CO2 Deactivation of Supported Amines: Does the Nature of Amine

Feb 9, 2012 - CO2 Deactivation of Supported Amines: Does the Nature of Amine ... Biological Engineering, University of Ottawa, Ottawa, Ontario K1N 6N5...
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CO2 Deactivation of Supported Amines: Does the Nature of Amine Matter? Abdelhamid Sayari,*,†,‡ Youssef Belmabkhout,† and Enshirah Da’na‡ †

Department of Chemistry, Centre for Catalysis Research and Innovation, University of Ottawa, Ottawa, Ontario K1N 6N5, Canada Department of Chemical and Biological Engineering, University of Ottawa, Ottawa, Ontario K1N 6N5, Canada

‡

S Supporting Information *

ABSTRACT: Adsorption of CO2 was investigated on a series of primary, secondary, and tertiary monoamine-grafted pore-expanded mesoporous MCM-41 silicas, referred to as pMONO, sMONO, and tMONO, respectively. The pMONO adsorbent showed the highest CO2 adsorption capacity, followed by sMONO, whereas tMONO exhibited hardly any CO2 uptake. As for the stability in the presence of dry CO2, we showed in a previous contribution [J. Am. Chem. Soc. 2010, 132, 6312−6314] that amine-supported materials deactivate in the presence of dry CO2 via the formation of urea linkages. Here, we showed that only primary amines suffered extensive loss in CO2 uptake, whereas secondary and tertiary amines were stable even at temperature as high as 200 °C. The difference in the stability of primary vs secondary and tertiary amines was associated with the occurrence of isocyanate as intermediate species toward the formation of urea groups, since only primary amines can be precursors to isocyanate in the presence of CO2. However, using a grafted propyldiethylenetriamine containing both primary and secondary amines, we demonstrated that while primary amines gave rise to isocyanate, the latter can react with either primary or secondary amines to generate di- and trisubstituted ureas, leading to deactivation of secondary amines as well.

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INTRODUCTION Amine-supported materials gained tremendous popularity in recent years as adsorbents for acid gases removal, particularly CO2.1,2 Properly designed, amine-functionalized materials exhibit high adsorption capacity, fast CO2 adsorption and desorption, and low-energy requirements for recycling compared to amine solutions.3,4 Because of the high reactivity of primary and secondary amines toward CO2, most aminesupported materials reported in the literature as CO 2 adsorbents consisted of primary amines or combinations of primary and secondary amines.5−10 In some instances where the amine-containing material is obtained via in-situ11 or exsitu12−14 polymerization, or via iterative reactions,15 tertiary amines were also included in the CO2 adsorbing species. Of particular interest, grafting propyldiethylenetriamine, which contains one primary and two secondary amine groups, on large-pore mesoporous silica such as pore-expanded MCM-41 (PE-MCM-41) gave rise to an adsorbent (TRI-PE-MCM-41) with high CO2 capacity at low partial pressure,3,9,16 albeit with lower efficiency (CO2/N = 0.33) compared to grafted primary monoamine, which exhibits a CO2/N ratio of 0.5,17,18 corresponding to quantitative formation of carbamate. In earlier reports,10,16,19 we also showed that TRI-PE-MCM-41 exhibits extremely high CO2 selectivity over N2, O2, CH4, and H2. With regard to the stability of amine-containing adsorbents in the presence of CO2, Drage et al.20 found that beyond 135 °C CO2 reacted with amine groups, giving rise to urea linkages. © 2012 American Chemical Society

We also demonstrated that the CO2 capacity of grafted propyldiethylenetriamine and primary monoamine as well as impregnated polyethylenimine is considerably reduced after a few tens of CO2 adsorption−desorption cycles in dry conditions.21 Although it is possible to regenerate CO2deactivated materials via hydrolysis of urea at high temperature, e.g. 200 °C, for any practical purpose such deactivation may be regarded as permanent.21 The formation and accumulation of urea according to Scheme 1 was found to be at the origin of Scheme 1. Reaction Pathway for the Formation and Accumulation of Urea in Dry Conditions

such deactivation. It was also shown that the formation of urea takes place only under dry conditions and is inhibited when moisture is added to the gas feed. On the basis of Scheme 1, it may be inferred at first glance that primary and secondary amines would be equally prone to deactivation in dry CO2 via Received: November 26, 2011 Revised: January 31, 2012 Published: February 9, 2012 4241

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groups.9 The nitrogen content was calculated based on the weight loss beyond 200 °C. The nitrogen content was also determined by elemental analysis using a TF Flash 1112 analyzer (Thermo Finnigan). Weighed samples placed into capsules were flash-combusted at about 1800 °C, and the gases were carried by helium through a reduction/ oxidation column to yield N2, CO2, and H2O. The gases were separated and quantified as they pass through a gas chromatograph column. The amine loading was calculated from the nitrogen content. The 13C CP/MAS NMR experiments were conducted on a Bruker AVANCE 500. The spinning frequency was set to 10 kHz. The contact time was 2 ms, with recycle delays of 2 s. Equilibrium and Cyclic Nonequilibrium Adsorption Measurements. Adsorption equilibrium measurements for single component CO2 as well as near-equilibrium adsorption−desorption cyclic measurements were performed using a Rubotherm gravimetric apparatus (Bochum, Germany). Details about the experimental setup as well as the procedure for measurements of the equilibrium isotherms may be found elsewhere.3 Prior to measurements, the material was activated under vacuum (5 × 10−4 mbar) at 120 °C for 2 h. For cyclic adsorption−desorption measurements, the sample was first exposed for 30 min at 120 °C to UHP nitrogen flowing at 50 mL/ min under atmospheric pressure. Subsequently, the sample was cooled to 55 °C at isobaric condition (1 bar) and the feed gas was switched to pure CO2 or CO2:N2 = 10:90 at 50 mL/min. The working adsorption capacity (nonequilibrium) corresponds to the weight gain of the sample after 30 min exposure. The material was then regenerated at 120 °C under flowing nitrogen (50 mL/min) for 30 min. This adsorption−desorption procedure was repeated over 60 cycles for pMONO and sMONO. Because dry CO2 was hardly adsorbed on tMONO (vide infra), only limited measurements were carried out in the presence of this material. In-Situ Diffuse Reflectance Infrared Fourier Transform (DRIFT) Spectroscopy Measurements. A Nicolet Magna-IR 550 spectrometer equipped with a MCT detector and a Thermo diffuse reflectance cell was used to collect DRIFT spectra. About 15 mg of powder sample was placed into the cell and pretreated in flowing ultrahigh purity He at 120 °C for 2 h. The DRIFT spectra were then recorded under He atmosphere for fresh materials. Then, 60 cycles of CO2 adsorption at 55 °C and desorption at 120 °C were carried out before recording DRIFT spectra. The spectrum for KBr was used as background.

urea formation at the expense of amine groups. However, this needs to be demonstrated. Thus, because of its relevance to the understanding of CO2−amine interactions and its fundamental and practical implications, we found it compelling to carry out a comprehensive comparative study of the CO2 adsorption properties of primary, secondary, and tertiary monoaminegrafted adsorbents using pore-expanded MCM-41 (PE-MCM41) silica as support. Moreover, a propyldiethylenetriaminegrafted PE-MCM-41, which contains both primary and secondary amines, was also used to investigate the possible occurrence of a cooperative deactivation mechanism involving both types of amines.

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EXPERIMENTAL SECTION

Materials. The mesoporous PE-MCM-41 silica support was prepared using Cab-O-Sil fumed silica from Cabot as the silica source, cetyltrimethylammonium bromide (CTAB, Aldrich) as template, tetramethylammonium hydroxide (TMAOH 25% in water, Aldrich) for pH adjustment, and dimethyldecylamine (DMDA 97% purity, Aldrich) as postsynthesis pore-expanding agent. Surface grafting of amine-containing species was carried out using four different aminosilanes, namely, 3-aminopropyltrimethoxysilane (pMONOsilane, 99%, Aldrich), N-methylaminoproyltrimethoxysilane (sMONO-silane, 99%, Aldrich), N,N-dimethylaminopropyltrimethoxysilane (tMONO-silane, >95% purity, Gelest), and N 1 -(3trimethoxysilylpropyl)diethylenetriamine (TRI-silane, technical grade, Aldrich). Carbon dioxide (99.99%), nitrogen (99.999%), helium (99.999%), and certified 0.1, 1, 5, and 10% carbon dioxide in nitrogen were supplied by Linde Ltd. Canada. All reagents and gases were used as-received. The pore-expanded MCM-41 silica support was prepared in two steps as described elsewhere.3 Briefly, MCM-41 silica was synthesized at 100 °C using CTAB in the presence of TMAOH. This was followed by pore expansion via hydrothermal treatment of as-synthesized MCM-41 in the presence of an aqueous emulsion of DMDA at 120 °C for 3 days. After removing the surfactant template and pore expander by air calcination, the obtained product was labeled PE-MCM-41. Incorporation of different amine-containing species was achieved via surface grafting as described earlier.17 After drying in a vacuum oven at 120 °C, a sample of PE-MCM-41 was loaded into a multineck glass flask containing 150 mL of toluene. Once a homogeneous suspension was produced, 0.4 mL of distilled deionized water per gram of PEMCM-41 was added and left stirring for 30 min. The glass flask with a condenser was then submerged in a silicon oil bath set at 85 °C using a temperature-controlled stirring hot plate with an external temperature probe. The aminosilane (2 mL per gram of silica) was subsequently added to the mixture and left stirring for 16 h. The material was then filtered and washed with copious amounts of toluene and then pentane. Finally, the recovered solid was dried at 100 °C in a natural convection oven for 1 h. The monoamine-containing adsorbents will be labeled xMONO, where x represents the type of amine grafted with (p) for primary, (s) for secondary, and (t) for tertiary. The propyldiethylenetriamine-derived material will be referred to as TRI. All samples were characterized by nitrogen adsorption at 77 K, using a Micromeritics 2020 instrument. The BET specific surface areas were determined from the adsorption data in the relative pressure (P/ P0) range from 0.05 to 0.2. The pore size distributions (PSDs) were calculated from the nitrogen adsorption branch using the Barrett− Joyner−Halenda (BJH) method with Kruk−Jaroniec−Sayari (KJS) correction,22 and the maximum of the PSD was considered as the average pore size. The pore volume was considered as the volume of liquid nitrogen adsorbed at P/P0 = ca. 1. The organic content was determined by thermogravimetric analysis (TGA) using a TA Q-500 instrument. The sample was heated at 10 °C min −1 under flowing nitrogen up to 800 °C followed by decomposition in the presence of air at the same heating rate, up to 1000 °C. The weight loss below 200 °C was attributed to adsorbed water, solvent, and any alcohol associated with nonhydrolyzed alkoxy

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RESULTS AND DISCUSSION Material Characterization. The N2 adsorption−desorption isotherms for xMONO at 77 K are shown in Figure 1. The

Figure 1. N2 adsorption−desorption isotherms at 77 K.

isotherms for the PE-MCM-41 support and TRI may be found elsewhere.3 All samples exhibited the type IV adsorption isotherm, typical of mesoporous materials. Table 1 shows the 4242

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Table 1. Structural Properties of the Materials material

SBET (m2/g)

Vp (cm3/g)

Dp (nm)

amine content (mmol/g)

amine content (molecules/nm2)

CO2 capacitya (mmol/g)

CO2/N (molar ratio)

PE-MCM-41 pMONO sMONO tMONO TRI

1153 233 471 519 367

2.81 0.46 0.72 1.05 0.87

11 8.2 8.3 7.5 9.4

4.6b (4.88c) 3.67b (3.71c) 3.25b 7.9b

11.8 4.7 3.7

2.3 1.24 0.2 2.7

0.5b (0.47c) 0.34b (0.34c) 0.061b 0.33b (0.38c)

a

At 5% CO2 in N2 at 25 °C and atmospheric pressure. bThermogravimetric analysis. cElemental analysis.

textural properties (surface area, pore volume, and pore diameter), amine content, amine surface density, and the molar amount of CO2 adsorbed per amine group (CO2/N). The equilibrium molar CO2/N ratio was determined using 5% CO2 in N2 at 25 °C and 1 atm, which are the conditions corresponding to saturation of the amine sites via chemical adsorption, while physisorption is negligible.17 As seen, for the xMONO series, the surface area and pore volume decreased as the amine content increased, whereas the pore size remained almost constant. Moreover, the decreasing pore volume and surface area upon amine loading is consistent with amine grafting on the internal surface of PE-MCM-41 silica. It is widely accepted that the mechanism of CO2 adsorption on primary and secondary amines in dry conditions involves the formation of zwitterion (eq 1), followed by the deprotonation of the zwitterion by a base to produce carbamate, possibly via formation of unstable carbamic acid,23,24 according to eq 2. R1R2NH + CO2 ⇌ R1R2NH+COO−

(1)

R1R2NH+COO− + b ⇌ R1R2NCOO− + bH+

(2)

Figure 2. CO2/N ratios vs amine loading for propylamine-grafted PEMCM-41 silicas (adsorption conditions: 5% CO2 in N2 at 25 °C and atmospheric pressure).

Under dry conditions, b is typically an amine species. Tertiary amines do not form carbamate and are generally not reactive with CO2 in nonaqueous conditions. As shown in Table 1, primary and tertiary amines exhibited the highest (0.5) and the lowest (0.06) CO2/N ratio, respectively. Adsorption of CO2 on primary amine corresponded to quantitative formation of carbamate with a stoichiometric ratio of 0.5, in line with the well-known high reactivity of such amines with CO2.25 Interestingly, the CO2/N efficiency of secondary amine was only 0.34. This may reflect the difference in the affinity of primary and secondary amines toward CO2 but may also be associated with the lower surface density of amines in sMONO compared to pMONO as the formation of carbamate requires two neighboring amines groups.26,27 These findings are in general agreement with recent literature reports.27,28 Notice that CO2 absorption in aqueous amine solutions showed comparable reactivity for primary and secondary amines resulting in similar absorption performances.29 To discriminate between the effect of amine surface density and the amine−CO2 affinity on CO2/N ratio, a series of propylamine-grafted PE-MCM-41 materials with different loadings were prepared by varying the synthesis conditions (aminosilane/silica ratio, amount of water added), and their adsorption capacity was measured in the presence of 5% CO2/ N2 at 25 °C. Consistent with our inference, Figure 2 shows that the CO2/N ratio increases with the amine surface density up to the stoichiometric ratio of 0.5. As seen, the data point representing sMONO is located significantly below the curve, indicating a weaker interaction with CO2 than pMONO. CO2 Adsorption Isotherms. Figure 3 and Figure S1 show the CO2 adsorption capacity of pMONO and sMONO up to 1

Figure 3. CO2 adsorption isotherms for pMONO at 298, 313, 328, and 343 K.

bar at four temperatures in the range of 25−70 °C. All isotherms were characterized by a strong increase in CO2 uptake versus pressure up to ca. 0.05 bar, followed by a much slower increase. Notice that the initial slopes were more pronounced for the primary amine compared to the secondary amine, indicating that the interactions of CO2 with pMONO are stronger compared to sMONO. This is consistent with the high heat of CO2 adsorption in aqueous solutions of primary amines.30 Based on the CO2 adsorption isotherms at different temperatures for pMONO and sMONO, the isosteric heat of adsorption (Qisos) was calculated (Figure S2 in Supporting Information) using the Clausius−Clapeyron equation.31 As 4243

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shown in Figure 3 and Figure S1, the stronger amine−CO2 interaction for pMONO was consistent with the higher isosteric heat of CO2 adsorption on pMONO in comparison to sMONO (89 vs 60 kJ/mol at low loading) over the whole range of CO2 loading (Figure S2). Interestingly, Qisos decreased at higher CO2 uptake for both pMONO and sMONO, reaching values typical of physical adsorption (i.e., 20 kJ/mol) at high loading, indicating that CO2 probes both materials as energetically heterogeneous surfaces.32 In particular, the plateau observed in Figure S2 at CO2 uptake around 2 mmol/g for pMONO at room temperature, which may be indicative of the onset of physical adsorption, while the highly exothermic interactions of CO2 with primary amine sites still occurring at this stage. The degree of heterogeneity of adsorption sites seems to be less significant for sMONO than pMONO, leading to a smoother transition between chemical and physical regimes. At room temperature, the onset of physical adsorption occurred at ca. 1.5 mmol/g without a plateau as the chemical adsorption was completely over. In an earlier report,3 we also demonstrated that the isosteric heat of CO2 adsorption over TRI decreased from ca. 90 kJ/mol to ca. 20 kJ/mol as the CO2 uptake increased. Long-Term Behavior of the Adsorbents in the Presence of CO2. The lifetime of adsorbents is a critical property of equal importance as CO2 adsorption capacity, selectivity, and kinetics with direct impact on the economics of the process. Depending on the CO2 separation application and the regeneration mode applied (e.g., pressure, temperature, or concentration gradient), amine supported materials may be adversely affected by different types of deactivation pathways, including degradation in the presence of dry CO220,21 and oxidative degradation33−35 when O2 is present in the process. As shown earlier,21 degradation of amine-supported materials in the presence of dry CO2 may occur even under mild adsorption and desorption conditions. In this work, we investigated the effect of dry CO2 at temperatures typical of flue gas treatment, i.e., 55 °C for adsorption and 120 °C for regeneration. Because most adsorption processes use a purge step in the presence of an inert gas or the purified gaseous product itself, a purge flow of N2 was used in the cyclic experiments. This temperature swing procedure will be referred to as TS(55-1-120), where the three numbers indicate the adsorption temperature (°C) and the desorption pressure (bar) and temperature (°C), respectively. Adsorption was carried out in the presence of pure CO2 to accelerate the deactivation process, if any. Materials that underwent TS(55-1-120) cycling under dry conditions will be referred to as xMONO-d, with x = p or s. Similar experiments for TRI were reported earlier.21 Figure 4 shows the CO2 working capacity over 60 adsorption−desorption cycles carried out in the presence of pMONO and sMONO under dry conditions. As seen, pMONO deactivated gradually and the percentage loss of adsorption capacity was 21% over 60 cycles. Under similar conditions, the degree of deactivation of TRI over 40 cycles was 14%.21 In contrast, the adsorption capacity for sMONO was stable throughout all cycles. At first glance, in light of Scheme 1, the striking difference in the stability of pMONO and sMONO is rather surprising. 13 C CP MAS NMR data for pMONO-d and sMONO-d samples after the cycling experiments and the corresponding fresh materials are shown in Figure S3 of the Supporting Information. As indicated, pMONO exhibited three peaks at 11, 23, and 44 ppm associated with carbon atoms of the propyl

Figure 4. Adsorption−desorption cycling of pure CO2 over pMONO and sMONO in dry conditions using TS(55-1-120) regeneration mode with N2 as purge gas.

chain, while sMONO exhibited three signals at 11, 23, and 54 ppm also associated with the propyl moiety, in addition to a signal at 34 ppm attributable to −NH−CH3. The small peak at 44 ppm is associated with propylamine impurity, whereas the very small signal at 63 ppm corresponds to unreacted methoxy groups. Notice that the ca. 23 ppm peak attributable to Si− CH2−CH2− is often broader than the other peaks or even split in two peaks at ca. 21 and 27 ppm. The carbon two bonds removed from the amine group is known to be very sensitive to the electronic properties of the nitrogen in aliphatic amines.36 Similar broadening or peak splitting of the C2 peak in adsorbed or grafted propylamine was shown to occur in partially hydrated materials.36,37 Interestingly, the fresh samples exhibited a signal at 164.6 ppm attributable to carbamate formed via adsorption of CO2 from ambient air.16 The deactivated pMONO-d (Figure 4) exhibited an additional peak at 160.5 ppm, assigned to the carbon atom of newly formed urea group (Scheme 1), which is responsible for the observed loss of adsorption capacity.21 Interestingly, the urea peak did not occur in the case of sMONO-d, although it underwent the same CO2 cycling as pMONO-d, but without any loss of adsorption capacity (Figure 4). It is thus inferred that sMONO is less prone to the CO2-mediated amine coupling into urea than pMONO. It may be argued that because of the lower amine surface density in sMONO compared to pMONO (4.71 vs 11.84 molecules/nm2, based on actual surface area of the samples), the formation of urea does not take place. However, based on the CO2/N chemisorption ratio of 0.34 obtained in the presence of sMONO, at least 68% of the amine groups are close enough to each other to form carbamate; thus they would be close enough to form urea. The contention that deactivation of pMONO is associated with the formation of stable urea groups was further substantiated by in-situ DRIFTS. As seen in Figure 5, the spectral region associated with NH2 scissoring for fresh pMONO is dominated by a band at 1600 cm−1. Cycling CO2 over this material at 55 °C for adsorption and 120 °C for desorption under dry conditions led to the development of two bands at 1560 and 1658 cm−1 assigned to urea.21 In contrast, the FTIR spectrum for sMONO remained virtually unchanged. The behavior of xMONO and TRI in pure dry CO2 was further investigated using thermogravimetric analysis up to 200 °C, which is slightly below the onset of thermal degradation of grafted amines under nitrogen.9 The materials were first 4244

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physisorption, the CO2/N ratio dropped to 0.25 versus 0.5 for the fresh material. This may be attributable to a decrease in the surface density of remaining free amine groups as the formation of carbamate requires the presence of two neighboring amine groups. However, assuming that the surface area did not change, the actual amine surface density of CO2-deactivated pMONO would be 4.5 molecules/nm2. According to Figure 2, the deactivated material exhibits half the adsorption efficiency of a fresh material with the same average amine surface density. It is inferred that the fresh material has a higher local amine density. This finding is reminiscent of recent work by Young and Notestein,38 who demonstrated the importance of local vs average amine density. Because the free amine catalyzes the silane condensation in solution or as rafts, aminosilane grafting affords high local surface densities of amine groups.38 As for sMONO, the material underwent, under flowing dry CO2, a gradual weight loss up to 200 °C (Figure 7). After cooling to 25 °C under N2 purge, the material adsorbed 7.1 wt

Figure 5. DRIFT spectra for pMONO and sMONO before and after CO2 cycling under dry conditions.

pretreated in dry nitrogen at 150 °C for 2 h, cooled to 25 °C (stage i), and then exposed to pure dry CO2 for 30 min (Figure 6). The temperature was then increased stepwise under flowing CO2 up to 200 °C. It was maintained constant for 2 h at 105,

Figure 7. Weight change vs time for sMONO in pure CO2 under different conditions (see Figure 6). Figure 6. Weight change vs time for pMONO in pure CO2 under different conditions: (i) pretreatment in N2 at 150 °C; (ii) exposure to pure CO2 for 30 min at 25 °C, then for 2 h at different temperatures, up to 200 °C; (iii) cooling under nitrogen to 25 °C and exposure to pure CO2 for 30 min.

% CO2, which is identical to the adsorption capacity of the fresh material. This indicates that, in contrast to pMONO, no change in the organic layer took place during exposure to pure CO2 at temperatures as high as 200 °C. The striking difference in the stability of pMONO and sMONO in the presence of dry CO2 is also consistent with the FTIR and NMR data reported in Figure S3 and Figure 5. Likewise, tMONO (data not shown) exhibited a similar behavior as sMONO, with a small CO2 uptake (0.31%) most likely due to physisorption. These findings indicate that as far as monoamine-containing adsorbents are concerned, only pMONO undergoes deactivation in the presence of dry CO2 due to urea formation. Thus, in contrast to polyethylenimine20,21 and other primary aminecontaining adsorbents,21 secondary monoamine-containing materials may be employed for the separation of CO2 in dry gases using temperature swing adsorption. In particular, it would be very attractive to use pure CO2 as purge gas for the production of a pure CO2 stream, ready for compression and sequestration or other usages, without any adverse effect on the adsorbent such as the formation of urea. This finding is particularly significant, since the other alternative being contemplated, which is the use of low quality steam is facing severe material stability limitations.34,39 With regards to adsorbents with mixed amines, the question arises as to whether only primary amines deactivate or there are

125, 150, 175, and 200 °C (stage ii). Subsequently, the materials were cooled to room temperature under nitrogen (stage iii) and exposed again to pure CO2 for 30 min at 25 °C. This procedure simulates accelerated degradation of grafted amine adsorbents in the presence of dry CO2. Figure 6 shows that as the temperature increased up to 125 °C in stage ii, pMONO lost weight, presumably because of CO2 desorption. However, beyond 125 °C, the sample weight increased gradually with temperature and with time. At the end of stage ii, the overall weight gain based on the material obtained after the pretreatment (stage i) was 4 wt %. This weight gain corresponds to extensive formation of urea.20,21 Upon regeneration and cooling under N 2 , followed by CO 2 adsorption at 25 °C, the weight gain (adsorption capacity) was only 1.9 wt % vs 9.9 wt % for the freshly pretreated sample. The 4% weight gain at 200 °C stems from the formation of 1.43 mmol of urea/g, corresponding to the consumption of 2.86 mmol of amine. The CO2 uptake on deactivated pMONO was 1.9 wt %, i.e., 0.43 mmol/g, whereas the amount of available amine was 1.76 mmol/g. Thus, ignoring the limited 4245

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secondary and tertiary amines. However, similarly to CO2induced deactivation, in the presence of mixed amines, cooperative mechanisms seem to occur, although they are yet to be delineated on the molecular scale.33−35 In addition to being consistent with all experimental findings, the proposed interpretation is supported by literature data. As far as primary amines are concerned, they were reported to interact with CO2 to generate carbamic acids in situ, which dehydrate in the presence of Mitsunobu reagents (PBu3 + dibenzyl or dialkyl azodicarboxylate) into isocyanate, which can react further with primary or secondary amines to afford different ureas in high yields.41−43 Alternatively, primary amine reacts with CO2 to generate carbamate, which transforms into urea in the presence of “dehydrating reagents” such as POCl3 or P4O10.44 In contrast, not only secondary amines cannot generate isocyanate, but “dehydration” of the corresponding carbamate is very difficult.45 The aforementioned literature data to support our findings are associated with organic reactions that take place in liquid phase, typically in the presence of catalysts and/or additives. Nonetheless, it is possible that such processes occur on solid surfaces, without catalysts/additives, albeit at a much slower rate. Recent work by Wu et al.46 lends strong support for the above interpretation. They found that under high temperature and CO2 pressure (180 °C, 10 MPa), but without any catalyst, nor solvent, primary amines yielded the corresponding disubstituted ureas, whereas mixtures of primary and secondary amines gave rise to primary aminederived disubstituted ureas and trisubstituted ureas originating from both amines, but no secondary amine-derived tetrasubstituted ureas. The authors used the intermediacy of isocyanate to interpret their data, although they did not detect it. Our results are fully consistent with such observations. Notice that in the case of adsorbents with multiple amine-containing species, while the occurrence of isocyanate offers a potential route to different ureas, there may be additional mechanisms leading to the formation of cyclic ureas such as imidazolidinones.47

deactivation pathways that involve different types of amines. To this end, we run the same experiment using TRI adsorbent, which contains primary and secondary amines in a 1 to 2 ratio. Figure S4 in the Supporting Information shows that the weight gain at 200 °C was 5.5 wt %. However, since TRI contains only 2.63 mmol of primary amine, the maximum weight gain stemming from complete and selective transformation of primary amine into urea should not exceed 3.4 wt %, whereas a 5.5 wt % increase corresponds to the consumption of 4.22 mmol amine. This indicates that contrary to sMONO, in TRI and most likely in other mixed-amine adsorbents, secondary amines participate into the deactivation process (formation of urea). Adsorption of CO2 on the remaining amine groups of TRI corresponds to a CO2/N ratio of only 0.14 versus 0.34 for the fresh material. Considering the overall reaction for urea formation proposed earlier (Scheme 1),21 it would be difficult to explain the significantly higher stability of secondary amine compared to primary amine. A possible explanation for the above findings is that in the presence of primary amine urea may occur not only via dehydration of carbamate but also through isocyanate, which forms by dehydration of the carbamic acid intermediate (Scheme 2). Recent work using solid state NMR and 13CO2 adsorption on amine-grafted silica demonstrated that carbamic acid is an intermediate toward the formation of carbamate.40 Scheme 2. Reaction Pathways for the Formation of Urea from Primary Amine

Notice that only carbamic acids stemming from primary amines are able to dehydrate into isocyanate. However, in the presence of secondary monoamine, not only the corresponding carbamate does not dehydrate into urea (vide infra), but the corresponding carbamic acid does not generate isocyanate (eq 3). Therefore, consistent with our findings, no urea is formed in the presence of adsorbents containing only a secondary monoamine.

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CONCLUSION Comparative investigation of CO2 adsorption on propyl-, Nmethylpropyl- and N,N-dimethylpropylamine grafted on poreexpanded MCM-41 silica was carried out over a wide range of temperature and CO2 concentration. The primary monoaminecontaining PE-MCM-41 exhibited the highest CO2 adsorption capacity regardless of CO2 concentration. The secondary monoamine exhibited lower CO2 capacity, while the tertiary monoamine adsorbed hardly any CO2. In terms of stability in the presence of dry CO2, the primary monoamine was found to be the least stable due to the formation of urea, most likely through isocyanate intermediate and possibly carbamate dehydration. The secondary monoamine showed exceedingly high stability in the presence of dry CO2 at temperature as high as 200 °C, most likely because (i) no isocyanate intermediate is formed and (ii) dehydration of carbamate does not occur. Thus, the formation of urea does not take place. However, when combined with primary amines, secondary amines may actually deactivate through reaction with primary amine-derived isocyanate to form trisubstituted ureas.

In the presence of both types of amine as in TRI or in supported polyethylenimine, once isocyanate is formed using the primary amine, it can react with either primary or secondary amine to generate di- or trisubstituted ureas (Scheme 3). Thus, in contrast to sMONO, provided that there are primary amines, secondary amines may also deactivate through the formation of Scheme 3. Urea Formation from Isocyanate in the Presence of Primary and Secondary Amines

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ASSOCIATED CONTENT

S Supporting Information *

mixed ureas. Notice that with regard to oxidative degradation primary amine was found to be much more stable than

CO2 adsorption isotherms for sMONO, isosteric heat of CO2 adsorption on pMONO and sMONO as a function of CO2 4246

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loading, 13C CP MAS NMR spectra for pMONO and sMONO before and after CO2 adsorption−desorption cycling under dry conditions, weight change vs time for TRI in pure CO2 under different conditions. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS We gratefully acknowledge the financial support of the Natural Science and Engineering Council of Canada (NSERC) and Carbon Management Canada is acknowledged. Y.B. thanks NSERC for a postdoctoral fellowship. A.S. thanks the Federal Government for the Canada Research Chair in Nanostructured Materials for Catalysis and Separation (2001−2015).

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dx.doi.org/10.1021/la204667v | Langmuir 2012, 28, 4241−4247