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• W2912476701 abstract "•Crystallization of an aqueous guanidine sorbent leads to effective CO2 capture•X-ray structural analysis reveals hydrogen-bonded bicarbonate dimers in the crystals•Mild heating of the bicarbonate crystals releases the CO2 via carbonic acid dimers•The CO2-separation cycle requires substantially less energy than benchmark sorbent Human activities in the last one and a half centuries have perturbed the natural carbon cycle, shifting massive amounts of carbon from the geosphere into the atmosphere and leading to climate change at an unprecedented pace. Carbon capture and storage, consisting of capturing CO2 from fossil fuel emissions and sequestering it deep underground, offer the prospect of limiting the increase in the atmospheric CO2 concentration and the global temperature. This requires the development and large-scale deployment of energy-efficient carbon-capture technologies. Here, we demonstrate a promising approach to CO2 capture based on crystallization of bicarbonate-water clusters with a simple guanidine compound. The CO2 separation cycle involves a unique proton-transfer mechanism via the formation of a carbonic acid dimer, leading to efficient CO2 release and quantitative regeneration of the guanidine compound and requiring significantly less energy than state-of-the-art carbon-capture technologies. Limiting global temperature rises is increasingly dependent on the development of energy-efficient carbon-capture methods. Here, we report a simple CO2-separation cycle using an aqueous bis(iminoguanidine) (BIG) sorbent that reacts with CO2 and crystallizes into an insoluble bicarbonate salt. X-ray diffraction analysis of the bicarbonate crystals revealed “anti-electrostatic” hydrogen-bonded (HCO3−)2 dimers, stabilized by guanidinium cations and water. Mild heating of the crystals releases the CO2 and regenerates the BIG sorbent quantitatively, thereby closing the CO2-separation cycle. Experimental and computational investigations support a CO2-release mechanism consisting of surface-initiated low-barrier proton transfer from guanidinium groups to bicarbonate anions with the formation of carbonic acid dimers, followed by CO2 and H2O release in the rate-limiting step, with a measured activation energy of 102 ± 12 kJ/mol. The minimum energy required for sorbent regeneration is 151.5 kJ/mol CO2, which is 24% lower than the regeneration energy of monoethanolamine, a benchmark industrial sorbent. Limiting global temperature rises is increasingly dependent on the development of energy-efficient carbon-capture methods. Here, we report a simple CO2-separation cycle using an aqueous bis(iminoguanidine) (BIG) sorbent that reacts with CO2 and crystallizes into an insoluble bicarbonate salt. X-ray diffraction analysis of the bicarbonate crystals revealed “anti-electrostatic” hydrogen-bonded (HCO3−)2 dimers, stabilized by guanidinium cations and water. Mild heating of the crystals releases the CO2 and regenerates the BIG sorbent quantitatively, thereby closing the CO2-separation cycle. Experimental and computational investigations support a CO2-release mechanism consisting of surface-initiated low-barrier proton transfer from guanidinium groups to bicarbonate anions with the formation of carbonic acid dimers, followed by CO2 and H2O release in the rate-limiting step, with a measured activation energy of 102 ± 12 kJ/mol. The minimum energy required for sorbent regeneration is 151.5 kJ/mol CO2, which is 24% lower than the regeneration energy of monoethanolamine, a benchmark industrial sorbent. (Bi)carbonate anions play a critical role in the carbon cycle, taking part in chemical equilibria between atmospheric CO2, carbonic acid, and bicarbonate salts dissolved in natural waters and carbonate minerals (mainly CaCO3).1Lower S.K. Carbonate Equilibria in Natural Waters. Simon Fraser University Press, 1996Google Scholar Human activity since the industrial revolution has perturbed the natural carbon cycle, mostly through emissions from fossil fuels, which shifted a massive amount of carbon from the geosphere into the atmosphere in a relatively short period, leading to an increase in global temperature.2Emanuel K.A. Climate Science and Climate Risk: A Primer. Massachusetts Institute of Technology Press, 2016Google Scholar Significant efforts in the last couple of decades have been directed toward reversing this trend and re-establishing the carbon balance through carbon capture and storage (CCS), which aims to reduce carbon emissions by capturing the CO2 generated by the energy sector and other industrial sectors and permanently storing it deep underground.3International Energy Agency 20 years of carbon capture and storage.www.iea.org/publications/freepublications/publication/20YearsofCarbonCaptureandStorage_WEB.pdfDate: 2016Google Scholar One possible approach to CCS is to employ the same chemical equilibria involving (bi)carbonate anions in the natural carbon cycle for CO2 separation from flue gas or other industrial sources. In such a process, an aqueous alkaline solution absorbs the CO2 and converts it into (bi)carbonate, which is then removed from solution by precipitation as an insoluble carbonate salt, akin to CaCO3 sedimentation in the oceans. Heating the solid carbonate salt releases the CO2 and regenerates the sorbent, thus closing the CO2-separation cycle. The carbonate precipitation is an important step, as it minimizes the overall energetic requirement by avoiding energy-demanding processes such as heating the aqueous sorbent and evaporating water. The simplest implementation of such a (bi)carbonate cycle is to use aqueous KOH or NaOH sorbents to scrub the carbon dioxide and then precipitate the resulting (bi)carbonate anions with Ca2+. Calcination of the resulting CaCO3 at high temperatures (∼800°C–900°C) then releases the CO2, which can be sent to geological storage. This simple cycle is at the basis of one of the oldest CO2-scrubbing technologies, and it has been recently employed in a process for capturing CO2 from the atmosphere.4Keith D.W. Holmes G. St. Angelo D. Heidel K. A process for capturing CO2 from the atmosphere.Joule. 2018; 2: 1573-1594Abstract Full Text Full Text PDF Scopus (591) Google Scholar Though the inorganic (bi)carbonate cycle offers a simple solution to the carbon-capture problem, it is very energy intensive due to the exceedingly high temperatures required for CO2 release. Organic alkaline sorbents, on the other hand, present a thermodynamically more favorable alternative to CO2 separation. Compared with their inorganic counterparts, organic (bi)carbonates may involve different mechanisms of CO2 binding and release, especially when the organic cation is capable of hydrogen bonding with the (bi)carbonate anions.5Baldwin D.A. Denner L. Egan T.J. Markwell A.J. Structure of guanidinium bicarbonate: A model for the bicarbonate anion binding site of the transferrins.Acta Crystallogr. C Cryst. Struct. Commun. 1986; 42: 1197-1199Crossref Google Scholar, 6Tiritiris I. Mezger J. Stoyanov E.V. Kantlehner W. Orthoamides and iminium salts, LXXI. Capturing of carbon dioxide with organic bases (part 2) – Reactions of guanidines and ω-aminoalkyl-guanidines with carbon dioxide.Z. Naturforsch. 2011; 66b: 407-418Google Scholar, 7Pérez E.R. Santos R.H. Gambardella M.T.P. de Macedo L.G.M. Rodrigues-Filho U.P. Launay J.C. Franco D.W. Activation of carbon dioxide by bicyclic amidines.J. Org. Chem. 2004; 69: 8005-8011Crossref PubMed Scopus (171) Google Scholar, 8Chutia R. Das G. Hydrogen and halogen bonding in a concerted act of anion recognition: F– induced atmospheric CO2 uptake by an iodophenyl functionalized simple urea receptor.Dalton Trans. 2014; 43: 15628-15637Crossref PubMed Google Scholar, 9Dalapati S. Jana S. Saha R. Alam M.A. Guchhait N. Reusable amine-based structural motifs for greenhouse gas (CO2) fixation.Org. Lett. 2012; 14: 3244-3247Crossref PubMed Scopus (13) Google Scholar, 10Mulugeta E. He Q. Sareen D. Hong S.J. Oh J.H. Lynch V.M. Sessler J.L. Kim S.K. Lee C.H. Recognition, sensing, and trapping of bicarbonate anions with a dicationic meso-bis(benzimidazolium) calix [4] pyrrole.Chem. 2017; 3: 1008-1020Abstract Full Text Full Text PDF Scopus (39) Google Scholar, 11McNally J.S. Wang X.P. Hoffmann C. Wilson A.D. Self-assembly of molecular ions via like-charge ion interactions and through-space defined organic domains.Chem. Commun. 2017; 53: 10934-10937Crossref PubMed Google Scholar Such hydrogen-bonded structures not only ensure effective CO2 binding but also provide a low-temperature path for CO2 release via proton transfer along the hydrogen bonds. In addition to its strong hydrogen-bond-accepting ability, the bicarbonate anion is also capable of acting as a proton donor through its O–H group. Furthermore, like other protonated oxyanions, such as HSO4− or H2PO4−,12He Q. Tu P.Y. Sessler J.L. Supramolecular chemistry of anionic dimers, trimers, tetramers, and clusters.Chem. 2018; 4: 46-93Abstract Full Text Full Text PDF PubMed Scopus (107) Google Scholar, 13Fatila E.M. Pink M. Twum E.B. Karty J.A. Flood A.H. Phosphate-phosphate oligomerization drives higher order co-assemblies with stacks of cyanostar macrocycles.Chem. Sci. 2018; 9: 2863-2872Crossref PubMed Google Scholar, 14He Q. Kelliher M. Bähring S. Lynch V.M. Sessler J.L. A bis-calix [4] pyrrole enzyme mimic that constrains two oxoanions in close proximity.J. Am. Chem. Soc. 2017; 139: 7140-7143Crossref PubMed Scopus (56) Google Scholar, 15Mungalpara D. Valkonen A. Rissanen K. Kubik S. Efficient stabilization of a dihydrogenphosphate tetramer and a dihydrogenpyrophosphate dimer by a cyclic pseudopeptide containing 1,4-disubstituted 1,2,3-triazole moieties.Chem. Sci. 2017; 8: 6005-6013Crossref PubMed Google Scholar, 16Fatila E.M. Twum E.B. Sengupta A. Pink M. Karty J.A. Raghavachari K. Flood A.H. Anions stabilize each other inside macrocyclic hosts.Angew. Chem. Int. Ed. 2016; 55: 14057-14062Crossref PubMed Scopus (96) Google Scholar, 17Rajbanshi A. Wan S. Custelcean R. Dihydrogen phosphate clusters: Trapping H2PO4– tetramers and hexamers in urea-functionalized molecular crystals.Cryst. Growth Des. 2013; 13: 2233-2237Crossref Scopus (37) Google Scholar the bicarbonate anion can hydrogen bond to itself and form (HCO3−)2 dimers, both in solution and the crystalline state. Though electrostatically repulsive, such inter-anion interactions,18Steiner T. Inter-anion O–H⋅⋅⋅O interactions are classical hydrogen bonds.Chem. Commun. 1999; : 2299-2300Crossref Scopus (47) Google Scholar, 19Macchi P. Iversen B.B. Sironi A. Chakoumakos B.C. Larsen F.K. Interanionic O–H⋅⋅⋅O interactions: The charge density point of view.Angew. Chem. Int. Ed. 2000; 39: 2719-2722Crossref PubMed Scopus (72) Google Scholar, 20Prohens R. Portell A. Font-Bardia M. Bauzá A. Frontera A. H-bonded anion-anion complex trapped in a squaramido-based receptor.Chem. Commun. 2018; 54: 1841-1844Crossref PubMed Google Scholar also called “anti-electrostatic” hydrogen bonds,21Weinhold F. Klein R.A. Anti-electrostatic hydrogen bonds.Angew. Chem. Int. Ed. 2014; 53: 11214-11217Crossref PubMed Scopus (151) Google Scholar, 22Weinhold F. Theoretical prediction of robust second-row oxyanion clusters in the metastable domain of antielectrostatic hydrogen bonding.Inorg. Chem. 2018; 57: 2035-2044Crossref PubMed Scopus (32) Google Scholar behave like classical hydrogen bonds and gain stabilization through localized O–H⋅⋅⋅O interactions with partial charge-transfer or covalent character. These inter-anion complexes can be further stabilized by interactions with organic ligands and macrocycles, counter-cations, or solvents, thereby allowing for their observation and isolation in the gas phase, solution, or the solid state.12He Q. Tu P.Y. Sessler J.L. Supramolecular chemistry of anionic dimers, trimers, tetramers, and clusters.Chem. 2018; 4: 46-93Abstract Full Text Full Text PDF PubMed Scopus (107) Google Scholar, 13Fatila E.M. Pink M. Twum E.B. Karty J.A. Flood A.H. Phosphate-phosphate oligomerization drives higher order co-assemblies with stacks of cyanostar macrocycles.Chem. Sci. 2018; 9: 2863-2872Crossref PubMed Google Scholar, 14He Q. Kelliher M. Bähring S. Lynch V.M. Sessler J.L. A bis-calix [4] pyrrole enzyme mimic that constrains two oxoanions in close proximity.J. Am. Chem. Soc. 2017; 139: 7140-7143Crossref PubMed Scopus (56) Google Scholar, 15Mungalpara D. Valkonen A. Rissanen K. Kubik S. Efficient stabilization of a dihydrogenphosphate tetramer and a dihydrogenpyrophosphate dimer by a cyclic pseudopeptide containing 1,4-disubstituted 1,2,3-triazole moieties.Chem. Sci. 2017; 8: 6005-6013Crossref PubMed Google Scholar, 16Fatila E.M. Twum E.B. Sengupta A. Pink M. Karty J.A. Raghavachari K. Flood A.H. Anions stabilize each other inside macrocyclic hosts.Angew. Chem. Int. Ed. 2016; 55: 14057-14062Crossref PubMed Scopus (96) Google Scholar, 17Rajbanshi A. Wan S. Custelcean R. Dihydrogen phosphate clusters: Trapping H2PO4– tetramers and hexamers in urea-functionalized molecular crystals.Cryst. Growth Des. 2013; 13: 2233-2237Crossref Scopus (37) Google Scholar While providing stabilization against dissociation into separate anion monomers, a hydrogen-bond donor may be able to activate the anions toward neutralization by proton transfer and formation of the conjugated acid dimers. In this respect, the (HCO3−)2 dimer is unique among other oxyanion analogs, as such a proton-transfer reaction from a hydrogen-bonding ligand would lead to the formation of carbonic acid dimers, which are predisposed to decompose into carbon dioxide and water (Scheme 1).23de Marothy S.A. Autocatalytic decomposition of carbonic acid.Int. J. Quantum Chem. 2013; 113: 2306-2311Crossref Scopus (19) Google Scholar, 24Ghoshal S. Hazra M.K. New mechanism for autocatalytic decomposition of H2CO3 in the vapor phase.J. Phys. Chem. A. 2014; 118: 2385-2392Crossref PubMed Scopus (20) Google Scholar Here, we explore this concept in a crystalline guanidinium bicarbonate salt, formed by the reaction of an aqueous solution of glyoxal-bis(iminoguanidine) (GBIG) with CO2. The resulting crystals contain hydrated (HCO3−)2 dimers, hydrogen bonded by guanidinium groups and water molecules, and undergo facile CO2 and H2O loss under relatively mild heating, regenerating the GBIG ligand quantitatively and closing the CO2 separation cycle (Scheme 1). GBIG is readily synthesized by the imine condensation of glyoxal with aminoguanidinium chloride, followed by neutralization with NaOH (see Supplemental Experimental Procedures). GBIG and its various salts were first reported by Thiele and Dralle in 1898.25Thiele J. Dralle E. Condensationproducte des aminoguanidins mit aldehyden und ketonen der fettreihe.Justus Liebigs Ann. Chem. 1898; 302: 275-299Crossref Scopus (39) Google Scholar More recently, GBIG and other bis(iminoguanidinium) analogs have been found to form extremely insoluble crystalline salts with oxyanions (e.g., sulfate and nitrate), which allows for the selective separation of this class of anions from aqueous solutions by crystallization.26Custelcean R. Williams N.J. Seipp C.A. Aqueous sulfate separation by crystallization of sulfate-water clusters.Angew. Chem. Int. Ed. 2015; 54: 10525-10529Crossref PubMed Scopus (40) Google Scholar, 27Custelcean R. Williams N.J. Seipp C.A. Ivanov A.S. Bryantsev V.S. Aqueous sulfate separation by sequestration of [(SO4)2(H2O)4]4– cluster within highly insoluble imine-linked bis-guanidinium crystals.Chem. Eur. J. 2016; 22: 1997-2003Crossref PubMed Scopus (30) Google Scholar In the particular case of 2,6-pyridine-bis(iminoguanidine) (PyBIG), an aqueous solution of the ligand was found to react with atmospheric CO2 and lead to crystallization of the corresponding carbonate salt.28Seipp C.A. Williams N.J. Kidder M.K. Custelcean R. CO2 capture from ambient air by crystallization with a guanidine sorbent.Angew. Chem. Int. Ed. 2017; 56: 1042-1045Crossref PubMed Scopus (71) Google Scholar The very low aqueous solubility of PyBIG carbonate (Ksp = 1.0 × 10−9) drives the equilibrium toward the carbonate crystallization, despite the very low concentration of CO2 in the air.29Brethomé F.M. Williams N.J. Seipp C.A. Kidder M.K. Custelcean R. Direct air capture of CO2 via aqueous-phase absorption and crystalline-phase release using concentrated solar power.Nat. Energy. 2018; 3: 553-559Crossref Scopus (90) Google Scholar This prompted the question whether the simpler GBIG congener would exhibit similar chemistry in the presence of carbon dioxide. Surprisingly, to our knowledge, more than 120 years after GBIG was first synthesized, there is no report of a (bi)carbonate salt of this simple guanidine ligand in the literature. An aqueous solution of GBIG left overnight under a CO2 atmosphere led to the formation of blade-shaped crystals with an elemental composition consistent with the dihydrated bicarbonate salt of GBIG [GBIGH2(HCO3)2(H2O)2]. The same product was obtained by reacting solid GBIG with a saturated solution of NaHCO3 (see Supplemental Experimental Procedures). Single-crystal X-ray diffraction analysis of the GBIG bicarbonate salt (Figure 1) revealed the existence of centrosymmetric (HCO3−)2 dimers in the crystals (Figure 1B), with H⋅⋅⋅O contact distances of 1.71 Å.30A Cambridge Structural Database (CSD version 5.39, November 2017) search found 78 examples of (HCO3–)2 with a mean OH⋅⋅⋅O contact distance of 1.71 ± 0.11 Å.Google Scholar Each bicarbonate anion additionally accepts three guanidinium hydrogen bonds, with N–H⋅⋅⋅O contact distances of 1.93, 1.93, and 2.17 Å, and two water hydrogen bonds, with O–H⋅⋅⋅O contact distances of 1.92 and 2.24 Å (Figure 1B). The bicarbonate dimers are linked by the water molecules into one-dimensional ladder-shaped clusters along the crystallographic [1 0 0] direction (Figure 1C). On the other hand, the essentially planar GBIGH22+ cations stack along the same direction, with a mean interplanar distance of 3.25 Å (Figure 1D). Finally, the cationic stacks flank the anionic clusters in a close-packed arrangement (Figure 1E). A critical part in the development of any carbon-capture cycle is the accurate measurement of the thermodynamics of CO2 absorption and release, which allows for determination of the minimum energy requirements for the overall process and comparison with state-of-the-art benchmarks. The reactions involved in the CO2 absorption and release with GBIG are shown in Scheme 2, and the corresponding thermodynamic parameters are listed in Table 1.Table 1Equilibrium Constants (K) and Enthalpies (ΔH) for the Reactions Involved in the CO2 Capture by GBIG and the CO2 Release from Its Crystalline Bicarbonate SaltEquationReactionKaDetermined at 25°C unless otherwise specified.ΔH (kJ/mol)aDetermined at 25°C unless otherwise specified.Reference1GBIG dissolution1.15 × 10−2bK1 = Ksp(GBIG).38.3hDetermined by van 't Hoff analysis of solubility values measured in the 15°C–35°C range.this study (Figure S1)2GBIG protonation10−12cK2 = (Kw)2/(Ka1 × Ka2), where Ka1 = 10−7.33 and Ka2 = 10−8.65.40.3iDetermined by van 't Hoff analysis of pKa values measured in the 15°C–35°C range. The enthalpy for reaction 2 is composed of the sum of the two protonation enthalpies of GBIG (−27.5; −43.8 kJ/mol) plus twice the enthalpy of deprotonation of water (55.8 kJ/mol).this study (Figure S2)3CO2 dissolution3.4 × 10−2dK3 = KH (Henry’s law constant for CO2, M atm−1).−19.4Carroll et al.34Carroll J.J. Slupsky J.D. Mather A.E. The solubility of carbon dioxide in water at low pressure.J. Phys. Chem. Ref. Data. 1991; 20: 1201-1209Crossref Scopus (482) Google Scholar4HCO3− formation3 × 107−50Wang et al.35Wang X. Conway W. Burns R. McCann N. Maeder M. Comprehensive study of the hydration and dehydration reactions of carbon dioxide in aqueous solutions.J. Phys. Chem. A. 2010; 114: 1734-1740Crossref PubMed Scopus (178) Google Scholar5GBIGH2(HCO3)2(H2O)2 crystallization5.4 × 108eK5 = 1/Ksp(GBIG-bicarb).−24.1jDetermined by van 't Hoff analysis of solubility values measured in the 15°C–35°C range.this study (Figure S1)6overall CO2 absorption6.5 × 106fK6 = K1 × K2 × (K3)2 × (K4)2 × K5 (atm−2).−84.3kΔH6 = ΔH1 + ΔH2 + 2 × ΔH3 + 2 × ΔH4 + ΔH5.this study7CO2 release–gCO2 release was not measured under equilibrium conditions.243lDetermined from the endotherm observed in the DSC at 122°C ± 6°C.this studya Determined at 25°C unless otherwise specified.b K1 = Ksp(GBIG).c K2 = (Kw)2/(Ka1 × Ka2), where Ka1 = 10−7.33 and Ka2 = 10−8.65.d K3 = KH (Henry’s law constant for CO2, M atm−1).e K5 = 1/Ksp(GBIG-bicarb).f K6 = K1 × K2 × (K3)2 × (K4)2 × K5 (atm−2).g CO2 release was not measured under equilibrium conditions.h Determined by van 't Hoff analysis of solubility values measured in the 15°C–35°C range.i Determined by van 't Hoff analysis of pKa values measured in the 15°C–35°C range. The enthalpy for reaction 2 is composed of the sum of the two protonation enthalpies of GBIG (−27.5; −43.8 kJ/mol) plus twice the enthalpy of deprotonation of water (55.8 kJ/mol).j Determined by van 't Hoff analysis of solubility values measured in the 15°C–35°C range.k ΔH6 = ΔH1 + ΔH2 + 2 × ΔH3 + 2 × ΔH4 + ΔH5.l Determined from the endotherm observed in the DSC at 122°C ± 6°C. Open table in a new tab The CO2 absorption consists of five elementary steps: dissolution of GBIG in water (equation 1), proton transfers from water to GBIG to generate the GBIGH22+ cations and HO− anions (equation 2), CO2 transport from air into the aqueous solution (equation 3), reaction of CO2 with HO− to generate HCO3− (equation 4), and crystallization of GBIGH22+ and HCO3− ions into GBIGH2(HCO3)2(H2O)2 (equation 5). The overall reaction is represented by equation 6 and has an equilibrium constant of 6.5 × 106 atm−2. The CO2 absorption is driven mostly by the last two steps, the HCO3− formation and the GBIGH2(HCO3)2(H2O)2 crystallization, which are highly favorable in comparison with GBIG protonation and dissolution of GBIG and CO2. The overall enthalpy of CO2 absorption, relative to solid GBIG (equation 6), is −84 kJ/mol or −42 kJ/mol CO2. For direct comparison with other aqueous sorbents, we may consider aqueous GBIG as the reference state, resulting in an enthalpy of CO2 absorption of −123 kJ/mol or −61 kJ/mol CO2. This falls close to the lower end of the 60–80 kJ/mol range of absorption enthalpies for most amine-based sorbents.31Rochelle G.T. Conventional amine scrubbing for CO2 capture.in: Feron P.H.M. Absorption-Based Post-Combustion Capture of Carbon Dioxide. Woodhead Publishing, 2016: 35-67Crossref Scopus (51) Google Scholar The CO2 release from crystalline GBIGH2(HCO3)2(H2O)2 (equation 7) was investigated by thermogravimetric analysis coupled with mass spectrometry (TGA-MS) and differential scanning calorimetry (DSC). The TGA-MS revealed a 48.3% mass loss around 112°C, with concomitant evolution of CO2 and H2O in about a 1:2 ratio (Figure S3). These results are consistent with the loss of 2 equiv of CO2 and 4 equiv of H2O (48.5% theoretical mass loss), as expected from equation 7. Periodic structure calculations based on density functional theory (DFT) confirmed that the concomitant release of water and H2CO3 (as CO2 + H2O) becomes the thermodynamically most favorable path with increasing temperatures, mostly because of its more favorable entropy than that of alternative paths involving stepwise release of H2O and H2CO3 (Figure S5) The DSC analysis (Figure S4) revealed an endotherm in the same temperature range corresponding to the loss of CO2 and H2O in the TGA, with a measured enthalpy of reaction of 243 ± 14 kJ/mol, which corresponds to 121.5 ± 7 kJ/mol CO2. The minimum energy requirement for the regeneration of GBIG can be obtained as the sum of the enthalpy of CO2 and H2O release (121.5 kJ/mol CO2) and the sensible heat corresponding to increasing the temperature of the GBIGH2(HCO3)2(H2O)2 crystals from ambient to about 120°C. The latter was estimated at 30 kJ/mol CO2 from the specific heat capacity of GBIGH2(HCO3)2(H2O)2 measured by DSC (Figure S6). Thus, the total regeneration energy for GBIG is estimated at 151.5 kJ/mol CO2 (3,443 kJ/kg CO2), which corresponds to a 24% reduction compared with monoethanolamine (MEA), a benchmark sorbent employed in industrial CO2 scrubbing (4,503 kJ/kg regeneration energy).32Gottlicher G. The energetics of carbon dioxide capture in power plants. US Department of Energy, Office of Fossil Energy, 2004Google Scholar In direct contrast to MEA and other aqueous amine sorbents,31Rochelle G.T. Conventional amine scrubbing for CO2 capture.in: Feron P.H.M. Absorption-Based Post-Combustion Capture of Carbon Dioxide. Woodhead Publishing, 2016: 35-67Crossref Scopus (51) Google Scholar the regeneration of GBIG is done in the solid state, thereby avoiding the large energetic penalties associated with the sensible heat and latent heat of vaporization of aqueous solutions, which may account for as much as two-thirds of their total regeneration energies.32Gottlicher G. The energetics of carbon dioxide capture in power plants. US Department of Energy, Office of Fossil Energy, 2004Google Scholar Isothermal TGA measurements of the GBIGH2(HCO3)2(H2O)2 crystals at T = 80°C, 90°C, 100°C, and 110°C (Figure S7) allowed for the kinetic analysis of the CO2 and H2O release (Figure 2). After plotting the fractional conversion (α) as a function of time (Figure 2A), the most common solid-state reaction kinetics models were tested, including Avrami-Erofeev, Prout-Tompkins, and geometrical phase-boundary (Figure S8).33Brown M.E. Introduction to Thermal Analysis. Kluwer Academic Publishers, 2001Google Scholar The best fit was found for the contracting volume model (Equation 1; Figure 2B), corresponding to the reaction’s initiation on the crystal’s surfaces, followed by inward advancement of the reaction interface, which results in a deceleratory α-t curve as the interfacial area decreases.33Brown M.E. Introduction to Thermal Analysis. Kluwer Academic Publishers, 2001Google Scholar1−(1−α)1/3=kt(Equation 1) The measured rate constants (k) increased with temperature, yielding a linear Arrhenius plot (Figure 2C), from which an activation barrier (Ea) of 102 ± 12 kJ/mol was obtained. Notably, the measured activation energy is very similar to the calculated barrier for the autocatalytic decomposition of carbonic acid of 92–109 kJ/mol (Figure S14), wherein the reaction is facilitated by the formation of a carbonic acid dimer that undergoes concerted double proton transfer concomitant with CO2 elimination in the rate-limiting step.23de Marothy S.A. Autocatalytic decomposition of carbonic acid.Int. J. Quantum Chem. 2013; 113: 2306-2311Crossref Scopus (19) Google Scholar, 24Ghoshal S. Hazra M.K. New mechanism for autocatalytic decomposition of H2CO3 in the vapor phase.J. Phys. Chem. A. 2014; 118: 2385-2392Crossref PubMed Scopus (20) Google Scholar Further evidence that the CO2 release reaction is initiated on the crystal’s surface was provided by multimodal all-optical imaging measurements at different temperatures. Specifically, confocal reflectance (CR) and transmission (TR) microscopies were applied to probe the changes induced by the temperature increase on the crystal surface and bulk, respectively. These measurements were selected according to the premise that at its early stage, a reaction starting on the crystal’s surface should have a profound effect on the surface—but not the bulk-dominated modalities—simply because the resulting surface changes should have only a limited effect on the optical transmission through the bulk crystal. Representative CR and TR images acquired from the (001) face of a GBIGH2(HCO3)2(H2O)2 crystal at 22°C and 109°C are shown in Figure 3, and additional images are provided in Figures S9–S11. The increase in temperature leads to clear changes in the CR image measured at 109°C, whereas the corresponding TR image exhibits noticeable but less significant changes, providing direct evidence for the reaction initiation on the crystal surface. To gain atomic-scale mechanistic insights into the initial steps of CO2 release, we performed ab initio molecular dynamics (MD) simulations for a 3 × 2 supercell slab model of the (001) surface of GBIGH2(HCO3)2(H2O)2, consisting of three molecular layers (Figure S12). The (001) surface was chosen because it is the dominant face in the observed crystal morphology, consistent with its small calculated surface energy in the gas phase (0.38 J/m2). As expected, high temperature conditions (450 K used in the MD simulations) forced the water molecules at the solid-vacuum interface to move to the vacuum region within one picosecond of the MD run. A direct proton transfer from GBIGH22+ to HCO3− was observed after 1.5 ps (Figure 4), followed by many proton transfer events back and forth between the bicarbonate oxygen and the imine nitrogen of the ligand in the next 3.5 ps. Under the MD simulation conditions, the proton transfer occurred only in the topmost molecular layer, in agreement with the experimental evidence that the reaction starts on the crystal surface. Removing the interfacial water molecules in a separate MD simulation (Figure S13) facilitated proton transfer and H2CO3 formation, as indicated by the longer residence times of the protons on the bicarbonate oxygens. These calculations support a reaction mechanism in which initial water evaporation from the crystal surface promotes proton transfer from the iminoguanidinium group to the bicarbonate anion to yield carbonic acid as a transient intermediate on the path to CO2 release. Extending the ab initio MD simulations to longer times to track the subsequent reaction steps and capture the full dynamics of CO2 release from the surface would be computationally prohibitive. However, on the basis of the foregoing results and prior investigations of carbonic acid chemistry,23de Marothy S.A. Autocatalytic decomposition of carbonic acid.Int. J. Quantum Chem. 2013; 113: 2306-2311Crossref Scopus (19) Google Scholar, 24Ghoshal S. Hazra M.K. New mechanism for autocatalytic decomposition of H2CO3 in the vapor phase.J. Phys. Chem. A. 2014; 118: 2385-2392Crossref PubMed Scopus (20) Google Scholar it is reasonable to assume that the next step in the CO2 release process involves the formation of hydrogen-bonded carbonic acid dimers, followed by autocatalytic decomposition into carbon dioxide and water. This conjecture is supported by additional DFT calculations on a simplified iminoguanidinium-bicarbonate cluster model (Figure S14) that found very low activation barriers of 10.5 and 16.3 kJ/mol for consecutive proton transfer events with the formation of a carbonic acid dimer. Furthermore, quantum chemical calculations at different levels of theory predict a barrier height of 92–109 kJ/mol for the concerted double proton transfer and CO2 release from an isolated carbonic acid dimer in the gas phase (Figure S14). This is in excellent agreement with the experimentally obtained activation energy of 102 ± 12 kJ/mol for the GBIGH2(HCO3)2(H2O)2 decomposition. Overall, the combined molecular dynamics and static DFT calculations strongly support a reaction mechanism involving low-barrier proton transfers from GBIGH22+ to HCO3− with the formation of a carbonic acid dimer intermediate, followed by CO2 and H2O release in the rate-limiting step. To demonstrate the practical utility of GBIG in capturing carbon dioxide, we ran a full CO2 separation cycle using flue gas simulant containing 12.8% CO2 (EPA protocol standard). The flue gas mixture was bubbled through an aqueous solution of GBIG, leading to the formation of an off-white precipitate within minutes (Figure 5; Video S1). The crystalline solid formed after 1 h of bubbling was collected by filtration and confirmed by powder X-ray diffraction (PXRD) and Fourier-transformed infrared spectroscopy (FTIR) to be pure GBIGH2(HCO3)2(H2O)2 (Figure S15). The filtrate was analyzed by ultraviolet-visible spectroscopy to determine the concentration of GBIG left in solution, which found that 89% ± 2% of the initial GBIG was converted to crystalline GBIGH2(HCO3)2(H2O)2. This corresponds to 1.78 ± 0.04 mol of CO2 captured per mol of GBIG. eyJraWQiOiI4ZjUxYWNhY2IzYjhiNjNlNzFlYmIzYWFmYTU5NmZmYyIsImFsZyI6IlJTMjU2In0.eyJzdWIiOiJiNGRlNjY0ZmQwOTcwMTFhZmRmNWFiYTk3OGQ0NjY1NSIsImtpZCI6IjhmNTFhY2FjYjNiOGI2M2U3MWViYjNhYWZhNTk2ZmZjIiwiZXhwIjoxNjc4NDEzNjU2fQ.gfpAFWGKTur1HmSiKnMG2Wj7hB2-Qh1N3vdfI2CTTnrUzB3CwZhGSMU7OHDktHFj3_XDRK-3rCn5K_OFacZjAUN0ThsmeQhpbOBuLUKu_M10Q83cl-psi5vLKoWcVFC6wd1VGehjYMy3RY-NXTnZfhR5eQumjHlO-b9BOUBVsOk3R9q195xJMcLGQOwjYrB7lcpjfuoPhLsxiUYodkgeEuEBmsC_ADLcj30KKoU6vOO4AgaIeTVX30d53nq-cWm9_LptN2EUfAsSHYel3wcb7eiCzm9zfGp6L2ZeQZL6sEU-Dn_9hIHnxIxuwZEqJKg-JWtQrMHMivNX-mMItgIHLg Download .mp4 (5.32 MB) Help with .mp4 files Video S1. CO2 Capture ExperimentTime-lapse video of flue gas simulant (EPA protocol standard, 12.8% CO2 + 87.2% N2) bubbling through an aqueous solution of GBIG (10 mM) at a flow rate of 0.4 L/min. The onset of GBIGH2(HCO3)2(H2O)2 crystallization is clearly seen after about 5 min Heating the isolated GBIGH2(HCO3)2(H2O)2 solid at 120°C for 2 h led to a quantitative release of CO2 and water, as determined by gravimetric and elemental analyses, and the resulting solid was confirmed by PXRD (Figure S16) and FTIR (Figure S17) to be anhydrous GBIG. The recovered GBIG solid was redissolved into the filtrate and the sorbent was recycled. Overall, ten consecutive CO2 capture-release cycles were conducted (Figure 6), with only 3% observed decrease in the CO2 absorption capacity. Although the long-term stability and recyclability of GBIG remain to be tested over many more CO2-capture cycles under real-world conditions, our initial results indicate that this simple bisiminoguanidine sorbent is remarkably robust. On the basis of isothermal TGA measurements (Figure S18), GBIG showed no signs of decomposition even after 1 week of continuous heating at 120°C in air. We demonstrated efficient CO2 scrubbing from flue gas mixtures by using a remarkably simple and robust bis-iminoguanidine compound (GBIG), whose synthesis by the imine condensation of glyoxal with aminoguanidinium salts was first reported in the literature 120 years ago. CO2 absorption by the aqueous GBIG sorbent leads to crystallization of a relatively insoluble bicarbonate salt, with an estimated reaction enthalpy of −61 kJ/mol CO2. Single-crystal X-ray diffraction analysis of the GBIG bicarbonate salt revealed the formation of “anti-electrostatic” hydrogen-bonded (HCO3−)2 dimers, linked by water molecules into extended ladder-shaped clusters, and stabilized by multiple hydrogen bonds from the iminoguanidinium cations. The CO2 can be released, with concomitant loss of water, by mild heating of the bicarbonate crystals at 80°C–120°C, resulting in quantitative regeneration of GBIG, which can be recycled multiple times. Kinetic measurements and optical imaging, together with ab initio MD simulations and DFT calculations, provided experimental and computational evidence for a CO2 release mechanism consisting of surface-initiated low-barrier proton transfer from the iminoguanidinium cations to the bicarbonate anions with the formation of carbonic acid dimers, followed by CO2 and H2O release in the rate-limiting step, with a measured activation energy of 102 ± 12 kJ/mol. The CO2 capture cycle was tested with a flue gas simulant, showing excellent performance (89% crystallization yield, quantitative GBIG recovery) over ten consecutive cycles. The minimum energy required for the regeneration of GBIG is 151.5 kJ/mol CO2, which is 24% lower than the regeneration energy of MEA, a benchmark industrial sorbent. The significant energy savings are a benefit of GBIG regeneration being done in the solid state, which avoids heating and evaporation of aqueous solutions, energy intensive processes typically involved in the regeneration of aqueous sorbents. With further optimization and scale-up, this simple CO2 separation cycle offers the prospect for an energy-efficient and cost-effective carbon-capture technology that could help mitigate climate change." @default.
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• W2912476701 title "CO2 Capture via Crystalline Hydrogen-Bonded Bicarbonate Dimers" @default.
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