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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 Pavel Valerievich Kortunov, Michael Siskin, Lisa Saunders Baugh, and David C. Calabro Energy Fuels, Just Accepted Manuscript • DOI: 10.1021/acs.energyfuels.5b00985 • Publication Date (Web): 10 Aug 2015 Downloaded from http://pubs.acs.org on August 24, 2015

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IN SITU NMR MECHANISTIC STUDIES OF CARBON DIOXIDE REACTIONS WITH LIQUID AMINES IN NONAQUEOUS SYSTEMS: EVIDENCE FOR FORMATION OF CARBAMIC ACIDS AND ZWITTERIONIC SPECIES 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

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Abstract In a previous study, we reported the use of in situ 1H- and 13C-NMR to elucidate mechanistic pathways for the reaction of carbon dioxide with a broad range of amines (pKa ~4.5-15.5), including alkanolamines of commercial interest, in water. In the aqueous systems of that study, water most importantly functions as a Brønsted acid/Lewis base and as the amine is consumed and pH decreases hydrolyzes the initially formed carbamate species (1:2 CO2:amine stoichiometry), into the alkyl ammonium bicarbonate with a more beneficial 1:1 CO2:amine stoichiometry. This study has been extended herein to amines, amidines and guanidines dissolved in non-aqueous solvent systems such as dimethylsulfoxide, sulfolane, toluene, 1-methyl-2-pyrrolidinone and the ionic liquid 1-ethyl-3-methyl-imidazolium acetate. The use of non-aqueous organic solvents shuts off some CO2 reaction pathways available in aqueous solution. However, more importantly, it opens up new possibilities and reaction pathways for amine based carbon capture. Two important aqueous-system pathways are eliminated: the direct hydration of CO2 with tertiary amines or guanidines to form bicarbonates, and the hydrolysis of carbamates at lower pH to form bicarbonates. In non-aqueous solution, the initial step for the reaction of primary and secondary amines with CO2 is the same as in aqueous solution - nucleophilic attack by the amine nitrogen on CO2. However, additional mechanistic pathways are enabled in non-aqueous solvents, particularly the stabilization of carbamic acid(s) (rather than carbamates) products in certain organic solvents. The formation of carbamates requires no water and is favored by higher amine concentrations and basicities (higher amine pKa). In contrast, carbamic acid/zwitterion formation is favored by lower amine concentrations, higher CO2 partial pressures, lower amine pKa, and selection of more polar organic solvents that promote hydrogen bonding. The new amine- CO2 reaction pathways enabled here by the use of non-aqueous solvents introduce stabilizing interactions between the non-aqueous solvent and the amine- CO2 reaction products, facilitating higher capacity and selectivity for carbon capture than in water solutions. The effects of temperature, amine basicity, solvent electronic structures, and concentration on amine- CO2 reaction products (carbamic acid/zwitterion/carbamate and equilibria between neutral and ion-paired forms) are discussed in detail herein.

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Introduction and Background been reported.15-21 Additional non-aqueous studies on the kinetics of CO2 absorption have been reported with aniline (AN) in carbon tetrachloride and chloroform22, in acetonitrile, methyl ethyl ketone, toluene and xylene23, and in aniline-mixed amine systems24. The kinetics of reactions between CO2 and alkanolamines in both aqueous and nonaqueous solutions have been reviewed by Versteeg25 and modeled by Jamal26.

Continued growth of the global economy, industry, and population, all lead to increasing demands for energy. Traditional fossil fuels (e.g., oil, natural gas and coal) will remain the main sources of energy on the planet. This is likely for the next several decades. The development and utilization of efficient CO2 capture and sequestration technologies could reduce global emissions of CO2 into the atmosphere. All viable technology options that offer improved process efficiencies, such as greater sorbent capacities, lower energy requirements, scale up ease, smaller footprint and reduced cost, are therefore of interest. In particular, liquid sorbent-based technologies may offer short term advantages as drop-in replacements to current technologies, whereas the technologies of solid absorbents are less mature and considered longer term options.

This paper focuses on non-aqueous liquid amine systems for carbon capture and presents an indepth fundamental discussion of the reaction mechanisms of CO2 with various nitrogenous bases having a wide range of basicities (pKa ~ 4.5-15.5; ranging from anilines to guanidines) in nonaqueous solution. Recently developed NMR techniques have provided new fundamental insights into CO2-removal reaction mechanisms27-34. The use of titration35, 32 and FTIR for mechanistic studies has been reported29,36-38 but quantification and real time / flow unit identification of unexpected species is cumbersome. Herein, 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. Our novel use of this technique for the purpose of mechanistic investigation has allowed us to target the identification of more economically efficient and highly selective CO2 scrubbing processes for potential industrial applications, e.g., CO2 scrubbing from flue gas. The mechanistic understanding gained on new amine-CO2 reaction chemistries has enabled the selection and finetuning of absorbent structure, reaction mechanism, and product distribution. These findings represent potentially significant advances for non-aqueous liquid-phase amine-based carbon capture approaches, including increases in overall efficiency, reduced energy, and lower operating cost.

We have previously reported in situ NMR-derived details for the mechanistic pathways associated with the reaction of a broad range of amines (pKa ~4.515.5), including commercially relevant alkanolamines, with carbon dioxide in water1. The existing literature on carbon capture with amines in non-aqueous solvents is sparse and fragmentary, and in almost all cases focuses on kinetics as opposed to the elucidation of mechanistic pathways. Reports include the reaction of CO2 with secondary alkanolamines in polyethylene glycol2,3 and in ethanol and butanol.4 Use of Nmethyldiethanolamine (MDEA)5,6, disopropylamine (DIPA)7, 2-amino-2-methyl-1-propanol (AMP)8, and triethanolamine (TEA) in alcohols, glycols and propylene carbonate9,10; ethylenediamine (EDA) and 3-amino-1-propanol (3-AP) in methanol and ethanol11; diethanolamine (DEA) in ethanol12,13 and isopropanol; cyclohexylamine (CH) in alcohol and toluene solutions14 alkylamines in alcohol, octane, and toluene, and DEA and monoethanolamine (MEA) in sulfolane and alcoholic solutions have

Reaction Chemistry of Amine/CO2 Sorption The chemistry of CO2 derives from the high electron deficiency of the carbon atom bonded to two highly electronegative oxygens. The electrophilicity of this carbon makes it susceptible to nucleophilic attack by various N- and O-donors.

Traditional reaction chemistry for non-aqueous amine/ CO2 reactions with primary and secondary amines is summarized in Scheme 1 for primary amines.25

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Scheme 1. Reaction pathways for CO2 with primary amines in non-aqueous solvents. The amine functions as a nucleophile and attacks the carbon atom of the CO2 molecule to form a zwitterion39, which is the charge-separated resonance form of the isomeric carbamic acid (both species having a 1:1 amine: CO2 ratio). 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: CO2 ratio). In non-aqueous solution, since neither carbonic acid nor free H2O is present to facilitate (bi)carbonate-forming reaction pathways, only carbamate and carbamic acid products are formed.1,25 Aromatic carbamic acids of pyrroles and indoles are known,40-42 and a handful of spectral studies on aliphatic and other aromatic carbamic acids formed from free amines and CO2 under nonaqueous conditions have been reported.43-48 This study is the first extensive structure-property NMR investigation of these species and the relevance of their equilibria with other reaction products for CO2 capture processes.

consequence of these different reaction pathways, and the motivation for the amine work described in this and related publications,1,49,50 is that theoretical CO2 sorption capacity for an amine system is a function of the products that are formed. When ammonium carbamates are the preferred reaction product under non-aqueous conditions, the theoretical CO2 sorption capacity is only 50 mol% per amine (2:1 amine:CO2 ratio). However, if carbamic acids or zwitterions are preferred, the theoretical CO2 sorption capacity doubles to 100 mol% per amine (1:1 amine: CO2 ratio). Since carbamate and carbamic acid species have different relative stabilities and energies of formation, the balance between these two species is tunable by tailoring amine structure and reaction parameters (such as temperature) to provide amine scrubbing systems with the highest potential CO2 capture capacities.

Tertiary amines, unlike primary and secondary amines, are not able to rearrange from zwitterion to carbamic acid via intramolecular proton transfer since there is no proton available on the amine. Thus, they do not generate stable ammonium carbamate products with CO2 in aqueous or non-aqueous solution, although a zwitterion can theoretically be formed.∗ An important practical

Experimental Section Experimental Design Using in situ NMR monitoring, we describe the temporal evolution of product formation for the reaction of CO2 with a series of primary and secondary amines in non-aqueous solution as a function of reaction temperature, amine concentration, basicity, nucleophilicity, CO2 partial pressure, and effect of the solvent. The quantitative results illustrate continuous product formation and decomposition under absorption and desorption conditions, respectively.

∗

The tertiary amine 1,5-bis[2-(N,Ndimethylamino)ethyl[ether, structurally analogous to the secondary and primary etheramines1 and 2 discussed below, showed no reaction by 13C or 1H NMR with 1 atm CO2 at 30 °C under the conditions of this study.

The aminoethers 2-ethoxyethylamine (EEA) and 1,5-bis(methylamino)-3-oxapentane (DMAOP), as 4


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well as several other primary and secondary amines with similar basicity, were evaluated in order to study the effects of amine structural type on CO2 reaction chemistry. Less basic amines, such as amino-nitriles, and stronger bases, such as piperidine, amidines, and guanidines, were also studied (Table 1). These specific amines were chosen for their commercial availability, generally high solubility in organic solvents, and relatively high boiling points. For amines bearing two identical amine groups, the presence of the second amine did not appear to affect behavior at low concentrations compared to monofunctional amines. However, concentrated non-aqueous solutions of difunctional amines showed amine

cross-linking (or gel formation) after reaction with CO2.51,52 The pKa values used herein as a measure of basicity are 25 °C predicted values obtained using ACD/Labs (Advanced Chemistry Development) software V8.14 for Solaris© as reported by Chemical Abstracts Service (SciFinder).53 These values do not specifically represent basicity in non-aqueous solution (as opposed to selection of experimental pKas measured in dimethylsulfoxide (DMSO), which were not available for all of the compounds in Table 1). However, they provide a universal, consistent, and general measurement of basicity across the wide range of substrates used. 6], 98%). 1,1,3-Trimethyl-3-propylguanidine (TMPG, no CAS no. assigned), N-butylguanidine (BG, [46269-1], see text), and N-propylguanidine (PG, [46225-9], see text) were prepared at the University of Florida. Epoxylated polyethylenimine (PEI-MEH, see text) was prepared by treating hyperbranched, ethylenediamine-endcapped polyethylenimine (Sigma-Aldrich Chemical Co. Catalog No. 408719, Mw 800 g/mol, Mn 600 g/mol, [25987-06-8]) with ca. 6 eq. of 1,2-epoxyhexane per amino proton dropwise at 50-90 °C. Materials not described above were purchased from common commercial sources in high purity and used as received.

Materials Materials were purchased from commercial sources (if not noted otherwise, from Sigma-Aldrich Chemical Co.) and used as received: 2(methylamino)ethanol (MAE, [109-83-1], 98+%), 1,5bis(methylamino)-3-oxapentane (DMAOP, [262027-1], 98%, Acros), ), 1,5-diamino-3-oxapentane (DAOP, [2752-17-2], 95%, TCI), 2-ethoxyethylamine (EEA, [110-76-9], 99%, TCI), 2-aminoethanol (monoethanolamine, MEA, [141-43-5], 99+%), 3aminopropionitrile (APN, [151-18-8], 98%, stabilized with K2CO3, TCI), 3,3’-iminopropionitrile (IDPN, [111-94-4], 98%, TCI), aminoacetonitrile (AAN, [54061-4], 95+%, American Custom Chemicals Corp.), aniline (AN, [62-53-3], 99.5+%), piperidine (PP, [11089-4], 99+%), piperazine (PZ, [110-85-0], anhydrous, 99%, Fluka), 1,4,5,6-tetrahydropyrimidine (THP, [1606-49-1], 97%), 1,1,3,3-tetramethylguanidine (TMG, [80-70-6], 99%), 1,1-dimethylguanidine (DMG, [6145-42-2], 96%, Chiron), N-(2aminoethyl)acetamide (N-acetylethylenediamine, AEDA, [1001-53-2], 90%), N-(3-aminopropyl)-2pyrrolidinone (APP, [7663-77-6], Sigma-Aldrich Technical Grade), bis[(2-(N,Ndimethylamino)ethyl]ether ([3033-62-3], 97%), 1ethyl-3-methylimidazolium acetate (EMIM-OAC, [143314-17-4], 97%), catechol (1,2-benzenediol, [12080-9], 99%), pinacol (2,3-dimethyl-2,3-butanediol, [76-09-5], 98%), 2,4-dimethyl-2,3-butanediol ([24892-49-7], 99%), triethanolamine (TEA, [102-71-

Materials in Supplementary Material: 1methylimidazole (1MI, [616-47-7], 99%), 2cyanoguanidine (dicyandiamide, [461-58-5], 99+%), 1-acetylguanidine ([5699-40-1], 98%), 1,5,7triazabicyclo[4.4.0]dec-5-ene (TBD, [5807-14-7], 98%), 1-phenylbiguanide ([102-02-3], 98%), and 1-(otolyl)biguanide ([93-69-6], 98%) were purchased from Sigma-Aldrich Chemical Co. 2-Aminoimidazoline ([19437-45-7]), 2-amino-1-ethylimidazoline (N,N-dimethyl-1piperidinecarboximidamide ([690616-76-3]), N,Ndimethyl-4-morpholinecarboximidamide ([128627600-3]), 4-morpholinecarboximideamide ([17238-663]), 1-piperidinecarboximidamide ([4705-39-9]), N(3-hydroxypropyl)guanidine ([4362-87-2]), were prepared at the University of Florida).

Table 1. Physical properties of amine, guanidine, and amidine sorbents.

Compound No.

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

Structure

5


ACD predicted pKa of a conj. acid

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Compound No.

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Structure

75.11 0.935 158

1 Methylaminoethanol (MAE)

2

3,3’-Iminodipropionitrile (IDPN)

Aminoacetonitrile (AAN)

56.04 0.956 58-66/8-15 mm

3 1,5-Diamino-3-oxapentane (DAOP) 4 2-Ethoxyethylamine (EEA) 5 Monoethanolamine (MEA) 6

7

8

9.40

132.21 0.872 76-78/20 mm 104.15 0.98 64/4 mm 89.14 0.85 105 61.08 1.012 170 70.09 0.9584 185 123.16 1.02 173/10 mm

1,5-Bis(methylamino)-3-oxapentane (DMAOP)

3-Aminopropionitrile (APN)

93.13 9

10

11

1.022

Aniline (AN)


184

9.87

9.07 8.92 9.16 7.14

6.13

5.43

4.61

10.45

Piperidine (PP)

85.15 0.862 106

Piperazine (PZ)

86.14 1.10 (solid) 145-146

9.55 5.41

84.12 1.024 88-89/1 mm

12.21

115.18 0.918 52-54/11 mm

15.20

87.12 unknown unknown

14.54

12 1,4,5,6-Tetrahydropyrimidine (THP)

NH 13

N

N

1,1,3,3-Tetramethylguanidine (TMG)

14 1,1-Dimethylguanidine (DMG)


15

6


Not reported

b

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Compound No.


Structure


1,1,3-Trimethyl-3-propylguanidine (TMPG)


16

14.04

N-Butylguanidine (BG) 101.15 unknown unknown

17

14.05

N-Propylguanidine (PG) 102.14 1.07 123/3 mm

8.92

142.20 1.014 120-123/1 mm

9.85

18 N-Acetylethylenediamine (AEDA)

19 N-(3-Aminopropyl)-2-pyrrolidinone (APP) a

Predicted using ACD/Labs (Advanced Chemistry Development) software V8.14 for Solaris© as reported by Chemical Abstracts Service (SciFinder), 25 °C. bNot found in Chemical Abstracts database.

General NMR Procedure for CO2 Uptake and Desorption laboratory ventilation that set an ambient, e.g. 0 psig (or 1.0 bar absolute) pressure in the NMR tube.

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 as an external standard for accurate temperature and heat release 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

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. Premixed gas sources were used to study effects of CO2 partial pressures below PCO2 = 1.0 bar (e.g., 10 mol% / 90 mol% CO2/N2 mixture, purchased from Matheson Tri-Gas, for PCO2 = 0.1 bar). Amine solutions (typically 15-35 wt% amine) were prepared directly in 10 mm diameter, 8-inch NMR tubes using DMSO-d6 (taken from a recently opened ampule and/or stored in a glove box), toluene-d8, sulfolane (taken from a newly opened bottle stored and handled in a dry box), Nmethylpyrrolidinone (NMP), or 1-ethyl-3methylimidazolium acetate (EMIM-OAC). 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. 0.7-1.5 g of amine and ca. 4 g of solvent. In order to eliminate solvent 7


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influences completely, CO2 uptake was also studied using several neat liquid amines. Results qualitatively showed the presence of similar reaction products for neat amines as seen in the non-aqueous solvents listed above.

C natural abundance variability on analytical results and interpretation, we performed carbon isotope characterization on EEA (4), MEA (5), and CO2. The corresponding δ13C values of -30.9‰, 27.0‰, and -11.0‰ (with respect to a Pee Dee Belemnite standard, 1.111% 13C) indicate that 13C abundance in selected compounds varies over the range 1.077% (for EEA) - 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 the total amount of reacted CO2 versus the total amount of amine present in the solution.

After heating the solution to the desired temperature, CO2 flow was initiated. The solution temperature was controlled by a pre-heated N2 purge (house N2 or liquid N2 vapor) flowing at 1200 L/h through the probe. A thermocouple was mounted 10 mm below the sample. 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 1H and 13C 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.

Per cent CO2 loading was calculated by integration of the 13C NMR carbonyl resonance(s) (usually at ca. 165.0-155.0 ppm) versus –CH3, –CH2–, –CH=, etc. resonances representing the total amine methylene and (if present) methyl or methine 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 per cent CO2 present per molecule and CO2/amine group ratio (correcting when necessary for 2 amines per molecule).

Spectral Acquisition For 13C NMR quantitative analysis of the starting solution and final product(s) of CO2 sorption, 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 starting solutions, 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

1

H NMR spectra were used for monitoring amine concentration in the solvent (amine/DMSO ratio). Both amine and solvent can gradually evaporate in the described flow-through system, especially during desorption at elevated temperatures. 1H NMR spectra were also used to calculate the relative concentrations of carbamic acid and carbamate species in CO2 sorption product mixtures as subsequently described. In principle, 1H NMR can give quantitative product information via the splitting of amine backbone resonances (methylene CH2, etc.) into multiple peaks associated with different CO2 reaction products. In this study, such methods were not used for quantification of reaction products, since they provide only an indirect measurement whereas 13C NMR can directly detect reacted CO2 and all reaction products. However, in water-free non-aqueous solutions, 1H NMR provides key information about CO2-amine reaction products and their evolution and equilibrium.

Spectral Analysis

1

H NMR Spectral Analysis of Carbamate and Carbamic Acid

13

Only the C 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 ~1.11 %. To eliminate a potential effect of

The quantification of carbamate and carbamic acid species in solution using 1H NMR data is not completely straightforward since several assumptions are required. The 1H NMR resonance 8


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of the free amine -NH2 or -NRH before CO2 addition is initially detected at low ppm shift (1.0 – 3.0 ppm), but significantly shifts and splits (for primary amine -NH2) into two peaks as CO2 is added. These peaks typically appear above 8 ppm and between 6-7 ppm at equilibrium. The most downfield peak represents –CO2H protons of the carbamic acid in rapid exchange with the –NH3+ (or -NRH2+) protons of the carbamate counter-ion. The second peak at 6-7 ppm is assigned to the –NH– protons of carbamic acid and carbamate, which may or may not be in a fast exchange with the –CO2H/–NH3+ protons. Calculation of the relative amounts of carbamic acid) and carbamate assumes that no free amine is present to contribute exchanging protons with carbamic acid and carbamate species. It is also assumed that the –CO2H/–NH3+ protons and the – NH– protons are not exchanging on the NMR

timescale, and that no more than two product species (carbamate / carbamic acid) are present for calculation purposes. The relative concentrations of carbamic acid and carbamate in the solution at equilibrium, assuming complete reaction of the amine with CO2 to give either CO2-bearing species or ammonium counter-ions and the absence of free amines in the solution, was calculated based on quantitative analysis of these two peaks. 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 real kinetics enables careful monitoring of reaction sequences and relative kinetics as a function of concentration and temperature.

RESULTS AND DISCUSSION 1.

Secondary Amines in Non-Aqueous Solution

CO2 Absorption 164 ppm. Existing spectral studies of carbamic acids referenced earlier place the carbamic acid C=O resonance at 157-160 ppm.24, 26-28 Thus, based on the 13 C NMR resonance of the equilibrium peak at 158.6 ppm, the DMAOP/CO2 reaction product at the given conditions can be interpreted as a carbamic acid species in fast exchange (through proton transfer) with the carbamate (expected at 165-164 ppm) present in the solution at lower concentration. Integration of the C=O peak versus the DMAOP backbone carbons (Fig. 2, bottom) shows an equilibrium CO2 loading of approximately 0.74 CO2 per each amino group (or 1.48 CO2 per DMAOP). This CO2/amine loading is significantly higher than the theoretical maximum for carbamate formation (0.5 CO2 per amine). In the absence of water, we rationalize exceeding of the carbamate maximum by formation of a carbamate/carbamic acid equilibrium mixture. A CO2 /amine mole ratio of 0.74 corresponds to 50.1 wt % CO2 loading based on the weight of amine and a 7.5 wt % loading based on the combined weight of the amine and solvent.

The secondary alkanolamine methylaminoethanol (MAE (1), pKa 9.40) has a simple molecular structure and a low molecular weight (75.11 g/mol), which render it beneficial for efficient CO2 capture on a weight basis and a prime secondary amine target for this study. However, the 1H NMR resonance of its –OH proton can interfere with the 1 H NMR signals of the –NH proton, whose movement is critical for monitoring during reaction with CO2. Therefore, for initial experiments we instead chose to primarily study its ether dimer, 1,5bis(methylamino)-3-oxapentane (DMAOP (2), pKa 9.87) as a representative secondary amine. The absence of an –OH proton allows analysis of the 1H NMR signals from the –NH proton of the free amine, the –COOH proton of the carbamic acid product, and the –NH2+ proton of the carbamate product cation. After confirmation of the amine purity (98+%) with 13C/1H NMR, a 15 wt% (1.2 M) solution of DMAOP in DMSO-d6 was treated with pure (1.0 bar) CO2 at 30 °C at a flow rate of 5.5 cc/min. At equilibrium, one new 13C NMR CO2 resonance, which was integrated to determine uptake, appeared at 158.6 ppm (Figure 2). Our study of this amine in aqueous solution1 showed a 13C resonance for the carbamate product species at approximately

The 1H NMR -OCH2CH2NHCH3 resonance of the free amine was detected at 1.7 ppm (Figure 2, top). As CO2 was introduced into the solution, the resonance of this amino-proton significantly shifts downfield as a single resonance (Fig. 2, bottom). This peak approaches 10.2 ppm at equilibrium and 9


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represents –OH protons of the carbamic acid in rapid exchange with the –NH2+ protons of the ammonium group cation of the carbamate species. C-H and H-H cross-polarization experiments did not show any correlations of this proton with carbons and protons of the amine structure. Absence of the correlation might be caused by the fast proton migration between molecules and/or by the weak polarization transfer through nitrogen and oxygen of the amine and carbonyl groups. Therefore, the configuration of the reaction products was not established. Based on the total CO2 loading of 0.74 per each amine, we estimate 54 mol% of the amines formed carbamates (1:2

CO2:amine ratio, i.e., 27 mol% of amines as ammonium species and 27 mol% as carboxylate anions) and 46 mol% formed carbamic acids (1:1 CO2:amine ratio). We thus envision the following pathway for the reaction of CO2 with DMAOP in non-aqueous solution: the secondary amine directly attacks free CO2 to form a zwitterion, which may rapidly transfer a proton intramolecularly to form a carbamic acid, or remain in solution in low concentration as a zwitterion or zwitterion pair (Scheme 2):

Scheme 2. Pathways for CO2 reactions with secondary amine DMAOP in non-aqueous solution: direct formation of the zwitterion followed by rapid rearrangement to the carbamic acid.

The carbamic acid would be a desirable reaction product for CO2 capture purposes due to the 1:1 mole ratio of captured CO2 per amine utilized in its formation. A portion of the zwitterion and/or carbamic acid in the solution can then be

deprotonated by reaction with a second equivalent of free amine to produce a carbamate salt having the overall stoichiometric requirement of two moles of amine per one mole of carbon dioxide absorbed (0.5:1 CO2:amine) (Scheme 3):

Scheme 3. Deprotonation of a carbamic acid to a carbamate salt.

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1. Primary Amines Solution

In the reaction described above, the second molecule of amine functions as a Brønsted base by accepting a proton from carbamic acid. The equilibrium for this second reaction can be advantageously controlled to stabilize the carbamic acid product over the ammonium carbamate (vide infra). Employing this strategy in a commercial non-aqueous acid gas removal system would afford the benefit of increasing the molar loading capacity of the sorbent amine above the theoretical maximum of 0.5:1 for a system providing the exclusive formation of ammonium carbamates. The CO2/amine ratios in this work vary between 0.5 and 1.0 depending on the relative concentration of carbamate and carbamic acid species. In a further study, we also show that primary amines have the potential to form dicarboxylated, dicarbamic acid products, and that zwitterion/carbamic acid deprotonation can be accomplished by a separately added, chemically distinct non-nucleophilic nitrogeous bases to give “mixed carbamates” composed of an amine-based carboxylate anion and the protonated form of the non-nucleophilc base as the counterion.49

in

Non-Aqueous

Absorption The CO2 adsorption chemistry of primary amines was initially studied using solutions of the diaminoether 1,5-diamino-3-oxapentane (DAOP) (3, pKa 9.07) in DMSO-d6. However, at DAOP concentrations above 15 wt%, the CO2/DAOP reaction products formed a viscous gel or solid precipitate51,52 which could not be analyzed by solution NMR spectroscopy. The monoaminoether 2-ethoxyethylamine (EEA (4), pKa 8.92) was thus selected instead. The commercially important alkanolamine monoethanolamine (MEA) (5, pKa 9.16), was also investigated briefly. However, EEA is more soluble in organic solvents and lacks the complication of–OH group interference with the – NH2 1H NMR signal. A 3M (26.1 wt%) solution of EEA in DMSO-d6 was treated with pure CO2 at 1.0 bar at 30 °C under conditions similar to those described for DMAOP. Figure 3 shows evolution of the chemical reaction and intermediate products as monitored by in-situ 13 C and 1H NMR. As pure CO2 was introduced into the amine solution, one product was initially seen in the C=O region of the 13C NMR spectrum at 161.5 ppm (Fig. 3, top). Since the predicted 13C NMR resonances of carbamate and carbamic acid species are at approximately 163 and 158 ppm, respectively, the initial EEA/ CO2 reaction product observed at 161.5 ppm can be interpreted as carbamic acid and carbamate species in fast exchange through rapid proton transfer. The close proximity of the observed peak to the expected carbamate value of 163.0 ppm suggests a higher concentration of carbamate species at early reaction times. Initial direct nucleophilic attack of the amino group on the electron deficient carbon of CO2 gives a zwitterion/carbamic acid equilibrium mixture; deprotonation of either of these species by the initial high concentration of free amines produces carbamate. Introduction of more CO2 into the EEA solution increased the intensity of the 161.5 ppm peak with time; it initially shifted downfield to 162.5 ppm (indicating a higher concentration of carbamate species) and then shifted upfield to 159.8 ppm, (suggesting a higher concentration of carbamic acid species). The solution, once depleted of free amine, becomes less basic, thereby allowing the equilibrium to be shifted away from carbamate species to favor formation of the more acidic carbamic acid species.

CO2 Desorption Desorption of the CO2 from the CO2-loaded DMAOP solution described above was studied in two ways. In the first case, the amine sample saturated with CO2 at 30 °C was regenerated to free amine and free CO2 by simple thermal desorption in the absence of fresh CO2, without use of any purge gas, by simply holding at elevated temperature (50 °C, 70 °C, and 90 °C) until equilibrium was reached and no changes in solution composition were observed. The reaction products are stable at ambient temperature and begin to decompose at 50 °C. Complete desorption of CO2 was observed at 90°C (Figs. S5.1, S5.2). In the second case, the CO2-saturated amine sample was desorbed by a N2 purge (10 cc/min) at 30 °C until equilibrium was reached. The solution was then purged to equilibrium at 50 °C, and then 70 °C, at which point all reaction products were gone. Dropping the CO2 partial pressure by purging with nitrogen allows for complete amine regeneration (CO2 desorption) by 70 °C, whereas thermal regeneration leaves a residual CO2 loading of approximately 10% at 70 °C in the presence of a finite partial pressure of CO2.

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We envision the mechanism shown in Scheme 4 for the reaction of CO2 with the primary amine EEA in DMSO or other non-aqueous solvent: The primary amine (4) attacks free CO2 to directly form an unstable zwitterion (B), some fraction of which rapidly rearranges into carbamic acid (C) via

intramolecular proton transfer, but is still in equilibrium with the zwitterion. The carbamic acid/zwitterion species may exist in monomeric form or as ion pairs. Carbamate (D) is formed by deprotonation of either component of this equilibrium mixture by free amine.

Scheme 4. The proposed mechanism of the reaction of CO2 with the primary amine EEA in non-aqueous solvents.

The backbone carbons of the EEA molecule show related shifting in the 13C NMR spectrum, as shown in Figure 3 (middle) for the CH3CH2OCH2CH2NH2 methylene carbons. Both peaks shift upfield while the -OCH2CH2NH2 peak at 73.1 ppm, which is more sensitive to amine reaction, splits into two peaks representing carboxylated amine -OCH2CH2NHCOO- and -OCH2CH2NH-COOH (69.5 ppm) and protonated amine -OCH2CH2NH3+ in fast exchange with remaining free amine -OCH2CH2NH2 (66.9 ppm). The latter peak shifts upfield during the reaction, which indicates a higher concentration of protonated amines and a lower concentration of free amines in solution. The less sensitive CH3CH2O- carbon resonance does not show similar peak splitting.

In contrast to the preferential carbamate formation pathway described in the literature,25, 55 we observe a significant fraction of amines forming a stable carbamic acid and/or zwitterion. The preferred reaction pathway and fraction of amines forming stable carbamic acid (C) or carbamate (D) depends on many parameters (vide infra), such as the nature of the solvent, amine concentration (concentration of free amine in solution), temperature, and CO2 partial pressure. Integration of the CO2 C=O peak at 159.8 ppm versus the EEA backbone carbons (Figure 4, bottom right; numbering scheme is explained in Scheme 4) shows an equilibrium CO2/amine ratio of approximately 0.77, corresponding to a 38.1 wt% CO2 loading based on the weight of amine and a 11.4 wt% CO2 loading based on the combined weight of the amine and solvent. The 1H NMR of the saturated solution (Figure 4, bottom left) shows downfield peaks at 9.8 ppm and 6.6 ppm, which correspond to –OH / –NH3+ and –NH– groups of the reaction products, respectively.

The 1H NMR resonance of the free amine CH3CH2OCH2CH2NH2 was initially detected at 1.30 ppm (Fig. 4, top left). As CO2 is introduced, it significantly shifts and splits into two peaks (Fig. 3, bottom left). These peaks approach 9.8 ppm and 6.6 ppm at reaction equilibrium. The most downfield peak at 9.8 ppm represents –CO2H protons of the carbamic acid in rapid exchange with 12


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the –NH3+ protons of the carbamate counterion. The peak at 6.6 ppm is assigned to the –NH– protons of carbamic acid and carbamate, which may or may not be in a fast exchange with the –CO2H/– NH3+ protons. These discrete, non-exchanging peak assignments allow for calculation of the relative amounts of carbamic acid and ammonium carbamate products in solution (assuming no residual free amine present), using the ratio of the 9.8 ppm peak to the 6.6 ppm peak. Ratios of 1:1 and 3:1 reflect pure carbamic acid and carbamate, respectively. For the 1 H NMR spectrum shown in Figure 4, a peak ratio of 1.75 gives a calculated concentration of carbamic acid amine products of ~45.3 mol% and a concentration of carbamate amine products of 54.7% (27.35% carboxylate anions and 27.35% ammonium cations) for a total of ~72.6 molar % loading of CO2-bearing products, which is in a good agreement with the CO2 loading determined directly from 13C NMR spectra (77.4%).

Scheme 5. Hydrogen-bonded pairs of EEA-CO2 carbamic acids and zwitterions. Such pairs can be partially stabilized by the organic solvent present in the solution (vide infra). 13 C-1H NMR cross-polarization spectroscopy (HSQC single-bond quantum correlation and HMBC multibond carbon-proton correlation experiments) was used to help define the reaction products through detection of the proton location relative to the carbons of EEA. The carbon of reacted CO2 (C=O peak at 159.8 ppm) shows a weak long-range crosspolarization with -O-CH2-CH2-NH- protons (3.4 and 3.1 ppm, respectively), but does not show interactions with –NH- and -OH protons (Fig. S2.5). In fact, neither of the –NH- and -OH protons at 9.8 / 6.6 ppm showed cross-coupling with amine carbons. This can be explained either by the fast migration of –NH and –OH protons between molecules, or by very weak carbon-proton crosspolarization through nitrogen and oxygen. The proton –NH2 of the free amine also did not show long-range correlation through nitrogen or with – CH2-NH2. However, 1H-1H COSY experiments detected a weak cross-polarization of the protons appearing at 3.1 and 6.6 ppm. Based on assignment of the 3.1 ppm peak to the –CH2-NH- proton, we interpret the 6.6 ppm peak as –NH– protons (Fig. S2.4).

To summarize, we observe that EEA in organic solvent at low concentration typically exceeds the theoretical maximum CO2 uptake capacity for the formation of carbamate species, instead forming products approaching a 1:1 CO2:amine ratio. This is potentially significant for applications of such nonaqueous amine solutions in high-capacity CO2 capture. The carbamic acid products may be in equilibrium with the analogous zwitterions, either as free or hydrogen-bonded pair forms (Scheme 5).

The concentration of carbamic acid and carbamate species, both during the non-aqueous CO2/amine reaction period and at equilibrium, depends on many parameters, including solvent properties, CO2 partial pressure, amine concentration, and solution temperature. Whereas the effects of the latter three parameters can be interpreted in a straightforward manner, the effects of solvent on the relative concentrations of carbamic acid and carbamate are unexpected and not fully understood. EEA in DMSO-d6 (3M concentration) forms a high concentration of carbamic acid with CO2 at 30 °C. However, a 3M solution of EEA in toluene preferentially forms carbamate species with CO2 under identical conditions. The initial reaction product in toluene was detected at 162.8 ppm (close to the predicted carbamate shift) and slightly shifted upfield during 13


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the reaction, reaching 162.1 ppm at equilibrium (Figure 5, top and Supplementary Material). The total CO2 loading in toluene solution of 0.58 CO2 per EEA also supports the high carbamate concentration observed. These results are reproduced with other amines having pKa >9 (vide infra). Without exception, highly polar DMSO solvent gives product mixtures richer in the neutral acid product, while less polar toluene gives predominantly carbamate salt. Thus, the composition of the carbamic acid/carbamate product mixture is counter to that expected based purely on solvent polarity (as will be discussed subsequently, sulfolane exhibits a product composition much closer to that observed for toluene than DMSO).

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2. WEAKLY BASIC AMINES (pKa 70 wt%. In order to analyze CO2 loading at 70 wt% APN with high resolution solution NMR, a small amount of fresh DMSO-d6 was added to dissolve the solid precipitate. CO2 loading was not affected by this procedure since gas release was not detected (gas bubbles and pressure build-up were not observed). No attempts were made to dissolve the precipitated CO2 reaction products at 90 wt% and 100 wt% since final CO2 loading may have been affected by changes in solvent concentration. Amine- CO2 product precipitation is likely due to the ionic nature of the carbamate species and lesser solubility provided by the nitrile group of APN (in comparison to the EEA ether group).

Figure 9 shows CO2 loading curves and speciation for all other amine concentrations studied. At low amine concentrations (e.g., higher concentrations of DMSO), the majority of the EEA forms carbamic acid species, with a small amount of carbamate detected. At amine concentrations above 50 wt%, carbamate is the predominant reaction product. This observation is confirmed by (i) an equilibrium CO2/amine ratio of 0.50-0.70, (ii) a 13C NMR C=O resonance towards 162.0 ppm (downfield from the carbamic acid region towards the carbamate region), and (iii) a high ratio, approaching 3.0, of OH/-NH3+ protons to –NH- protons in the 1H NMR spectrum. In neat EEA, the majority (82 mol%) of the amine forms carbamate products, leading to a CO2 loading only slightly greater than 0.5 per amine. Therefore, at higher concentrations of EEA, a bimolecular reaction with second order kinetics is favored and the carbamate product, with a 1:2 CO2:amine ratio, is favored. At low amine concentration, however, the equilibrium of the bimolecular reaction, which is controlled by concentration, is slower and carbamic acid predominates. These results confirm the crucial role of amine concentration in solution on CO2-amine reaction pathways. The DMSO stabilizes the carbamic acid and favors pathway (C) (Scheme 4, top) over carbamate (D) (Scheme 4, bottom). At high solvent concentration (e.g., low amine concentration), a greater fraction of amine can be stabilized as carbamic acid.

The range of APN concentrations at which reaction products are soluble in DMSO-d6 characterizes the trend of product equilibrium. Increasing the APN concentration from 10 to 50 wt % dropped carbamic acid content from ~ 84 to 36 mol%. As indicated previously, the presence of excess free amine in the solution drives proton transfer from carbamic acid and allows for subsequent bimolecular (2nd order) reaction to form carbamate as amine concentration increases. Table 2 summarizes the above concentration studies for EEA and APN.

A similar study was performed with APN at 10-100 wt% (Figures S6.10-S6.12). Results are qualitatively

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Table 2. Summary of the reaction of CO2 with EEA and APN dissolved in DMSO-d6 at various concentrations. 13

EEA conc. in DMSO (wt%)

C NMR C=O Peak (ppm)

10 15 30 50 70 90 100

158.5 158.8 159.9 160.8 161.4 161.8 162.0 13

APN conc. in DMSO (wt%) 10 15 30 50 70 90 100

1

CO2/EEA Molar Ratio

H NMR + -OH/-NH3 Peak (ppm)

H NMR -NHPeak (ppm)

Integral Ratio of Downfield 1 H peaks

Carbamic acid mol%

Carbamate mol%

0.88 0.85 0.75 0.69 0.65 0.63 0.58

9.97 9.35 9.97 9.63 9.82 9.63 9.19

6.66 6.60 6.60 6.46 6.35 6.15 5.76

1.31 1.60 1.69 2.05 2.17 2.43 2.78

73 54 49 31 26 17 6

27 46 51 69 74 83 94

H NMR -NHPeak (ppm)

Integral Ratio of Downfield 1 H peaks

Carbamic acid mol%

Carbamate mol%

7.01 7.29 6.85 6.61 6.55

1.27 1.52 1.73 2.02 2.53

76 59 47 32 13

24 41 53 68 87

1

1

C NMR C=O Peak (ppm)

CO2/EEA Molar Ratio

H NMR + -OH/-NH3 Peak (ppm)

158.0 158.9 159.7 160.8 161.2

0.93 0.86 0.75 0.67 0.58

10.0 10.21 9.91 9.69 8.56

1

precipitation of reaction products

showed that CO2 pressure exerts only a minor effect on the CO2-amine reaction pathway but significantly affects the CO2-amine equilibrium.

6. Effects of CO2 Partial Pressure on Amine-CO2 Reaction Equilibrium Lower CO2 pressure is expected to provide a lower driving force for CO2 to form a chemical bond with the amine. To screen the magnitude of partial pressure effects, we treated three identical solutions of 3M EEA in DMSO-d6 with either 0.20, 1.00, or 10.04 bar of CO2 at (30 0C). At the lower 0.20 and 1.00 bar pressures of CO2, as described above and shown in Figs. 3 and S7.4-S7.6, EEA initially formed a mixture of carbamic acid and carbamate, which was preferentially transformed into carbamic acid at reaction equilibrium with a total loading of 0.730.77 CO2 per EEA.

A more detailed study of CO2 partial pressure effects was conducted with MEA (of similar basicity to EEA), and the less basic APN. At a given amine concentration, the CO2/amine equilibrium loading in MEA/DMSO solution changes over the range of 0.45-0.72 at 45 0C and 0.33-0.52 at 90 0C as the CO2 partial pressure is changed from 0.01 to 1.00 bar (Figure 10). At low CO2 pressure, carbamate is the predominant reaction product, whereas at higher partial pressure of CO2 the reaction equilibrium shifts toward carbamic acid. Thus, at 0.01 bar CO2 and 45 0C, approximately 90 mol% of the MEA forms carbamates, with no carbamic acid detected, resulting in a total CO2 loading of ~0.45. At 0.1 bar of CO2, only 64 mol% of the amines are in the carbamate form; another 36 mol% of the MEA forms carbamic acid for a total loading of 0.68. At these conditions, the CO2/amine reaction equilibrium is controlled by other factors besides pressure (temperature, amine concentration, and, likely, the ability of the organic solvent to stabilize particular reaction products). The maximum possible CO2 capacity for MEA may be achievable at 45 0C, but at pressures significantly higher than 1.0 bar.

When CO2 was introduced at 10.04 bar, the initial reaction product was observed at 161.0 ppm, confirming the presence of both carbamate and carbamic acid (Fig. 5 (middle) and Figs. S7.1-S7.3). However, the equilibrium CO2 peak was detected at 159.1 ppm (vs. 159.8 ppm at 1.0 bar of CO2), suggesting a higher concentration of carbamic acid species. At equilibrium, approximately 80 mol% of EEA formed carbamic acids and another 20 mol% formed carbamates (10 mol% as carboxylate anions and 10 mol% as ammonium cations) with a total loading of 0.90 CO2 per EEA. This initial evaluation 21


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The pressure dependence of the CO2-amine thermodynamic equilibrium for APN/DMSO was significantly larger than for MEA/DMSO. At low partial pressures of CO2, the lower driving force for reaction affects equilibrium, particularly affecting less stable reaction products.

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comparably polar solvent to DMSO, yields a product distribution much closer to that of toluene than DMSO, suggests that observed solvent effect cannot be straightforwardly attributed to polarity. To directly probe solvent effects on the CO2amine reaction, 3 M solutions of EEA in DMSO-d6 and toluene were treated with 1.0 bar CO2 at 30 0C. As previously described and shown in Figure 3, EEA/DMSO-d6 forms a mixture of carbamic acid and carbamate at early stages of the reaction (161.5 ppm) and preferentially forms carbamic acid at reaction equilibrium (159.8 ppm; 54 mol% of amine as carbamic acid and 46 mol% as ammonium carbamate; total loading ~0.77 CO2 per EEA). In contrast, the EEA/toluene solution preferentially forms carbamate species with CO2 (Figure 5, top, Figs. S8.1-S8.3). The initial product was detected at 162.8 ppm (in the carbamate region) and shifted only slightly upfield during reaction to reach 162.1 ppm at equilibrium, indicating a high concentration of carbamate species. The total CO2 loading of 0.58 CO2 per EEA also indicates high carbamate concentration (~84 total mol% of amines as carbamates) and low carbamic acid concentration (~16 mol% of amines).

As shown in Figure 10, the CO2/APN equilibrium loading (3 M solution) changes over the broad range of 0.15-0.90 at 45 0C and 0.05-0.40 at 90 0C as the partial pressure of CO2 is varied from 0.01 to 1.00 bar. At low CO2 pressure, both amine-CO2 reaction products (carbamate and carbamic acid) are very unstable. At higher partial pressures of CO2, the reaction yield is higher and CO2 capacity significantly increases. At a CO2 partial pressure of 0.01 bar and 45 0C, the equilibrium CO2 loading was ~0.15 per amine, and the 13C NMR product resonance (161.3 ppm), suggested a carbamate/carbamic acid product mixture. At 0.1 bar CO2, the 13C NMR product resonance shifts upfield to 160.5 ppm, suggesting a higher concentration of carbamic acid, with an equilibrium loading of 0.50 CO2 per amine. At 1.0 bar equilibrium is shifted further towards carbamic acid (159.3 ppm) and shows, as a result, a high CO2 loading of 0.90 CO2 per amine. Again, under these conditions, the CO2/amine reaction equilibrium is controlled by other factors. It may be possible to achieve 45 0C maximum CO2 capacity for APN at relatively low pressures.

EEA solutions were also prepared in sulfolane (tetramethylenesulfone), 1-methyl-2-pyrrolidinone (NMP), toluene-d8, and the ionic liquid 1-ethyl-3methylimidazolium acetate (EMIM-OAc) (Table 3) at concentrations from 10 wt% to 90 wt% (Figs. S8.5-S8.7). If non-aqueous solvents influence the CO2-amine equilibrium, the greatest effect would be expected at low amine concentration (e.g., high solvent concentration). Table 4 summarizes details of these reactions under dilute conditions (10-15 wt% amine) (see Supplementary Material for further details). EEA/DMSO exhibited the highest CO2 capacity of these studies, (~0.88 CO2), with the majority of the amine in the carbamic acid form. EEA/toluene-d8 and EEA/sulfolane show lower capacities of approximately 0.58 and 0.64, respectively, with the majority of the amine in the carbamate form. CO2 loading for EEA/NMP was 0.86 CO2 per amine, with approximately 72 mol% of the amine as carbamic acid and 28 total mol% as carbamate.

7. Effects of Organic Solvent on AmineCO2 Reaction Equilibrium Given the numerous sorbent properties and reaction conditions affecting CO2-amine reaction equilibria, changing the nature of the non-aqueous solvent (in contrast to aqueous solutions) offers additional opportunities to manipulate amine-CO2 chemistry. Solvent effects on the stabilization of certain reaction products were anticipated. Organic solvents with higher dipole moments, such as DMSO, were expected to preferentially stabilize charged ionic species such as carbamate, while less polar solvents such as toluene were expected to favor neutral carbamic acid species. However, we observe the opposite product speciation trend with DMSO and toluene. As stated above, combined with the observation that sulfolane, which is a

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Table 3. Physical properties of non-aqueous solvents.a Mol. Wt. (g/mol), Density (g/mL), 0 Boiling Pt. ( C)

Dipole b Moment (D)

Dielectric Constant

Polarity

Solubility Parameter

78.13 1.10 189

3.96

46.6

44.4

13.0

99.13 1.028 202

4.1

32.2

36

11.0

4.69

44

41

N/A

Sulfolane

120.17 1.261 285

0.4

2.38

9.9

8.9

Toluene

92.14 0.865 110-111

N/A

N/A

N/A

N/A

Solvent

DMSO

NMP

1-Ethyl-3-methylimidazolium acetate (EMIM-OAc)

170.21 1.03 ~250 (decomposes)

a

c

Smallwood, I. M. Handbook of Organic Solvent Properties. Halsted, New York: 1996. b25 °C. cWhere water = 100.

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Table 4. Summary of the reaction of EEA and APN with CO2 in various non-aqueous solvents at low concentrations. 13

EEA conc. (wt %)

Solvent

C NMR C=O peak (ppm)

10

DMSO-d6

158.5

10

NMP

158.9

30

sulfolane

161.5

15

toluene

162.6

15

EMIM-OAc

160.2

50

EMIM-OAc

160.5 13

APN conc. (wt %)

Solvent


10

DMSO-d6

158.1

10

NMP

158.3

10

sulfolane

160.6

10

toluene

160.7

10

EMIM-OAc

160.6

30

EMIM-OAc

160.6

CO2/EEA molar ratio (acid:carbamate ratio) 0.88

1

(76:24) 0.86 (72:28) 0.64 (28:72) 0.58 (16:84) 1.20 (120:0) 0.90 (80:20) CO2/EEA molar ratio (acid:carbamate ratio) 0.93 (86:14) 0.86 (72:28) 0.64 (28:72) 0.56 (12:88) 1.07 (107:0) 1.02 (102:0)

1

H NMR H NMR Integral ratio Carbamic Carbamate + -OH/-NH3 -NH- peak of downfield acid mol% mol% 1 peak (ppm) (ppm) H peaks 9.97

6.66

1.31

73

27

9.31

6.70

1.50

60

40

9.21

6.05

2.5

13

87

9.95

6.18

2.62

11

89

14.46

7.43

2.81

120

0

13.21

6.93

1.24

80

20

1

1

H NMR H NMR Integral ratio Carbamic Carbamate + -OH/-NH3 -NH- peak of downfield acid mol% mol% 1 peak (ppm) (ppm) H peaks 10.00

7.01

1.27

76

24

9.17

7.07

1.50

60

40

7.09

6.24

2.50

14

86

8.78

6.93

2.07

30

70

14.94

7.61

1.01

107

0

14.57

7.37

1.43

102

0

The 1H NMR spectroscopic methodology used to determine product ratios in the previous sections (see Experimental Section and Supplementary Material) breaks down when EMIM-OAC is used as a solvent, requiring alternative methodology. Data for EMIM-OAC is discussed below.

These results clearly indicate a strong solvent effect on the observed product distribution. If this effect were simply a function of solvent polarity, all of the polar solvents would be expected to give a product distribution similar to that of sulfolane, i.e., favoring the charge-paired carbamate salt over the uncharged acid. Indeed, three of the four polar solvents (DMSO, NMP, EMIM-OAC) exhibit the opposite trend, and the least polar solvent (toluene) gives a product distribution comparable to sulfolane.

The carbamic acid stabilization mechanism operative for DMSO and NMP may involve hydrogen-bonded interactions between the –CO2H group of the carbamic acid and the DMSO/NMP solvent X=O group (X = C or S, Scheme 13).

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N

O HO C HN

N O

O

O N

OH

N

O Scheme 13. Stabilization of a carbamic acid via hydrogen-bonded interactions with the X=O groups of DMSO and NMP.

O

O C

O N H

Scheme 14. Stabilization of EEA carbamic acid product with EMIM-OAC.

In Figure 11 (50 wt% EEA in EMIM-OAc), integration of the 13C NMR C=O peak at 160.5 ppm versus the amine backbone carbon backbone peaks (69.5, 65.7, 40.8, 15.1 ppm) gives a high loading of 0.90 CO2 per EEA. At lower concentrations of EEA in the ionic liquid (15-30 wt%), additional mechanisms for CO2 uptake and C=O product peaks were observed (dicarbamic acid formation and imidazolium C-carboxylation). These are discussed in related report.50 As compared to EEA/EMIM-OAc, EEA at similar or higher concentrations in other non-aqueous solvents shows lower CO2 loadings with greater concentrations of carbamate species. Figure 12 shows a comparison of 13C NMR-derived CO2 loading capacities and product speciation for EEA in toluene-d8 and EMIM-OAc at various amine concentrations (1.0 bar CO2, 30 °C). Concentrationvariable data for EEA/NMP is quantitatively similar to that for EEA/DMSO-d6 (Figure 9), whereas EEA/sulfolane behavior resembles that of EEA/toluene.

Clearly, toluene cannot stabilize the carbamic acid in this fashion. Surprisingly, sulfolane does not appear to stabilize the carbamic acid via the hydrogen-bonding interaction shown above. The reason for this unexpected result is not immediately obvious, but may be due to the electronic structure of the sulfolane group. Ionic liquids can be considered to function as polar organic solvents. For EEA in EMIM-OAc (15 wt%), the 1.20:1 CO2/amine ratio calculated by 13C NMR of 1.20:1 (vide infra), indicates that each molecule of EEA fully reacts with at least one CO2 molecule. Under these conditions it is assumed that all of the EEA forms carbamic acid products, and that the most downfield proton in the 1H NMR spectrum represents the -CO2H proton of carbamic acid only. Interestingly, the carbamic acid proton in the 1H NMR spectrum of the EEA/EMIM-OAc solutions is shifted further downfield than for other solvents to 13.2-14.5 ppm (shown for 50 wt% in Figure 11). This can be explained by the strong Hbonding interaction between the carbamic acid – CO2H proton and the negatively charged oxygen atom of the EMIM-OAC acetate anion. In this case, the acetate anion can be considered as a promoter for carbamic acid stabilization and formation (Scheme 14). If the conjugate anion of the carbamic acid is a stronger base than the acetate anion, it can exchange/displace it and form a stronger ion pair with the imidazolium cation.

Generally speaking, at higher amine concentrations, EEA and similar amines tend to form carbamate species through a second order/bimolecular reaction step, rather than forming carbamic acids (which result from direct nucleophilic attack on CO2 or rapid proton transfer from the zwitterion). Interestingly, sulfolane and toluene, which do not support stabilization of the carbamic acid at low amine concentrations in the manner described for DMSO and NMP, show a very weak dependence of the CO2-amine reaction equilibrium on amine concentration. In fact, neat liquid EEA shows approximately the same CO2:amine loading as EEA in toluene or sulfolane under similar conditions (and, for the case of

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toluene, approximately the same product distribution). This observation suggests that nonpolar, non-aqueous solvents such as toluene do not influence CO2-amine reaction chemistry, although they play a practical role in maintaining low solution viscosity. In contrast, polar non-aqueous solvents, such as DMSO, NMP, and in particular EMIM-OAC, significantly affect CO2-amine reaction chemistry via selective stabilization of certain reaction products. The effect is especially pronounced at high solvent concentration (e.g., low amine concentration) and vanishes at high amine concentration. As described above, curiously, sulfolane, which is similar in polarity to DMSO and NMP, behaves like toluene.

Page 26 of 55

benchmark chemical shift difference at high levels of intermolecular H-bonding interactions of OH groups of ethanol. As shown in Fig. 13, decreasing the EtOH concentration in cyclohexane to as low as 5 wt% (1:10 EtOH:CHX molar ratio) only modestly decreases the δ(OH)-δ(CH2) chemical shift difference to 1.02 ppm (-OH group is still downfield). Only at approximately 2 wt% of EtOH (1:50 EtOH:CHX molar ratio) were ethanol molecules sufficiently diluted by cyclohexane to significantly reduce intermolecular Hbonding. Under these conditions, the 1H NMR of the δ(OH) is shifted upfield relative to δ(CH2). The δ(OH)-δ(CH2) value of -0.52 ppm is close to the predicted value (CambridgeSoft ChemDrawTM Ultra 11.0) for non-interacting ethanol molecules and can serve as a benchmark chemical shift difference for δ(OH)-δ(CH2) at low levels of intermolecular Hbonding interactions of the OH groups of ethanol.

Hydrogen Bonding The studies described above in organic solvents show a pronounced influence of the solvent on equilibrium reaction products. Again, at high concentration of DMSO or NMP in solution (when the effect of the solvent is maximized), amines tend to preferentially form the carbamic acid on reaction with CO2. A significantly smaller fraction of the amines form the carbamic acid when sulfolane or toluene is used as a solvent.

After benchmarking of the δ(OH)-δ(CH2) shift of EtOH at high and low levels of H-bonding interactions, solutions containing EtOH in DMSOd6, toluene-d8, sulfolane-d8, and NMP over the molar ratio range of 1:50 - 10:1 were prepared. A small amount of non-interacting cyclohexane was used to calibrate the 1H NMR spectra. Figure 13 shows the 1H NMR resonance of the δ(OH) group relative to δ(CH2) of EtOH dissolved in the various organic solvents. At low concentrations of EtOH (minimal intermolecular ethanol H-bonding), the 1H NMR shift of the δ(OH)-δ(CH2) is 0.88-0.90 ppm for EtOH in DMSO-d6 and 0.99-1.02 ppm for EtOH in NMP. These values are similar to the δ(OH)-δ(CH2) shift of ethanol experiencing a high level of Hbonding interactions. As such, this result shows that diluted EtOH molecules are still influenced by H-bonding interactions, this time with DMSO and NMP. On the other hand, toluene-d8 reduces H-bonding interactions with ethanol molecules. At a 1:10 molar ratio of EtOH:toluene-d8, the δ(OH)-δ(CH2) was 0.43 ppm, which is very close to the benchmark δ(OH)-δ(CH2) value determined for low level Hbonding interactions described above. Interestingly, at lower EtOH concentration in toluene-d8, the 1H NMR of the δ(OH) further shifts upfield and the δ(OH)-δ(CH2) chemical shift reaches -2.27 ppm, implying no interaction with toluene or between ethanol molecules. Ethanol in sulfolane shows reduced H-bonding at even higher ethanol concentration. For example, the 1H NMR shift δ(OH)−δ(CH2) is 0.45 ppm at 1:1 EtOH:sulfolane and -0.28 ppm at 1:4 EtOH:sulfolane, and further drops to -2.12 ppm at

It was reasoned that organic solvents capable of hydrogen bonding can stabilize a carbamic acid. 1 H NMR was used to screen the hydrogen bonding interactions of the organic solvents under study. First the 1H NMR of ethanol (EtOH) dissolved in selected organic solvents at various concentrations was measured, specifically the δ(OH) and δ(CH2) proton chemical shifts referenced against tetramethylsilane (TMS), and the difference of δ(OH) and δ(CH2) proton chemical shifts. Changes of the δ(OH)-δ(CH2) reflect the extent of intermolecular interactions of the OH group with other molecules in the solution. Cyclohexane (CHX), incapable of hydrogen bonding, was used as an internal 1H NMR reference standard. To estimate intermolecular hydrogen bonding interactions between EtOH molecules, several solutions containing different concentrations of ethanol in cyclohexane were prepared. At high ethanol concentration, hydrogen bonding interactions between EtOH molecules are anticipated. For the solution containing 95 wt% EtOH in cyclohexane (38:1 EtOH:CHX molar ratio), the δ(OH) proton chemical shift is downshifted relative to the δ(CH2) protons of EtOH by 1.73 ppm (Figure 13). The δ(OH)-δ(CH2) value of 1.73 is a 26


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absorption capacity are less pronounced than effects of amine concentration. The weight-percent CO2 capacities of EEA dissolved in DMSO, NMP, sulfolane and toluene are similar and primarily depend on amine concentration. Compared to the other non-aqueous solvents, the ionic liquid EMIM-OAC has a greater enhancing effect on CO2 loading at a fixed amine concentration. Its ability to stabilize a high concentration of the carbamic acid product leads to significantly higher CO2 loading capacities. Thus, with an ionic liquid solvent, less amine is required to capture the same amount of CO2, offering potential beneficial reductions in viscosity, corrosivity, and process expense. For example, a 30 wt% EEA ionic liquid solution captures approximately 15 wt% CO2 per solution weight, whereas this level of CO2 loading can only be achieved at amine concentrations of 45-50 wt% in the other non-aqueous solvents. Qualitatively similar results were obtained with the less basic amine APN in DMSO, NMP, sulfolane, toluene and EMIM-OAc at various concentrations. The APN/CO2 reaction products remained in solution at low amine concentrations. Under these conditions, the more polar non-aqueous solvents DMSO, NMP, and particularly EMIM-OAc significantly affected CO2-amine reaction chemistry via stabilization of the carbamic acid reaction products. In a similar fashion to the results for EEA as discussed above, this effect for APN is especially pronounced at high solvent concentration (e.g., low amine concentration) but vanished at high amine concentration.

lower ethanol concentration. Sulfolane tends to disrupt intermolecular H-bonding in ethanol molecules and does not H-bond with ethanol itself. Explanation of the observed difference in carbon capture product distribution is not straightforward, but can be rationalized as follows: NMP and DMSO show H-bonding interactions with the OH group of ethanol. There is no potential for H-bonding of ethanol with non-polar toluene, so the observed Hbonding of ethanol in toluene at high concentration is due to self-associated ethanol attributed to the “hydrophobic effect”.63-65 The ethanol molecules avoid the non-polar environment of the toluene by self-association. Sulfolane does not show Hbonding interactions with the OH group of ethanol. This cannot be attributed to the hydrophobic effect, but is attributed to the semipolar nature of the sulfolane bonding structure.66 It does not follow the octet rule and as such (based on the divalency of oxygen) the sulfur atom of the sulfone, with twelve valence electrons, has an effective +2 charge on the S atom and has a -1 charge on each oxygen; a “semipolar bond”. This ionic-type polarity stabilizes the carbamate salt more effectively than the uncharged carbamic acid. The observed similarity in stabilization of sulfolane and toluene of the carbamate vs. NMP and DMSO stabilization of the carbamic acid is therefore not rationalized by simple solvent polarity arguments (dielectric constant-dipole moment; see above), but can be explained by the relative ability of these solvents to H-bond with the carbamic acid product. Therefore, DMSO and NMP hydrogen bond with the OH group of ethanol to a much greater extent than sulfolane or toluene. Consistent with this finding, DMSO and NMP stabilize the carbamic acid via H-bonding, much more so than sulfolane and toluene. Moreover, sulfolane and toluene seem to repel molecules containing OH-groups, such as ethanol and carbamic acids. Sulfolane helps to stabilize and isolate ionic carbamate species, and toluene, by its lack of interaction, allows preferential formation of carbamates. The weight percent CO2 uptake capacity for the various solutions depicted in Figures 9 and 12 (weight CO2 per weight of combined amine + solvent) is shown in Figure 14 and indicates that solvent identity has a lesser effect on CO2-amine reaction equilibrium at higher amine concentration (e.g., lower solvent concentration). Amines dissolved in all of the non-aqueous solvents absorb more CO2 than the carbamate maximum anticipated from theory (0.5 mole% per amine). However, effects of solvent identity on solution

Except with EMIM-OAc, precipitation of the APN/CO2 reaction products was observed for neat APN and solutions containing >70 wt% APN. The conditions for which precipitation were observed correspond to the high-carbamate region of the carbamic acid-carbamate equilibrium. Intentional precipitation of reaction products between gaseous CO2 and a liquid amine could potentially be used for the development of CO2-capture processes based on phase separation.

8.

Comparison with Aqueous Solution

Figures 15 and S9.1 show the CO2 loading capacity and speciation of EEA in aqueous solution at various concentrations after treatment with 1.0 bar CO2 to equilibrium at 30 °C. Although different reaction products ((bi)carbonates and carbamates rather than carbamates and carbamic acids) are formed in aqueous solution1, total CO2 loading is comparable to that of EEA in DMSO or NMP 27


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solution (Figure 9) and lower than for EEA in EMIM-OAc (Figure 12, bottom). In a previous report1, bicarbonate products were shown to be less thermally stable than carbamates. In this study, carbamates (and carbamic acids) are found to be less stable in non-aqueous solvents than in conventional aqueous amine systems. Thus, the non-aqueous systems offer the benefit of lower regeneration energy. However, aqueous amine solutions have the benefit of low viscosity of the CO2-rich solution and low solvent cost. These alternative options must be weighed when selecting a carbon capture approach.

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concentrations (Fig. S9.2), with a total CO2 capacity significantly lower than that achievable in nonaqueous solutions at the same amine concentration. Table 5 illustrates the loading advantages of anhydrous CO2 reactions with weak bases compared to reactions performed in the presence of water. This can be explained by a relatively low affinity of weakly basic amines (pKa < 8) to function as Brønsted bases, i.e., to accept a proton and form a carbamate and/or bicarbonate product with CO2 in aqueous solution. However, these amines may have a relatively larger affinity to function as Lewis bases (nucleophiles) towards CO2 under non-aqueous conditions, and form zwitterion/carbamic acid products with higher total CO2 loading.

APN in aqueous solution preferentially forms carbamate species over a broad range of amine

Table 5. Comparison of CO2 sorption results for EEA and APN in DMSO-d6 vs. H2O at various concentrations (1.0 bar CO2, 30 °C). Amine conc. (wt%)

10 15 30 50 70 90

CO2/EEA

CO2/APN

molar ratio

molar ratio

in DMSO-d6

in H2O

in DMSO-d6

in H2O

(carbamate/

(carbamate/

(carbamate/

(carbamate/

carbamic acid)

(bi)carbonate)

carbamic acid)

(bi)carbonate)

0.88

0.86

0.93

0.54

(0.12/0.76)

(0.14/0.74)

(0.07/0.86)

(0.46/0.08)

0.85

0.78

0.86

0.52

(0.15/0.70)

(0.22/0.56)

(0.14/0.74)

(0.48/0.04)

0.75

0.69

0.75

0.52

(0.25/0.50)

(0.31/0.38)

(0.25/0.50)

(0.48/0.04)

0.69

0.59

0.67

0.52

(0.31/0.38)

(0.41/0.18)

(0.33/0.34)

(0.48/0.04)

0.65

0.57

0.58

0.52

(0.35/0.30)

(0.43/0.14)

(0.42/0.12)

(0.48/0.04)

0.63

0.53

precipit.

precipit.

(0.37/0.26)

(0.47/0.06)

28


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herein tend to form a separate liquid phase from the unreacted amine. Solubility can be controlled by proper selection of amine structure and organic solvent.

Non-aqueous CO2 reactions in DMSO were studied for other primary amines having basicities similar to EEA (MEA, DAOP, Nacetylethylenediamine (AEDA, 18), and aminopropylpyrrolidinone (APP, 19) (Table 1)). The CO2-amine reaction equilibria for these amines were very similar to that for EEA under similar conditions. Multifunctional amines such as DAOP tend to precipitate from solution or form gels after reaction with CO2, even at low concentration in solution. Large agglomerates can be formed through the intermolecular association of carbamate products or carbamic acid hydrogenbonded pairs when multifunctional molecules are used.51,52

9. Intramolecular Stabilization Carbamic Acids by Functional Groups

of

Given the above findings, a next logical step for the design of highly efficient non-aqueous amine sorbents is the incorporation of stabilizing functional groups with high dipole moments (e.g., C=O or S=O) into amine molecular structures. In this situation, the amino group will react with CO2 via nucleophilic addition to form a carbamic acid (or zwitterion) that is stabilized by the functional group. To evaluate this approach, the behavior of AEDA and APP was studied further. The following reaction schemes are expected for AEDA (top) and APP (middle/bottom), respectively (Scheme 15):

At high concentrations (typically above 50 wt%) in non-aqueous solution, the reaction products of some monofunctional primary amines studied

Or

Scheme 15. Intramolecular stabilization of a carbamic acid by hydrogen-bonded interactions in AEDA (top) and APP (middle, bottom).

amine, the stabilizing polar functional group, and the organic solvent.

CO2 reaction for both amines was probed in DMSO-d6 over the concentration range of 10 - 100 wt% amine (1.0 bar CO2, 2-3 h to equilibrium, 30 °C). Reaction of CO2 with neat AEDA and APP was expected to reflect the intrinsic tendency of amines to preferentially form carbamate at high concentration, while reaction in dilute solution was expected to reflect the combined effect of the

In both cases, at low amine concentrations the majority of amines react with CO2 forming carbamic acid species, with a small amount of carbamate detected. However, CO2 loading for AEDA/DMSO solution (Supplementary Material, Figs. S10.1, S10.3, top) is very similar to CO2 loading for EEA/DMSO solution (Figure 9) over the broad concentration

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range. Thus, reaction equilibrium appears primarily controlled by the DMSO solvent, and the acetyl group of AEDA is not providing the necessary properties to intramolecularly stabilize a product having 1 CO2 per amine molecule.

dicarbamic acids.49 The dicarbamic acid species may have the first acid group stabilized by the pyrrolidinone carbonyl oxygen, and the second by the solvent, which should be absent in toluene or sulfolane (Scheme 16).

Interestingly, the APP/DMSO solution shows significantly higher molar CO2 loading than the EEA/DMSO solution (Figures 16 and 9, respectively). At 10 wt% APP concentration, CO2 loading exceeds 1 CO2 per APP and reaches 1.06 (Table 6, Figs. S10.2, S10.3 bottom). In addition to the two “regular” -CO2H and -NH downfield protons, usually observed around 9.90 and 6.60 ppm, a 1H NMR peak at 18.57 ppm (not accounted for in product ratio calculations) was detected and confirms a very complex structure for the CO2amine-solvent reaction product. A total loading of 1.06 CO2 per amine may be explained by invoking a product mixture having 94% of the amines present as monocarbamic acids, with two carbonyl groups attached to the remaining 6 mol% of amines to give

CO2 loadings above 1.0 per amine were also observed with EEA in EMIM-OAc at low amine concentrations.28 However, at higher EEA concentrations in EMIM-OAc, total CO2/amine loading decreases, with the carbamic acid decomposing to carbamate species. For APP/DMSO, total CO2/amine loading and carbamic acid concentration remains higher than for the reference system EEA/DMSO at higher amine concentrations. This confirms that the pyrrolidinone group of APP exerts an extra stabilizing influence on the carbamic acid reaction product. Table 6 summarizes the CO2/DMSO-d6 system at various concentrations.

Scheme 16. Proposed mechanism for dicarboxylation of the primary amine APP facilitated by intramolecular and intermolecular stabilization of carbamic acid groups by H-bonding interactions.

Table 6. Summary of the reaction of CO2 with APP in DMSO-d6 at various concentrations. APP conc. in DMSO (wt%)

13

1

1


CO2/APP molar ratio 13 ( C NMR)

H NMR + -OH/-NH3 peak (ppm)

158.31

1.06

8.64

30

159.41

0.88

9.17

6.74

18.39

75

25

50

160.35

0.85

8.51

6.68

none

70

30

70

160.70

0.81

9.44

6.65

none

63

27

90

161.46

0.77

7.31

6.30

none

54

46

10

H NMR -NHpeak (ppm)

Downfield 1 H NMR peak (ppm)

Carbamic acid mol% 1 ( H NMR)

Carbamate mol% 1 ( H NMR)

6.70

18.57

100

12

30


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separated by a distance well-suited to the size of the linear O=C=O molecule, such as catechol, pinacol (2,3-dimethyl-2,3-butanediol), and 2,4-dimethyl-2,4pentanediol (Scheme 18).

10. CO2 Physisorption in Non-Aqueous Solution During the above-described investigations of CO2 reaction with the dinitrile amine IDPN in DMSO-d6, a new 13C NMR resonance at 125.05 ppm was noted. Based on the ppm shift value, this signal cannot be interpreted as a chemically bound carbonyl group, but rather as free CO2 weakly interacting with certain functional groups via Van der Waals forces (13C NMR shift of free CO2 = ~124.5 ppm). Therefore, we associate the signal at 125.05 ppm with physisorbed CO2 in solution. Based on the absence of physisorbed CO2 in pure DMSO-d6 solutions and with the more basic mono-nitrile amine APN, the physisorption mechanism is proposed to involve the two nitrile groups of the IDPN molecule (Scheme 17). The loading of physisorbed CO2 reached 0.08 CO2 per IDPN at equilibrium (Supplementary Material).

N C

O C O

At 1.00 bar of CO2, 30 0C, and ~10 wt% diol solution concentration, the observed molar CO2 loadings were 0.09 per molecule for catechol (3.3 wt%), 0.10 CO2 per pinacol (3.8 wt%), and 0.11 CO2 per pentanediol (3.3 wt%). A solution of the tertiary amine triethanolamine (TEA, N(CH2CH2OH)3, pKa 7.77) in DMSO-d6 did not react with CO2, but physisorbed approximately 0.13 CO2 per TEA (13C NMR peak at 125.18 ppm). Very high CO2 physisorption was detected with hyperbranched poly(ethylenimine) (repeat unit –CH2CH2NHnominal Mn= 600) in which the majority of the primary and secondary amines were capped with 1,2-epoxyhexane to form tertiary amines bearing one or two –OH groups per amine (PEI-MEH). After treating PEI-MEH solutions (15 wt% in DMSO-d6) with 1.0 bar CO2 at 30 °C, up to 0.24 moles of physisorbed CO2 per polymer –OH group was detected by 13C NMR (125.21 ppm). This is in addition to regular CO2 chemisorption at the remaining uncapped primary and secondary amines of the PEI, 162-160 ppm).

N C

N H

In these experiments, the sorbent molecules were dissolved in DMSO-d6 solvent before exposure to CO2 gas, and the observed physisorbed CO2 cannot be explained as gaseous CO2 dissolved in the organic solvent (CO2 was not physisorbed in pure DMSO-d6). The CO2 detected in solution may be captured via Van der Waals interactions or hydrogen bonding of both CO2 oxygen atoms to two –OH groups of the abovementioned diol molecules. Similar interactions may exist between CO2 and molecules with a single –OH group; however, the sorbing effect is maximized when two –OH groups are present on one molecule and separated by the appropriate length to effect bidentate chelation.

Scheme 17. Proposed mechanism for CO2 physisorption by hydrogen-bonded interactions with C≡N groups. Relatively high amounts of CO2 physisorption were also detected in DMSO-d6 with other molecules having two –CN or –OH groups per molecule. In contrast, no physisorption was observed in non-aqueous solutions of sorbents having only one –CN or -OH group. The highest physisorption yield was achieved with diols having a fairly inflexible structure and two –OH groups

OH OH

O C O

Scheme 18. Proposed mechanism for CO2 physisorption with diols via H-bond interactions with –OH groups.

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aqueous solution.1 In the absence of water, the formation of bicarbonate or carbonate salts is not possible, and the exclusive formation of ammonium carbamate is expected as the sole reaction product of CO2 with amines capable of acting as both a Lewis and Brønsted base.

SUMMARY In this study, in-situ NMR techniques have been used to confirm the commonly accepted primary mechanism for the formation of carbamates from primary or secondary amines and CO2 under nonaqueous conditions. This process proceeds through an unstable zwitterion intermediate which rapidly rearranges via proton transfer to a carbamic acid. The zwitterion and/or carbamic acid is then rapidly deprotonated by a second equivalent of free amine to produce a carbamate salt with the overall stoichiometric requirement of two moles of amine per CO2 (0.5:1 CO2:amine). In the absence of excess free amine in the solution, the presence of excess CO2, or the use of weak Brønsted base amines (pKa < 8), carbamic acids with a 1:1 CO2:amine ratio predominate and increase CO2 loading capacity. Zwitterion/carbamic acid species were not independently detected in this study as spectrally isolated from carbamates due to fast reaction equilibria. However, spectral analysis of equilibrium product mixtures indicates the presence of carbamic acids, based on the high CO2/amine molar ratios and peak shifts seen in 13C NMR in conjunction with additional speciation carried out via 1H NMR. Carbamic acids have not been observed under commercial aqueous amine scrubbing conditions.

In a water-free environment, sufficiently stable carbamic acid products are formed which, under appropriate conditions of concentration, pressure, molecular structure and temperature, only undergo a limited degree of subsequent Brønsted reactivity to form carbamates. This unexpected chemistry provides a beneficial increase in the amount of absorbed CO2 (molar uptakes in excess of the theoretical maximum predicted from complete carbamate formation). It was shown herein that amines with low basicity (pKa < 8) typically react with CO2 to preferentially form equilibrium mixtures favoring carbamic acid products in non-aqueous media. The product mixtures for these weakly basic amines contain both carbamic acid and carbamate species, and are significantly less thermally stable than carbamate products formed with stronger primary and secondary amine bases, e.g. EEA (pKa 8.92). This is explained by a lower affinity of the less basic amine nitrogen for the electron deficient carbon of CO2. More basic piperidine (PP, pKa 10.45) forms carbamates in higher concentration than EEA and APN. For even more basic tetrahydropyrimidine (THP, pKa 12.21), an equilibrium CO2 loading of approximately 0.95 CO2 per THP molecule is observed at 30 °C. This high CO2:THP ratio can only be explained by the predominant formation of carbamic acid or zwitterion species. Combined with the high concentration of protonated amidines observed in the 13C NMR spectrum (Figure 6), this provides powerful first time evidence for the formation of a stable zwitterion in this carbon capture system.

The use of primary and secondary amine sorbents in non-aqueous solvents can produce CO2:amine molar ratios exceeding the theoretical ammonium carbamate maximum of 0.5:1. In addition to this benefit, non-aqueous solvents provide reduced corrosivity and amine regeneration energy needs. Thus, such non-aqueous systems can potentially provide operational and economic benefits for future commercial carbon capture technology. Conventional thinking regarding the nature of amine-CO2 reaction chemistry focuses on Brønsted acid-base chemistry, largely ignoring the fact that CO2, as a thermodynamic endpoint, can only become reactive in the presence of a nucleophile (Lewis base). Water, alcohols and amines are known to be sufficiently nucleophilic to activate CO2. Water activation leads to the well-known carbonic acid/bicarbonate/carbonate pH-dependent equilibrium mixture. Since most amines are stronger nucleophiles than water, they preferentially attack the carbon atom of CO2, forming a zwitterion intermediate which can rapidly tautomerize to carbamic acid, which is subsequently neutralized to ammonium carbamate, even in

The CO2-amine thermodynamic equilibrium is significantly influenced by solution temperature. At elevated temperatures, the CO2-amine reaction equilibrium is shifted towards carbamate, and then further towards free CO2 and free amine. The temperature range at which the CO2-amine absorption reaction is favorable depends on the type, structure, and basicity of the amine as well as the thermal stability of its CO2 reaction products. CO2 loading of primary and secondary amines drops with increasing reaction temperature, which is associated with the decomposition of carbamic acid products to lower-capacity carbamates. 32


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Weaker bases generally have lower loading capacities at all temperatures. As amine basicity is increased, more stable carbamate species are formed, and CO2 loadings at elevated temperature may still be higher than 0.5 CO2 per amine. This presents opportunities to carry out carbon capture at elevated temperatures.

the exception of sulfolane, the propensity for stabilizing the acid product is generally enhanced with polar solvents, as expected. The inability of sulfolane to stabilize the acid is attributed to a less favorable electronic structure that inhibits hydrogen bonding.

At low amine concentrations (e.g., higher concentrations of non-aqueous solvent), carbamic acid products dominate with a minor amount of carbamate formed. At amine concentrations above about 50 wt%, carbamate is the predominant product in non-aqueous solution; however, above this concentration, the CO2-rich solutions become very viscous and at concentrations of 70 wt% or higher precipitation is frequently observed.

CONCLUSIONS In summary, secondary and primary amines at lower amine concentrations in non-aqueous solution favor carbamic acid products. These acids are the product of a 1:1 stoichiometry with CO2 with the amine vs. the traditional 0.5:1 stoichiometry of carbamate formation. The formation of carbamic acids is inherently advantageous for CO2 capture because it provides higher CO2 capacity. However, at higher amine concentration and/or basicity, the carbamic acid product is deprotonated by reaction with a second equivalent of free amine to produce a carbamate. The formation of carbamates from carbamic acids can be partially circumvented by the use of appropriate organic solvents that preferentially stabilize the carbamic acid through H-bonding interactions. Stabilization of a carbamate is achieved by use of an appropriate solvent that repels H-bonding interactions. Ionic liquid solvents are most effective in this respect. Carbamic acid stabilization may alternatively be achieved using an amine sorbent bearing a second stabilizing functionality.

At low CO2 pressures, both types of amine-CO2 products (carbamates and carbamic acids) are very unstable. Higher partial pressures of CO2 increase reaction yield, and thus significantly increase the CO2 capacity of the amine solution. Generally, the CO2/amine reaction equilibrium is controlled by non-pressure factors, specifically temperature and amine concentration, and likely also by the ability of the organic solvent to stabilize particular reaction products. At higher concentrations in solution, EEA and similar amines tend to form carbamate species with a 1:2 CO2/amine ratio via a second order/bimolecular reaction process (rather than carbamic acids which form by direct nucleophilic attack or rapid proton transfer from a zwitterion intermediate). Interestingly, sulfolane and toluene, which do not support stabilization of the carbamic acid at low amine concentrations in the manner observed for DMSO and NMP, show a very weak dependence of CO2-amine reaction equilibrium on amine concentration. Neat EEA sorbent shows approximately the same CO2:amine loading as EEA/toluene or EEA/sulfolane (and, in the case of toluene, approximately the same product distribution).

Lower sorption temperatures also favor carbamic acid formation because carbamic acids are less stable than carbamates. Carbamic acid products thus also may provide potential process advantages of easier CO2 desorption and lower energy regeneration of free amine. Lower energy desorption for carbamic acid-rich sorption systems has been confirmed via both pure thermal degradation in the absence of a purge gas, and in conjunction with an inert purge gas. Therefore, a combined Pressure and Temperature swing (PSATSA) approach may be beneficial for non-aqueous amine regeneration.

More polar non-aqueous solvents, such as DMSO, NMP, and ionic liquids, significantly affect CO2amine reaction chemistry via the stabilization of certain reaction products. This effect is especially pronounced at high solvent concentration (e.g., low amine concentration) and vanishes at high amine concentration.

The CO2-amine thermodynamic equilibrium varies significantly with amine concentration and feed gas CO2 partial pressure. At low partial pressures of CO2, the reaction driving force is lower, favoring carbamic acid over carbamate species.

We attribute these solvent effects to the ability of the non-aqueous solvent to shift the carbamic acidcarbamate equilibrium in the direction of the acid via hydrogen-bonding to the acidic proton. With

Weakly basic amines react with CO2 in nonaqueous solutions to form primarily carbamic acids which have particularly low thermal stability. These features of high CO2/amine ratio and relatively low33


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energy CO2 regeneration may provide advantages for low-temperature CO2 capture. In contrast, strong amidine and guanidine bases react effectively with CO2 above 50 0C and therefore can be used for CO2 capture at elevated temperatures. These bases also react with CO2 forming a stable zwitterion at ambient conditions.

Supplementary Material 1

H and 13C NMR spectra of unreacted and reacted amines discussed in the paper as well as results of CO2 reaction with other primary and secondary amines, guanidines and biguanides are included in the supplementary material.

Acknowledgement. We gratefully acknowledge Prof. Alan R. Katritzky and Dr. Ekaterina Todadze, University of Florida, for preparation of some of the guanidines used in this study.

Author Information Corresponding Author

[email protected]

Key Words: Carbon capture; non-aqueous solution; organic solvents; ionic liquids; CO2 absorption; carbamate; carbamic acid; stabilization of CO2 reaction products; zwitterion; guanidines; hydrogen bonding; 13 1 in situ C and H NMR

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References 1.

Kortunov, P.; Siskin, M; Baugh, L. S.; Calabro, D. C. “In Situ NMR mechanistic studies of carbon dioxide reactions with liquid amines in aqueous systems: New insights on carbon capture reaction pathways.” Energy and Fuels, accepted for publication.

2.

Davis, R.A.; Sandall, O.C. “Kinetics of the reaction of carbon dioxide with secondary amines in polyethylene glycol.” Chem. Eng. Sci. 1993, 48, 3187-3193.

3.

Li, J.; You, C.; Chen, L.; Ye, Y.; Qi, Z.; Sundmacher, K. “Dynamics of CO2 absorption and desorption processes in alkanolamine with cosolvent polyethylene glycol.” Ind. Eng. Chem. Res. 2012, 51, 12081-12088.

4.

Versteeg, G.F.; Van Swaaij, W.P.M. “On the kinetics between CO2 and alkanolamines both in aqueous and non-aqueous solutions – I. Primary and secondary amines.” Chem. Eng. Sci.1988, 43, 573-585.

5.

Park, S.-W.; Lee, J.-W.; Choi, B.-S.; Lee, J.-W. “Absorption of carbon dioxide into non-aqueous solutions of Nmethyldiethanolamine.” Korean J. Chem. Eng. 2006, 23, 806-811.

6.

Versteeg, G. F.; van Swaaij, W. P. M. “On the kinetics between CO2 and alkanolamines both in aqueous and non-aqueous solutions - II. Tertiary amines.” Chem. Eng. Sci. 1988, 43, 587-591.

7.

Hwang, K.-S.; Park, S.-W.; Park, D.-W.; Oh, K.-J.; Kim, S.-S. “Absorption of carbon dioxide into diisopropanolamine solutions of polar organic solvents.” J. Taiwan Inst. Chem. Eng. 2010, 41, 16-21.

8.

Xu, S.; Wang, Y.-W.; Otto, F. D.; Mather, A. E. “Kinetics of the Reaction of Carbon Dioxide with 2-Amino-2-methyl-1propanol Solutions.” Chem. Eng. Sci. 1996, 51, 841-850.

9.

Park, S.-W.; Choi, B.-S.; Lee, J.-W. “Chemical absorption of carbon dioxide with triethanolamine in non-aqueous solutions.” Korean J. Chem. Eng. 2006, 23, 138-143.

10.

Sada, E.; Kumazawa, H.; Ikehara, Y.; Han, Z.Q "Chemical kinetics of the reaction of carbon dioxide with triethanolamine." Chem. Eng. J. 1989, 40, 7-12.

11.

Kadiwala, S.; Rayer, A. V.; Henni, A. “Kinetics of carbon dioxide (CO2) with ethylenediamine, 3-amino-1-propanol in methanol and ethanol, and with 1-dimethylamino-2-propanol and 3-dimethylamino-1-propanol in water using stoppedflow technique.” Chem. Eng. J. 2012, 179, 262-271.

12.

Crooks, J. E.; Donnellan, J. P.; “Kinetics of the formation of N,N-dialkylcarbamate from diethanolamine and carbon dioxide in anhydrous ethanol.” J. Chem. Soc. Perkin, 1988, 191-194.

13.

Littel, R. J.; Versteeg, G. F.; van Swaaij, W. P. M. “Physical absorption into non-aqueous solutions in a stirred cell reactor.” Chem. Eng. Sci. 1991, 46, 3308-3313.

14.

Imaishi, N.; Iguchi, H.; Hozawa, M.; Fujinawa, K. “Chemical absorption of carbon dioxide by non-aqueous solutions of cyclohexylamine.” Kagaku Kogaku Ronbunshu 1981, 7, 261-266.

15.

Sada, E.; Kumazawa, H.; Osawa, Y.; Matsuura, M.; Han, Z.Q. "Reaction kinetics of carbon dioxide with amines in nonaqueous solvents." Chem. Eng. J. 1986, 33, 87-95.

16.

Takesita, K.; Kitamoto, A. “Relation between separation factor of carbon isotope and chemical reaction of CO2 with amine in nonaqueous solvent.” J. Chem. Eng. Jpn. 1989, 22, 447-454.

17.

Takeshita, K.; Kitamoto, A. “Chemical equilibria of absorption of CO2 into nonaqueous solution of amin.e” J. Chem. Eng. Jpn. 1988, 21, 411-417.

18.

Takeshita, K. “Separation of carbon isotopes by using the chemical reaction of carbon dioxide with amines in nonaqueous solution.” Kagaku Kogaku 1991, 55, 426-428.

19. Yushko, V. L.; Sergienko, I. D.; Kosyakov, N. E.; Khokhlov, S. F.; Pushkin, A. G.; Ivanova, A. E.; Kosenko, E. D. “Effect of water content on the solubility of carbon dioxide in solutions of monoethanolamine in sulfolane.” Voprosy Khimii i Khimecheskoi Tekhnologii 1973, 30, 3-5.

35


Energy & Fuels

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

Page 36 of 55

20. Sada, E.; Kumazawa, H.; Han, Z. Q.; Matsuyama, H. “Chemical kinetics of the reaction of carbon dioxide with ethanolamines in nonaqueous solution.” AIChE J. 1985, 31, 1297-1303. 21.

Sada, E.; Kumazawa, H.; Han, Z. Q. “Chemical absorption of carbon dioxide into ethanolamine solutions of polar solvent.” AIChE J. 1986, 32, 347-349.

22. Dinda, S.; Patwardhan, A.V.; Panda, S. R.; Pradhan, N. C. "Kinetics of the reaction of carbon dioxide with solutions of aniline in carbon tetrachloride and chloroform." Chem. Eng. J. 2008, 136, 349-357. 23. Dinda, S.; Patwardhan, A. V.; Pradhan, N. C. "Kinetics of reactive absorption of carbon dioxide with solutions of aniline in nonaqueous aprotic solvents." Ind. Eng. Chem. Res. 2006, 45, 6632-6639. 24. Ali, S. H.; Merchant, S. Q.; Fahim, M. A. "Kinetic study of reactive absorption of some primary amines with carbon dioxide in ethanol solution." Separation and Purification Tech. 2000, 18, 163-175. 25. Versteeg, G. F.; Van Dijck, L. A. J.; Van Swaaij, W. P. M. "On the kinetics between CO2 and alkanolamines both in aqueous and non-aqueous solutions. An overview." Chem. Eng. Comm. 1996, 144, 113-158 and references therein. 26. Jamal, A.; Meisen, A.; Lim, C. J. “Kinetics of carbon dioxide absorption and desorption in aqueous alkanolamine solutions using a novel hemispherical contactor. I. Experimental apparatus and mathematical modeling.” Chem. Eng. Sci., 2006, 61, 6571-6589. 27. 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. 28. Ballard, M.; Brown, M.; James, S.; Yang, Q. ”NMR studies of mixed amines.” Energy Procedia, 2011, 4, 291-298. 29. 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, 1009710106. 30. Rainbolt, J.E.; Koech, P.K.; Yonker, C.R.; Zheng, F.; Main, D.; Weaver, M.L.; Linehan, J.C.; Heldebrant, D.J. “Anhydrous tertiary alkanolamines as hybrid chemical and physical CO2 capture reagents with pressure-swing regeneration.” Energy Env. Sci. 2011, 4, 480-484. 31.

Lail, M.; Coleman, L.; Jamal, A.; Amato, K.; Gupta, R. “Separation of carbon dioxide from gas mixtures using nonaqueous solvent systems.” Abstracts of Papers, 241st National Meeting of the American Chemical Society, Anaheim, CA, March 2731, 2011; American Chemical Society: Washington, DC, 2011; FUEL 43.

32. McCann, N.; Phan, D.; Wang, X.; Conway, W.; Burns, R.; Attalla, M.; Puxty, G.; Maeder, M. “Kinetics and mechanism of carbamate formation from CO2(aq), carbonate species and monoethanolamine in aqueous solution.” J. Phys. Chem. A 2009, 113, 5022-5029. 33. Jakobsen, J.P. da Silva, E. F. Krane, J.; Svendsen, H. F. “NMR study and quantum mechanical calculations on the 2-[(2aminoethyl)amino]-ethanol-H2O-CO2 system.” J. Magn. Res. 2008, 191, 304-314. 34. Garcia-Abuin, A.; Gomez-Diaza, D.; Navaza, J. M.; Rumbo, A. “NMR studies in carbon dioxide - amine chemical absorption.” Procedia Engineering 2013, 42, 1242-1249. 35. 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. 36. Richner, G.; Puxty, G. “Assessing the chemical speciation during CO2 absorption by aqueous amines using in situ FTIR.” Ind. Eng. Chem. Res. 2012, 51, 14317−14324. 37. 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. 38. 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.

36


Page 37 of 55

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

Energy & Fuels

39. Robinson, K.; McCluskey, A.; Attalla, M. I. “The effect molecular structural variations has on the CO2 absorption characteristics of heterocyclic amines.” In: Recent Advances in Post-Combustion CO2 Capture Chemistry; Attalla, M. I., Ed. ACS Symposium Series 1097; American Chemical Society: Washington, DC, 2013; Chapter 1, pp 1-27. 40. 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. 41.

Boger, D.L.; Patel, M. “Activation and coupling of pyrrole-1-carboxylic acid in the formation of pyrrole N-carbonyl compounds: pyrrole-1-carboxylic acid anhydride.” J. Org. Chem. 1987, 52, 2319-2323.

42. Boger, D.L.; Patel, M. “Indole N-carbonyl compounds: preparation and coupling of indole-1-carboxylic acid anhydride.” J. Org. Chem. 1987, 52, 3934-3936. 43. Katritzky, A. R.; Faid-Allah, H.; Marson, C. M. “The N-carboxylic acids of nitrogen heterocycles.” Heterocycles 1987, 26, 1333-1344. 44. Ito, Y. “Formation of carbamic acids and their photochemistry.” Kokagaku 2002, 33, 205-212. 45. Kayaki, Y.; Suzuki, T.; Ikariya, T. “Utilization of N,N-dialkylcarbamic acid derived from secondary amines and supercritical carbon dioxide: stereoselective synthesis of Z-alkenyl carbamates with a CO2-soluble ruthenium-P(OC2H5)3 catalyst.” Chem. Asian J. 2008, 3, 1865-1870. 46. (a) Hampe, E. M.; Rudkevich, D. M. “Reversible covalent chemistry of CO2.” Chem. Commun. 2002, 1450-1451. (b) Hampe, E. M. Rudkevich, D. M. “Exploring reversible reactions between CO2 and amines.” Tetrahedron 2003, 59, 96199625. 47. Dijkstra, Z. J.; Doornbos, A. R.; Weyten, H.; Ernsting, J. M.; Elsevier, C. J.; Keurentjes, J. T. F. “Formation of carbamic acid in organic solvents and in supercritical carbon dioxide.” J. Supercritical Fluids 2007, 41, 109-114. 48. Masuda, K.; Ito, Y.; Horiguchi, M.; Fujita, H. “Studies on the solvent dependence of the carbamic acid formation from ω-1(naphthyl)alkylamines and carbon dioxide.” Tetrahedron 2005, 61, 213-229. 49. Kortunov, P.; Baugh, L. S.; Calabro, D. C.; Siskin, M. “In Situ NMR mechanistic studies of carbon dioxide reactions with liquid amines in mixed base systems: the interplay of Lewis and Brønsted basicity.” Energy and Fuels, accepted for publication. 50. Kortunov, P.; Baugh, L. S.; Siskin, M. “Mechanism of the chemical reaction of carbon dioxide with ionic liquids and amines in ionic liquid solution.” Energy and Fuels, accepted for publication. 51.

Aronu, U.E.; Svendsen, H. F.; Hoff; K. A.; Juliussen, O. “Solvent selection for carbon dioxide absorption.” Energy Procedia 2009, 1, 1051-1057.

52. Carretti, E.; Dei, L; Baglioni, P.; Weiss, R. G. “Synthesis and characterization of gels from polyamine and carbon dioxide as gellant.” J. Am. Chem. Soc. 2003, 125, 5121-5129. 53. See, for example: (a) Bates, R. G.; Pinching, G. D. “Acidic dissociation 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 K.” J. Chem. Eng. Data 1990, 35, 276277. (c) Antelo, J. M.; Arce, F.; Casado, J.; Sastre, M.; Varela, A. “Protonation constants of mono-, di-, and triethanolamine. Influence of the ionic composition 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) K.” 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. Eng. Chem. Res. 2012, 51, 11924-11930. 54. For some experimentally determined aqueous and non-aqueous 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. 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 conversions to sulfonamides at high pH.” Can. J. Chem. 2005, 83, 1525-1535. (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) 37


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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. (i) Ishikawa, T., Ed. “Superbases for organic synthesis: guanidines, amidines, phosphazenes and related oganocatalysts.” JohnWiley and Sons, Ltd.,: Chichester, UK, 2009. (j) Pahadi, N. P.; Ube, H.; Terada, M. “Aza-Henry reaction of ketimines catalyzed by guanidine and phosphazene bases.” Tetrahedron Lett. 2007, 8700-8703 and references therein. (k) Kaljurand, I.; Kutt, A.; Soovali, L.; Rodima, T.; Maemets, V.; Leito, I.; Koppel, I. A. “Extension of the self-consistent spectrophotometric basicity scale in acetonitrile to a full span of 28 pKa units: unification of different basicity scales.” J. Org. Chem. 2005, 70, 1019-1028. (l) Kovacevic, B.; Maksic, Z. B. “Basicity of some organic superbases in acetonitrile.” Org. Lett. 2001, 3, 1523-1526. 55. Astarita, G; Savage, D. W.; Bisio, A. “Gas treating with chemical solvents.” Wiley-Interscience Publications, John Wiley & Sons: New York, 1983, pp. 208-210. 56. Laurence, C.; Gal, J.-F. “Lewis basicity and affinity Scales – data and measurement.” John Wiley & Sons: New York, 2010. 57. (a) Orth, J.H. “Preparation of aminocarbamates from diamines and carbon dioxide.” PCT Int. Appl. (1997), WO 9729083 A1 19970814 CAN 127:190473 AN1997:542425. (b) Yamada, T.; Lukac, P.J.; George, M.; Weiss, R.G. “Reversible, roomtemperature ionic liquids. Amidinium carbamates derived from amidines and aliphatic primary amines with carbon dioxide.” Chem. Mat. 2007, 19, 967-969. (c) Yamada, T.; Lukac, P.J.; Yu, T.; Weiss, R.G. “Reversible, room-temperature, chiral ionic liquids. Amidinium carbamates derived from amidines and amino-acid esters with carbon dioxide.” Chem. Mat. 2007, 19, 4761-4768. (d) 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.” Chem. Mat. 2008, 20, 5337-5344. (e) Yu, T.; Weiss, R.G.; Yamada, T.; George, M. “Reversible room-temperature ionic liquids for gas sequestration.” PCT Int. Appl. (2008), WO 2008094846 A1 20080807 CAN 149:249600 AN 2008:942211. 58. (a) 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 mono- and dicarbamic acids/carbamtes of piperazine at 25.0 °C.” J. Phys. Chem. A 2013, 117, 806-813. (b) Bishnoi, S.; Rochelle, G. T. “Absorption of carbon dioxide in aqueous piperazine / methyldiethanolamine.” AICHE Journal 2002, 48, 2788. (c) Bishnoi, S.; Rochelle, G. T. “Thermodynamincs of piperazine / methyldiethanolamine / water / carbon dioxide.” Ind. Eng. Chem. Res. 2002, 41, 604-612. (d) 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. 59. Endo, T.; Nagai, D.; Monma, T.; Yamaguchi, H.; Ochiai, B. “A novel construction of a reversible fixation-release system of carbon dioxide by amidines and their polymers.” Macromolecules 2004, 37, 2007-2009. 60. Gund, P. “Guanidine, trimethylenemethane, and Y-delocalization: can acyclic compounds have “aromatic” stability.” J. Chem. Ed. 1972, 49, 100-103. 61.

Anderson, M. L.; Hamner, R. N. “Properties of 1,1,3,3-tetramethylguanidine as a nonaqueous solvent.” J. Chem. Eng. Data 1967, 12, 442.

62. Minami, S.; Shono, T.; Matsumoto, J. “Pyrido[2,3-d]pyrimidine antibacterial agents. II. Piromidic acid and related compound.s” Chemical & Pharmaceutical Bulletin 1971, 19, 1426-32. 63. Breslow, R. “Hydrophobic effects on simple organic reactions in water.” Acc. Chem. Res. 1991, 24, 159-164. 64. Breslow, R.; Maitra, U.; Rideout, D. “Selective Diels-Alder reactions in aqueous solutions and suspensions.” Tetrahedron Lett. 1983, 24, 1901-1904. 65. Greieco, P. A. “Organic chemistry in unconventional solvents.” Aldrichimica Acta 1991, 24, 59. 66. Price, C.C.; Oae, S. “Sulfur bonding.” The Ronald Press Co.: New York, 1962, pp. 61-62.

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Figures:

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. 13C and 1H NMR spectra of 15 wt % DMAOP in DMSO-d6 after CO2 saturation at 30 °C. The reaction product is associated with the 13C peak at 158.6 ppm and 1H peak at 10.2 ppm; equilibrium CO2 loading per each amine of DMAOP)= 0.74.

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Figure 3. Evolution of the amine-CO2 reaction of 30 wt% EEA in DMSO-d6 at 1.0 bar of CO2 at 30 0C with time monitored by 13C and 1H NMR. Formation of carbamate/carbamic acid species (top), evolution of –CH2-O-CH2- carbons (middle), and evolution of 1H NMR spectra (bottom).

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Figure 4. 13C and 1H NMR spectra of 3 molar EEA in DMSO-d6 before (top) and after CO2 saturation at 30 °C (below). The reaction product is associated with the 13C peak at 159.8 ppm and 1H peaks at 9.8 and 6.6 ppm; CO2/EEA mole ratio = 0.77.

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Figure 5. Evolution of carbamate/carbamic acid species during the amine-CO2 reaction monitored by 13C NMR for 3 M EEA in toluene solution at 1.0 bar of CO2 at 30 0C (top); for 3 M EEA in DMSO-d6 at 10.0 bar of CO2 at 30 0C (middle); for pure EEA at 1.0 bar of CO2 at 30 0C (bottom).

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Figure 6. Evolution of the amine-CO2 reaction for 3 M THP in DMSO-d6 at 1.0 bar of CO2 at 30 C with time monitored by 13C NMR. Formation of zwitterion/carbamic acid species (top), evolution of –CH2-CH2-CH2- carbon (bottom). 0

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Figure 7. 13C and 1H NMR spectra of 10 wt % 3-butylguanidine in DMSO-d6 before (top) and after CO2 treatment (bottom) at 30 °C. Two types of reaction products were detected at 160.3 ppm and 154.9 ppm. The total CO2 loading is 0.785 CO2 per guanidine. The peak at 124.9 ppm indicates physisorbed CO2 (0.205 CO2 per guanidine).

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Figure 8. CO2 uptake of 15 wt % solutions of IDPN, APN, EEA and PP in DMSO-d6 solution at various temperatures at 1.0 bar of CO2.

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Figure. 9. CO2 loading capacity and product speciation for EEA in DMSO-d6 solution at various amine concentrations and a fixed CO2 partial pressure of 1.0 bar at 30 °C.

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Figure 10. CO2 loading capacity of 28 wt% (5 M) MEA in DMSO and 20 wt% (3 M) APN in DMSO as a function of CO2 partial pressure at 45 and 90 0C.

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Energy & Fuels

Figure 11. 13C and 1H NMR spectra of 50 wt% EEA in EMIM-OAc with drop of DMSO-d6 (as reference) before (top) and after CO2 treatment (bottom) at 30 °C. The reaction product is associated with the 13C peak at 160.5 ppm and 1H peaks at 13.2 and 6.9 ppm; CO2/EEA mole ratio = 0.90.

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Figure 12. Comparison of 13C NMR CO2 loading capacities and product speciation for EEA dissolved in various solvents at various concentrations at 1.0 bar CO2 and 30 °C: in toluene-d8 (top) and EMIM-OAc ionic liquid (bottom).

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Energy & Fuels

Figure 13. 1H NMR of OH relative to adjacent CH2 of ethanol dissolved in various organic solvents as a function ethanol concentration. Changes of the δ(OH)-δ(CH2) reflect the extent of intermolecular interactions (specifically, hydrogen bonding) of the OH group with other

molecules in the solution.

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Figure 14. Comparison of CO2 loading capacities in solution (per combined weight of amine and solvent) for EEA in various solvents at various amine concentrations (1.0 bar CO2, 30 °C; IL = EMIM-OAc). Dashed line shows theoretical CO2 capacity for exclusive ammonium carbamate formation.

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Figure 15. Comparison of 13C NMR CO2 loading capacities and product speciation for EEA aqueous solution at various amine concentrations (1.0 bar CO2, 30 °C).

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Figure 16. CO2 loading capacity and product speciation for APP in DMSO-d6 at various amine concentrations (1.0 bar CO2, 30 °C).

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TOC

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In Situ Nuclear Magnetic Resonance Mechanistic ... - ACS Publications

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