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Jan 20, 2016 - The mechanisms of the reaction of CO2 with hindered alkanolamines are described based on direct reaction monitoring with in situ 13C NM...
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CO2 Reaction Mechanisms with Hindered Alkanolamines: Control and Promotion of Reaction Pathways Pavel V. Kortunov,* Michael Siskin, Michele Paccagnini, and Hans Thomann Corporate Strategic Research Laboratory, ExxonMobil Research and Engineering Company, 1545 Route 22 East, Annandale, New Jersey 08801, United States S Supporting Information *

ABSTRACT: The mechanisms of the reaction of CO2 with hindered alkanolamines are described based on direct reaction monitoring with in situ 13C NMR spectroscopy. Four amines of increasing steric hindrance were studied and compared with unhindered primary, secondary, and tertiary alkanolamines. The number and location of additional methyl (or other) groups on the nitrogen or adjacent carbon atoms create different degrees of steric hindrance. As steric hindrance increases, the Lewis basicity (nucleophilicity), as an affinity of the amine nitrogen to directly attack the electrophilic carbon of CO2 and form a carbamate, decreases. At the same time, methyl (or other) groups present on nitrogen or adjacent carbon atoms do not change the Brønsted basicity of hindered amines. Coupled with lower Lewis basicity, the CO2−amine−water reaction equilibrium favors formation of bicarbonate, with higher CO2 loading capacity and lower thermal stability, both of which are favorable properties for cyclic CO2 capture. Due to lower stability of the hindered N−C bond of the carbamate, hindered amines AMP and MAP in aqueous solution show a phenomenon of steric acceleration: the more rapid transfer of CO2 from the nitrogen of the amine (e.g., carbamate) to form the bicarbonate anion, maintaining relatively high reaction rates. A strong nucleophilic base such as piperazine present in the amine solution at low concentration maintains the high reaction rate by attacking free CO2 and transferring it to the amine as bicarbonate. This described mechanism helps to improve CO2 capture rates with severely hindered amines such as in the secondary amine MAMP, which without a promoter acts as a tertiary amine slowly reacting with CO2. The mechanism of chemical reaction between CO2 and amines dissolved in methanol is similar to that in watermethanol attacks CO2 and forms an O−C bond and methylbicarbonate anion (analogue of bicarbonate) with CO2, which is stabilized by a protonated amine. The effects of amine concentration, CO2 partial pressure, and temperature are discussed.

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INTRODUCTION AND BACKGROUND Amine−CO2 Reaction Chemistry in Aqueous Solution. In situ 1H and 13C NMR mechanistic studies reported earlier (see ref 1 and background references cited therein) elucidated the pathways of carbon capture with primary, secondary, and tertiary amines as well as amidines and guanidines in aqueous media. Largely traditional chemistry with formation of carbamates, carbonates, and bicarbonates was observed. The electrophilic nature of the carbon atom in CO2 makes it susceptible to nucleophilic attack by various N-, O-, and Cdonors, e.g., of amines, hydroxyl groups, and certain ionic liquids, respectively.2−4 Primary and secondary amines can act as a nucleophile (Lewis base) by direct attack on free CO2 to form a zwitterion, which rapidly rearranges to the carbamic acid via intramolecular proton transfer. In the presence of another free amine, which now acts as a Brønsted base, the carbamic acid may be converted into a carbamate via intermolecular proton transfer. Additionally, primary and secondary amines can form bicarbonate salts by deprotonation of carbonic acid or by hydrolysis of carbamate. Based on detailed monitoring of the monoethanolamine (MEA) reaction intermediates, we observed that at an early reaction stage the primary amine acts as a strong Lewis base forming exclusively the carbamate at a relatively fast reaction rate.1 Only after reaching the theoretical maximum of carbamate formation (CO2/N = 0.5) and significant reduction of solution pH does the less nucleophilic water hydrolyze the © 2016 American Chemical Society

primary carbamate product into bicarbonate, ultimately generating a carbamate/bicarbonate equilibrium mixture (Scheme 1). Carbamate hydrolysis liberates a free amine which, with continued introduction of fresh CO2, forms either carbamate, which quickly decomposes into bicarbonate, or directly forms the bicarbonate salt via aqueous carbonic acid (in a fashion similar to reaction with tertiary amines discussed below) at this reduced pH. The resulting carbamate−bicarbonate equilibrium depends on many parameters, such as amine basicity, amine nucleophilicity, amine concentration, solution temperature, and CO2 partial pressure.1 Unlike primary and secondary amines, tertiary amines are weak Lewis bases but can be relatively strong Brønsted bases. As such, the nitrogen of tertiary amines can act exclusively as proton acceptors by forming the two types of stable products with CO2: a bicarbonate and/or carbonate (dependent on the pH of the solution). In our previous work1 we confirmed that tertiary amines react with CO2 (as carbonic acid) by simultaneously forming carbonate and bicarbonate products at an early stage of the reaction (Scheme 2). Bicarbonate appears to be the favored reaction product at all stages of the reaction, surprisingly even Received: November 3, 2015 Revised: January 15, 2016 Published: January 20, 2016 1223

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Scheme 1. General Mechanism (Left) and 13C NMR Spectroscopic Evidence (Right) of the Reaction of CO2 with an Aqueous Solution of Primary (or Secondary) Amine: Nucleophilic Attack on Free CO2 with Formation of Zwitterion/Carbamic Acid (Step 1) Followed by Deprotonation by Another Basic Amine To Form a Carbamate (Step 2) and Partial or Complete Hydrolysis into (Bi)carbonate (Step 3)

Scheme 2. General Mechanism (Left) and 13C NMR Spectroscopic Evidence (Right) of the Reaction of CO2 with an Aqueous Solution of Tertiary Amine: Formation of Carbonic Acid between CO2 and H2O Followed by Deprotonation by Another Basic Tertiary Amine into Carbonate (at High pH) or Bicarbonate (at Lower pH), Which Exist in Equilibrium

following series: 2-amino-1-propanol (AP), 2-amino-2-methyl1-propanol (AMP), 2-(methylamino)-2-methyl-1-propanol (MAMP), and 2-(methylamino)-1-propanol (MAP) under aqueous and nonaqueous (e.g., DMSO) conditions.

at high solution pH. With increasing protonation of the tertiary amine, the basic amine is consumed and the solution pH drops favoring the formation of bicarbonate from carbonic acid even more. The carbonate concentration initially increases at high pH and then declines as pH decreases. A second route to bicarbonate formation becomes possible at low pH. After a majority of amines are protonated and form carbonate and bicarbonate species with CO2, carbonate anions (conjugate base) can deprotonate a carbonic acid molecule to form two bicarbonate anions. Insights gained for the previous aqueous studies and the new pathways gleaned from the nonaqueous work2,3 made us realize that the reactivity of hindered amines described in this work needs to be elucidated and differentiated from other primary, secondary, and tertiary amines/alkanolamines because they form bicarbonates vs carbamates which doubles the capacity of CO2 absorption and are stable enough to be compatible with a typical 45 °C absorption followed by a 110−120 °C desorption/regeneration scheme. Bougie5 and others6−10 have written a comprehensive review of sterically hindered amine absorbents for carbon capture. AMP (2-amino-2-methyl-1-propanol) is a hindered primary amine favored for H2S selectivity in acid gas scrubbing processes.10−12 Because of its steric hindrance, kinetic formation of the hydrosulfide salt is favored over CO2 capture. Aronu,13 in a publication on carbon capture properties of tetraethylenepentamine (TEPA), disclosed AMP as one component of the Mitsubishi KS-1 carbon capture process. This work will elucidate the effects of the stepwise increase in steric hindrance on carbon capture chemistry by studying the

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EXPERIMENTAL DESIGN

Using in situ 13C NMR monitoring, we describe the temporal evolution of product formation for the reaction of CO2 with a series of primary and secondary hindered amines (Table 1) in aqueous solution as a function of reaction temperature, CO2 partial pressure, and amine concentration. The quantitative results illustrate stepwise product formation and decomposition under absorption and desorption conditions, respectively. Subsequently, we will discuss the behavior of amine sorbents with increasing steric hindrance around the amine group. These studies employ two series of monoalkanolamines having one or two methyl groups on the α-carbon (carbon atom adjacent to the amine group) to study the effect of amine structural type on reactivity. Alkanolamines are preferred over the use of amines lacking polar functional groups in order to facilitate high water solubility of the amine and its reaction products with CO2. Properties of sterically hindered amines (4−7) to react with CO2 were compared with unhindered amines (1−3). Results of additional reactions studied (but not discussed in detail in the text) are tabulated in the Supporting Information (SI). Hindered amines AP (4) and AMP (5) were purchased from SigmaAldrich while MAP (6) and MAMP (7) were synthesized in the laboratory; their structures were confirmed by 1H and 13C NMR spectroscopy and elemental analysis. The general NMR procedure for CO2 absorption and desorption as well as spectral acquisition and analysis is explained in detail in our previous work.1 Amine saturation with CO2 inside the NMR 1224

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Energy & Fuels Table 1. Physical Properties of Amines Used in Aqueous CO2 Sorption Experiments

a

Predicted using ACD/Laboratories (Advanced Chemistry Development) software V8.14 for Solaris.

bar of CO2 at 30 °C, the aqueous solution (3 M MEA) forms a mixture of carbamate (61 mol %) and bicarbonate (39 mol %) with the total loading of approximately 0.7 CO2 per amine.1 It is very important to again note that MEA exclusively forms carbamate species with CO2 at early reaction stages while bicarbonate products appear as a result of carbamate hydrolysis after a majority of amines form carbamate and solution pH dropped. Such reaction mechanism and product equilibrium suggest that MEA can be considered as a strong nucleophile (Lewis base) with high affinity of its nitrogen to the carbon of CO2 and as a strong Brønsted base to accept a proton and stabilize either carbamate or bicarbonate species. SI Figures S1.1 and S1.2 show the evolution of CO2−MEA reaction in aqueous solution. A methyl group attached to nitrogen in secondary amine (methylamino)ethanol (MAE, 2) creates steric hindrance in the amine, reduces its Lewis basicity, and increases Brønsted basicity. Under the conditions applied previously to MEA (1.0 bar of CO2 at 30 °C), an aqueous solution of 3 M MAE forms predominantly bicarbonate species (72 mol %) and a smaller amount of carbamate (28 mol %). Such product equilibrium increases the total CO2 loading to approximately 0.86 CO2 per amine at the same pressure, temperature, and amine concentration.1 SI Figures S1.3 and S1.4 show the evolution of CO2−MAE reaction in aqueous solution.

spectrometer was organized in two ways. The experiments at CO2 partial pressure of 1.0 bar and below were conducted in a flow-through mode (SI Figure S1.0). CO2 containing gas (or pure CO2) was bubbled through the amine solution at a total pressure of about 1.0 bar. Although the majority of experiments reported here were performed by purging pure CO2 gas (PCO2 = 1.0 bar), special mixtures of gases were also used to study effects of CO2 partial pressure. For example, a 10/90 CO2/N2 mixture (mol %), purchased from Matheson Tri-Gas, was used to monitor reaction at a CO2 partial pressure of 0.1 bar. Experiments at CO2 pressure above 1.0 bar, or using more volatile hindered amines (such as methoxylated amines), or solvents (such as methanol) were performed in a batch mode. In this case, the amine solution is pressurized by pure CO2 at a given pressure, which changes only modestly during the gas absorption experiments due to a high volume ballast reservoir connected to the NMR cell. It is important to note that, under the conditions used for these measurements, the gas−liquid reaction is often limited by the gas delivery rate and the time-dependent data are not reflective of the intrinsic reaction kinetics. Rather than a detriment, this attenuation of the real kinetics has enabled careful monitoring of the sequence of reactions and relative kinetics occurring as a function of concentration and temperature.

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RESULTS AND DISCUSSION We have chosen the primary amine monoethanolamine (MEA, 1) as the benchmark for this study. After the reaction with 1.0 1225

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Energy & Fuels Two methyl groups attached to the nitrogen in tertiary amine (dimethylamino)ethanol (DMAE, 3) further reduces the Lewis basicity of the nitrogen atom while maintaining a similar or even higher Brønsted basicity. As a result of the extremely low affinity of the nitrogen atom to the carbon of CO2, DMAE does not form an N−C bond with CO2 and stabilizes as a mixture of carbonate (10 mol %) and bicarbonate (90 mol %) by accepting a proton. The total CO2 loading on DMAE further increases to approximately 0.95 CO2 per amine at conditions comparable to those for MEA and MAE previously discussed.1 The drawback of the CO2 reaction with a tertiary amine is the slow rate of chemical reaction, which in contrast to primary and secondary amines is controlled by the low concentration of carbonic acid in the solution. SI Figures S1.5 and S1.6 show the evolution of the CO2−DMAE reaction in aqueous solution. 1. CO2 Reaction with Aqueous Solutions of Hindered Amines. Amines with high Brønsted basicity at moderate to low Lewis basicity are beneficial for CO2 capture because they preferentially form bicarbonate salts with high CO2 capacity and have lower thermal stabilities. Additionally, carbamate species formed with weaker Lewis bases such as MAE are less stable than carbamate species with strong Lewis bases such as MEA. The Lewis basicity of the amine can also be modified by placing bulky groups around the amine which creates steric hindrance and reduces the affinity of the amino nitrogen to the carbon of CO2. To study potential changes of the CO2−amine reaction mechanism as well as product equilibrium, we studied several moderately hindered aminesAP (4), AMP (5), MAP (6), and MAMP (7):

2-Amino-1-propanol. The primary amine AP has a single methyl group on the carbon adjacent to the amine group and represents the simplest chemical structure of a moderately sterically hindered amine. Based on monitoring of the AP reaction intermediates, we conclude that the CO2 reaction mechanism with AP is similar to reaction with primary amine MEA and secondary amine MAE: initial carbamate formation followed by partial carbamate hydrolysis into bicarbonate and quantitative overall conversion of the amine groups with CO2 into carbamate and bicarbonate salts (Figure 1a and SI Figures S1.7 and S1.8). However, equilibrium CO2 loading and concentration of reaction products of primary hindered amine AP is close to secondary amine MAE (SI Figures S1.4 and S1.5). This result confirms that a methyl group present on the carbon adjacent to the nitrogen atom of AP creates a moderate steric hindrance and reduces the nucleophilicity (Lewis basicity) of AP. Due to a now weaker N−C bond, the carbamate of AP more easily hydrolyzes to the bicarbonate with a 1:1 CO2/amine mole ratio, which leads to increased CO2 capacity of the aqueous AP solution relative to more nucleophilic MEA. 2-Amino-2-methyl-1-propanol. The primary amine AMP has two methyl groups on the carbon α to the amine group. These two methyl groups are expected to create more significant hindrance at the amine nitrogen and further reduce its Lewis basicity relative to AP.

Figure 1. Evolution of the reaction of CO2 with 3 M aqueous solutions of AP (A), AMP (B), MAP (C), and MAMP (D) at 30 °C monitored by in situ 13C NMR. Pure CO2 at 1.0 bar was purged at 5 cm3(STP) min−1.

As the aqueous AMP solution at 30 °C was treated with 5 cm3(STP) min−1 of pure CO2 at 1.0 bar, two 13C resonances at approximately 165.2 and 164.4 ppm were initially detected in the CO region (Figure 1b and SI Figures S1.9 and S1.10). The second resonance at 164.4 ppm represents the carbamate species similar to that of AP described earlier (Figure 2, top). It is remarkable that the concentration of carbamate species in AMP solution remains low during the reaction and does not exceed 20 mol % (or 0.1 CO2 per AMP, SI Figure S1.9) while 1226

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Scheme 4. Second Mechanism of Bicarbonate Formation with Hindered Amine AMP by Carbamate Hydrolysis

carbonation3 of AMP (NH2C(CH3)2CH2O−COO−), which is likely stabilized by another AMP molecule. Although concentration of this product is low under the given experimental conditions (1.0 bar of CO2 and 30 °C), it may increase at either higher CO2 pressure or lower solution temperature. 2-(Methylamino)-2-methyl-1-propanol. Secondary amines have lower Lewis basicity relative to the corresponding primary amines.1 In the previous section we described an example of hindered primary amine AMP, which already shows very weak Lewis basicity toward the carbon of CO2 due to two methyl groups on the carbon α to the nitrogen but still retains high Brønsted basicity. To examine the behavior of an even weaker, more hindered, Lewis base and its reaction mechanism with CO2 in aqueous solution, we synthesized hindered secondary amine MAMP, which represents an N-methylated version of the hindered primary amine AMP, and prepared its 3 M solution in water (SI Figure S1.13, top). It is remarkable that, unlike AMP discussed earlier, no carbamate species were detected at 164−165 ppm as pure CO2 at 1.0 bar was introduced into the aqueous solution of secondary hindered amine MAMP at 30 °C and 10 cm3(STP) min−1 (Figure 1d and Scheme 5). This result implies that

Figure 2. Evolution of the reaction of CO2 with 3 M MAMP in H2O at 30 °C monitored by 13C NMR spectroscopy. Quantitative CO2 product speciation and solution pH. Pure CO2 at 1.0 bar was purged at 10 cm3(STP) min−1.

90 mol % of less sterically hindered AP molecules form an N− C bond (as carbamate species) under similar conditions (SI Figure S1.7). The low concentration of AMP carbamate suggests lower stability of its N−C bond. In analogy with tertiary amines (Scheme 2), we interpret the first resonance at 165.2 ppm as a mixture of carbonate and bicarbonate with higher concentration of carbonate species at the high pH of the starting solution. With more CO2 introduced into the solution, this CO peak increases in intensity and shifts upfield to 160.7 ppm indicating higher concentration of bicarbonate in the solution. This is explained by the two mechanisms. With increasing protonation of amines, the solution pH decreases favoring the formation of bicarbonate species from either carbonic acid or from carbonate.1 Additionally, the less stable AMP carbamate (peak at 164.4 ppm) is more prone to hydrolysis into bicarbonate at low pH. Due to its low stability, no carbamate was detected at equilibrium, which was reached after approximately 3 h of CO2 purge at 5 cm3(STP) min−1. The total amount of reacted CO2 achieved 0.99 CO 2 per AMP and represents mainly bicarbonate. This finding implies that Brønsted basicity of AMP dominates its reaction properties with CO2 in aqueous solution. It is important to stress that the steric hindrance results in formation of an unstable carbamate that facilitates the increased rate of CO2 transfer to form the bicarbonate. Effectively, (low) Lewis basicity of AMP promotes the bicarbonate formation through direct nucleophilic attack of a free CO2 (Scheme 4). This route is expected to be faster than direct carbonic acid deprotonation (Scheme 3), which reaction is limited by the (low) concentration of the carbonic acid at low CO2 pressure. Figure 1b (top) also indicates a minor CO resonance at 158.3 ppm which represents the trace amount of O-

Scheme 5. Hindered Secondary Amine MAMP Not Carboxylated by CO2 in Aqueous Solution

nucleophilicity of the amino group of MAMP is further reduced relative to the already weak Lewis base AMP. Three methyl groups around the amine nitrogen create a barrier and inhibit the approach of a CO2 molecule to the nitrogen atom to create an N−C bond. Methyl groups on nitrogen and adjacent carbon donate electrons increasing Brønsted basicity of the amine. Hindered secondary amine MAMP behaves similarly to tertiary amines as a proton acceptor by forming exclusively carbonate and bicarbonate species with CO2 in water. The evolution of the CO2 reaction with MAMP monitored by 13C NMR (Figure 1d) confirms that, at an early reaction stage when solution pH is high (Figure 2), MAMP preferentially forms carbonate species (167 ppm). With more CO2 introduced into the solution, C O resonance increases in intensity and shifts upfield implying that the MAMP solution absorbs CO2 in the form of bicarbonate by deprotonation of carbonic acid (Scheme 6). At lower pH, the second route is possible; the carbonate reacts with carbonic acid and forms two molecules of bicarbonate. At equilibrium, CO2 loading achieved 0.95 CO2 per MAMP molecule predominantly as bicarbonate represented by a resonance at 160.67 ppm (Figure 1d). In the described reaction pathway, the rate limiting step is the formation of carbonic acid by nucleophilic attack of water (a weak Lewis base) on acidic CO 2 . The subsequent carbonic acid

Scheme 3. First Mechanism of Bicarbonate Formation with Hindered Amine AMP by Deprotonation of Carbonic Acid

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Scheme 6. Reaction Pathway of Hindered Secondary Amine MAMP with CO2 in Aqueous Solution: Formation of Carbonate at High pH and Bicarbonate (Directly by Carbonic Acid Deprotonation or through Carbonate) at Lower pH

Scheme 7. Pathway of the Reaction of CO2 with an Aqueous Solution of Hindered Secondary Amine MAP: Nucleophilic Attack on Free CO2 with Formation of Carbamate Followed by Partial or Complete Hydrolysis into (Bi)carbonate

splitting in the 13C NMR spectrum into two peak clusters: the first representing a MAP−carbamate anion structure (HOCH2CH(CH3)NH(CH3)−COO−) which correlates with the CO peak at 164.4 ppm; the second an ammonium−MAP counterion structure (HOCH2CH(CH3)NH2+(CH3)). With more CO2 purged through the solution, the intensity of the peak at 164.4 ppm (e.g., concentration of carbamate) increased with time and reached its maximum after approximately 10 min of CO2 flow at 10 cm3(STP) min−1. 13 C NMR indicates that approximately 50 mol % of MAP formed carbamate with CO2 and a total loading reached approximately 0.25 CO2 groups per all amines in the solution, as shown in SI Figure S1.15. At this point, the solution pH dropped to 11.0 from 12.3 and the carbamate peak at 164.4 ppm began to decline as a second resonance appeared at 164.2 ppm and grew larger, shifting upfield (Figure 1c and SI Figure S1.15). We assign this peak as a mixture of bicarbonate and carbonate in equilibrium. These species are likely formed by hydrolysis of the carbamate, which is confirmed by the reduction of carbamate concentration in the solution. With more CO2 introduced into the solution and more amine protonated, solution pH continues to decrease favoring formation of bicarbonate. Based on the observed evolution of CO2 reaction with MAP, we conclude that Lewis basicity of MAP is unlocked. MAP can directly attack free CO2 delivered into the solution and forms carbamate (Scheme 7). However, MAP−carbamate is less stable than MEA−carbamate or AP−carbamate, and it hydrolyses into bicarbonate after the pH of the solution is reduced. The intermediate formation of carbamate is very important for the overall reaction rate. In contrast to MAMP under similar conditions, which also formed bicarbonate through carbonic acid, the bicarbonate from MAP is formed through the MAP− carbamate, which can exist in the solution at significantly higher concentration (up to 60 mol % in the described example). The CO2 reaction mechanism of MAP is similar to AMP, which also forms an unstable carbamate as an intermediate product (Figure 1b). However, MAP appears to be slightly more nucleophilic and captures the majority of purged CO2 as carbamate while the slightly less nucleophilic AMP forms bicarbonate more directly, predominantly though carbonic acid. 2. Effect of Hindered Amine Concentration, Temperature, and CO 2 Partial Pressure on Amine−CO 2

deprotonation and carbonate−bicarbonate interchange are considerably faster. Since the carbonic acid concentration is low and the rate of formation at 1.0 bar of CO2 and 30 °C is slow, the overall rate of bicarbonate formation represented by Scheme 6 is slower than the formation of bicarbonate through hydrolysis of carbamate (Scheme 4). For example, 3 M AP and 3 M AMP reached equilibrium after approximately 180 min of purging pure CO2 through the solution at 5 cm3(STP) min−1 (SI Figures S1.7 and S1.9) while 3 M MAMP did not reach an equilibrium after 360 min of the CO2 flow under the same conditions (SI Figure S1.14). When CO2 was bubbled at a higher flow rate of 10 cm3(STP) min−1, the equilibrium with 3 M MAMP solution in water was achieved after approximately 180 min (Figure 2). To summarize, the hindered secondary amine MAMP effectively reacts with CO2 reaching a high loading of 0.95 CO2 per MAMP at 1.0 bar of CO2 and 30 °C. Due to predominant formation of bicarbonate species, CO2 loading capacity of MAMP is significantly higher than that of unhindered primary and secondary amines MEA and MAE, which act as strong Lewis bases and mainly form carbamate species. Due to severe steric hindrance of the amine nitrogen and resulting lack of Lewis basicity of MAMP, it is unable to directly attack a free CO2 and form carbamate species with fast reaction rates. Instead, MAMP deprotonates carbonic acid present in the solution at low concentration. The latter pathway slows the overall reaction rate and makes CO2 capture with MAMP less effective unless MAMP is used with promoters (discussed later). 2-(Methylamino)-1-propanol. The previously described hindered secondary amine MAMP cannot form a stable N−C bond in aqueous solution. MAMP can be considered as a too weak Lewis base due to severe hindrance of the amine nitrogen with three surrounding methyl groups. To reduce the steric hindrance around the amine, we synthesized the hindered secondary amine MAP, which represents an N-methylated version of the moderately hindered primary amine AP, and prepared its 3 M solution in water. SI Figure S1.16, top, shows that the MAP solution contains approximately 3 mol % AP as a side product of the MAP synthesis. As pure CO2 was introduced into the MAP aqueous solution at 10 cm3(STP) min−1, a peak at 164.4 ppm was detected (Figure 1c). We interpret this peak as carbamate because the backbone carbons of the MAP molecule show related peak 1228

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Energy & Fuels Reaction. In our previous work,1,2 we showed that the concentration of CO2−amine reaction products and total CO2 loading in the solution varies significantly with changing CO2 partial pressure and solution temperature. This effect is especially important for less stable reaction products such as bicarbonate. For tertiary amines in aqueous solution, bicarbonate−carbonate equilibrium is shifted toward carbonate at lower partial pressure of CO2 while the carbamate− bicarbonate equilibrium with primary and secondary amines in aqueous solution is shifted toward carbamate. The highest CO2 loading capacity of primary, secondary, and tertiary amines was achieved at the lowest temperature studied. As solution temperature increases, measured solution pH decreases indicating the drop of the effective amine Brønsted basicity (and potentially Lewis basicity), which leads to lower amines−CO2 reaction yield. Particularly, primary and secondary amines keep acting as strong Lewis bases and form predominantly carbamate species at elevated temperature. The effect of amine concentration in aqueous solution on reaction yield was less intuitive to expect. Tertiary amines at low concentrations (7 M for DMAE) showed a lower sorption capacity due to the formation of less bicarbonate and more carbonate species. The different behavior of the more concentrated solution can be explained by exceeding the critical amine/solvent ratio above which the equilibrium formation of divalent carbonate species is favored over monovalent bicarbonate. Similar effects were detected with primary and secondary amines, which tend to form carbamate species (2:1 amine/carbon dioxide) over monovalent bicarbonates at high amine concentration: higher pH. Detailed information on behavior of strong Lewis bases MEA and MAE and weak Lewis base DMAE is described in ref 1. To study the effects of solution temperature, amine concentration, and CO2 partial pressure on CO2 reaction with hindered amines, we prepared aqueous solutions of hindered primary amines AP and AMP, hindered secondary amines MAP and MAMP, and primary amine MEA as an example of a strong Lewis base. These solutions were treated with a gas containing 10 mol % CO2/90 mol % N2 (0.1 bar CO2 in N2), then switching the reaction gas to pure (1 bar) CO2, and then finally switching to 1 mol % CO2/99 mol % N2 (0.01 bar of CO2). Experiments above ambient pressure were performed in a batch mode by introducing CO2 in a head space of a specially designed high pressure NMR tube, which was connected to the ballast volume. 1H and 13C NMR samples were taken after every experiment. The described experiments were conducted at 45, 60, 90, and 120 °C to evaluate a temperature effect on CO2−amine reaction. CO2 Partial Pressure. Figure 3, top, shows the CO2−amine thermodynamic equilibrium as a function of CO2 partial pressure for selected amines at a fixed amine concentration (3 M) and fixed solution temperature (45 °C). The effect of CO2 partial pressure is more pronounced for hindered amines, which predominantly make (bi)carbonate reaction products (AMP and MAMP), and less pronounced for nucleophilic amines, which make more stable carbamate species (MEA and AP). As shown in Figure 3, top, at a given amine concentration and temperature (3 M and 45 °C), the CO2/amine equilibrium loading for the hindered amines AMP and MAMP changes over the range of 0.3−0.5 (in carbonate) at 0.01 bar of CO2 to 0.95

Figure 3. Comparison of CO2/amine equilibrium for the series of hindered primary and secondary amines AP, AMP, MAP, and MAMP with MEA at 3 M in water as a function of CO2 partial pressure at 45 °C (top) and as a function of temperature at a fixed CO2 partial pressure of 1.0 bar (middle) and CO2 uptake capacity as carbamate (open squares) and bicarbonate (open circles) of MAP dissolved in H2O at various concentrations with CO2 at 1.0 bar and 45 °C (bottom).

at 1.0 bar of CO2 and to 1.00 (in bicarbonate) at 10.0 bar of CO2. The nucleophilic amines AP and MEA show a smaller change of CO2 loading at higher CO2 pressure. At low CO2 pressure, carbamate is the predominant reaction product with nucleophilic amines AP and MEA while (bi)carbonate formation is less favorable. At higher partial pressure of CO2, 1229

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less hindered primary amine AP behaved similarly to nucleophilic secondary amine MAE, which we reported earlier.1 We describe the details of the amine concentration effect on CO2−amine equilibrium using a hindered secondary amine MAP with lower Lewis and high Brønsted basicity as an example. We showed that MAP is nucleophilic enough to attack a free CO2 and form an N−C bond. The resulting MAP− carbamate is less stable than the carbamate formed with a strong Lewis base (e.g., MEA) and can be either hydrolyzed into bicarbonate by water or decomposed into free amine and free CO2 at higher temperature. Figure 3, bottom, shows the CO2−MAP product speciation after reaction of CO2 at 1.0 bar with MAP dissolved in water over the range of concentrations at 45 °C. The highest CO2 loading of 0.95 per amine was achieved at lowest MAP concentration (3 M) with 90 mol % MAP as bicarbonate and 10 mol % as carbamate. Higher amine concentrations favor carbamate formation because higher solution pH is maintained, disfavoring bicarbonate formation. We observe the decrease of MAP−bicarbonate at higher amine concentrations while the fraction of MAP−carbamate increases resulting in lower total CO2 loading (Figure 3, bottom). In more dilute amine solution, the amine is titrated with CO2 more rapidly, causing a more rapid rate of pH decrease and shifting the reaction equilibrium toward bicarbonate. The concentration effect can be explained in terms of amine concentration range which favors the equilibrium formation of bicarbonate with a single MAP molecule solvated by water (at lower concentration) versus formation of carbamate involving two MAP molecules solvated by water (at higher concentration). 3. Reaction of CO2 with Hindered Amines in Methanol Solution. The hindered amines form weaker N−C bonds with CO2 and preferentially act as proton acceptors by stabilization of bicarbonate with CO2. Therefore, it is important to understand the mechanism of bicarbonate formation. We showed that water plays a dual role in the reaction of CO2 with amines. First, water forms carbonic acid with CO2, which can be deprotonated by the basic amine to form a carbonate or bicarbonate (Scheme 2). Alternatively, water can hydrolyze carbamate into a bicarbonate (Scheme 1). In both cases, water acts as a nucleophile by attacking the carbon of CO2 and creating an O−C bond with CO2. A similar mechanism can work with other molecules containing hydroxyl groups, for example alcohols.27−31 The oxygen atom of alcohols can be turned into nucleophiles by removing the hydrogen with, for example, a basic amine. The resulting oxygen nucleophile can now attack the carbon of CO2 to form an alkylbicarbonate CH3O−COO− +H−amine. In our previous work we showed that the hydroxyl group of an alkanolamine can act as a nucleophile for CO2 in nonaqueous solution and be stabilized in the presence of a strong Brønsted base (pKa > 12).3 We studied CO2 reactions with the hindered amine MAP in alcohol solution and compared results with an aqueous solution of MAP. Negligible acidity and low steric hindrance makes methanol a preferred choice to mimic properties of water. Due to the high volatility of methanol, these experiments were conducted in a batch mode by pressurization of the amine solution with CO2 at 10.0 bar in the NMR tube located inside the NMR spectrometer and recording 13C NMR spectra every 2 min to monitor the pathway of the chemical reaction. Due to the absence of solution mixing and lower gas−liquid contact area, the total time to reach equilibrium in a batch mode is longer than in flow-through mode described in the previous

the carbamate/(bi)carbonate equilibrium is able to shift toward (bi)carbonate to increase CO2 loading in the amine solution. At low amine concentration and low temperature (3 M and 45 °C), CO2 loading capacity of hindered amines AMP, MAMP, and MAP is higher than that of nucleophilic amines MEA and AP due to formation of bicarbonate species. Tables S1−S3 of the Supporting Information give details of CO2 loading capacity of aqueous solutions of AP, AMP, and MAP at various amine CO2 partial pressures. Solution Temperature. Figure 3, middle, compares the temperature effect on the CO2−amine thermodynamic equilibrium for weak Lewis bases AMP, MAMP, and MAP with stronger Lewis bases AP and MEA at a fixed amine concentration (3 M) and CO2 pressure (1.0 bar). The highest CO2 loading capacity of all solutions was achieved at the lowest temperature studied, 30 °C. It is worth noting, that hindered amines AMP, MAMP, and MAP form predominantly bicarbonate species at these conditions with loading of approximately 1.0 CO2 per amine, while stronger Lewis bases AP and MEA form bicarbonate at lower concentrations60 and 40 mol %, respectively. At higher solution temperature, CO2 loading of all solutions decreases implying the shift of bicarbonate−carbonate equilibrium to carbonate with hindered amines AMP and MAMP and shift of bicarbonate−carbamate equilibrium to carbamate with nucleophilic amines AP and MEA. This shift is affected by the changes of solution pH at elevated temperatures. The pH of the starting (CO2-free) amine solution decreases as a function of temperature.1 The lower pH value of the starting solution at elevated temperature leads to even lower solution pH after initial carbonate (with hindered amines) or carbamate (with nucleophilic amines) formation. At conditions when solution pH is approaching a neutral value after initial carbonate or carbamate formation, hydrolysis into the bicarbonate is more favorable. As a result of the lower thermal stability of bicarbonate, CO2 capacity of amines which predominantly form bicarbonate (AMP, MAMP, and MAP) is reduced at elevated temperature making them desirable candidates for a cyclic CO2 capture process. At 90 °C, all amines under study show comparable CO2 loading of about 0.5−0.7 with reaction products as either carbonate (AMP and MAMP) or carbamate (AP, MAP, and MEA). However, carbamate with weaker Lewis base MAP is significantly less stable than carbamate with a strong Lewis base MEA. Tables S1−S3 of the Supporting Information give details of CO2 loading capacity of aqueous solution of AP, AMP, and MAP at various solution temperatures. Amine Concentration. We studied the effect of amine concentration on the CO2−amine reaction equilibrium and found that hindered amines AMP and MAMP precipitate from the solution at concentrations above 3 and 5 M, respectively, after reaction with CO2. Precipitation can be caused by the high concentration of a less soluble bicarbonate salt in aqueous solution. Phase diagrams for solid/slurry formation with AMP dissolved in aqueous and nonaqueous (DMSO) solutions at various concentrations after reaction with CO2 at fixed partial pressure (0.1 and 1.0 bar) are included in the Supporting Information (Figures S2.3 and S2.4). While precipitation can interfere with liquid amine approaches to carbon capture, it is being applied by others: e.g., by Alstom using chilled ammonia, by IFP using solution phase separation approaches, and by GE using a solids formation approach.14−26 On the other hand, the 1230

DOI: 10.1021/acs.energyfuels.5b02582 Energy Fuels 2016, 30, 1223−1236

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Energy & Fuels

bonate. The rate of MAP−carbamate formation is significantly higher due to a direct nucleophilic attack of amine on either gaseous CO2, in the head space, or with CO2 diffused into the solution. Initially, the growing concentration of MAP− carbamate reached 40 mol % after approximately 1 h of reaction and began to decline due to decomposition of MAP− carbamate into MAP−bicarbonate and MAP−methylbicarbonate. It is remarkable that the two latter species were formed with similar relative rates and existed at similar concentrations in the solution. After 3 h, CO2 loading achieved approximately 0.8 CO2 per MAP with 38 mol % MAP in bicarbonate form, 38 mol % in methylbicarbonate, and the remaining 14 mol % as carbamate. Complete equilibrium was achieved after 9 h with approximately 50 mol % MAP forming bicarbonate with water, 47 mol % MAP forming methylbicarbonate with methanol, while only 3 mol % MAP remained as carbamate. 13C NMR spectroscopy confirms the participation of the methanol molecules in the reaction. The initial methanol peak splits into two peaks representing unreacted (CH3OH, 49.19 ppm) and reacted methanol (CH3O−COO−, 52.90 ppm). In the solution 17 mol % methanol molecules participated in the reaction (SI Figure S4.1, bottom). The experiment confirmed the ability of methanol to act similarly to water by creating an O−C bond and forming methylbicarbonate with CO2 and amine (Scheme 8). We

sections. Equilibration time in a batch mode primarily depends on CO2 solubility, CO2 diffusion rate into the liquid solution, and solution mixing due to chemical and density gradients. Results obtained in batch mode for solutions at similar conditions (e.g., concentration, temperature, and CO 2 pressure) can provide relative comparisons. Figure 4, top, shows the evolution of chemical reaction of 3 M MAP in a 1:1 mixture of methanol and water with CO2 at

Scheme 8. Proposed Mechanisms of Methylbicarbonate Formation by CO2 Reaction with Methanol Solution of Hindered Secondary Amine MAP: Nucleophilic Attack on Carbamate (Top) or Deprotonation of Methylcarbonic Acid (Bottom)

Figure 4. Evolution of the reaction of CO2 with 3 M MAP in 1:1 MeOH/H2O mixture at 45 °C monitored by 13C NMR spectroscopy. The solution was pressurized by pure CO2 at 10.0 bar. Formation of carbamate, (bi)carbonate, and methylbicarbonate species (top) and quantitative CO2 product speciation (bottom).

expect other alcohols such as ethanol, propanol, and 2propanol, etc., to show even lower affinity toward the carbon of CO2 due to their lower acidity and some steric hindrance. 4. Promotion of the Reaction of CO2 with Hindered Amines. Tertiary and hindered amines with reduced Lewis basicity and relatively high Brønsted basicity such as DMAE, AMP, MAMP, and MAP exclusively or preferentially form bicarbonate with a more beneficial 1:1 CO2/amine ratio. Stronger Lewis bases such as MEA preferentially form carbamate with lower CO2 capacity1:2 CO2/amine ratio. However, the reaction of weaker Lewis bases with CO2 is slower because their ability to directly react with gaseous CO2 is either lower (for AMP and MAP) or negligible (for DMAE or MAMP) and the total rate of bicarbonate formation depends on (low) carbonic acid concentration and rate of formation in basic aqueous solution. The presence of a strong Lewis and weaker Brønsted base in the solution can speed up the CO2 capture rate by direct nucleophilic attack on gaseous CO2 and its transformation to a stronger Brønsted base (hindered) or tertiary amine. To check

10.0 bar monitored by 13C NMR spectroscopy. Similar to reaction in water solution (SI Figures S4.2 and S4.3), MAP simultaneously forms carbamate and carbonate/bicarbonate (peaks at 164.8 and 162.2 ppm). 13C NMR spectroscopy detects the third peak at 160.0 ppm, which was never observed in water solution. Unlike the peak at 162.2 ppm, which represents the carbonate/bicarbonate mixture and shifts upfield due to higher concentration of bicarbonate, the peak at 160.0 ppm does not shift implying that the associated reaction product is not in exchange with other species. We interpret this third peak at 160.0 ppm as methylbicarbonate, which is formed by either a nucleophilic attack of MeOH on MAP−carbamate or by deprotonation of methyl carbonic acid formed by reaction of CO2 with methanol: Figure 4, bottom, shows the concentrations and relative rates of formation of carbamate, (bi)carbonate, and methylbicar1231

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Energy & Fuels this hypothesis, we studied the CO2 reaction mechanism with solutions containing 1 M piperazine (PZ, 8) as a “promoter” and 3 M of hindered amines AMP, MAP, and MAMP or tertiary amine DMAE, which showed slower equilibration time relative to MEA. PZ. The CO2 reaction with PZ (pKa, 9.55, 5.41) in aqueous solution was studied in detail by Bishnoi et al. and Cullinane and Rochelle.32−34 In our previous work1 we also confirmed that PZ acts as a strong nucleophile preferentially forming carbamate with CO2, some of which is hydrolyzed by water into bicarbonate (Scheme 9). At the conditions discussed in this

Scheme 10. Proposed PZ Promoted Mechanism of the Reaction of CO2 with an Aqueous Solution of Hindered Amine MAMP and Stronger Lewis Base PZ: Nucleophilic Attack of PZ on Free CO2 with Stabilization by MAMP Followed by Either Hydrolysis into Bicarbonate or Reaction with Another Molecule of CO2

Scheme 9. Mechanism of the Reaction of CO2 with an Aqueous Solution of Piperazine: Nucleophilic Attack on Free CO2 with Formation of Carbamate Followed by Partial Hydrolysis into Bicarbonate

work (1.0 bar of CO2, 30 °C), the amount of CO2 in carbamate form (0.60−0.65 CO2/PZ) was approximately double the CO2 as bicarbonate (0.30−0.35 CO2/PZ). A total loading of approximately 1.0 CO2 per PZ molecule with 0.65 CO2/PZ as carbamate (SI Figures S5.1 and S5.2) also implies that two nitrogen atoms of the same PZ molecule can act as Lewis base to accept a carbon of CO2 and as a Brønsted base to accept a proton and stabilize either bicarbonate or carbamate species. After protonation of the first nitrogen, the second nitrogen is less basic (second pKa value of 5.41), but may remain nucleophilic. In this case, one piperazine molecule can potentially participate in the formation of intermolecular reaction species in combination with CO2, bicarbonate and carbamate units on neighboring molecules. MAMP/PZ. In section 1 we showed that hindered secondary amine MAMP is a very weak Lewis base. Lacking nucleophilic properties, MAMP is unable to directly attack the carbon of CO2 and form carbamate species. Instead, MAMP relies on its Brønsted basicity and deprotonates carbonic acid into carbonate and bicarbonate. Due to the low concentration of carbonic acid, the total equilibration time of the MAMP−CO2 chemical reaction was the longest within our series of hindered amines. To compare the CO2−MAMP reaction mechanism with and without a strong nucleophile PZ, we prepared an aqueous solution containing 3 M MAMP and 1 M PZ (SI Figure S5.4, top), measured pH ∼ 12.5, and treated it with CO2. Summarizing the observed chemical reaction between CO2, MAMP, and PZ in aqueous solution, we observed the following reaction mechanism (Scheme 10). Strong Lewis base PZ attacks CO2 by initially forming a zwitterion/carbamic acid, which is immediately deprotonated by either another PZ molecule or more likely by a stronger Brønsted base MAMP (Figure 5). The part of the formed mixed carbamate between PZ, CO2, and MAMP can be hydrolyzed by water molecules into the MAMP−bicarbonate and free PZ, which can capture the next gaseous CO2 molecule. After carboxylation of one nitrogen atom of PZ, the second nitrogen remains a sufficiently strong Lewis base and is able to attack another CO2 molecule, which will again be stabilized by the MAMP. Actually, the presence of more basic MAMP in the solution helps PZ to

Figure 5. Evolution of the reaction of CO2 with an aqueous solution containing 3 M MAMP and 1 M PZ at 30 °C monitored by 13C NMR spectroscopy. Pure CO2 at 1.0 bar was purged at 10 cm3(STP) min−1. Formation of carbamate with PZ and (bi)carbonate with MAMP.

attack CO2 rather than stabilize the already captured CO2. Brønsted base MAMP exclusively acts as a proton acceptor by stabilizing mixed carbamate with PZ and deprotonation of carbonic acid into carbonate (at high pH) or bicarbonate (at lower pH). Additionally, the hydroxyl group of a few MAMP molecules reacts with CO2 forming the O-carbonated form, which is then stabilized by another MAMP molecule. Due to the ability of PZ to directly react with gaseous CO2, the total equilibration time of the CO2−PZ−MAMP−H2O reaction is about 2 times shorter than the total reaction time of CO2−MAMP−H2O solution at similar conditions (SI Figure S5.5 vs Figure 2). PZ acts as a CO2 transfer agent from the gas phase into a mixed carbamate or bicarbonate with an amine. Higher CO2 loading capacity represents the additional benefit of a solution containing PZ to act as a promoter of chemical reaction with CO2. It is important to note that integration of the PZ−carbamate peak cluster at 162−163 ppm over PZ structural carbons at 45−40 ppm (SI Figure S5.4, bottom) gives 1.28 CO2 per PZ (in PZ−carbamate form) and suggests that approximately 71 mol % PZ molecules are monocarboxylated by CO2 (1 CO2 per PZ) while the remaining 29% PZ are doubly carboxylated with CO2 located on both nitrogen atoms (2 CO2 per PZ) and stabilized by protonated MAMP. Two configurations of PZ carboxylation give two overlapping 13C peaks at 162.76 and 162.03 ppm. Double carboxylation of PZ molecules also implies 1232

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Energy & Fuels that at least a part of the CO2 that reacted with PZ is stabilized with MAMP and formed a mixed carbamate PZ−CO2− +H− MAMP. AMP/PZ and MAP/PZ. Primary amine AMP is less sterically hindered relative to MAMP and able to directly attack CO2 by forming carbamate species (Figure 1b). However, the reaction mechanism of the aqueous solution of AMP with a small amount of stronger Lewis base PZ with CO2 is qualitatively similar to the MAMP/PZ mixture (SI Figures S5.7−S5.9). In section 1 we showed that MAP is a stronger Lewis base relative to AMP (and MAMP). MAP directly attacks gaseous CO2 delivered to the solution and forms a relatively stable carbamate, which only partially hydrolyzes in aqueous solution into bicarbonate (Figure 1c), unlike AMP−carbamate, which is completely hydrolyzed into bicarbonate (Figure 1b). In terms of reaction rates with CO2, the solution containing nucleophilic MAP does not benefit from an additional nucleophilic amine (such as PZ) as much as AMP and especially MAMP, which does not attack CO2. With the exception of carbamate formation by the hindered amine, the overall reaction mechanism and total amount of captured CO2 by MAP−PZ solution is similar to AMP−PZ and MAMP−PZ described earlier (Scheme 10). A strong nucleophilic base such as PZ present in the amine solution at low concentration helps to speed up the reaction rate by preferentially attacking the gaseous CO2 and transferring it to the amine (as bicarbonate), but does not consume the amine capacity by forming a mixed carbamate with it. The best reaction promoters are those amines with the highest possible Lewis basicity but Brønsted basicity similar to or even lower than Brønsted basicity of the promoted amine. The results of the reaction of CO2 with nonhindered and hindered amines promoted by strong Lewis base PZ at various concentrations in the solution are summarized in SI Tables S4 and S5.

Scheme 11. Mechanism of the Reaction of CO2 with Aqueous Solution of Hindered Amines: Nucleophilic Attack on Free CO2 with Formation of Carbamate Followed by Partial or Complete Hydrolysis into (Bi)carbonate

basicities of hindered amines favors a higher concentration of bicarbonate and reduced concentration of carbamate in the solution, which under certain conditions doubles the capacity of CO2 absorption. Second, bicarbonates are generally less stable than carbamates and offer an option to reduce the energy consumption associated with amine regeneration (e.g., CO2 desorption). Third, carbamates with hindered amines are less stable than carbamates with unhindered amines such as MEA. Due to their high basicity and lower nucleophilicity, hindered amines allow fine-tuning of the distribution of CO2−amine reaction products. Thus, formation of carbamates, which are undesirable from low capacity (1 CO2 per 2 amines) and higher ΔHf (which requires higher temperature regeneration) points of view, can be partially or completely suppressed using selected hindered primary or secondary amines. In this case, CO2 will be captured in the form of a bicarbonate, which is beneficial due to high CO2 capacity (1 CO2 per 1 amine), lower ΔHf, and lower viscosity and corrosivity. Lower stability of bicarbonate formed from hindered amines also offers amine regeneration (e.g., CO2 desorption) at lower temperature or higher CO2 pressure than the carbamate with unhindered amines, substantially reducing process energy requirements and regeneration costs. Unlike tertiary amines, which also offer exclusive formation of carbonates and bicarbonates in aqueous solution at low rates, the advantage of hindered amines is their relatively high reaction rates driven by an intermediate/ transient step of carbamate formation. The number and location of additional methyl (or other) groups on nitrogen or adjacent carbon atoms create different degrees of steric hindrance. The presence of one methyl group on the carbon adjacent to the nitrogen atom of AP creates a moderate steric hindrance and slightly reduces the nucleophilicity of the amine. The reaction mechanism of AP with CO2 in aqueous solution is similar to nonhindered amines such as primary amine MEA. However, the carbamate of primary amine AP more easily hydrolyzes to the bicarbonate with increased CO2 loading, which is close to secondary amine MAE rather than to primary amine MEA. Two methyl groups on the α-carbon significantly reduce the Lewis basicity while maintaining the high Brønsted basicity of the primary amine group in AMP and shift the reaction equilibrium more toward bicarbonate. Carbamate with AMP was detected at low concentrations at early stages of reaction, but it quickly hydrolyzes into bicarbonate in the presence of water. Effectively, bicarbonate formation with AMP is promoted through direct nucleophilic attack of AMP on a free CO2 (to form carbamate) followed by fast hydrolysis. This route is expected to be faster than direct carbonic acid deprotonation, which reaction is limited by the (low) concentration of the carbonic acid at low CO2 pressure. No carbamate was detected at equilibrium on reaction of AMP with CO2 at 30 °C due to the low stability of the AMP

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SUMMARY AND CONCLUSIONS In situ 13C NMR mechanistic studies elucidated the pathways of the reaction of CO2 with hindered alkanolamines in aqueous and nonaqueous media. Insights gained from the previous aqueous studies1 and the new pathways gleaned from the nonaqueous work2,3 made us realize that Lewis basicity (nucleophilicity) of the amine is as important for reaction with CO2 as traditionally discussed Brønsted basicity. The reactivity of hindered alkanolamines described in this work needs to be elucidated and differentiated from other primary and secondary amines/alkanolamines in terms of reduced affinity of the amine nitrogen for the carbon atom of CO2. Similar to tertiary amines, hindered amines tend to form bicarbonates with more favorable CO2 capacity: 1:1 CO2 per amine. In contrast to tertiary amines and similar to primary and secondary unhindered amines such as MEA, the reaction of CO 2 with certain hindered amines proceeds through intermediate carbamate formation followed by partial or complete hydrolysis into bicarbonate (Scheme 11). Such a reaction mechanism is favorable in terms of relatively high reaction rates and leads to high CO2 loading capacity. Why does the reactivity of hindered amines need to be differentiated from other primary, secondary, and tertiary amines/alkanolamines? First, hindered amines have reduced Lewis basicity maintaining similar or even higher Brønsted basicity when compared with corresponding MEA-like unhindered amines. Such interplay of Lewis and Brønsted 1233

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Scheme 12. Mechanism of the Reaction of CO2 with an Aqueous Solution of Hindered Amine and Stronger Lewis Base PZ: Nucleophilic Attack of PZ on free CO2 with Amine Stabilization Followed by Hydrolysis into Bicarbonate

carbamate. This finding implies that Brønsted basicity of AMP dominates its reaction properties with CO2 in aqueous solution. The highly hindered secondary amine MAMP with one methyl group on nitrogen and two methyl groups on the αcarbon can be considered as a too weak Lewis base to directly attack CO2 and form carbamate species with fast reaction rates. Instead, secondary amine MAMP reacts similarly to tertiary amines such as DMAE by deprotonation of carbonic acid present in the solution at low concentration. The latter pathway slows the overall reaction rate and makes CO2 capture with MAMP less effective unless MAMP is used with the promoters. Due to predominant formation of bicarbonate species, CO2 loading capacity of MAMP is significantly higher than that of unhindered primary and secondary amines MEA and MAE, which act as strong Lewis bases and mainly form carbamate species. The presence of two methyl groups, one on nitrogen and a second on the α-carbon, unlocks Lewis basicity of secondary amine MAP. MAP acts similarly to primary hindered amine AMP (with two methyl groups on the carbon α to the nitrogen) via a direct attack of free CO2 delivered into the solution and forms an intermediate carbamate. However, MAP−carbamate is more stable than AMP−carbamate and only partially hydrolyzes into bicarbonate after the pH of the solution is reduced. MAP appears to be slightly more nucleophilic and captures a majority of the purged CO2 as carbamate while the slightly less nucleophilic AMP forms bicarbonate more directly, predominantly though carbonic acid. The intermediate formation of carbamate is very important for the overall reaction rate. In contrast to MAMP under similar conditions, which also formed bicarbonate through carbonic acid, the bicarbonate from MAP is formed through the MAP− carbamate, which can exist in the solution at significantly higher concentration. At higher solution temperature, CO2 loading decreases implying the shift of the bicarbonate−carbonate equilibrium to carbonate with more hindered amines AMP and MAMP and shift of the bicarbonate−carbamate equilibrium to carbamate with less hindered amines AP and MAP. This shift is affected by the changes of solution pH at elevated temperatures. As a result of the lower thermal stability of bicarbonate, amines which predominantly form bicarbonate (AMP and MAMP) are mostly affected. MAP also reduces CO2 loading capacity at elevated temperature, which reflects the low stability of MAP− bicarbonate and the N−C bond of MAP−carbamate. The effects of CO2 partial pressure on the reaction of CO2 with hindered amines is more pronounced for hindered amines, which predominantly make (bi)carbonate reaction products (AMP and MAMP), and less pronounced for hindered amines, which make more stable carbamate species (MAP and AP). At low CO2 pressure, carbamate is the predominant reaction product with nucleophilic amines AP and MAP while (bi)carbonate formation is less favorable. At higher partial

pressure of CO2, the carbamate/(bi)carbonate equilibrium is able to shift toward (bi)carbonate to increase CO2 loading in the amine solution. The CO2−amine reaction equilibrium and concentration of reaction products also depends upon amine concentration in the solution. For less hindered amines AP and MAP, higher amine concentrations favor carbamate formation because higher solution pH is maintained, disfavoring bicarbonate formation. In more dilute amine solution, the amine is titrated with CO2 more rapidly, causing a more rapid rate of pH decrease and shifting the reaction equilibrium toward bicarbonate. Hindered amines AMP and MAMP precipitate from the solution at concentrations above 3 and 5 M, respectively, after formation of bicarbonate salt with CO2. Methanol and other alcohols can be considered as a solvent for amines to replace water. Similarly to water, methanol can attack CO2 and form an O−C bond and methylbicarbonate with CO2 and amine. The mechanism of chemical reaction between CO2 and MAP dissolved in methanol is similar to that in water and involves formation of carbamate with amine, which is followed by partial or complete hydrolysis into the methylbicarbonate salt. The overall rates of the CO2 reaction with amine in methanol solution are also similar. However, nucleophilic properties of methanol to attack CO2 and form an O−C bond are weaker relative to water. We expect other alcohols such as ethanol, propanol, and 2-propanol, etc., to show even lower affinity to the carbon of CO2 due to their lower acidity and steric hindrance. A strong nucleophilic base such as PZ present in the amine solution at low concentration helps to speed up the reaction rate by attacking the free CO2 and transferring it to the amine (as bicarbonate) but does not utilize the amine capacity by forming a mixed carbamate with it (Scheme 12). The best reaction promoters are those amines with the highest possible Lewis basicity but Brønsted basicity similar to, or even lower than, the Brønsted basicity of the promoted amine. The viscosity of the amine solution is affected by hydrogen bonding interactions between amino sites (-NH2 or -NH-) and -OH groups with neighboring amine molecules and water. Addition of N-methyl groups to the nitrogen atom and converting a primary amine into a secondary and tertiary amine reduce solution viscosity and are explained by the reduction of amine−OH hydrogen bonding interactions due to steric hindrance of the amine group. In contrast, an α-methyl group adjacent to the amine carbon does not prevent an -OH group from approaching and hydrogen bonding with a primary amine but increases the kinetic diameter of the molecule, which leads to higher solution viscosity. Solution viscosity can be further reduced by turning off hydrogen bonding interactions between hydroxyl groups of amines. Methoxylated versions of alkanolamines are significantly less viscous with viscosities even lower than water. Such amines can be beneficial for a carbon capture approach because capping a hydroxyl group with a methyl (or 1234

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Energy & Fuels

(4) Kortunov, P. V.; Baugh, L. S.; Siskin. Mechanism of the Chemical Reaction of Carbon Dioxide with Ionic Liquids and Amines in Ionic Liquid Solution. Energy Fuels 2015, 29 (9), 5990−6007. (5) Bougie, F.; Iliuta, M. C. Sterically Hindered Amine-Bases Absorbents for Removal of CO2 from Gas Streams. J. Chem. Eng. Data 2012, 57, 635−669. (6) Versteeg, G. F.; Van Dijck, L. A. J.; Van Swaaij, W. P. M. On the kinetics between CO2 and alkanolamines both in aqueous and nonaqueous solitions. An overview. Chem. Eng. Commun. 1996, 144, 113− 158 and references cited therein.. (7) Park, S.-W.; Lee, J.-W.; Choi, B.-S.; Lee, J.-W. Absorption of carbon dioxide into non-aqueous solutions of N-methyldiethanolamine. Korean J. Chem. Eng. 2006, 23, 806−811. (8) 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. (9) Xu, S.; Wang, Y.-W.; Otto, F. D.; Mather, A. E. Kinetics of the Reaction of Carbon Dioxide with 2-Amino-2-methyl-1-propanol Solutions. Chem. Eng. Sci. 1996, 51, 841−850. (10) Sartori, G.; Savage, D. W. Sterically Hindered Amines for CO2 Removal from Gases. Ind. Eng. Chem. Fundam. 1983, 22, 239−249. (11) Say, G. R.; Heinzelmann, F. J.; Iyengar, J. N.; Savage, D. W.; Bisio, A.; Sartori, G. Treating Acid and Sour Gas: A New, Hindered Amine Concept for Simultaneous Removal of CO2 and H2S from Gases. Chem. Eng. Progress 1984, 72−77. (12) Goldstein, A. M.; Edelman, A. M.; Beisner, W. D.; Ruziska, P. A. Hindered amines yield improved gas treating. Oil Gas J. 1984, 70−76. (13) Aronu, U. E.; Svendsen, H. F.; Hoff, K. A.; Juliussen, O. Solvent selection for carbon dioxide absorption. Energy Procedia 2009, 1, 1051−1057 (see p 1052).. (14) Kozak, F.; Petig, A.; Morris, E.; Rhudy, R.; Thimsen, D. Chilled Ammonia Process for CO2 Capture. Energy Procedia 2009, 1, 1419− 1426. (15) 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. (16) Aleixo, M.; Prigent, M.; Gibert, A.; Porcheron, F.; Mokbel, I.; Jose, J.; Jacquin, M. Physical and Chemical Properties of DMX Solvents. Energy Procedia 2011, 4, 148−155. (17) Raynal, L.; Alix, P.; Bouillon, P.-A.; Gomez, A.; le Febvre de Nailly, M.; Jacquin, M.; Kittel, J.; di Lella, A.; Mougin, P.; Trapy, J. The DMX process: An original solution for lowering the cost of postcombustion carbon capture. Energy Procedia 2011, 4, 779−786. (18) Hu, L. Methods and Systems for Deacidizing Gaseous Mixtures. U.S. Patent 7,718,151, May 18, 2010. (19) Cadours, R.; LeComte, F.; Magna, L.; Barrere-Tricca, C. Method of Decarbonating a Combustion Fume with Extraction of the Solvent Contained in the Purified Fume. U.S. Patent 2006/0188423, Aug. 24, 2006. (20) Lemaire, E.; Raynal, L. IFP Novel Concepts for Postcombustion Carbon Capture. Capture and Geological Storage of CO23rd International Symposium, France; November 2009. (21) Lemaire, E.; Bouillon, P. A.; Gomez, A.; Kittel, J.; Gonzalez, S.; Carrette, P. L.; Delfort, B.; Mougin, P.; Alix, P.; Normand, L. New IFP optimized first generation process for post-combustion carbon capture: HiCapt+TM. Energy Procedia 2011, 4, 1361−1368. (22) Porcheron, F.; Gibert, A.; Jacquin, M.; Mougin, P.; Faraj, A.; Goulon, A.; Bouillon, P.-A.; Delfort, B.; Le Pennec, D.; Raynal, L. High throughput screening of amine thermodynamic properties applied to post-combustion CO2 capture process evaluation. Energy Procedia 2011, 4, 15−22. (23) Perry, B.; Grocela-Rocha, T.; Genovese, S.; Wood, B.; Westendorph, T.; O’Brien, M.; Meketa, M.; Draper, S.; Vipperla, R.K.; Enick, B.; Tapriyal, D. CO2 Capture Process Using PhaseChanging Absorbents, ARPA-E Project DE-AR0000084, Sep. 21, 2010; Advanced Research Projects Agency, U.S. Department of Energy: Washington, DC, USA.

ethyl) group does not change the chemical properties of amine to absorb and desorb CO2. The viscosities of carbamate and bicarbonate species differ. The higher viscosity of carbamate is shown and explained by both hydrogen bonding capability and doubling in effective molecular weight of the two interacted amines. On the other hand, bicarbonate species are less viscous because bicarbonate does not participate in hydrogen bonding and represents the product of a single amine reaction. The weaker Lewis base MAP and stronger Lewis base MEA after saturation with CO2 and formation of carbamate and bicarbonate at different concentrations still show comparable solution viscosity at similar concentration on an amine weight basis. Although hindered amine MAP shows CO2 loading capacity comparable with MAE, the lower stability of MAP−CO2 reaction products relative to MAE is the main advantage of MAP as a molecule for the cyclic CO2 capture. Easier desorption of MAP−CO2 reaction products is a reflection of a less stable N−C bond of MAP−carbamate. Lower carbamate stability became possible by reduction of Lewis basicity (or amine nucleophilicity) of MAP created by a methyl group on the nitrogen and another on the α-carbon atom. Weaker Lewis bases AMP and MAMP show 2-fold advantage over MEAgreater CO2 capacity and lower thermal stability of reaction products, which suggest benefits of decreased circulation rates and lower regeneration energy. However, the disadvantage of AMP and MAMP is the low solubility of their bicarbonate reaction products and relatively low reaction rates, which can be increased by the use of a stronger Lewis base promoter.

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.energyfuels.5b02582. 1 H and 13C NMR spectra of unreacted and reacted amines discussed in the work and experimental results with other hindered amines and promoter as well as viscosity of CO2-lean and CO2-rich solutions (PDF)

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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REFERENCES

(1) Kortunov, P. V.; 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 Fuels 2015, 29 (9), 5919−5939. (2) Kortunov, P. V.; Siskin, M.; Baugh, L. S.; Calabro, D. C. 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. Energy Fuels 2015, 29 (9), 5940−5966. (3) Kortunov, P. V.; Baugh, L. S.; Siskin, M.; Calabro, D. C. In Situ NMR Mechanistic Studies of Carbon Dioxide Reactions with Liquid Amines in Aqueous Systems: The Interplay of Lewis and Brønsted Basicity. Energy Fuels 2015, 29 (9), 5967−5989. 1235

DOI: 10.1021/acs.energyfuels.5b02582 Energy Fuels 2016, 30, 1223−1236

Article

Energy & Fuels (24) Perry, R. J.; O’Brien, M. J.; Lam, T. H.; Soloveichik, G. L.; Kniajanski, S.; Lewis, L. N.; Rubinsztajn, M. I.; Hancu, D. Liquid Carbon Dioxide Absorbent and Methods of Using the Same. U.S. Patent 2010/0154639 Jun. 10, 2010. (25) Perry, R. J.; Grocela-Rocha, T. A.; O’Brien, M. J.; Genovese, S.; Wood, B. R.; Lewis, L. N.; Lam, H.; Soloveichik, G.; Rubinsztajn, M.; Kniajanski, S.; Draper, S.; Enick, R. M.; Johnson, J. K.; Xie, H.; Tapriyal, D. Aminosilicone Solvents for CO2 Capture. ChemSusChem 2010, 3, 919−930. (26) Jou, F.-Y.; Mather, A. E.; Otto, F. D. The solubility of CO2 in a 30 mass percent of monoethanolamine solution. Can. J. Chem. Eng. 1995, 73, 140−147. (27) Heldebrant, D. J.; Yonker, C. R.; Jessop, P. G.; Phan, L. Organic liquid CO2 capture with high gravimetric CO2 capacity. Energy Environ. Sci. 2008, 1, 487−493. (28) Heldebrant, D. J.; Yonker, C. R.; Jessop, P. G.; Phan, L. CO2binding organic liquids (CO2BOLs) for post-combustion CO2 capture. Energy Procedia 2009, 1, 1187−1195. (29) Jessop, P. G.; Heldebrant, D. J.; Li, X.; Eckert, C. A.; Liotta, C. L. Green chemistry: Reversible nonpolar-to-polar solvent. Nature 2005, 436, 1102. (30) Heldebrant, D. J.; Yonker, C. R. Capture and release of mixed acid gasses with binding organic liquids. U.S. Patent Appl.2009/ 0136402 (Battelle), May 28, 2009. (31) Phan, L.; Chiu, D.; Heldebrant, D. J.; Huttenhower, H.; John, E.; Li, X.; Pollet, P.; Wang, R.; Eckert, C. A.; Liotta, C. L.; Jessop, P. G. Switchable solvents consisting of amidine/alcohol or guanidine/ alcohol mixtures. Ind. Eng. Chem. Res. 2008, 47, 539−545. (32) Bishnoi, S.; Rochelle, G. T. Absorption of Carbon Dioxide in Aqueous Piperazine/Methyldiethanolamine. AIChE J. 2002, 48, 2788. (33) Bishnoi, S.; Rochelle, G. T. Thermodynamincs of Piperazine.Methyldiethanolamine/Water/Carbon Dioxide. Ind. Eng. Chem. Res. 2002, 41, 604−612. (34) Cullinane, J. T.; Rochelle, G. T. Kinetics of Carbon Dioxide Absorption into Aqueous Potassium Carbonate and Piperazine. Ind. Eng. Chem. Res. 2006, 45, 2531−2545.

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DOI: 10.1021/acs.energyfuels.5b02582 Energy Fuels 2016, 30, 1223−1236