and tertiary mono- and multi-functional amines are included in the supplementary material.

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References 12. Einbu, A.; Ciftja, A.F.; Grimstveldt, A.; Zakeri, A.; Svendsen, H.F. “Online analysis of amine concentration and CO2 loading by ATR-FTIR spectroscopy” Energy Procedia 2012, 23, 55-63. 13. Park, H.; Jung, Y. M.; You, J. K.; Hong, W. H.; Kim, J.-N. “Analysis of the CO2 and NH3 reaction in an aqueous solution by 2D IR COS: Formation of bicarbonate and carbamate” J. Phys. Chem. A 2008, 112, 6558-6562. 14. (a) Kortunov, P.V.; Baugh, L.S.; Calabro, D.C.; Siskin, M. “In Situ NMR Mechanistic Studies of Carbon Dioxide Reactions with Liquid Amines in Non-Aqueous Systems: Evidence for Formation of Carbamic Acids and Zwitterionic Species” Manuscript in preparation, (b) Kortunov, P.; Baugh, L.; Calabro, D.C.; Siskin, M. “In Situ NMR Mechanistic Studies of the Reactions of Carbon Dioxide with Liquid Amines: Mixed Base Systems” Manuscript in preparation, (c) Kortunov, P.; Baugh, L.; Siskin, M. “In Situ NMR Mechanistic Studies of the Reactions of Carbon Dioxide with Liquid Amines: Ionic Liquids” Manuscript in preparation. 15. Johnson, S. L.; Morrison, D. L. Kinetics and Mechanism of Decarboxylation of N-Arylcarbamates. Evidence for Kinetically Important Zwitterionic Carbamic Acid Species of Short Lifetime. J. Am. Chem. Soc. 1972, 94, 1323-1334. 16. There are numerous and detailed studies in the literature regarding acidity constants for alkanolamines commonly used in CO2 capture processes. See, for example: (a) Bates, R. G.; Pinching, G. D. “Acidic Disociation Constant and Related Thermodynamic Quantities for monoethanolammonium Ion in Water From 0° to 50 °C” J. Res. Nat. Bur. Std. 1951, 46, 349-352. (b) Littel, R. J.; Bos, M.; Knoop, G. “Dissociation Constants of Some Alkanolamines at 293, 303, 318, and 333 ºK” J. Chem. Eng. Data 1990, 35, 276-277. (c) Antelo, J. M.; Arce, F.; Casado, J.; Sastre, M.; Varela, A. “Protonation Constants of Mono-, Di-, and Triethanolamine. Influence of the Ionic Compositon of the Medium” J. Chem. Eng. Data 1984, 29, 10-11. (d) Hamborg, E. S.; Versteeg, G. F. “Dissociation Constants and Thermodynamic Properties of Amines and Alkanolamines from (293 to 353) ºK” J. Chem. Eng. Data 2009, 54, 1318-1328. (e) Sumon, K. Z.; Henni, A.; East, A. L. L. “Predicting pKa of Amines for CO2 Capture: Computer versus Pencil-and-Paper” Ind.

1. Nirula, S.C.; Ashraf, M. “Carbon Dioxide Separation”, SRI International Report No. 180. 2. Versteeg, G.F.; Van Dijct, L.A.J.; Van Swaaij, W.P.M. “On the kinetics between CO2 and alkanolamines in both aqueous and non-aqueous solutions: An overview” Chem. Eng. Comm. 1996, 144, 113-158 and references cited therein. 3. Jamal, A., Meisen, A., and Lim, C.J., “Kinetics of carbon dioxide absorption and desorption in aqueous alkaonolamine solutions using a novel hemispherical contactor I. Experimental apparatus and mathematical modeling”, Chemical Engineering Science, 2006, 61, 65716589. 4. McCann, N., Phan, D., Wang, X., Conway, W., Burns, R., Attalla, M., Puxty, G., and 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. 5. Jakobsen, J.P., da Silva, E.F., Krane, J., and Svendsen, H.F., “NMR study and quantum mechanical calculations on the 2-[(2-aminoethyl)amino]-ethanolH2O-CO2 system”, J. Magn. Res., 2008, 191, 304-314. 6. Mani, F.; Peruzzini, M.; Stoppioni, P. “CO2 absorption by aqueous NH3 solutions: speciation of ammonium carbamate, bicarbonate and carbonate by a 13C NMR study” Green. Chem. 2006, 8, 995-1000. 7. Ballard, M.; Brown, M.; James, S.; Yang, Q.”NMR studies of mixed amines” Energy Procedia, 2011, 4, 291-298. 8. Garcia-Abuin 1249 Garcia-Abuin, A.; GOmezDiaz, D.; Navaza, J. M.; Rumbo, A. “NMR Studies in Carbon Dioxide - Amine Chemical Absorption” Procedia Engineering 2013, 42, 1242-1249. 9. Pereira, F.S.; Ribeiro de Azevedo, E.; da Silva, E.F.; Bonagamba, T.J.; da Silva, A., Deuber L.; Magalhaes, A.; Job, A.E.; Perez G.; Eduardo R. “Study of the carbon dioxide chemical fixation-activation by guanidines” Tetrahedron 2008, 64, 10097-10106. 10. Matin, N. S.; Remias, J. E.; Neathery, J. K.; Liu, K. “Facile method for Determination of Amine Speciation in CO2 Capture” Ind. Eng. Chem. Res. 2012, 51, 6613-6618. 11. Richner, G.; and Puxty, G. “Assessing the Chemical Speciation during CO2 Absorption by Aqueous Amines Using in Situ FTIR” Ind. Eng. Chem. Res. 2012, 51, 14317−14324.


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Eng. Chem. Res. 2012, 51, 11924-11930. (f) March, J. Advanced Organic Chemistry, 3rd Ed.; John Wiley & Sons: New York, 1985, p. 221 17. For some experimentally determined pKa values for amines used in this study, see: (a) Soloway, S.; Lipschitz, A. “Basicity of Some Nitrilated Amines” J. Org. Chem. 1958, 23, 613-615. (b) Stevenson, G. W.; Williamson, D. “Base Strengths of Cyanoamines” J. Am. Chem. Soc. 1958, 80, 5943-5947. (c) King, J. F.; Gill, M. S.; Ciubotaru, P. “Benzenesulfonyl Chloride with Primary and Secondary Amines in Aqueous Media – Unexpected High Conversion to Sulfonamides at High pH” Can. J. Chem. 2005, 83, 15251535. (d) Hall, H. K. “Correlation of the Base Strengths of Amines” J. Am. Chem. Soc. 1957, 79, 5441-5444. (e) Hall, H. K. “Steric Effects on the Base Strengths of Cyclic Amines” J. Am. Chem. Soc. 1957, 79, 5444-5447. (f) Searles, S.; Tamres, M.; Block, F.; Quarterman, L. “Hydrogen Bonding and Basicity of Cyclic Imines” J. Am. Chem. Soc. 1956, 78, 4917-4920. (g) Xu, S.; Otto, F. D.; Mather, A. E. “Dissociation Constants of Some Alkanolamines” Can. J. Chem. 1993, 71, 1048-1050. (h) Hetzer, H. B.; Robinson, R.A.; Bates, R.G. “Dissociation Constants of Piperazinium Ion and Related Thermodynamic Quantities from 0 to 50 °C” J. Phys. Chem. 1968, 72, 2081-2086. 18. Astarita, G; Savage, D.W.; Bisio, A. “Gas Treating with Chemical Solvents” Wiley-Interscience Publications, John Wiley& Sons, New York 1983, pp. 208-210. 19. Laurence, C.; Gal, J-F. “Lewis Basicity and Affinity scales – Data and Measurement”, Wiley, pg. 7, 2010. 20. (a) Yamada, T.; Lukac, P.J.; George, M.; Weiss, R.G. “Reversible, room-temperature ionic liquids. Amidinium carbamates derived from amidines and aliphatic primary amines with carbon dioxide” Chemistry of Materials 2007, 19, 967-969. (b) Yamada, T.; Lukac, P.J.; Yu, T.; Weiss, R.G. “Reversible, roomtemperature, chiral ionic liquids. Amidinium carbamates derived from amidines and amino-acid esters with carbon dioxide” Chemistry of Materials 2007, 19, 4761-4768. (c) Yu, T.; Yamada, T.; Gaviola, G.C.; Weiss, R.G.. “Carbon Dioxide and Molecular Nitrogen as Switches between Ionic and Uncharged Room-Temperature Liquids Comprised of Amidines and Chiral Amino Alcohols” Chemistry of Materials 2008, 20, 5337-5344. (d) Yu, T.; Weiss, R.G.; Yamada, T.; George, M. “Reversible roomtemperature ionic liquids for gas sequestration” PCT Int. Appl. (2008), WO 2008094846A1 20080807 CAN 149:249600 AN 2008:942211. (e) Orth, J.H. “Preparation of aminocarbamates from diamines and carbon dioxide” PCT Int. Appl. (1997), WO 9729083 A1 19970814 CAN 127:190473 AN 1997:542425 21. Bishnoi, S.; Rochelle, G.T. “Absorption of Carbon Dioxide in Aqueous Piperazine/Methyldiethanolamine” AICHE Journal 2002, 48, 2788. 22. Bishnoi, S.; Rochelle, G.T. “Thermodynamincs of Piperazine. Methyldiethanolamine/Water/Carbon Dioxide” Ind. Eng. Chem. Res. 2002, 41, 604-612. 23. Cullinane, J. T.; Rochelle, G.T. “Kinetics of Carbon Dioxide Absorption into Aqueous Potassium Carbonate and Piperazine” Ind. Eng. Chem. Res. 2006, 45, 2531-2545

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24. Conway, W.; Fernandes, D.; Beyad, Y.; Burns, R.; Lawrence, G.; Puxty, G.; Maeder, M. Reactions of CO2 with Aqueous Piperazine Solutions: Formation and Decomposition of Monoand Dicarbamic Acids/Carbamtes of Piperazine at 25.0 °C. J. Phys. Chem. A 2013, 117, 806-813. 25. Kortunov, P.; Siskin, M.; Sysyn, D.A. “Hydrolysis of Guanidines and Amidines” Unpublished results.


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Figure 1. In-situ NMR setup for studying amine-CO2 reaction chemistry and schematic of NMR probe used for in-situ CO2 uptake studies

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Figure 2.1.

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C NMR spectra of DMAE before (top) and after (bottom) chemical reaction with CO2 at 30 °C

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Figure 2.2. Evolution of CO2/DMAE reaction over time in H2O at 30 0C monitored by 13C NMR. Formation of carbonate/bicarbonate species (top), contour plot of evolution of structural carbons – NCH2CH2OH (bottom right) and carbonate/bicarbonate species (bottom left). Blue and red curves represent the best fit of (bi)carbonate and structural carbon peak trajectories.

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Figure 3. Analysis of CO2 reaction versus time for 3 M aqueous DMAE at 30 °C based on insitu NMR results on Figure 2. Fraction of unreacted free base DMAE (dashed red curve), fraction of protonated DMAE (solid red curve), CO2/amine ratio in carbonate and bicarbonate species (dash and dot-dash blue curves, respectively), total CO2 uptake (solid blue curve), and solution pH (filled spheres).

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Figure 4. Evolution of amine-CO2 reaction for 3 M aqueous MEA, 1.0 bar of CO2 at 30 0C monitored by 13C NMR. Formation of carbamate/(bi)carbonate species (top), evolution of HOCH2- carbon (bottom).

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Figure 5. CO2 uptake, reaction product speciation and solution pH versus time for aqueous solutions of primary amines: 18 wt% of monofunctional MEA at 30 °C (top) and 15 wt% of bifunctional DAOP at 10 °C (bottom). CO2 was introduced at 1.0 bar at a flow rate of 5.5 cc/min.

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Figure 6. CO2 uptake, reaction product speciation and pH in aqueous solutions of 15 wt% APN (top) and 15 wt% of AAN (bottom) as a function of time at 10 °C and 1.0 bar of CO2.

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Figure 7. CO2 reaction with aqueous solutions of 15 wt% piperidine at 10 0C and 1.0 bar of CO2. Formation of carbamate/(bi)carbonate species monitored by 13C NMR (top), reaction product speciation and pH as a function of reaction time (bottom).

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Figure 8. Evolution of CO2/Piperazine reaction products over time in H2O at 300C monitored by 13C NMR. Formation of carbamate/(bi)carbonate species (top), evolution of structural carbons (bottom).

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Figure 9. CO2 uptake, reaction product speciation and pH in aqueous solutions of 1 molar piperazine as a function of time at 30 °C

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Figure 10. Formation of carbonate-bicarbonate species and evolution of structural carbon C=N monitored by 13C NMR for CO2/TMG reaction in H2O at 300C and 1.0 bar of CO2 (top). Same plot in contour form with fitted carbonyl resonance of reacted CO2 and C=N resonance of TMG by blue and red curves, respectively (bottom).

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Figure 11. Analysis of CO2 reaction with TMG in aqueous solution versus time based on in-situ NMR results on Figure 10. Top figure shows the fractions of unreacted free base TMG (dashed red curve), fraction of protonated TMG (solid red curve), CO2/amine ratio in carbonate and bicarbonate species (dash and dot-dash blue curves, respectively), total CO2 uptake (solid blue curve). The bottom figure shows CO2 uptake and pH in aqueous solutions of 30 wt% TMG as a function of time at 45 °C at 0.1 bar of CO2.

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Figure 12. Comparison of CO2/amine equilibrium for series of amines and guanidines as a function of temperature at a fixed CO2 partial pressure (1.0 bar) and concentration in water (15 wt%) with extrapolation shown by black arrows

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Figure 13. Comparison of CO2/amine equilibrium for an analogous series of primary, secondary and tertiary amines as a function of CO2 partial pressure at 45 0C: 3 M in water (top), 5 M in water (bottom)

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Figure 14. Comparison of CO2/amine equilibrium for analogous series of primary, secondary and tertiary amines as a function of temperature at fixed CO2 partial pressure: 3 M in water (top), 5 M in water (bottom)

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Table of Content Graphic

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In Situ Nuclear Magnetic Resonance Mechanistic Studies of Carbon

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