Adsorptive separation of carbon dioxide: From conventional porous materials to metal–organic frameworks

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Adsorptive separation of carbon dioxide: from conventional porous materials to metal–organic frameworks Dong-Dong Zhou , Xue-Wen Zhang , Zong-Wen Mo , Yu-Zhi Xu , Xiao-Yun Tian , Yun Li , Xiao-Ming Chen , Jie-Peng Zhang PII: DOI: Reference:

S2589-7780(19)30019-3 https://doi.org/10.1016/j.enchem.2019.100016 ENCHEM 100016

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EnergyChem

Received date: Revised date: Accepted date:

24 June 2019 8 September 2019 9 September 2019

Please cite this article as: Dong-Dong Zhou , Xue-Wen Zhang , Zong-Wen Mo , Yu-Zhi Xu , Xiao-Yun Tian , Yun Li , Xiao-Ming Chen , Jie-Peng Zhang , Adsorptive separation of carbon dioxide: from conventional porous materials to metal–organic frameworks, EnergyChem (2019), doi: https://doi.org/10.1016/j.enchem.2019.100016

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Highlights   

Key points for CO2 adsorptive separation are comprehensively analyzed. Many active sites and strategies can be used to improve CO2 adsorption. Meal–organic frameworks are ideal for studying CO2 adsorptive separation.

Adsorptive separation of carbon dioxide: from conventional porous materials to metal–organic frameworks Dong-Dong Zhou,a Xue-Wen Zhang,a Zong-Wen Mo,a Yu-Zhi Xu,a,b Xiao-Yun Tian,a Yun Li,a Xiao-Ming Chen,a and Jie-Peng Zhang*,a

a

MOE Key Laboratory of Bioinorganic and Synthetic Chemistry, School of Chemistry, Sun Yat-Sen

University, Guangzhou 510275, China b

School of Materials Science and Engineering, South China University of Technology, Guangzhou

510640, China * Correspondence: [email protected] (J.-P. Z.)

Abstract: Excessive CO2 emission is the main cause of global greenhouse effect. Separating CO2 from relevant sources such as flue gas, natural gas, and syngas may reduce emissions directly or indirectly by increasing the energy efficiency. Compared with amine absorption, adsorptive separation offers merits such as a low energy cost, no corrosion, and many other advantages, and has attracted increasing attention, especially in the past 10 years. Different types of adsorbents such as activated carbons, zeolites, porous silica, porous polymers or porous organic polymers, and porous coordination polymers or metal–organic frameworks have demonstrated many unique characteristics for CO2 adsorptive separation. This review discusses the main challenges in CO2 adsorptive separation and analyses the strategies developed for the typical types of porous materials, especially for metal–organic frameworks.

Keywords: Carbon dioxide; Adsorption; Separation; Porous materials; Metal–organic frameworks

1. INTRODUCTION The global greenhouse effect has caused many crises, including rising temperatures, melting of glaciers, rising sea levels, accelerating the extinction of species, and so on. CO2 is the main greenhouse gas, and its emission reached 36.8 Gt in 2017, which is five times higher than that in 1950.1,2 The atmosphere CO2 concentration in April 2018 reached 410 ppm, which is the highest in the past 800000 years; it was only 280 ppm in 1780.3 Reducing energy consumption is the most promising way to reduce CO2 emission, since fossil fuels are still the main sources of energy. Chemical separation in industry contributes ca. 1/6th of the global energy consumption.4 Therefore, developing more efficient separation methods is very meaningful in the aspects of both energy and the environment. Before sufficient optimisation of the global energy structure and development of energy-efficient technologies, carbon capture and sequestration/storage can serve as a short-term strategy for lowering the CO2 emission. The CO2 captured may be consequently utilised or stored by many methods.5,6 CO2 can be used in the food industry (dry ice, carbonated beverages, etc.) or can serve as a raw material in the chemical industry for synthesising value-added products (polycarbonate, urea, etc.). High-concentration CO2 can improve the photosynthesis efficiency of plants. Electrochemical reduction of CO2 to produce low-carbon fuels and other chemical products is an effective approach to making full use of renewable energies. Direct air capture (DAC) of CO2, being somewhat similar to the carbon fixation process of photosynthesis, has attracted great academic interest.7,8 However, it is obvious that capturing CO2 from high-concentration environments/sources such as flue gas, biogas, and natural gas is much easier than DAC. The CO2 emissions from fossil fuel combustion exceed 80% of the total emissions.9,10 With regard to the combustion industry, the CO2 separation technologies can be divided into three types: pre-combustion, post-combustion, and oxy-combustion. Post-combustion CO2 capture involves the fuel gas, which mainly contains ca. 15% CO2, ca. 75% N2, enriched from air, and ca. 3−4% O2, as well as saturated water vapour and small amounts of acid gases such as SOx and NOx, with a total pressure of ca. 1 atm and a temperature of ca. 40−80 °C.11 In pre-combustion CO2 capture, the fuel is converted to H2 (ca. 70−80%) and CO2 (ca. 15−25%).12,13 Oxyfuel combustion uses pure O2 for fuel conversion, producing mainly water vapour and CO2 as the flue gas, which is easy to separate, but efficient production of pure O2 is still a great challenge. Most efforts

have been devoted to post-combustion CO2 capture. Separating the CO2 present in biogas, natural gas, and syngas is also beneficial for subsequent uses of these gases. Depending on the sources, they display very different CO2 concentrations. Biogas mainly contains ca. 60−70% CH4 and ca. 30−40% CO2, and is produced at ambient temperature and pressure.14 Natural gas mainly consists of CH4 and generally contains ca. 1% CO2 (in some cases >40%), and is produced at ca. 30−60 °C and 50−70 bar.15 Syngas mainly consists of 10−30% CO2, 20−40% CO, and 30−40% H2, and is usually produced above 400 °C at ca. 20−40 bar.16 The separation of C2 hydrocarbons and CO2 is also worth studying. For example, in addition to CH4 and CO2, natural gas often contains C2H6. In modern industry, C2H2 is produced by selective oxidation of CH4 and hydrocarbon cracking, with CO2 being one of the by-products.

1.1 CO2 separation methods Compared with those of other relevant gases (H2: 20 K, N2: 77 K, CO: 82 K, O2: 90 K, CH4: 112 K), CO2 (195 K) exhibits a much higher boiling point. This means CO2 can develop much stronger intermolecular interactions with other molecules, substrates, or materials. In fact, CO2 also displays larger quadrupole moment and molecular weight, which permit stronger multipole–multipole (electrostatic or Keesom) and dispersion interactions. Therefore, CO2 is generally the strongest absorbed/adsorbed gas in a mixture. By reducing the pressure/concentration and/or elevating the temperature, the CO2 absorbed/adsorbed can be released to regenerate the absorbent/adsorbent. Inert gases are usually (especially for academic studies used to promote desorption by reducing the partial pressure of the desorbed gas, behaving like the VSA process. Nevertheless, to obtain pure CO2, the desorption process should be carried out in pure CO2 atmosphere. In practical applications, the selection of desorption method/condition mainly depends on the characteristics of the absorbent/adsorbent, especially the CO2 affinity and water/heat stability. Note that to save energy and time, it is not necessary to desorb all the absorbed/adsorbed gases. Therefore, the working capacity for CO2 capture is the difference of uptakes between absorption/adsorption and desorption conditions, just the same as that for gas storage. Absorption by liquids. Owing to its high CO2 binding affinity and selectivity, an aliphatic amine (mainly monoethanolamide; MEA) solution has been applied in industrial CO2 separation. Amines, except tertiary amines (owing to their large steric hindrances and the absence of active

protons), react with CO2 to yield carbamate and ammonium, and the carbamate is hydrolysed in water to form carbonate/bicarbonate and amine. Tertiary amines in water also form carbonate/bicarbonate and ammonium (Fig. 1). The main drawbacks of amine solutions are low CO2 capacity, high energy consumption (high regeneration temperature and high heat capacity of water), equipment corrosion, and degradation of the amine by evaporation, oxidation (by O2), and formation of stable salts (SO2 and NO2).17 For example, the CO2 uptake of a typical MEA solution (30 wt%) is only ca. 1.2 mmol g–1 for flue gas.18,19 Based on the relatively high solubility of CO2 in water, methanol, polyethylene glycol dimethyl ether, and other liquids, physical absorption with low corrosion and heat consumption but of a low efficiency has a few industrial applications for separating high-concentration CO2.17 Adsorption by solids. Gas molecules can attach onto solid surfaces by intermolecular interactions, physically and/or chemically. Porous materials exhibit much larger surface areas than common solids. Moreover, the pore sizes, shapes, and surface functionalities of porous materials are additional parameters for the modification of solid–gas interactions. Except for H2 (289 pm), CO2 displays the smallest kinetic diameter (330 pm), compared with those of other relevant gases (O2: 346 pm, N2: 364 pm, CO: 376 pm, CH4: 380 pm), which means that it has a kinetic advantage in diffusing through and adsorbing in the pores. Membrane separation is another application of porous materials; it offers quite different characteristics and challenges compared with adsorptive separation,20,21 and is not discussed in this review. Adsorption by using solid porous materials can overcome the many shortcomings of liquid absorption such as corrosion and energy consumption. Besides temperature swing adsorption (TSA), pressure swing adsorption (PSA), and their combination, vacuum swing adsorption (VSA) can also be facilely applied to regenerate the adsorbent, because a solid adsorbent does not vaporise like a liquid does. The main problem with the conventional porous materials for CO2 adsorptive separation is the relatively weak CO2 binding and selectivity (compared with the case of chemical absorption). However, some porous materials can be functionalised by aliphatic amines, and physical adsorption of ultra-micropores may be as strong as chemisorption.22 Compared with CO2, H2O is smaller and exhibits much higher polarity and boiling point, therefore, it can be adsorbed more strongly by polar functional groups, which is another problem of adsorptive separation of CO2. Moreover, water may destroy porous materials, especially for those based on non-covalent bonds. Note that, H2 exhibits an

extremely low boiling point and can hardly be adsorbed, therefore, CO2/H2 can be well separated by using commercial adsorbents such as activated carbons. On the other hand, it is very difficult to separate C2 hydrocarbons (the boiling points are C2H2: 189 K, C2H4: 169 K, C2H6: 184 K) and CO2, owing to the similarities in the physical properties and molecular shapes/sizes.

1.2 Evaluation of CO2 adsorptive separation The CO2 separation performance of an adsorbent should be evaluated based on many aspects. Pore volume, surface area (usually calculated by the BET (Brunauer-Emmett-Teller) and Langmuir models), and pore size are not only basic porosity parameters but also necessary for predicting the gas adsorption and separation properties. These parameters are generally calculated from the N2/Ar/CO2 isotherm measured at or near the adsorbate boiling point. Note that neither adsorptive CO2 separation nor storage is practiced at 195 K. However, if an ultra-microporous adsorbent can adsorb CO2 195 K and cannot adsorb N2 at 77 K, it might display a kinetic separation effect at room temperature for gases. Pore volume determines the saturated guest uptake which can be achieved at low temperatures and/or high pressures, which are generally far from the conditions for practical CO2 separation. Surface area usually determines the guest uptake at medium temperatures and pressures, which can be related to the separation of CO2 at a high concentration/pressure, such as from biogas and syngas. Pore size relates to the guest binding affinity, and small pores can provide stronger overlaps of the potential energy to increase the adsorbent–adsorbate binding. However, conventional adsorbents consisting of small pores usually display small pore volumes and surface areas.

1.2.1 Single-component adsorption and desorption Using commercial instruments, single-component gas adsorption and desorption isotherms can be conveniently and accurately measured around room temperature and at pressures ranging from near vacuum to ca. 1 bar. When adsorbent regeneration is considered, the isotherms can be measured at higher temperatures. CO2/N2 selectivity is generally very high. Therefore, when the CO2 uptake of an adsorbent for pure CO2 at ca. 0.15 bar is very high, it can be approximately considered as the CO2 uptake of this material for flue gas (though there is some overestimation), which can be true in the case of chemisorption involving aliphatic amines. Although 1 bar is not the concentration of CO2 in

air or flue gas, the CO2 uptakes at this pressure have been frequently reported and compared. The gas uptakes at ambient temperature and pressure are usually similar to those measured at somewhat higher temperatures and pressures, which can be useful in predicting the separation performances for natural gas and syngas. High-pressure isotherms can be directly measured by using high-pressure apparatus, but the measurement is less convenient and the results are less accurate than those obtained using the common apparatus which only works below ambient pressure. Note that, because the saturated vapour pressure of CO2 at ambient temperature is not very high (64.5 atm at 298 K), the CO2 uptake at a very high pressure essentially reflects the adsorbent porosity, rather than the storage or separation performance, which is similar to the case of the uptake measured at 1 bar and 195 K. Owing to their very low boiling points, H2, N2, CO, and O2 generally show very low uptakes at ambient temperature and pressure, which suggest that the CO2 separation performances are mainly dependent on the CO2 adsorption behaviours rather than those of the weakly adsorbed gases. In other words, the CO2 separation performance of an adsorbent can be qualitatively predicted with only the CO2 adsorption data. Nevertheless, without using large amounts of adsorbent samples and/or carefully checking and fitting their isotherms, large relative errors can easily occur for these weakly adsorbed gases. The gas uptake (for a given amount of adsorbent) can be expressed in many different units; the adsorbate/adsorbent amount can be expressed as a volume (cm3 at STP), weight (g), or in moles (mmol). For both volumetric and gravimetric measurement methods, only the weight of adsorbent can be directly and reliably measured, and the volume of adsorbent can be transduced from the weight, according to the crystal density, if available. For practical applications, the bulk density (which is always lower than the intrinsic/crystal density) of the adsorbent packed in a container volume is necessary. Both the gravimetric and volumetric uptakes are important parameters, but the volumetric uptake is more important for CO2 separation, because, in most cases, the separation unit filled with an adsorbent does not need to move during operation.22,23 For consistency and to avoid confusion, the gas uptakes in this review are expressed in mmol g–1 and/or mmol cm–3. Based on the single-component isotherms of relevant gases, the CO2 separation performances, such as the selectivity and working capacity, can be predicted through many methods. For example, the CO2 selectivity can be calculated (predicted) from the ratio of the initial isotherm slopes, from the ratio of the uptakes at relevant pressures, and by the ideal adsorption solution theory (IAST).

Note that, a small error (in the absolute value) for a weakly adsorbed gas can produce a very large error in the selectivity. In addition, the popularly used IAST not only requires accurate fitting of the single-component isotherms based on appropriate models, but also has many restrictions in its application. For example, flexible or heterogeneous adsorbents, adsorbates with significantly different adsorbing abilities (e.g. CO2/H2), and molecular sieves with different accessible surfaces for different adsorbates are not suitable for IAST.23 Therefore, the predicted selectivities should be regarded as qualitative measures rather than quantitative measures. Besides the thermodynamic difference, guest molecules can be kinetically separated based on the difference in the diffusion rate, which arises when the pore size is very small and similar to that of the guest molecules. In the ideal situation, the pore size is smaller than that of some of the guest molecules, which results in the molecular sieving effect with infinite selectivity (but, in practice, adsorption on a particle surface is unavoidable). Although molecular sieving displays a simple mechanism, precise control of the pore size is very challenging, especially for flexible adsorbents.

1.2.2 Mixture adsorption and separation Considering that the adsorption and separation behaviours of many adsorbent–adsorbate systems are still difficult to predict, direct measurement of the related data is obviously very important, even if it is only for academic purpose. Compared with the case of single-component isotherms, the measurement of mixture isotherms is very complicated and time-consuming. Among the various measurement methods, fixed-bed breakthrough is the most widely adopted, because it closely resembles the industrial separation operations. Using breakthrough setups, the adsorption and separation performance under a given temperature, pressure, and mixture composition (including water and other interfering components) can be readily controlled. The breakthrough times of the adsorbates directly reveal their relative adsorption tendencies, which are combinations of the thermodynamic and kinetic effects. When the outlet flow rate of each adsorbate is measured, the adsorbate uptake at any time can be obtained by integrating the breakthrough curve. However, measurement of the outlet flow rate is much more difficult than that of the relative concentration. Thermogravimetry (TG) is a powerful method for studying not only the thermal stability but also the gas adsorption property of adsorbents, because the temperature and gas compositions can be readily programmed, and the weight change of the sample directly corresponds to the total guest

uptake at any time. For example, heating the sample under pure CO2, purging the sample under inert gas, and heating the sample under inert gas during the desorption stage simulate the TSA, VSA, and the temperature-vacuum swing adsorption (TVSA) processes, respectively. In the absence of a strongly adsorbed species such as H2O, the total uptake can be approximately regarded as the sole contribution of CO2. For CO2/N2 separation, the weight change can be assigned to the change of CO2 uptake, because N2 adsorption at room temperature is generally too weak to measure. In this way, the working capacity and kinetic data can be conveniently obtained, which is useful for optimizing the desorption condition. In principle, fix-bed column breakthrough experiments can give comprehensive adsorption/desorption data including uptakes of all involved gases, adsorption selectivity, working capacity, adsorption/desorption kinetics, etc., but accurate measurement of these parameters are very challenging. In other words, in mixture adsorption/separation experiments, the gas uptakes (especially for weakly adsorbed gases) are still difficult to measure accurately. In fact, adsorbents showing very strong CO2 bindings, such as those functionalised by aliphatic amines, always exhibit very high CO2/N2 selectivities, which are very difficult and not required to measure. On the other hand, for adsorbents with relatively weak CO2 bindings, the CO2/N2 selectivities are of little practical significance. Therefore, although adsorption selectivity is important for academic and practical purposes,24 it is not emphasised in this review.

1.2.3 Adsorption and separation mechanism The structure–property relationship, including the structure, energy, and diffusion, is critical for designing/modifying an adsorbent. Because the relevant gases (H2, N2, CO, O2, and CH4) generally display very weak adsorptions, the CO2 adsorptive separation performance mainly depends on the CO2 adsorption behaviour. Besides the classic carbamate mechanism for aliphatic amines, which can be reliably predicted, CO2 can also be adsorbed by many other types of interactions which vary significantly in structure and energy. For example, CO2 molecule can coordinate with open metal sites (OMSs) through its oxygen atoms.25-28 CO2 can also form hydrogen bonds and various multipole–multipole interactions through its electronegative oxygen atoms and the electropositive carbon atom. The detailed host-CO2 structure can help elucidate the adsorption energy and separation performance. There are many experimental methods for determining the host–guest structure. For crystalline

porous materials, the atomic parameters of the host–guest systems, including the bond lengths and angles, can be straightforwardly visualised by using diffraction techniques.29-34 The high mobility and disorder of gas molecules in the crystal structures are the main problem,35,36 which may be overcome by using a low temperature and a high gas loading, although these deviate from the separation conditions. Fortunately, for adsorbents showing satisfactory CO 2 separation performances, the CO2 bindings are usually strong enough to clearly show the guest locations even at above room temperature for a low gas loading.32,33 Single-crystal X-ray diffraction (SCXRD) can provide very accurate and comprehensive results, but maintaining good single-crystallinity after the activation and adsorption processes is always challenging.32,37 Alternatively, powder X-ray diffraction (PXRD), with a lower accuracy, can be used, although the location and motion of the gas molecules are difficult to determine.35,38 Compared with the case of commercialised X-ray diffractometers, synchrotron and neutron radiations can provide a higher accuracy and more information,30,31,39-43 but they are only available in limited sources. Many other methods can be applied for both crystalline and amorphous samples. For example, nuclear magnetic resonance (NMR) spectroscopy can provide detailed molecular-level information on the local chemical environment, including molecular motion.44 Infrared/Raman spectroscopy can straightforwardly indicate the formation/cleavage or change in the covalent and coordination bonds. X-ray absorption spectroscopy and X-ray scattering can provide the local structural information pertaining to the geometry and interatomic distances.45-47 The CO2 adsorption enthalpy is closely related to the separation performance. A high adsorption enthalpy usually corresponds to large isotherm slope and adsorption selectivity. However, a high adsorption enthalpy also increases the thermal conduction load and requires a large amount of energy for adsorbent regeneration. In other words, a higher temperature and/or a lower pressure are required for desorption for the TSA/PSA and related separation operations. Obviously, an ideal adsorbent should be easy to desorb CO2. While PSA is a commercial technology for many practical applications for the low energy requirement and fast regeneration of adsorbents,20,48 TSA is considered to be energy efficient for post-combustion CO2 capture, because it is difficult to compress or apply a vacuum to a large volume of the low-pressure post-combustion gas stream.23,49 The energy consumption for CO2 desorption in a TSA process can be simply quantified by considering two parts: 1) the energy needed to heat up the adsorber to reach the desorption conditions (sensible heat), which

can be calculated using the specific heat capacity Cp of the adsorbent, the total mass of the adsorbent, and the temperature difference between adsorption and desorption conditions, and 2) the energy required to undo the adsorption process, which consists of the working capacities multiplied by the adsorption enthalpy for each flue gas component.50 The host–guest binding strength can be qualitatively judged from high-quality diffraction and spectroscopy data, such as the thermal motions of atoms determined by SCXRD and NMR.32,51 Because direct measurement (by calorimetric method) is relatively complicated, the adsorption enthalpy or isosteric heat (Qst) is usually calculated from the adsorption isotherms by the van't Hoff equation. Computational simulation using density functional theory (DFT), periodic density functional theory (PDFT), molecular mechanics (MM), grand canonical Monte Carlo (GCMC), and/or other methods is powerful for studying and predicting adsorption and separation behaviours. The energy consumption associated with the guest-induced structural transformation of an adsorbent is difficult to measure, and is usually calculated by computational simulation, although the available simulation methods still find it difficult to accurately handle flexible solids.52 Diffusion behaviours are also mainly studied by computational simulations.52,53 Compared with the case of amorphous solids, crystalline materials with periodic structures can be modelled more easily through computational simulations. Furthermore, simulating CO2 adsorption is much easier than simulating the adsorptions of other weak adsorbing gases. Adsorptive separation of CO2 has received significant research interest, especially in the past 10 years (Fig. 2). There have been many reviews focusing on CO2 separation by both adsorption and membranes,23,53-56 CO2 separation from flue gas only,20,57-60 and both CO2 separation and conversion.61-63 According to the composition and structure, porous materials can be categorised into many types. Since each type of adsorbent exhibits unique characteristics for CO2 adsorptive separation, most reviews concentrated on individual types or groups of porous materials.23,53,64-69 This review intends to comprehensively analyse the common strategies involving typical types of porous materials for CO2 adsorptive separation from not only flue gas but also other mixtures (air, natural gas, syngas, biogas, etc.). The representative types of porous materials, from the conventional ones to the modern ones, including activated carbons, zeolites, porous silica, porous polymers or porous organic polymers (POPs), and porous coordination polymers (PCPs) or metal–organic frameworks (MOFs), are sequentially discussed with select examples. Being consistent with the

research activities on different types of adsorbents, our discussions mainly focus on MOFs, because they have demonstrated an overwhelmingly rich chemistry for CO2 adsorption/separation (Scheme 1),30,34,70-77 compared with the other types of adsorbents.

Scheme 1. Key strategies and results for MOF-based CO2 adsorptive separation.

2. POROUS CARBONS FOR CO2 ADSORPTIVE SEPARATION Porous carbons, mostly referred to as activated carbons, have been widely used in adsorptive separation. Activated carbons can be produced by pyrolysis of various carbon-containing precursors.67 Coal, pitch, and biomass are the most commonly and commercially used precursors, but their carbon structures are difficult to control. As carbon precursors, synthetic polymers can provide better controls over the composition, structure, and morphology.78 MOFs offer similar advantages as synthetic polymers, but they display higher structural accuracies, incorporate metal ions, and reveal large porosities; they are emerging as intriguing carbon precursors and templates of carbon precursors.79,80 For CO2 separation applications, activated carbons exhibit the advantages of high stability, low cost, and relatively low water affinity, but suffer from the disadvantages of weak CO2 binding, non-uniformity, and hardly controllable structures. Several commercial activated carbons have been studied for CO2 separation from flue gas, but they display low CO2 uptakes (~0.6 mmol g−1 at ~0.15 bar and 298 K) and low CO2/N2 selectivities.81,82 Tuning the pore size and doping heteroatoms are the main strategies for improving the CO2 separation performances of activated carbons. As a classic example, Gogotsi et al. studied the correlation between the CO2 uptake and the pore

size of porous carbons.83 A series of porous carbons were obtained by the reaction of Cl2 with titanium carbide powders of different sizes at 200−1200 °C. The N2 adsorption isotherms showed that all the samples exhibit broad pore size distributions. The pore volume, BET surface area, and pore size distribution can be adjusted depending on the precursor size and the reaction condition. The CO2 adsorption isotherms obtained at 273 K showed that the CO2 uptakes at 0.1 bar and 1.0 bar are poorly correlated to the pore volume or surface area. Instead, the volume of pores smaller than 8 Å or 5 Å correlated linearly with the CO2 uptake at 1.0 bar or 0.1 bar, respectively (Fig. 3). In other words, at 1.0 bar and 0.1 bar, pores smaller than 8 Å and 5 Å, respectively, contribute the most to the CO2 uptake. Nevertheless, the volume of pores smaller than 5 Å is relatively small, which leads to a low CO2 uptake, which indicates that narrowing the pore size distribution can hardly improve the CO2 uptake at room temperature and low pressures. The pore sizes of porous carbons can play an extraordinary role for CO 2 separation. Carbon molecular sieves (CMSs) are a special class of porous carbons which display relatively uniform pore sizes. Compared with conventional activated carbons, CMSs require precise controls of the precursor structures and pyrolysis conditions, and/or the chemical vapour deposition of the additional carbons to reduce the pore sizes.84 Several CMSs, such as Bergbau Forschung and 3K-161, have been used in commercial CO2 capture from landfill gas.85,86 Besides the intrinsic thermodynamic selectivity, these carbons with narrow pores exhibit a high CO2 diffusivity, compared with that of CH4, which provides additional kinetic selectivity. Molecular sieving of CO2/CH4 is easier than that of CO2/N2, because the kinetic diameter of CH4 is greater than that of N2. For example, Silvestre-Albero et al. prepared a CMS named VR-93-M by pyrolysing a mixture of KOH and mesophase pitch, which showed a high CO2 uptake of 3.75 mmol g−1 and a low N2 uptake of 0.25 mmol g−1 at 298 K and 0.53 bar, but revealed complete exclusion of CH4 at the same condition (Fig. 4).87 Doping heteroatoms, especially nitrogen, is the main strategy for improving the CO2 bindings of porous carbons, which can be achieved by proper selection/design of precursors and/or post-synthetic

modification (PSM).59

Nitrogen

doping

is

usually performed

by

using

nitrogen-containing precursors and/or through a high-temperature reaction of porous carbons with NH3. Natural precursors with rich nitrogen contents, such as chitosan and pine cone, have been used to prepare nitrogen-doped activated carbons. Most precursors of activated carbons contain oxygen, which can be retained when carbonisation is incomplete. In porous carbons, the nitrogen species are

mainly graphitic, pyridinic, pyrrolic, etc., whereas the oxygen species mainly exist in the form of hydroxyls, carboxyls, carbonyls, etc. Metal azolate frameworks (MAFs), as a subclass of MOFs consisting of nitrogen-heterocycle ligands, are suitable precursors for preparing nitrogen-doped activated carbons. Banerjee et al. first studied the CO2 adsorption properties of nitrogen-doped porous carbons derived from MOFs.88 Isostructural metal imidazolate frameworks [Zn(nIM)(bIM)] (ZIF-68; HnIM = 2-nitroimidazole, HbIM

=

benzimidazole,

pore

size

10.3

Å),

[Zn(nIM)(5cbIM)]

(ZIF-69;

H5cbIM

=

5-chlorobenzimidazole, pore size 7.8 Å), and [Zn(nIM)0.87(IM)1.13] (ZIF-70; HIM = imidazole, pore size 15.9 Å) with GME topology and different porosities were used to host furfuryl alcohol and then serve as carbon precursors to prepare three nitrogen-doped activated carbons, which were named C-68, C-69, and C-70, respectively. BET surface areas of 1171−1510 m2 g−1, pore volumes of 0.725−1.749 cm3 g−1, nitrogen contents of 7.02−13.02 wt%, and CO2 uptakes of 4.54−5.45 mmol g−1 were obtained at 273 K and 1 bar for the three samples, and these showed the same orders as the nitrogen contents of their precursors i.e. ZIF-70 > ZIF-68 > ZIF-69. In general, the nitrogen content of activated carbon increases with that of the precursor. On the other hand, the structure of the precursor should also play an important role in determining the nitrogen content of the obtained activated carbon. However, carbon precursors including MOFs with different structures generally exhibit different nitrogen contents. By using isomeric MOFs (same composition, but different structures), Zhang et al. showed the simple structural relationships between carbon precursors and activated carbons.89 Three supramolecular isomers of [Zn(eim)2] (Heim = 2-ethylimidazole), namely MAF-5 (ANA, 38.6%), MAF-6 (RHO, 55.4%), and MAF-32 (qtz, 0%) with different network topologies and porosities (Fig. 5a), were pyrolysed to yield nitrogen-doped porous carbons, which were named CMAF-5, CMAF-6, and CMAF-32, respectively. The morphologies of the precursors were retained after carbonisation, and the orders of the pore volumes (0.35−0.65 cm3 g−1), BET surface areas (730−1453 m2 g−1), and pore sizes (3.7−7.5 Å) of the obtained nitrogen-doped activated carbons showed the same trends as those of the precursors i.e. CMAF-6 > CMAF-5 > CMAF-32 (Fig. 5b). The nitrogen contents showed the reverse order of CMAF-32 (5.45 wt%) > CMAF-5 (5.05 wt%) > CMAF-6 (2.55 wt%), which indicated that larger pores of the precursor benefit the removal of nitrogen species during carbonisation. The CO2 uptakes at 298 K and 1 bar followed the same trend as the pore volumes (Fig. 5c), namely CMAF-6 (3.53

mmol g−1) > CMAF-5 (3.39 mmol g−1) > CMAF-32 (3.26 mmol g−1). However, the uptakes at 0.15 bar revealed the opposite trend i.e. CMAF-32 (1.30 mmol g−1) > CMAF-5 (1.15 mmol g−1) > CMAF-6 (1.09 mmol g−1), which indicated that CMAF-32 adsorbed CO2 the strongest. Indeed, the Qst followed the order of CMAF-32 (56.1 kJ mol−1) > CMAF-5 (50.6 kJ mol−1) > CMAF-32 (41.2 kJ mol−1), which is consistent with the order of the nitrogen contents and the reversed order of the porosities. All the samples adsorbed a small amount of N2, and the IAST CO2/N2 selectivities were calculated as 37, 180, and 262 for 15:85 CO2/N2 mixture. Column breakthrough experiments of CMAF-32 using 15:85 CO2/N2 mixture yielded a selectivity of 234. Through non-uniform doping, activated carbons can form heterojunctions. Yuan et al. showed that the electron transfer across the heterojunction of activated carbon can improve the CO2 adsorption.90 Microporous carbon fibres (PCFs) were coated with nitrogen-containing poly(ionic liquid) (PIL) and then pyrolysed to yield the hybrid microporous carbon fibers C-PIL/PCF, with nitrogen-doped porous carbon (C-PIL) as the shell and the non-doped PCFs as the core. While the PCFs and C-PIL showed low CO2 uptakes of 2.3 mmol g−1 and 2.8 mmol g−1 at 273 K and 1 bar, respectively, C-PIL/PCF displayed uptakes of 6.9/4.7 mmol g−1 at 273/298 K and 1 bar, and 1.1 mmol g−1 at 298 K and 0.15 bar. Column breakthrough experiments (298 K) using 20:80 CO2/N2 mixture showed a CO2 adsorption capacity of 2.1 mmol g−1. Interestingly, because CO2 adsorption changes the electrical resistance, C-PIL/PCF can also be used as a CO2 sensor. Heteroatoms generally increase the hydrophilicity of porous carbons, which decreases the CO2 adsorption and separation performances. Interestingly, Tour et al. showed that pre-adsorption of H2O on activated carbon may increase the CO2 separation performance (Fig. 6a).91 Pyrolysis of asphalt/KOH mixture afforded the activated carbon uGilT, with a large BET surface area of 4200 m2 g−1 and a pore volume of 2.41 cm3 g−1. It showed CO2 and CH4 uptakes of 30 mmol g−1 and 12 mmol g−1, respectively, at 298 K and 35 bar. Adsorption of water resulted in the host–guest material uGilTH2O (containing 148 wt% H2O) with negligible porosity (BET surface area 7 m2 g−1). The CO2 uptake of uGilTH2O drastically increased from 4.6 mmol g−1 at 20 bar to 30 mmol g−1 at 40 bar (Fig. 6b), whereas the CH4 uptake remained at only 1.8 mmol g−1 up to 54 bar. The performance of uGilTH2O in the separation of 10:90 CO2/CH4 mixture was tested at 258 K, and the result showed a steep rise in the gas uptake at 15 bar. Infrared spectroscopy indicated that below 20 bar, the CO2 absorption peak (2348 cm−1) downshifted to 2331 cm−1, which suggested physical adsorption (Fig.

6c). Above 20 bar, a new CO2 peak appeared at 2341 cm−1, and the H2O absorption peaks at 3221 cm−1 and 3355 cm−1 also became stronger (Fig. 6d), which indicated the formation of CO2 hydrates. Carbon nanotubes (CNTs) and graphene have also been studied as porous carbon materials for CO2 separation. Owing to their simple, regular, and tuneable structures, CNTs have been widely used as models for studying gas diffusion and permeability.92 However, as was the case with activated carbons, CNTs exhibit a high CO2 adsorption capacity only at high pressures, owing to the lack of surface polarity or active sites.

3. ZEOLITES FOR CO2 ADSORPTIVE SEPARATION Zeolites, including natural and synthetic ones, are highly crystalline, stable, microporous aluminosilicates exhibiting moderate pore volumes and surface areas93 which have been widely used in adsorptive separation and purification. Owing to the high crystallinity, their host–guest structures can be readily analysed by diffraction techniques.94 However, single crystals of zeolites are hard to grow because the strong Si/Al-O covalent bonds suffer from poor reversibility.95 Owing to their well-defined structures, excellent stabilities, easily repeatable properties, as well as relatively low costs, zeolites are used as the benchmark for CO2 adsorption.96 Zeolites can be regarded as porous isomers of SiO2 which display low-density four-connected topologies, with some Si(IV) substituted by Al(III) non-periodically. The Si/Al ratio directly determines the amount of counter cations inside a zeolite pore and also controls many porosity parameters such as the pore volume, pore size, and hydrophilicity. The strongest adsorption sites in zeolites are generally the counter cations, namely alkaline and alkaline earth metal ions,97,98 which bind CO2 more weakly than aliphatic amines, but the bindings are usually stronger than those in the cases of many other polar active sites. On the other hand, they bind H2O more strongly than CO2. Framework topology is the main factor determining the pore structures of zeolites. A few types of zeolites, such as LTA (zeolite A) and FAU (zeolite X or Y), have been studied for CO2 adsorption by mainly focusing on the Si/Al ratio and the counter cations. A counter cation with a high charge density increases CO2 binding. For example, Long et al. used Mg2+ (Mg-A) and Ca2+ (Ca-A) to exchange the Na+ in Na-A and studied the effect on CO2 adsorption (Fig. 7a).99 At 313 K and 0.15 bar, Ca-A exhibits much higher CO2 uptake and Qst (3.7

mmol g−1 and 58 kJ mol−1), compared to those of Mg-A (1.6 mmol g−1 and 36 kJ mol−1) and Na-A (1.3 mmol g−1 and 30 kJ mol−1) (Fig. 7b). For 15:75 CO2/N2 mixture at 313 K, the IAST selectivity of Ca-A (250) is also much higher than those of Mg-A (90) and Na-A (200). Synchrotron and neutron powder diffraction (SPD and NPD, respectively) showed that CO2 molecules are located at neighbouring six-rings and at the centre of the eight-ring, with the latter revealing a higher CO2 occupancy, which indicated that the eight-ring is the preferred adsorption site (Fig. 7a). More importantly, the diffraction data showed that Ca-A contains no Na+, whereas Mg-A still contains Na+ at the eight-ring sites. Therefore, considering that the smaller Mg2+ should interact strongly with CO2, the weaker CO2 binding of Mg-A can be explained by the incomplete ion exchange. Owing to small pore apertures, the gas diffusion in zeolites is relatively slow, compared with that in other adsorbents.24,100,101 A small cation can increase not only CO2 binding but also CO2 diffusion. Recently, Webley et al. showed these effects by changing the Na+ in Na-ZSM-25 (MWF topology, Fig. 8a) to the smaller Li+.102 Na-ZSM-25 was treated with LiCl solution for different times to yield modified zeolites, which were named Li(x)-ZSM-25 (Li+/Na+ x = 0.063, 0.10, and 0.13). Note that the zeolite framework collapses for x > 0.13. The CO2 adsorption isotherms showed that both the adsorption rate and adsorption amount increase along with x, because Li+ provides a larger free volume and a higher CO2 affinity, compared with those for Na+. For example, Na-ZSM-25 reveals the lowest CO2 adsorption rate of 0.43×10−4 s−1 at 303 K and the adsorption capacity of 2.6 mmol g−1 at 353 K and 9.5 bar, whereas Li(0.13)-ZSM-25 displays the highest CO2 adsorption rate of 4.24×10−4 s−1 and an adsorption capacity of 3.4 mmol g−1 at the same conditions (Fig. 8b). Column breakthrough experiments (303 K) using 1:1 CO2/CH4 mixture at 2 bar and 7 bar were performed on Li(0.13)-ZSM-25 to obtain selectivities of 66.9 and 11.3, respectively. Molecular sieving effect is an important feature of zeolites. Larger/more counter cations can reduce the pore volume and aperture size of zeolites. For example, 3A, 4A, and 5A zeolites with gradually increasing aperture sizes exhibit the same LTA topology, but different cations (Fig. 9a). With suitable counter cations, the aperture size can allow the diffusion of CO2 and block other gases such as CH4 and N2. For example, the pore diameter of NaCs-RHO is 3.6 Å, which is between the kinetic diameters of CO2 and CH4. It reveals a moderate CO2 adsorption capacity of 3.3 mmol g−1 at 303 K and 1 bar, but negligible CH4 adsorption under the same conditions. The molecular sieving effect for CO2/CH4 was confirmed by PXRD, in which a phase change occurs in CO 2, instead of CH4,

atmosphere.103 By tuning the Na/K ratio of NaKA zeolite, the aperture size can be adjusted in the range 0.30−0.38 nm. For 17 at% K+, NaKA shows a CO2 uptake of 2.4/3.43 mmol g−1 at 0.15/0.85 bar and 298 K, and negligible N2 adsorption at the same conditions, which indicate a molecular sieving effect (Fig. 9b). Higher K+ contents lead to less CO2 adsorptions, whereas lower K+ contents result in obvious N2 adsorption.104 Tuning the Si/Al ratio can simultaneously change the amounts of counter cations. The Si/Al ratio of K-CHA can be hardly tuned below 2, because a high aluminium content induces the formation of amorphous alumina (Fig. 10a).105 Using coal fly ash as the raw material, Li et al. synthesised r1.9K-CHA (Si/Al = 1.9) and obtained CO2 uptake of 0.81/1.49/1.54 mmol g−1 at 273/303/333 K, respectively, and 1 bar (Fig. 10b).106 This abnormal trend indicates that gas diffusion is kinetically limited by the small aperture. Under the same conditions, the N2 isotherms showed negligible uptakes. Column breakthrough experiments using 50:50 CO2/N2 mixture revealed a selectivity of 688 (Fig. 10c). DFT calculations showed that lower Si/Al ratios result in higher barriers for the passage of gas through the eight-membered rings, because a high cation density increases the space hindrance. Protonated polyamines can be introduced as counter cations into zeolites. Choi et al. achieved stable chemical adsorption of CO2 by using a zeolite loaded with mono-protonated ethylenediamine (Heda+).107 The NH4-Y zeolite was calcined at 300 °C to yield H-Y, which subsequently adsorbed eda vapour to afford Heda-Y (Fig. 11a). At low pressures, Heda-Y exhibited a much steeper CO2 adsorption isotherm than the conventional zeolites NaX, ETS-10, and SGU-29. A saturation uptake of ~1.5 mmol g−1 was realised with Heda-Y below 20 mbar, which can be assigned to chemical adsorption (Fig. 11b). Column breakthrough experiments using a dry flue gas (15% CO2, 0% H2O) and a wet flue gas (15% CO2, 3% H2O) at 313 K were performed. The conventional zeolites exhibited much lower equilibrium CO2 uptakes (0.17−0.23 mmol g−1) in the wet flue gas than in the dry flue gas (1.3−2.1 mmol g−1), while Heda-Y displayed small differences (1.3 mmol g−1 and 1.9 mmol g−1) under the two conditions. Besides the Si/Al ratio and the counter cation, zeolites can be modified by other strategies. For example, aliphatic amines can also serve as neutral guests in zeolites or chemically graft on the crystal surface defects of zeolites, which is similar to the modification of porous silica. However, zeolites display much smaller pore sizes and lesser defects than porous silica with amorphous

structures. Other elements with tetrahedral coordination geometry, such as boron, phosphorus, germanium, and some transition metals, can be doped into the Si/Al frameworks of zeolites for tuning the pore size or directly used to construct new topologies.108-110

4. POROUS SILICA FOR CO2 ADSORPTIVE SEPARATION Besides nonporous polymorphs such as quartz and porous pure silica zeolites such as Silicate-I, SiO2 can also be amorphous, which are usually mesoporous. Silica gels have been widely used in many areas. Mesoporous silica is usually referred to as a relatively new type of porous material which exhibits ordered pore structures and disordered atomic locations.111 The typical porous silica, such as MCM-41 and SBA-15 with hexagonal structures and KIT-6 with cubic structure, reveal tuneable mesopores, high surface areas, and large pore volumes. Most porous silica show relatively high hydrothermal stabilities, whereas some MCM-41-type structures having thin pore walls of thicknesses ~1−2 nm display relatively poor stabilities. Pristine porous silica reveal low CO2 uptakes at low pressures, usually below 0.2 mmol g−1 at ambient conditions,112 which indicate that the silanol groups on the pore surface interact weakly with CO2 molecules. However, the silanol groups can readily condense with alkoxysilane groups, which can be used to chemically graft the active aliphatic amines. Moreover, the large pore size and pore volume of porous silica are suitable for holding aliphatic amines as guest molecules (Fig. 12).68,113 Because amine-functionalized porous silica chemically adsorbs CO2, Qst and selectivity are generally not measured/reported. Usually, only the CO2 uptake under a certain condition (25−75 °C and 0.15 bar or 1 bar) is measured by TG. Similar to the case of amine solutions, the stability of the grafted or adsorbed amines is more important, because they are prone to leaching, decomposition, and oxidation at relatively high operation temperatures. Note that, under dry high-temperature CO2-rich conditions, amines can react with CO2 to form urea derivatives, which reduces the CO2 adsorption performance of amine-modified porous silica. Through condensation of silanol and alkoxysilane, simple aliphatic amines can be post-synthetically

grafted

onto

the

pore

surface

of

porous

silica.114,115

Because

alkoxysilane-functionalised aliphatic amines usually contain one or two amine groups per molecule, the amine contents of modified porous silica are relatively low (2−4 mmol g−1), which result in

limited

CO2

adsorption

capacities

(0.4−1.8

mmol

g−1).17

By

using

[(3-trimethoxysilyl)propyl]diethylenetriamine, bearing three amine groups, as the modification reagent,116 the CO2 adsorption capacity may increase to ca. 2.6 mmol g−1. Besides PSM, the chemical grafting may also be accomplished during the synthesis of the amine-modified porous silica. For example, Russell et al. used 3-(aminopropyl)triethoxysilane and tetraethoxysilane as raw materials to perform a one-step sol-gel reaction which yielded aerogel type porous silica anchored with aliphatic amines, which showed a high CO2 uptake of 5.55 mmol g−1 at 25 °C and 0.01 bar CO2 partial pressure.117 Besides the Si-C linkages resulting from the condensation of silanol and alkoxysilane, aliphatic amines can also be covalently linked on the pore surface of porous silica by O-C bonds (Fig. 12). Jones et al. showed that ring-opening polymerisation of aziridine can covalently link polyamines on the pore surface of porous silica.118 During the reaction, the aziridine monomers not only polymerise to form hyperbranched polyalkylamines, but also react with the Si-OH groups on the surface of porous silica to form Si-O-C- linkages. With hyperbranched polyamines on the pore surface, the porous silica displays a high nitrogen content of 7.0 mmol g−1, which results in a high CO2 uptake of 3.11 mmol g−1 in 10:90 CO2/Ar saturated with water at 298 K, in the fixed bed breakthrough experiment.

When

using

the

conventional

chemical

grafting

reagents

N-[(3-trimethoxysilyl)propyl]ethylenediamine and 3-aminopropyltrimethoxysilane, the nitrogen contents were only 2.5 mmol g−1 and 1.9 mmol g−1, and the CO2 uptakes only 0.7 mmol g−1 and 0.4 mmol g−1, respectively, under the same conditions. Compared with chemical grafting, the physical impregnation method is usually more convenient and can load more active amine groups into porous silica. Therefore, physical impregnation has gained more attention in recent years. The aliphatic amine molecules adsorbed on porous silica can retain the advantages of aqueous amine solutions, but avoid its problems of corrosion and high energy consumption for regeneration. To reduce the amine evaporation/leaking at high temperatures, polyamines such as poly(ethylenimine) (PEI; molecular weight, M.W., > 400) are required. As a classic example, Song et al. used MCM-41 as a molecular basket for hosting branched PEI (M.W. 600) to take advantage of the large porosity of MCM-41 and the high CO2 binding affinity of PEI.119 The CO2 uptakes of a series of samples with PEI loadings from 0 wt% (pure MCM-41) to 100 wt% (pure PEI) were measured by TG at different temperatures in 1 atm CO2. The highest CO2

uptakes were obtained at 75 °C, because the kinetic factor or gas diffusion problem dominates the adsorption and the amine groups agglomerate at lower temperatures; CO2 adsorption is weaker at higher temperatures. The results showed that as the amine loading increased, the CO2 adsorption capacity first increased and then decreased, and the maximum was obtained at 50 wt% PEI loading, which represents a balance between PEI loading and porosity. When the PEI loading was 50 wt%, the CO2 uptake reached 2.55 mmol g−1, which was 12 times higher than that of MCM-41 (0.2 mmol g−1) and even slightly higher than that of pure PEI (2.48 mmol g−1). The CO2 adsorption efficiency of PEI in MCM-41 is much higher than that of pure PEI (Fig. 13). More importantly, the CO2 desorption efficiency (N2 at 75 °C and 1 atm) of the sample with 50 wt% PEI (99.8%) was almost the same as that of pure MCM-41 (100%) and significantly higher than that of pure PEI (56.4%). A further study showed that moisture promotes CO2 adsorption.120 The CO2 uptake of MCM-41 with 50 wt% PEI loading increased from 2.02 mmol g−1 in dry flue gas (14.9% CO2, 4.25% O2, and 80.85% N2) to 2.84 mmol g−1 in humid flue gas (13.55% CO2, 3.86% O2, 72.72% N2, and 9.87% H2O), which was explained by the further reaction of carbamate, CO2, and H2O to form bicarbonate and ammonium. As a commonly used polyamine, PEI tends to undergo oxidative degradation at the secondary amine sites in the presence of oxygen at high temperatures. Jones et al. showed that poly(propylenimine) (PPI) can be more stable than PEI with regard to oxidation degradation (Fig. 14).121 Mesoporous silica SBA-15 was impregnated with small molecular (M.W. 146−302) linear or dendritic PEI and PPI. It was proposed that, with longer alkyl spacers, the amine groups in PPI exhibit higher basicities than those of PEI, which weaken the hydrogen-bonding interaction between neighbouring amines. The amine efficiencies were calculated as 0.22−0.23 and 0.13−0.18 CO2 per “active” amine group for PPI and PEI, respectively. After 4 cycles of adsorption at 35 °C in 1 bar CO2/N2 mixed gas containing 400 ppm CO2 and desorption at 70 °C in pure N2, the CO2 adsorption capacity of PPI-based sorbents remained almost constant, but that of PEI-based sorbents decreased by approximately 36%. Infrared, NMR, and electrospray ionisation mass spectrometry showed that small amounts of PPI molecules underwent intramolecular thermal rearrangements rather than oxidation, while PEI molecules oxidised to form imines and amides. It is of particular interest to reduce the formation of urea by the amines loaded in porous silica. Choi et al. showed that PEI modified by 1,2-epoxybutane (EB) can greatly enhance the adsorbent

stability during cyclic separation and regeneration by inhibiting urea formation and oxidation (Fig. 15).122 The reaction between PEI (M.W. 1200) and EB yielded alkylated amines with 2-ethyl-hydroxyethyl groups, which can lower the basicity of the amine centres which act as electron-withdrawing groups. Meanwhile, the steric hindrance near the amine increases, which renders the carbamate species more unstable. The combination of lowered basicity and increased steric hindrance of the amine centres reduces CO2 adsorption and promotes CO2 desorption, which can be useful because the CO2 adsorption abilities of aliphatic amines are generally too strong. More importantly, functionalisation of PEI with EB suppresses the urea formation and oxidative degradation. TSA involving adsorption at 40 °C in 1 bar wet flue gas (15% CO2, 3% H2O, 2% argon, and balance N2) and desorption at 120 °C in 1 bar dry CO2 (100%) was performed. A steady CO2 separation working capacity of 2.2 mmol g−1 was observed for the sample loaded with EB-modified PEI after 50 adsorption-desorption cycles. In contrast, without EB modification, the CO 2 separation working capacity decreased from 2.9 mmol g−1 to 1.1 mmol g−1 under the same condition. The EB modification can also be applied to other polyamines.123 Besides mesoporous silica with ordered pore structures, other types of porous silica can also be functionalised by loading amines as guests.124,125 For example, loading PEI (M.W. 423) in commercial silica gel can yield a high CO2 adsorption capacity of 3.1 mmol g−1 in dry flue gas (15% CO2, 4.5% O2, and 71.5% N2) at 75 °C, which is comparable to that of SBA-15 loaded with 50 wt% PEI.126 Porous silica nanoparticles with hierarchical pores may be useful for achieving better distribution of the amines and diffusion of CO2.125 The loading procedures can also be optimised to obtain better CO2 separation performances. For example, thermally driven evaporation is generally used for the preparation of amine-loaded porous silica, but freeze-drying may produce a host–guest type adsorbent with a much higher structural homogeneity (distribution of the amines in porous silica).127 The polyamines can also be introduced during the synthesis of mesoporous silica, which greatly reduces the time, energy, and amounts of chemical reagents.128

5. POPs FOR CO2 ADSORPTIVE SEPARATION POPs are a special type of organic polymers with permanent porosity, and most of them are amorphous solids. Based on reversible covalent bonds and network topology principles, crystalline

and even single-crystalline POPs can be designed and synthesised.129-132 However, most POPs exist as very small crystallites and need to be characterised by PXRD; their crystallography has not been well adopted for studying their host–guest structures. Being organic materials, the structures of POPs can be easily designed and tuned (compared with those of conventional adsorbents) to obtain a variety of active sites for improving the CO2 adsorption/separation performances.133-135

5.1 Aliphatic amines Aliphatic amines can be post-synthetically introduced into POPs. PAF-1, an amorphous POP with dia topology and a very large surface area, was designed and synthesised by Qiu and Zhu et al. by Yamamoto–Ullman cross coupling of tetrakis(4-bromophenyl)methane.136 Several similar POPs, namely PPN-3, PPN-4, PPN-5, and PPN-6, based on tetrakis(4-bromophenyl)adamantane, tetrakis(4-bromophenyl)silane,

tetrakis(4-bromophenyl)germanium),

and

tetrakis(4-bromophenyl)methane, respectively, were later synthesised by Zhou et al. by using an optimised Yamamoto homo-coupling procedure.137 Using a two-step PSM method, Zhou et al. successfully prepared a series of PPN-6 derivatives of covalently linked aliphatic amines with different lengths and shapes, namely PPN-6-CH2EDA, PPN-6-CH2TAEA, PPN-6-CH2TETA, and PPN-6-CH2DETA (Fig. 16a).138 Elemental analysis yielded the following order of nitrogen contents: PPN-6-CH2DETA (11.95%) > PPN-6-CH2TAEA (9.31%) > PPN-6-CH2TETA (9.04%) > PPN-6-CH2EDA (7.53%), which corresponded to amine concentrations of ca. 5.3−8.5 mmol g−1. The N2 isotherms showed gradual decreases in the BET surface areas, along with the increases in the amine concentrations. The CO2 isotherms at 295 K confirmed that an increase in the nitrogen content enhances the CO2 binding and uptake (Fig. 16b). PPN-6-CH2DETA, with the highest nitrogen content, showed the highest CO2 uptakes of 3.08 mmol g−1 and 4.31 mmol g−1 at 0.15 bar and 1 bar, respectively. The IAST selectivities of PPN-6-CH2DETA and PPN-6-CH2EDA for 15:85 CO2/N2 mixture were calculated to be 442 and 115, respectively. Farha, Hupp, and Nguyen et al. directly synthesised the PAF-1 derivative appended with phthalimide by using (4-bromo-3-methylphenyl)tris(4-bromophenyl)methane as a building block, which can be transformed to the aliphatic amine derivative PAF-1–CH2–CH2NH2 and further to the imine derivative PAF-1–CH2–CH2N=CMe2.139 The nitrogen contents of PAF-1–CH2–CH2NH2 and PAF-1–CH2–CH2N=CMe2 were determined by elemental analysis to be 3.83 wt% and 3.80 wt%,

respectively, which correspond to a nitrogen concentration of ca. 2.7 mmol g−1. At 273 K and 1 bar, the CO2 uptake increases from 2.46 mmol g−1 (PAF-1) to 4.4 mmol g−1 (PAF-1-CH2NH2) and 3.13 mmol g−1 (PAF-1-CH2N=CMe2), and Qst at zero-loading increases from 18.1 kJ mol−1 (PAF-1) to 57.6 kJ mol−1 (PAF-1-CH2NH2) and 19.3 kJ mol−1 (PAF-1-CH2N=CMe2). Obviously, a primary amine binds CO2 more strongly than a ternary amine or an aromatic hydrocarbon.

5.2 Other organic functional groups Through proper design of the monomers, various kinds of polar functional groups can be introduced into POPs for enhancing the CO2 adsorption.140-142 For example, Bhaumik et al. synthesised a series of POPs which were simultaneously functionalised by aromatic amine, triazine, and phenolic hydroxyl groups, namely TrzPOP-1, TrzPOP-2, and TrzPOP-3, by Schiff-base condensation of 1,4-bis(4,6-diamino-s-triazin-2-yl)-benzene and three dialdehydes with gradually increasing numbers of phenolic hydroxyl groups (Fig. 17a).143 PXRD showed that these POPs are amorphous. The N2 isotherms revealed gradual decreases in the BET surface area from 995 m2 g−1 to 772 m2 g−1 (Fig. 17b), pore width from 1.7 nm to 1.4 nm, and pore volume from 0.93 cm2 g−1 to 0.62 cm2 g−1, along with an increase in the phenolic hydroxyl content. In contrast, the CO2 isotherms showed gradual increases in the uptake at 298 K and 1 bar from 3.51 mmol g-1 to 5.09 mmol g-1 and zero-loading Qst from 29 kJ mol−1 to 37 kJ mol−1. DFT calculations on the proposed repeating unit of TrzPOP-2 suggested strong N−H···O and O−H···O hydrogen bondings, as well as weak N···C, O···C, and C−H···O interactions. The IAST selectivities of TrzPOP-1, TrzPOP-2, and TrzPOP-3 for 15:85 CO2/N2 mixture were calculated as 108.4, 140.6, and 167.4, respectively. Fluorinated organic compounds usually exhibit high hydrophobicities, which may be used to improve the moisture resistance of the CO2 adsorption in POPs. For example, Han et al. synthesised two triazine-based 2D honeycomb-like POPs, namely CTF-1, with ordinary C−H bonds, and FCTF-1, functionalised by C−F bonds, using terephthalonitrile and tetrafluoroterephthalonitrile as the monomers, respectively (Fig. 18a).144 CTF-1 displays the characteristic PXRD peaks of a 2D honeycomb-like structure, whereas FCTF-1 is amorphous (Fig. 18b). At 298 K and 1 bar, the CO2 uptake of FCTF-1 (3.21 mmol g−1) is much higher than that of CTF-1 (1.41 mmol g−1) (Fig. 18c). In addition, the Qst value of FCTF-1 (35 kJ mol−1 ) is higher than that of CTF-1 (30 kJ mol−1) at zero-loading. Column breakthrough experiments using 10:90 CO2/N2 dry mixture yielded

selectivities of 77 and 18 for FCTF-1 and CTF-1, respectively. The CO2 uptake in the column breakthrough experiment slightly decreased from 0.73 mmol g−1 in the dry condition to 0.64 mmol g−1 in the presence of saturated H2O vapour. The stronger CO2 adsorption of FCTF-1 was attributed to the high polarity of the C−F bond, which increased the host–guest interaction, but the C−F bond is well known to exhibit weak intermolecular interactions. The much smaller pore size of the fluorinated POP should form the basis of an alternative explanation, and can be further verified by comparing the CO2 adsorption properties of the expanded versions of FCTF-1 and CTF-1.145 Although various types of organic functional groups have been introduced into POPs, their CO2 bindings are all much weaker than those of aliphatic amines.134,146 The relatively large pores and low structural regularity of POPs are the main issues preventing the cooperation of multiple weak interactions like MOFs.

5.3 Metal ions Besides organic functional groups, OMSs can also be rationally introduced into POPs. Owing to their suitable symmetries and functionalities, porphyrin, phthalocyanine, and Salen are popular building blocks of POPs. With these functional groups, metal ions can be incorporated inside the covalent backbones of POPs. In 2D POPs, these planar functional groups tend to stack via - interaction to block metal ions.140,147,148 Designing a 3D structure can be a viable strategy to reduce the - stacking for exposing the metal ions.149,150 Mastalerz et al. reported a series of Salen-based POPs named MaSOFn by Schiff-base condensation

of

2,3,6,7,12,13-hexa-(3'-carbonyl-4'-hydroxyphenyl)triptycene

with

o-phenylenediamine in the presence of different proportions (n in percentage, 50 or 100) of various metal ions (Fig. 19a).149 The polymerisation can be carried out without adding a metal salt to obtain a metal-free POP named H2-SOF. According to the non-coplanar orientations of the six salicylaldehyde groups, these POPs should exhibit 3D porous structures, in which the Salen groups would be difficult to stack. Therefore, both faces of the square-planarly coordinated metal ions of the metal-Salen plane may be accessible to CO2 molecules. TG measurements of the POPs in air yielded metal oxide contents of 17.7−35.8 wt%, which corresponded to a metal concentration of ca. 2 mmol g−1 and an OMS concentration of ca. 4 mmol g−1. PXRD showed that all these POPs are amorphous, but the N2 isotherms revealed BET surface areas of 116−816 m2 g−1 and pore volumes of 0.17−0.52 cm3 g−1. At

273 K and 1 bar, the CO2 uptakes of MaSOFn are almost twice that of H2-SOF (Fig. 19b). Ni-MaSOF100 exhibits the highest CO2 uptake of 4.83 mmol g−1, whereas that of H2-SOF is only 2.12 mmol g−1. Metal ions can also be introduced by post-synthetic ion exchange.150-152 However, POPs functionalised with metal ions generally displayed CO2 uptakes lower than those of MOFs and zeolites. This can be attributed to the relatively soft structure of the organic backbones of POPs, which can satisfy the coordination preferences of the metal ions, thus reducing their ability to coordinate with the guest molecules.

6. MOFs FOR CO2 ADSORPTIVE SEPARATION Being inorganic–organic hybrids, MOFs offer the advantages of high crystallinity or structural regularity, extremely rich structural diversity, and being easy to design (compared with other types of adsorbents), which are ideal for studying gas adsorption and separation.30-32 Various types of active sites, with CO2 bindings varying from the typical chemical bonds (such as aliphatic amines and monodentate hydroxide), semi-chemical/physical interactions (such as open metal sites) to pure physical interactions, have been successfully introduced into MOFs and their effectiveness of CO2 adsorption/separation demonstrated (Scheme 1).34,70-77,153,154 These strategies, together with representative examples, are hereafter introduced/discussed sequentially according to the order of their development through research i.e. OMS, aliphatic amine, aromatic amine (including amino and amide groups) and other polar functional groups (including aromatic N-heterocycles, acidic groups, hydroxide, etc.), ultra-micropores and cooperation of weak interactions, and flexibility of the framework for improving/tuning CO2 adsorptive separation.

6.1 OMSs Small and labile terminal ligands usually exist in the coordination networks of MOFs. When a MOF is stable enough to endure the removal of these terminal molecules, OMSs can be generated for coordination with the target guest. Because OMSs are quite common in MOFs, they represent the early strategy for enhancing CO2 adsorption, such as the Cu(II) ions in [Cu3(btc)2] (HKUST-1, H3btc = 1,3,5-benzenetricarboxylic acid).155 Because of the weak coordination ability of the axial position

of Cu(II), HKUST-1 displays a moderate CO2 uptake of 4.86 mmol g−1 at 298 K and 1 bar.156 Many other MOFs consisting of Cu2(RCOO)4 paddlewheel clusters were also studied for CO2 adsorption, and their performances are similar to or lower than that of HKUST-1.157,158 [M2(dobdc)] (H4dobdc = 2,5-dihydroxyl-1,4-benzenedicarboxylic acid) (MOF-74156 or CPO-27159) is a typical MOF which exhibits high concentrations (6.1−8.2 mmol g−1, 7.2−8.4 mmol mL−1) and diverse types of OMSs (Fig. 20a).70,159-163 Matzger et al. first studied MOF-74 for CO2 adsorption70. The zero-coverage CO2 Qst of MOF-74 lies in the range 22−47 kJ mol−1. As a hard Lewis acid with the smallest radius (0.65 Å), Mg(II) interacts with CO2 more strongly than the other metal ions which are capable of forming the MOF-74 structure. Owing to the relatively small radii of Co(II) (0.74 Å) and Ni(II) (0.72 Å), their CO2 bindings are also strong. DFT calculations showed that the hypothetical V(II)-MOF-74 or Ti(II)-MOF-74 structure could bind CO2 6−9 kJ mol−1 stronger than Mg-MOF-74, because the CO2 lone-pair orbitals overlap with the empty d-level orbitals of transition metals.164 At 298 K, the gravimetric CO2 uptakes of Mg-MOF-74 (4.95 mmol g−1 at 0.15 bar and 8.61 mmol g−1 at 1.0 bar) are significantly higher than those of its analogues based on other metal ions. Regarding the volumetric uptake, Co-MOF-74 (8.20 mmol cm−3 at 1.0 bar) displays a slightly higher value than Mg-MOF-74 (7.92 mmol cm−3 at 1.0 bar). Therefore, the high gravimetric CO2 uptakes of Mg-MOF-74 are a benefit of the low atomic mass of magnesium, to some extent. The IAST selectivities of MOF-74 for 15:75 CO2/N2 mixture were reported to be ca. 57−148.43,49,53 At 298 K and 0.4 mbar, Mg-MOF-74 also shows the highest CO2 adsorption among the MOF-74 series (Fig. 20b), but the uptake is only 0.088 mmol g−1, which suggests that the Qst is not high enough.49 The poor water stability of the MOF-74 structure can be a problem in CO2 separation. For example, PXRD showed that Mg-MOF-74 collapsed after water sorption measurement.165 At 70% RH, the CO2 adsorption amounts at 0.15 bar for Mg-MOF-74 and Zn-MOF-74 are only 16% and 22%, respectively, of those in the dry condition. The more stable structures Ni-MOF-74 and Co-MOF-74 can retain 61% and 85%, respectively.166 MOFs consisting of high-valence (+3 or +4) metal ions usually display high stability, and trivalent OMSs can be found in many MOFs consisting of M3(3-O/OH)(RCOO)6 clusters. Benefiting from the higher charge density, trivalent OMSs usually show stronger CO2 bindings than divalent OMSs. For example, high CO2 Qst have been reported for MOFs consisting of Cr(III) (62 kJ

mol−1), In(III) (67.8 kJ mol−1), and V(III) (79.6 kJ mol−1).167 Generally, each CO2 molecule coordinates with one OMS in an angular orientation,40 and each OMS only binds one CO2. In some special cases, each CO2 can coordinate with two OMSs and each OMS can coordinate with two CO2 molecules. Zhou et al. first proposed bidentate CO2 in a designed MOF [Cu2(ndpa)] (PCN-88; H4ndpa = 5,5'-(naphthalene-2,7-diyl)diisophthalic acid).168 In PCN-88, two OMSs of two adjacent Cu2(RCOO)4 clusters orient in a face-to-face manner with a separation of 7.4 Å, which is suitable for sandwiching a CO2 molecule (length ca. 2.3 Å) with two weak Cu−O coordination bonds (length 2.5 Å, normal to the z-axis of the square-pyramid coordinated Cu(II); Fig. 21a). GCMC simulation was used to visualise the expected host–guest binding structure. The CO2 Qst was calculated to be 27 kJ mol−1. Zhang et al. reported interesting metal–CO2 interactions in [Cu2(OH)2(bdim)] (MAF-35; H2bdim = 1,5-dihydrobenzo[1,2-d:4,5-d']diimidazole).169 MAF-35 is constructed based on unique four-connected planar Cu2(-OH)2(bdim)4 clusters which expose both sides of square-planarly coordinated Cu(II) ions on the pore surface to yield an ultrahigh concentration of OMSs (17.2 mmol cm–3 or 12.6 mmol g–1). Moreover, the distance between two adjacent copper sites is 7.4 Å and the dihedral angle of adjacent planar Cu2(-OH)2 clusters is 57°, which is more suitable than the coplanar configuration for chelating a CO2 molecule. MAF-35 shows high CO2 Qst (47 kJ mol−1) and capacity (4.46 mmol g−1 or 6.05 mmol cm−3 at 298 K and 1 bar), as well as rapid adsorption (3 min for 90% equilibrium uptake) and desorption (5 min for 90% equilibrium uptake) dynamics. GCMC simulation revealed that MAF-35 can develop multiple interactions with CO2 molecules and each copper can expose two sides to bind two CO2 molecules, and two adjacent copper can chelate one CO2 molecule (Fig. 21b). Chen et al. reported a supramolecular isomer of MOF-74, namely UTSA-74, which exhibits a similar honeycomb-type 3D framework and a 1D channel.38 Interestingly, in UTSA-74, one Zn(II) adopts tetrahedral coordination geometry and another adopts octahedral coordination geometry, with two opposite sites occupied by removable guest molecules. Further, the two OMSs with two adjacent Zn(II) ions display a separation of 6.3 Å and a dihedral angle of 71°. SCXRD showed that, at low CO2 loadings, one CO2 can coordinate with two adjacent zinc ions, with the Zn-O distance being 2.2 Å (Fig. 21c). The CO2 Qst of UTSA was calculated as 25 kJ mol−1, which is slightly lower than that of Zn-MOF-74 (27 kJ mol−1).40

OMSs can also originate from the counter cations,170,171 although they always attach onto the pore surface after the guest removal. It can be seen that, compared with other types of adsorbents, MOFs can accommodate more types of metal ions with various coordination environments for CO2 adsorption. Although MOFs are notable for their flexibility, some of them like MOF-74 are rigid enough to retain the OMS structures. Consequently, stronger CO2 bindings and higher CO2 uptakes can be achieved at low pressures with OMS-functionalised MOFs.

6.2 Aliphatic amines Self-assembly of organic ligands containing aliphatic amine groups with metal ions generally results in metal–amine coordination. Only a handful of MOFs with free aliphatic amine groups have been directly synthesised.172-174 PSM is a rational strategy for introducing free aliphatic amines on the pore surfaces of MOFs.71,153,175-178 Because covalent PSM is quite complicated and achieving high concentrations of amine groups is difficult,153 most studies focused on coordinative PSM. Many types of coordinative PSM have been developed to introduce free aliphatic amines on the pore surfaces of MOFs. By using the coordination ability of OMSs, diamines can be conveniently immobilised by using the metal–amine coordination bond. In 2009, Long et al. post-synthetically introduced

eda

onto

the

pore

of

H3[(Cu4Cl)3(BTTri)8]

(CuBTTri,

H3BTTri

=

1,3,5-tris(1H-1,2,3-triazol-5-yl)benzene) containing Cu(II) OMSs, and observed a high CO2 Qst of 90 kJ mol–1, which is characteristic of chemical adsorption.71 Subsequently, some other diamines were also successfully grafted on CuBTTri24,178 to realise strong CO2 bindings of the order of 96 kJ mol–1. The expanded versions of Mg-MOF-74 have been extensively used for grafting diamines.41,42,177,179-183 Note that the pore size (11 Å) of Mg-MOF-74 is too small for stoichiometric grafting of diamines, and 1/6th of the metal ions can be used experimentally.184 Hong and Long et al. grafted

N,N'-dimethylethylenediamine

(mmen)

4,4'-dihydroxy-(1,1'-biphenyl)-3,3'-dicarboxylic

acid,

in

[Mg2(dobpdc)]

(H4dobpdc

pore

diameter

Å)

18.4

to

=

obtain

[Mg2(dobpdc)(mmen)1.6].174 At 298 K, it reveals CO2 uptakes of 3.13 mmol g−1 and 3.86 mmol g−1 at 0.15 bar and 1 bar (Fig. 22), respectively, which are lower than those of [Mg2(dobpdc)] (4.85 mmol g−1 and 6.42 mmol g−1 at 0.15 bar and 1 bar, respectively) because of the reduced pore volume and increased sample weight. However, [Mg2(dobpdc)(mmen)1.6] displays a CO2 uptake of 2.0 mmol g−1 at 0.39 mbar and 298 K, which is 15 times higher than that of [Mg2(dobpdc)]. Infrared spectroscopy

showed that the N−H groups in mmen indeed disappeared after adsorbing CO2, but the formation of ammonium carbamate was questioned. Abnormally, [Mg2(dobpdc)(mmen)1.6] exhibited gate-opening type CO2 isotherms, which indicated that the amines, as the strongest adsorption sites, are initially inaccessible. This isotherm shape can significantly lower the CO2 uptake at a slightly lower temperature and/or higher pressure, which can reduce the energy consumption during adsorbent regeneration. Subsequently, Long et al. used SPD to reveal the unique CO2 adsorption mechanism.30 Various analogues of [Mg2(dobpdc)(mmen)1.6] consisting of other metal ions (M = Mn, Fe, Co, or Zn) also showed the gate-opening type isotherms, and the manganese analogue with a high crystallinity was used. After the adsorption of CO2, a carbamate group forms, but it coordinates with manganese (Mn−O 2.10(2) Å), rather than remaining exposed in the pores (Fig. 22). The amine group exposed in the pore interacts with the uncoordinated carbamate oxygen (N···O 2.61(2) Å). The gate-opening pressure follows the order Mg < Mn < Fe < Zn < Co, which is in good agreement with the stabilities of the octahedral metal complexes.185 The nickel analogue of [Mg2(dobpdc)(mmen)1.6] showed simple type-I weak CO2 adsorption, which was attributed to the exceptionally high stability of the Ni−N bond. Long et al. further found that the substituent size/number of diamines can be used to effectively tune the isotherm gating pressure or CO2 desorption temperature.180,186 Specifically, the isobaric step temperatures follow the order of secondary–secondary > primary–secondary > primary–ternary amines, because the amine groups with smaller steric hindrances bind the OMSs stronger, and the free amine groups with smaller steric hindrances form stronger ammonium carbamate hydrogen bonds or ion-pair interactions. SCXRD was successfully carried out with the zinc framework, which confirmed the correlations between the M−N, M−O, and N···O distances and the steric hindrances and isobaric step temperatures. By balancing the adsorption–desorption working capacity and the regeneration temperature, [Mg2(dobpdc)] loaded with N,N-dimethylethylenediamine (dmen) was evaluated to be the best (13.3 wt% or 3.0 mmol g–1 working capacity from 40 °C and 0.15 bar to 75 °C under argon) for capturing CO2 from coal flue gas.42 For DAC, Hong et al. showed an ultrahigh CO2 uptake of 2.85 mmol g−1 or 2.72 mmol cm−3 at 298 K and 0.4 mbar by using [Mg2(dobpdc)(eda)1.6].41 Besides the low molecular weight, the small size of eda, which reduces the steric hindrance and weakens the hydrogen-bonding interactions

between adjacent amines, may be responsible for the high CO2 uptakes of [Mg2(dobpdc)(eda)1.6], compared with those of [Mg2(dobpdc)] grafted with other diamines. However, the CO2 uptake of [Mg2(dobpdc)(eda)1.6] at 0.4 mbar is much lower than the saturation uptake of 4.6 mmol g−1, which indicates that the intermolecular hydrogen bonding is quite strong.30,187 Upon slightly increasing the temperature to 323 K, the CO2 uptake at 0.4 mbar decreases to 0.12 mmol g−1. Zhang et al. showed that hydrazine can serve as the shortest diamine which can be used for modifying pristine Mg-MOF-74 to afford [Mg2(dobdc)(N2H4)1.8].76 MM and PDFT calculations showed that the shortest N···N separation is 3.81 Å, which is much longer than that (3.06 Å) in [Mg2(dobpdc)(eda)1.6] (Fig. 23). By virtue of the small framework volume and low molecular weight, as well as the weak hydrogen bonding between adjacent amine groups, [Mg2(dobdc)(N2H4)1.8] exhibits a high concentration of free amine groups (6.01 mmol g −1 or 7.08 mmol cm−3), which results in an ultrahigh CO2 uptake of 3.89 mmol g−1 or 4.58 mmol cm−3 at 0.4 mbar and 298 K, which changes to 1.04 mmol g−1 or 1.22 mmol cm−3 at 328 K. From the CO2 adsorption–desorption capacity and Qst, as well as the experimental specific heat, the regeneration energy of [Mg2(dobdc)(N2H4)1.8] was calculated as 3.02 MJ per kilogram of CO2, which is higher than those of [Mg2(dobpdc)(mmen)1.6] (2.34 MJ kg−1) and the state-of-the-art amine-based solutions (2.6 MJ kg−1), but lower than that of MEA (3.5 MJ kg−1). Coordination grafting of neutral diamines has the risk of the diamine escaping at high temperatures and/or humid conditions.179 By using charge-assisted coordination bonds or anionic amines, this problem may be avoided. Zhang et al. introduced an anionic aliphatic amine, deprotonated

ethanolamine

(Hea),

in

[Zr6O4(OH)4(bdc)6]

(UiO-66,

H2bdc

=

benzene-1,4-dicarboxylic acid).188 Because Hea cannot directly react with UiO-66, UiO-66 was first heated to dehydrate the Zr6O4(OH)4(RCOO)12 clusters to form Zr6O6(RCOO)6 clusters, which were then reacted with Hea to yield [Zr6O4(OH)2(ea)2(bdc)6] (UiO-66-EA) (Fig. 24). Owing to the exposed aliphatic amine groups, UiO-66-EA displayed a high CO2 Qst of 66 kJ mol−1. The NMR spectra indicated carbamate species were generated after the CO2 adsorption. The column breakthrough curves were obtained by using 10:90 CO2/N2 mixture at 313 K and 1 bar. The CO2 adsorption capacity of UiO-66-EA was found to be 18 times higher than that of UiO-66, and the value was retained at 82% RH.

6.3 Aromatic amines and other polar functional groups Aromatic amines. Because of their similarities with aliphatic amines, aromatic amines have attracted special interest. However, the electron lone pair of an aromatic amine conjugates with the aromatic ring and exhibits a low basicity/reactivity. Compared with isostructural MOFs without an aromatic amine group, these MOFs generally show improved CO2 bindings and higher CO2 uptakes at low pressures.154,156 However, hydrogen bonding, dipole–dipole interactions, and even simply the pore size effect may be more important for such improvements. In 2009, Shimizu et al. reported that ultra-microporous [Zn2(ox)(atz)2] (H2ox = oxalic acid, Hatz = 3-amino-1,2,4-triazole) can reveal a CO2 uptake of 3.8 mmol g−1 at 293 K and 1 bar, and the uptake exceeds 3 mmol g−1 at 0.15 bar.189 The Qst is 40 kJ mol−1 at near zero coverage, which decreases to 35 kJ mol−1 at 1 mmol g−1, but rises again to 38.6 kJ mol−1 at 3.8 mmol g−1, which was explained in the form of strong host–guest interactions at low loadings and notable guest–guest interactions at high loadings. Subsequently, they successfully performed SCXRD for [Zn2(ox)(atz)] loaded with CO2 at 123–293 K.34 There are two independent CO2 adsorption sites, defined as the first (CO2-I) and second (CO2-II), based on the different occupancies (Fig. 25). DFT showed that their binding energies are close (39.6 kJ mol−1 and 38.1 kJ mol−1 for CO2-I and CO2-II, respectively). For CO2-I, the carbon atom of CO2 forms a contact with the nitrogen atom of the NH2 group (3.152(6) Å), the length of which is shorter than the sum of the corresponding van der Waals radii (3.25 Å). The carbon atom also interacts with an oxalate oxygen atom (C···O 3.155(8) Å). Further, one oxygen atom of CO2 interacts with one NH2 group to form two weak hydrogen bonds (N···O 3.039(4)−3.226(9) Å). CO2-II is located between two oxalate groups, and the oxygen atom of CO2 interacts with the carbon atom of ox2− (O···C 2.961(5) Å) and with an NH2 group with a strong hydrogen bond (N···O 2.783(8) Å). Obviously, the aromatic amine does not dominate the CO2 binding, and its electron lone pair makes a limited contribution. The two CO2 sites are in close contact (C···O 3.023(7) Å), which is used to explain the increase in Qst at high loadings. As a derivative of amines, amides have also been used to improve CO2 adsorption.190 Although the electron lone pair of the nitrogen atom in amides permits stronger -conjugations and weaker CO2 bindings than that of simple amines, the carbonyl group of amides may provide stronger dipole– dipole interactions and smaller pore sizes. Considering that most amide-functionalised MOFs also contain open Cu(II) sites with stronger CO2 affinities, Schröder et al. used an amide-functionalised

pyrimidyl-isophthalate

ligand

to

synthesise

[Cu2(piaip)4]

(MFM-136,

H2piaip

=

5-[4-(pyrimidin-5-yl)benzamido]-isophthalic acid) without any open Cu(II) sites, which is suitable for studying the role of amide groups in CO2 adsorption. The results showed that MFM-136 exhibits a CO2 Qst of 25.6 kJ mol−1 and a CO2/N2 selectivity of 27, which are similar to the values obtained for other MOFs functionalised with amide groups. However, NPD revealed that CO2 molecules interact with the phenyl rings and pyrimidyl rings (O···ring centroid 2.99(6)−3.22(9) Å), and none of the CO2 molecules develop an apparent interaction with the amide moiety. Therefore, the CO2 adsorption mechanism was attributed to a combination of pore size, geometry, and the functional group, rather than to the amide group.191 Aromatic N-heterocycles. Azole/azolate and pyridine can provide active nitrogen atoms with free electron lone pairs72,192-196 and also occupy less space than not only aromatic amines but also ordinary aromatic rings.197 By comparing a pair of isostructural MOFs bearing triazolate nitrogen atoms with different degrees of exposure and different pore sizes, Zhang et al. demonstrated that an uncoordinated aromatic nitrogen atom which is even partially exposed on the pore surface can provide a stronger binding than that with a smaller pore size for enhancing CO2 adsorption.72 As a result of the high activity of the electron lone pairs, aromatic nitrogen atoms also display strong coordination ability. The strategies for exposing aliphatic amines on the pore surfaces by direct MOF synthesis can also be applied to aromatic nitrogen atoms.197 As revealed by their relatively weak basicity, aromatic N-heterocycles bind CO2 weaker than aliphatic amines, which results in relatively low CO2 uptakes at low pressures and Qst values.72,192-194,198 However, when aromatic nitrogen atoms are arranged at appropriate positions, each CO2 may interact with more than one N-heterocycle to achieve much stronger adsorption. Zhang

et

al.

reported

a

flexible

porous

MAF

[Zn2(btm)2]

(MAF-23,

H2btm

=

bis(5-methyl-1H-1,2,4-triazol-3-yl)methane) bearing pairs of triazolate nitrogen atoms as CO2 chelating sites.33 In MAF-23, Zn(II) ions adopt the tetrahedral coordination mode, and all triazolate rings adopt the imidazolate coordination mode, leaving an uncoordinated nitrogen atom for each triazolate ring. Interestingly, two adjacent uncoordinated nitrogen atoms from different triazolate rings form a claw-like guest binding site which results in a high CO2 Qst of 43.6 kJ mol–1 and a CO2/N2 selectivity (0.15 atm CO2, 0.75 atm N2) of 87 at 298 K. SCXRD analysis of the host–guest structures at different CO2 loadings showed that the two nitrogen atoms indeed chelate one CO2

molecule in a dynamic manner (C···N 2.90(4)–3.26(1) Å) like a claw (Fig. 26). Other polar groups. Besides nitrogen-based active sites, many other polar groups, including acidic ones, have been used to improve the CO2 binding.199-201 For example, Serre, Weireld, and Maurin et al. constructed a UiO-66 derivative, namely UiO-66-(COOH)2, with two additional free carboxylic groups per organic ligand.202 Although the pore volume decreased from 0.45 cm3 g−1 to 0.21 cm3 g−1, the experimental CO2/N2 adsorption selectivity increased from 25 to 56. GCMC simulations showed that the CO2 molecules were primarily distributed in the regions close to the

3-OH and −COOH groups. It should be noted that UiO-66 is stable in both water and acid.203 Metal hydroxides, both soluble and insoluble, are well known for their extremely high affinities toward CO2 owing to the acid–base reactions (the heat of formation of HCO3− from OH− and CO2 in the gas phase is ca. 205 kJ mol−1); they are useful for removing trace CO2, but difficult to regenerate. Many MOFs consist of hydroxide ligands, which generally coordinate with two or three metal ions and exhibit negligible basicity and low CO2 affinity. For example, Schröder et al. compared the CO2 adsorption behaviours of [VIII2(OH)2(L)] (MFM-300(VIII), H4L = biphenyl-3,30,5,50-tetracarboxylic acid) and its oxidised isostructural analogue [VIV2 O2(L)] or MFM-300(VIV).204 At 298 K and 1 bar, MFM-300(VIII) showed a CO2 uptake (6.0 mmol g−1) which is much higher than that of MFM-300(VIV) (3.5 mmol g−1). NPD revealed that the CO2 molecules are located very differently (Fig. 27). In MFM-300(VIII), the -OH ligand forms strong hydrogen bonds with the O atoms of CO2 (H···O 1.863(1) Å). In MFM-300(VIV), the CO2 molecule is sandwiched between two phenyl groups (O···centres of the phenyl rings 3.069(2) and 3.146(3) Å), and the carbon atom of CO2 is far away from the -O ligand (C···O 4.86 Å). In a handful of MOFs with large separation distances of the metal ions, hydroxides can serve as monodentate ligands. Zhang et al. proposed monodentate hydroxide ligands for strong and reversible CO2 adsorption by considering the moderate basicity and structural similarity with the key active sites of carbonic anhydrase.74 They designed and synthesised [MII2Cl2(bbta)] (H2bbta = 1H,5H-benzo(1,2-d:4,5-d')bistriazole, M = Mn, Fe, Co, Ni, or Cu), which is isoreticular with MOF-74, but exhibits a much higher stability owing to the azolate coordination effect.205 [MnII2Cl2(bbta)] (MAF-X25) and [CoII2Cl2(bbta)] (MAF-X27) can be oxidised by H2O2 to form [MnIIMnIIICl2(bbta)(OH)] (MAF-X25ox) and [CoIICoIIICl2(bbta)(OH)] (MAF-X27ox), respectively, in which the counter anion hydroxides serve as monodentate ligands on the pore surface (Fig. 28).

MAF-X25ox and MAF-X27ox reveal ultrahigh zero-coverage Qst of 120 kJ mol−1 and 124 kJ mol−1, respectively. Infrared spectroscopy suggested the reversible formation of HCO3−. At 298 K, the CO2 uptakes of MAF-X25ox and MAF-X27ox reach 4.1/4.1 mmol g−1 or 5.0/5.5 mmol cm−3 at 0.15 bar, and 7.1/6.7 mmol g−1 or 8.7/9.1 mmol cm−3 at 1 bar, respectively. Adsorption/desorption cycling experiments on MAF-X27ox between 15:85 CO2/N2 mixture at 40 °C and pure N2 at 85 °C showed that 12.2 wt% (2.8 mmol g-1) working capacity can be achieved within 13/19 min during the adsorption/desorption process. More importantly, the breakthrough curves obtained by using dry and humid (82% RH) 10:90 CO2/N2 mixtures are almost the same, since OH− binds CO2 more strongly than it binds H2O. Recently, Wade et al. used HCO3− (0.1 M, pH 8.0−8.3) to exchange the OAc− ligands of [Zn5(OAc)4(bibta)3] (CFA-1, H2bibta = 1H,1'H-5,5'-bibenzo[d][1,2,3]triazole), which was then heated at 100 °C under vacuum to yield [Zn5(OH)4(bibta)3] functionalised with monodentate hydroxide ligands.206 This reaction is reversible and can be used to capture the CO2 in air. At 27 °C and a CO2 pressure of 0.4 mbar, [Zn5(OH)4(bibta)3] exhibits a high CO2 uptake of 2.20 mmol g−1. Adsorption/desorption cycling experiments on [Zn5(OH)4(bibta)3] under simulated air conditions (0.4 mbar CO2) between 27 °C and 100 °C showed a working capacity of 5.8 wt% (1.3 mmol g-1).

6.4 Ultra-micropores and cooperation of weak interactions While most functional groups bind CO2 more weakly than aliphatic amines, the cooperation of multiple binding sites may afford the desired adsorption strength. It is well known that smaller pores interact more strongly with the guests, but only MOFs have demonstrated the extraordinary effectiveness of this feature. Besides the examples discussed above which show clear crystal structures,33,34 many other ultra-microporous MOFs can reveal similar effects.35,207,208 For example, Zhou et al. observed strong CO2 adsorption in [Cu(tzc)(dpp)0.5] (PCN-200, H2tzc = tetrazole-5-carboxylic acid, dpp = 1,3-di(4-pyridyl)propane).35 PCN-200 contains small cavities with the diameter of 4.4 Å which are lined with C−H moieties and carboxylate oxygen atoms. A high zero-coverage CO2 Qst of 49 kJ mol−1 and an IAST CO2/N2 selectivity of 205 were calculated from the isotherms. SPD revealed framework distortion after CO2 adsorption and the existence of CO2 in the cavity, but the guest geometry and position could not be refined. GCMC and DFT simulations showed that there is one CO2 molecule in each cavity which forms multiple relatively long contacts

with the host framework (Fig. 29). [M(bpy)2(SiF6)] (bpy = 4,4′-bipyridine) is a pcu type network containing cationic M(bpy)2 square grids (sql topology) and inorganic SiF62– pillars.208,209 All the three components can be changed to finely adjust the pore size, thus yielding very different CO2 adsorption behaviours (Fig. 30).73,210-212 Zaworotko et al. first showed the usefulness of the [M(bpy)2(SiF6)] structure for systematic tuning of the pore size and CO2 adsorption.183 Subsequently, Eddaoudi and Zaworotko et al. used three isoreticular analogues, [Cu(dpa)2(SiF6)] (SIFSIX-2-Cu, dpa = 4,4'-dipyridylacetylene), [Cu(dpa)2(SiF6)] (SIFSIX-2-Cu-i, two-fold interpenetrated), and [Zn(pyr)2(SiF6)] (SIFSIX-3-Zn, pyr = pyrazine), with the pore sizes of 13.05 Å, 5.15 Å, and 3.84 Å, respectively, to further demonstrate the critical role of pore size in CO2 adsorption.207 The three materials with gradually decreasing pore sizes show gradually increasing CO2 Qst, from 22 kJ mol–1 to 45 kJ mol–1, IAST CO2/CH4 (50:50) selectivity, from 5.3 to 231, IAST CO2/N2 (10:90) selectivity, from 13.7 to 1818, and CO2 uptake, from 0.23 mmol g–1 to 2.4 mmol g–1, at 0.1 bar and 298 K. Column breakthrough experiments on SIFSIX-2-Cu-i using 50:50 CO2/CH4 and 10:90 CO2/N2 mixtures revealed CO2 uptakes which were consistent with those obtained from adsorption isotherms. PXRD showed that SIFSIX-2-Cu-i can retain its structure at 95% RH, but the CO2 uptake decreased 16% at 74% RH because of competitive adsorption of H2O. On the other hand, SIFSIX-3-Zn transformed to another unidentified structure above 35% RH. Eddaoudi et al. showed that [Cu(pyr)2(SiF6)] (SIFSIX-3-Cu),211 the Cu(II) analogue of SIFSIX-3-Zn, has a smaller pore size of 3.5 Å and a higher CO2 Qst of 54 kJ mol–1. SIFSIX-3-Cu showed a CO2 uptake of 1.24 mmol g -1 at 0.4 mbar and 298 K, which is 10 times that of SIFSIX-3-Zn (0.13 mmol g–1). Further, column breakthrough experiments using 1:999 CO2/N2 mixture at 298 K yielded a selectivity of 10500 for SIFSIX-3-Cu and 7259 for SIFSIX-3-Zn. The researchers further used Ni(II) and NbOF52– to design and construct a similar structure, [Ni(NbOF5)(pyr)2] (NbOFFIVE-1-Ni), with an even smaller pore size of 3.2 Å, which revealed an ultrahigh CO2 uptake of 2.3 mmol g–1 at 0.4 mbar and 298 K213. Zaworotko et al. showed that [Ni(pyr)2(SiF6)] (SIFSIX-3-Ni) can achieve a high IAST CO2/CO selectivity of >4000 for 1:99 CO2/CO at 298 K and 1 bar, which is much higher than those of zeolite 13X (99) and activated carbon (115). Column breakthrough experiments using 50:50 CO2/CO mixture showed a CO2 uptake which is consistent with that obtained from the single-component adsorption isotherm, and a CO

outlet purity >99.99% was obtained.212 Another function of ultra-micropores is the molecular sieving effect. For example, Zhang and Zaworotko et al. reported that [Cu(qc)2] (Hqc = quinoline-5-carboxylic acid) can crystallise as two isomers with dia and sql topologies.75 The sql isomer reveals an ultra-small pore aperture size of 3.3 Å, whereas the dia isomer displays a much larger aperture size of 4.8 Å. Gas sorption measurements at 293 K up to 1 bar showed that the dia isomer can adsorb CO2 (1.75 mmol g–1), CH4 (1.10 mmol g– 1

), and N2 (0.26 mmol g–1), and the uptakes are consistent with the order of the boiling points of the

gases. In contrast, the sql isomer shows a noteworthy CO2 uptake (2.16 mmol g–1) and negligible uptakes of CH4 (0.058 mmol g–1) and N2 (0.013 mmol g–1), which can be attributed to adsorption on the particle surface. Comparing their adsorption behaviours, the latter is likely to display molecular sieving for CO2/N2 and CO2/CH4. The IAST selectivities calculated for 15:85 CO2/N2 and 50:50 CO2/CH4 mixtures are 40000 and 3300, respectively, for the sql isomer. Ultra-micropores of MOFs can also realise some special properties. For instance, the directions of the quadrupole moments of CO2 and C2H2 are opposite, but the gas molecules can rotate freely to select the best orientation which fits the electrostatic fields on the pore surface, which results in a relatively low adsorption selectivity.32,214 However, when the pore is small enough, it can recognise the opposite directions of the gas quadrupole moments and allow only one to enter, which can result in a very high adsorption selectivity.215

6.5 Framework flexibility Flexibility is one of the most remarkable features of MOFs, which are among the known types of porous materials. Guest-induced structural transformations always occur differently for different adsorbates, and CO2 usually exhibits the lowest threshold pressure for structural transformation, compared with those of H2, N2, CO, O2, and CH4, which display much lower boiling points,216,217 and even that of C2H2, which reveals a similar boiling point and stronger adsorption.218,219 The gate-opening type structural transformation is attractive for realising molecular sieving like, ultrahigh CO2 selectivity,77,220-222 although co-adsorption of other guests after the gate-opening should be a concern.71 Long

et

al.

showed

that

the

gate-opening

behaviour

of

[Co(bdp)]

(H2bdp

=

1,4-benzenedipyrazole) can be used to achieve molecular sieving of CO2/CH4.77 The

single-component gas adsorption isotherms at 298 K indicated that the gate-opening occurs at ca. 2 bar for CO2 and ca. 18 bar for CH4 (Fig. 31). The gas adsorption isotherms of mixtures showed that below 25.3 bar (CO2 10.9 bar, CH4 14.4 bar), [Co(bdp)] exhibited a CO2 uptake of 11.4 mmol g–1 and a negligible CH4 uptake, which suggested that the gate-opening caused by CO2 did not result in co-adsorption of CH4. Under CH4-rich atmosphere e.g. when the equilibrium pressures of CH4 and CO2 were 54.9 bar and 3.7 bar, respectively, the CO2/CH4 adsorption selectivity was determined as 61. SPD confirmed that gate-opening in the mixed-gas atmosphere occurs, similar to that in the single-component gases. The existence of a gate-opening pressure for the strongly adsorbed guest can lead to a high residue concentration of the guest. For example, Kitagawa et al. reported the CO2/CH4 mixture separation behaviour of [Zn(5NO2-ip)(bpy)]n (CID-5, 5NO2-H2ip = 5-nitroisophathalic acid), which showed gate-opening adsorption of CO2 at 0.1 MPa and negligible adsorption of CH4 up to 1 MPa (Fig. 32). Column breakthrough experiments using 40:60 CO2/CH4 mixture at 0.8 MPa revealed that the CO2 concentration cannot decrease below 10% when CH4 outflows.223 Therefore, decreasing and even eliminating the gate-opening pressure for a strongly adsorbed guest such as CO2 should be important for fully utilising the exceptional property of flexible MOFs. The apparent adsorption enthalpy relies on not only the host–guest binding energy but also the energy consumptions of the structural transformations involved.224 Structural transformation of a MOF during adsorption can reduce Qst, which benefits CO2 desorption. By using a MOF with a unique flexibility, Zhang et al. showed that the very high CO2 uptake observed at ambient conditions can be correlated with a low Qst. [Zn(atz)2] (MAF-66) exhibits an exceptional dia topology with the highest possible crystal symmetry (Fd 3 m), in which the triazolate ligands serve as bent imidazolate linkers (Fig. 33).225 PXRD of MAF-66 in different guest inclusion forms showed reversible structural transformations involving complicated structural disorder, as indicated by the abnormal symmetry. MAF-66 reveals a high concentration of uncoordinated amino groups and azolate nitrogen donors (ca. 17 mmol g–1), but these form N–H···N hydrogen bonds, which limit them as active CO2 binding sites. Interestingly, MAF-66 can adsorb large amounts of CO2 at ambient conditions (6.26/4.41 mmol g–1 at 273/298 K and 1 bar), but its CO2 Qst is only 26 kJ mol–1 at zero loading and abnormally increases with the increase in CO2 loading. PXRD and infrared spectroscopy showed that, when the

CO2 pressure increases, the host gradually changes its structure and the host–guest interactions also gradually increase, which are in contrast with the conventional trends (wherein the host structure abruptly changes and the host–guest interactions gradually decrease). Therefore, the low Qst was attributed to the lack of active CO2 binding sites and to the energy consumption of the structural transformation of the host, and the gradual increase in the Qst was attributed to the gradual breaking of hydrogen bonds, which releases the binding sites (Fig. 33). Light can also act as a driving force which promotes desorption mainly by inducing structural transformation of the photoactive moieties of MOFs. For example, Zhou et al. observed the photo-modulated CO2 sorption behaviours of a MOF (PCN-123) functionalised by azobenzene side groups (Fig. 34).226 The MOF exhibited CO2 uptakes of 1.02 mmol g–1, 0.75 mmol g–1, and 0.47 mmol g–1 before UV irradiation, immediately after UV irradiation, and after five-hour UV irradiation, respectively, which were ascribed to the trans to cis conformation change in azobenzene after the UV irradiation, which significantly decreased the pore size. Lyndon and Hill et al. observed 64% instant release of the CO2 adsorbed in a photo-responsive MOF constructed with –N=N– and –C=C– bridged organic ligands upon exposure to broadband UV light.227 By virtue of the ring-opening/closing reactions of dithiophenylethene groups being triggered by UV/visible light, which can slightly change the bridging length and angle of the corresponding ligand, Guo et al. realised instant release and adsorption of CO2 by changing the irradiation between UV light (300 nm) and visible light (600 nm).228 Some crystalline porous materials without extended metal–ligand connectivities can behave like MOFs. For example, molecules/ions such as organic molecules and metal complexes can form porous structures which are sustained by hydrogen bonds, ionic bonds, and many other types of supramolecular

interactions.229-231

Some

molecules or

supermolecules can also

contain

intramolecular cavities for adsorbing CO2.232,233 Further, some nonporous molecular crystals may transform their structures to adsorb CO2.234 Sustained by intermolecular interactions which are weaker and less directional than coordination bonds, these molecular/ionic crystalline porous materials are usually more difficult to design, but they may be recycled by recrystallisation.

7. SUMMARY AND OUTLOOK As discussed in the above sections, various types of CO2 adsorption sites have been developed,

but aliphatic amines may still be the most promising materials for achieving satisfactory CO2 binding and selectivity at low pressures. Chemically linked/grafted amines on porous materials may avoid the leaking and corrosion problem of aqueous solutions. Without the need for large amounts of water, the adsorbent regeneration energy can also be reduced, but there are greater risks of decomposition, oxidation, and urea formation, which have been mainly studied for amine-modified porous silica, rather than for other types of porous materials. Owing to the richest structural diversity and finest tuneability among all types of porous materials, many unique strategies and extraordinary CO2 separation performances have been revealed with MOFs. For example, CO2 adsorption by hydrazine, monodentate hydroxide, or in ultra-micropores can be as strong as that by aliphatic amines. For the same adsorption enthalpy, cooperation of multiple weak interactions should result in a much lower energy barrier than that for a single covalent bond, which promotes the desorption of CO2 and reduces the energy consumption. Further, many interesting structure–property relationships have been clearly established for MOFs. The biggest challenge in the application of MOFs in CO2 separation should be their relatively low stability, especially in water and acids, as exemplified by most of the MOFs discussed above. Nevertheless, many water- and even acid-stable MOFs are known, which can be rationally designed and constructed by using high-valence metal ions, azolate ligands, and/or hydrophobic functional groups.167,197 Even for adsorbents with sufficiently high stabilities, the presence of H2O in the gas mixture generally reduces the CO2 adsorption. However, a few examples have shown the opposite effect, because H2O serves as a CO2 binding site,235-237 but controlling the humidity is a challenge. Hydrophobic adsorbents can display water-resistant CO2 adsorption, but they generally adsorb CO2 weakly owing to the lack of active sites.238-241 Accurate engineering of the pore size and shape, particularly based on ultra-microporous MOFs, might be a viable way to solve this problem. So far, a few MOFs have demonstrated high and stable CO2 capture performance under high humidity and even acidic conditions, but the testing time/cycle is was generally too short for evaluating the suitability in industrial applications. In the absence of a strong CO2 binding (low energy cost), molecular sieving may be the ideal separation mechanism, especially for CO2/CH4 and CO2/C2 hydrocarbon mixtures, because of the extremely high CO2 selectivity (high efficiency). With a relatively low CO2 capacity at low pressures and large pore volumes, molecular sieves can be useful for separating high-pressure mixtures.

Nevertheless, only a handful of materials, including zeolites and MOFs with highly regular structures, as well as porous carbons with amorphous structures, have demonstrated CO 2/N2 and CO2/CH4 molecular sieving effect. In this context, the gate-opening type dynamism of flexible MOFs or other systems is attractive for achieving the molecular sieving effect without precise design of the adsorbent.77,220-222 Slow adsorption/desorption kinetics can be a problem for molecular sieves and ultra-micropores, but the reported breakthrough behaviours imply that the diffusion rate is not a critical problem in the selected experimental conditions.103,106,207,211 However, it should be noted that most of the experiments about MOF adsorption/desorption kinetics were proceeded in relatively low flow rates (for studying the thermodynamic properties). And much higher flow rates (representing kinetic properties and other important parameters) are required in practical applications. In principle, the diffusion rate can be significantly improved by optimising the crystal morphology and/or introducing hierarchical porosity. Many adsorbents have revealed satisfactorily high CO2 adsorption capacities under relevant conditions (Tables 1 and 2), and further increases in the number of such adsorbents and the CO2 separation performances can be anticipated in the near future, especially for MOFs. However, for practical applications, one must further consider desorption method/condition, working capacity, energy consumption/optimisation, long-term/cycling stability in the presence of water and acid, and other important aspects. So far, most of the researches about CO2 capture/separation by using MOFs only simply reported single-component sorption isotherms, maybe together with selectivity calculated from the isotherms. The relatively high cost for MOFs may be one of the main reasons for hindering in-depth investigations. However, the materials cost is closely associated with the application scale, and the cost of MOFs can vary a lot dependent on the composition and synthetic method. At the present stage, stability should be the most important consideration for a MOF showing high CO2 capture performance, and this is indeed a trend of recent studies. For a MOF combining high CO2 adsorption performance and sufficiently high heat, water, and acid stabilities, it is valuable to measure its long-time stability under practical CO2 separation conditions. Note that the adsorption/desorption kinetics, working capacity and energy consumption can be facilely studied, although the desorption method and working capacity have been rarely studied/optimized for MOFs so far. Other industrially important criteria, such as sorbent shaping, cost, mechanical strength, and

so on, can be considered latter. Finally, compositing different types of adsorbents may be a possible approach to combining their advantages and overcoming the disadvantages.242,243 For example, compositing porous materials with light, electric, and/or magnetic responsive materials can help overcome the low heat conductivities of porous materials to promote CO2 desorption and reduce energy consumption.244-246 Considering that water attacks MOFs from the outer surface or at the defects of crystals, coating the MOF crystals with hydrophobic materials may be useful to increase the water stability of MOFs.247-249 Further investigations on different types of adsorbents, including MOFs and their composite materials, can definitely discover useful adsorbents for CO 2 adsorptive separation.

Fig. 1. Reversible chemical absorption of CO2 by primary, secondary, and tertiary amines under dry and humid conditions.

Numbers

2000

Other types of adsorbents MOFs

1500 1000 500 0

2000

2005

2010 Year

2015

2020

Fig. 2. The change of the number of publications per year on CO2 adsorptive separation related to MOFs and other types of adsorbents (Data source: Web of Science).

(a)

(e)

(b)

(f)

(c)

(g)

(d)

(h)

Fig. 3. Relationships of CO2 uptakes (273 K) at 1.0 bar (a−d) and 0.1 bar (e−h) and the pore volume of pores smaller than 1.5 nm (a, e), 1.0 nm (b, f), 0.8 nm (c, g), and 0.5 nm (d, h). Reproduced with permission.83 Copyright 2011, Royal Society of Chemistry.

Fig. 4. Adsorption kinetics of VR-93-M for CO2 (▲), CH4 (■), and N2 (●) at 298 K and 0.53 bar. Reproduced with permission.87 Copyright 2010, Wiley-VCH.

(a)

600 CMAF-6 MAF-6

500

3

-1

Gas Uptake (cm g )

(b)

400 CMAF-5 CMAF-32 MAF-5

300 200 100 0 0.0

0.2

0.4

0.6

0.8

1.0

1.2

P/Po

(c)

Fig. 5. (a) Topologies of MAF-5, MAF-6 and MAF-32. (b) Ar (77 K) and (c) CO2 (298 K) adsorption (solid) and desorption (open) isotherms of CMAF-5, CMAF-6, and CMAF-32. Reproduced with permission.89 Copyright 2017, Royal Society of Chemistry.

(a)

(c)

(b)

(d)

Fig. 6. (a) Illustration of the preparation from uGilT to uGilTH2O. (b) Single-component CO2 and CH4 adsorption isotherms of uGilT and uGilTH2O at 298 K. Infrared spectra of uGilT H2O in CO2 focusing on the (c) H2O absorption peak and (d) the CO2 absorption peak. Reproduced with permission.91 Copyright 2017, Nature Publishing Group.

(a)

(b)

Fig. 7. (a) Crystal structure of CO2-loaded Ca-A and (b) Qst profiles of related zeolites. Reproduced with permission.99 Copyright 2013, Royal Society of Chemistry.

(a)

(b)

Fig. 8. (a) Topology of ZSM-25 and (b) CO2 adsorption kinetic profiles of related zeolites at 303 K and 7.23 kPa. Reproduced with permission.102 Copyright 2018, American Chemical Society.

(a)

(b)

Fig. 9. (a) Topology of zeolite A and (b) the relationships of CO2 and N2 uptakes (298 K, 0.85 bar) with the K+ ratio. Reproduced with permission.104 Copyright 2010, Royal Society of Chemistry.

(a)

(b)

(c)

Fig. 10. (a) Topology, (b) CO2, N2 and CH4 adsorption isotherms at 303 K, and (c) breakthrough curves (50:50 CO2/N2 mixture at 348 K and 1 bar) of r1.9K-CHA. Reproduced with permission.106 Copyright 2018, Royal Society of Chemistry.

(a)

(b)

Fig. 11. (a) Preparation and structure of Heda-Y and (b) CO2 adsorption isotherms of related materials. Reproduced with permission.107 Copyright 2016, Royal Society of Chemistry.

(a)

(b)

(c)

Fig. 12. Structures of (a) conventional chemical grafted, (b) polymerization chemical grafted, and (c) physical adsorbed amines in porous silica.

(a)

(b)

(c)

Fig. 13. PEI status in MCM-41 at (a) low temperature and (b) high temperature. (c) CO2 adsorption (pure CO2 at 75 oC and 1 atm) efficency of PEI in MCM-41. Reproduced with permission.119 Copyright 2002, American Chemical Society.

Fig. 14. The different resistence of PEI and PPI toward oxidation. Reproduced with permission.121 Copyright 2017, American Chemical Society.

Fig. 15. EB modification of PEI for increasing the stability of amine-silica hybrid. Reproduced with permission.122 Copyright 2016, Nature Publishing Group.

(a)

(b)

Fig. 16. (a) Syntheses and (b) CO2 isotherms (295 K) of PPN-6 derivatives. Reproduced with permission.138 Copyright 2012, Wiley-VCH.

(a)

(b)

Fig. 17. (a) Syntheses and (b) CO2 isotherms of TrzPOP-n. Reproduced with permission.143 Copyright 2018, American Chemical Society.

(a)

(b)

(c)

Fig. 18. (a) Syntheses, (b) PXRD patterns, and (c) CO2 isotherms (298 K) of CTF-1 and FCTF-1. Reproduced with permission.144 Copyright 2013, Royal Society of Chemistry.

(a)

(b)

Fig. 19. (a) Syntheses and (b) CO2 isotherms (273 K) of MaSOFn. Reproduced with permission.149 Copyright 2018, American Chemical Society.

(a)

(b)

Fig. 20. (a) Framework and local coordination structures. Reproduced with permission.250 Copyright 2017, American Chemical Society. (b) CO2 adsorption isotherms (298 K) of MOF-74. Reproduced with permission.40 Copyright 2014, Royal Society of Chemistry.

(a)

(b)

(c)

Fig. 21. Unconventional metal-CO2 binding proposed/observed in MOFs. (a) GCMC snapshots of a 15:85 CO2/N2 mixture in PCN-88 showing CO2 chelated by two Cu(II) sites. Reproduced with permission.168 Copyright 2013, Nature Publishing Group. (b) Bidentate CO2 and two-coordinated Cu(II) in MAF-35 suggested by GCMC. Reproduced with permission.169 Copyright 2013, Royal Society of Chemistry. (c) Bidentate CO2 in UTSA-74 determined by SCXRD. Reproduced with permission.38 Copyright 2016, American Chemical Society.

(a)

(b)

Fig. 22. (a) The coordination environments of [Mn2(dobpdc)(mmen)1.6] before and after CO2 adsorption, as well as an ammonium carbamate hydrogen bonding chain. Reproduced with permission.30 Copyright 2015, Nature Publishing Group. (b) CO2 adsorption isotherms for [Mg2(dobpdc)(mmen)1.6] at different temperatures. Reproduced with permission.177 Copyright 2012, American Chemical Society.

(a)

(b)

Fig. 23. Comparison of the computation simulated hydrogen bonding interactions of (a) [Mg2(dobdc)(N2H4)2] and (b) [Mg2(dobpdc)(eda)2]. Reproduced with permission.76 Copyright 2016, Royal Society of Chemistry.

Fig. 24. Computation simulated structure of UiO-66-EA. Reproduced with permission.188 Copyright 2015, Royal Society of Chemistry.

(a)

(b)

(c)

Fig. 25. SCXRD structure of [Zn2(ox)(atz)] loaded with CO2. (a) Interaction of the electron lone pair of the aromatic amine group with the C atom of CO2-I, (b) other interactions of other parts with CO2-I, and (c) interaction between CO2-I and CO2-II. Reproduced with permission.34 Copyright 2010, the American Association for the Advancement of Science.

Fig. 26. SCXRD structure of MAF-23 loaded with CO2. Reproduced with permission.33 Copyright 2012, American Chemical Society.

(a)

(b)

Fig. 27. NPD structures of (a) MFM-300(VIII) and (b) MFM-300(VIV) loaded with CO2. Reproduced with permission.39 Copyright 2017, Nature Publishing Group.

(a)

(b)

Fig. 28. Local coordination structures and CO2 adsorption mechanisms of (a) MAF-X27 and (b) MAF-X27ox. Reproduced with permission.74 Copyright 2015, Royal Society of Chemistry.

Fig. 29. GCMC-DFT simulated structure of PCN-200 loaded with CO2. Reproduced with permission.35 Copyright 2012, Wiley-VCH.

Fig. 30. The topological structure of [M(bpy)2(SiF6)]. Reproduced with permission.208 Copyright 2016, Royal Society of Chemistry.

(a)

(b)

Fig. 31. (a) Gate-opening structural transformation of [Co(bdp)] and molecular sieving adsorption of CO2/CH4. (b) Single-component and mixture CO2/CH4 mixture adsorption isotherms of [Co(bdp)] at 298 K. Reproduced with permission.77 Copyright 2018, American Chemical Society.

(a)

(b)

Fig. 32. (a) Single-component CH4 (square) and CO2 (circle) sorption isotherms and (b) breakthrough curves (40:60 CO2/CH4 mixture at 0.80 MPa) for CID-5 at 273 K. Reproduced with permission.223 Copyright 2012, Royal Society of Chemistry.

(a)

(b)

(c)

Fig. 33. (a) The dia network structure, (b) coverage-dependent Qst profiles of CO2 and CH4, and (c) CO2, CH4, and N2 adsorption isotherms of MAF-66. Reproduced with permission.225 Copyright 2012, American Chemical Society.

Fig. 34. Schematic illustration of the light responsible CO2 adsorption behavior of PCN-123. Reproduced with permission.226 Copyright 2012, American Chemical Society.

Table 1. Benchmark CO2 adsorption properties for flue gas and other mixture with higher CO 2 concentrations. (g cm ) 0.48b

CMAF-32 (carbon)

Adsorbent

−3

CO2 uptakes (mmol g−1/mmol cm−3)

Comment

References

2.1/1.0

Hydrophobic

251,252

1.34/NA

3.26/NA

N species

89

40.0

NA/NA

4.70/NA

N species

90

NA

37.1

2.0/NA

4.96/NA

N & O species

253

NA

NA

NA/NA

4.59/NA

Molecular sieving

87

NaX (zeolite)

1.426

44

3.1/3.2

4.98/5.14

Na+

254, 49

Ca-A (zeolite)

1.514

58

4.1/6.21

5.0/7.57

OMS

99

17atom% NaKA (zeolite)

NA

NA

2.4/NA

3.5/NA

Molecular sieving

104

Zeolite Rho (zeolite)

NA

NA

1/NA

3.5/NA

Molecular sieving

103

3dd-TEPA60% (silica)

NA

NA

5.1#1/NA

NA/NA

Aliphatic amine

125

TEPA-SBA-15 (silica)

NA

NA

4.7#2/NA

NA/NA

Aliphatic amine

127

PEI-MCM-41 (silica)

NA

NA

2.0#1/NA

2.55d, #1/NA

Aliphatic amine

119

TrzPOP-3 (POP)

NA

37

NA/NA

5.09/NA

Multiple polar groups

143

PPN-6-CH2DETA (POP)

NA

55.0

3.04/NA

4.3/NA

Aliphatic amine

138

HKUST-1

0.879

35

2.64/2.32#3

4.86/4.27#3

OMS

255

[Mg2(dobdc)]

0.920

47

5.34/4.93#4, §

8.6/7.9 #4

OMS

49,70

[Co2(dobdc)]

1.177

37

2.66/3.13#4, §

7.5/8.8 #4

OMS

70,256

[Ni2(dobdc)]

1.194

41

2.64/3.15#4, §

7.1/8.5 #4

OMS

70,256

[Mg2(dobpdc)]

0.713

44

4.85/5.3#4, §

6.42/4.58 #4

OMS

177

MAF-35

1.357

47

2.08/2.83

4.46/6.05

OMS

169

[Mg2(dobpdc)(eda)1.6]

0.955

51

3.62/3.46

4.57/4.36

Aliphatic amine

41

[Mg2(dobpdc)(mmen)1.6]

1.073

71

3.13/4.1

3.86/4.14

Aliphatic amine

177

mmen-CuBTTri

1.059

96

2.38/2.52

4.2/4.4

Aliphatic amine

27

NA

NA

4.2/NA

5.0/NA

Aliphatic amine

257

UiO-66-EA

1.310

66

1.05/1.376

1.70/2.23

Alkylol amine

188

[Mg2(dobdc)(N2H4)1.8]

1.178

90

5.18/6.10

5.51/6.49

N2H4

76

[Zn2(ox)(atz)2]

1.156

40.8

1.9/2.2#3

3.8/4.4#3

Aromatic amine

189

MAF-66

1.132

26

1.29/1.46

5.0/5.6

Multiple N

225

MAF-X27ox

1.354

124

4.1/5.5

6.7/9.07

Monodentate OH

74

MAF-X25ox

1.227

120

4.1/5.0

7.14/8.76

Monodentate OH

74

USTA-16

1.116

35

1.37/1.53#4

6.40/7.14#4

Multiple weak interactions

43

PCN-200

1.645

49

0.62/1.02#4

1.78/2.93#4

Multiple weak interactions

35

SIFSIX-3-Cu

1.60

54

2.34/3.68§

2.58/4.13

Multiple weak interactions

211

SIFSIX-3-Zn

1.62

45

4.7/3.75§

2.55/4.1

Multiple weak interactions

207

Qc-5-Cu-sql

1.492

36

0.85/1.27#3

2.16/3.22#3

Molecular sieving

75

0.15 bar

1 bar

25.7

0.6/0.29

NA

56.1

NPCF-10 (carbon)

NA

PYDC-600-2 (carbon) VR-93-M (carbon)

BPL (carbon)

PEI-MIL-101-100

a



Qst (kJ mol−1)

a

Except MOFs, a note of the adsorbent type is given for other types of adsorbents. b Bulk density. #1 75 oC; #2 90 oC; #3 20 oC; #4 23 oC. § 0.1 bar. NA = Not Applicable or Not Available.

Table 2. Benchmark CO2 adsorption properties for DAC. Porous materials



CO2 uptakes at 0.4 mbar −3

−1

−1

Comment

T o

(mmol g /mmol cm )

[Mg2(dobdc)(N2H4)1.8]

1.178

3.89/4.58

Hydrazine grafted MOF

25

76

[Mg2(dobpdc)(eda)1.6]

0.955

2.83/2.72

Aliphatic amine grafted MOF

25

41

NA

2.5/NA

Aliphatic amine impregnated

35

258

25

259

Davisil 646-TEPA

( C)

References

(g cm )

commercial porous silica PEI/silica

NA

2.4/NA

Aliphatic amine impregnated commercial porous silica

[Zn5(OH)4(bibta)3]

0.697

2.20/1.53

MOF with monodentate hydroxide

27

206

[Mg2(dobpdc)(mmen)1.6]

1.073

2.0/2.15

Aliphatic amine grafted MOF

25

177

NbOFFIVE-1-Ni

1.762

1.3/2.29

MOF with ultramicropore

25

213

SIFSIX-3-Cu

1.60

1.24/1.96

MOF with ultramicropore

25

211

MAF-X27ox

1.354

1.13/1.53

MOF with monodentate hydroxide

25

74

NA

1.04/NA

Aliphatic amine grafted POP

22

260

MAF-X25ox

1.227

0.95/1.16

MOF with monodenate hydroxide

25

74

SIFSIX-3-Zn

1.62

0.13/0.21

MOF with ultramicropore

25

207

[Mg2(dobdc)]

0.92

0.088/0.081

MOF with OMS

25

70,256

PPN-6-CH2DETA

NA = Not Applicable or Not Available.

Acknowledgements This work was supported by NSFC (21731007, 21821003, and 21701191) and Guangdong Pearl River Talents Program (2017BT01C161).

Author contributions The manuscript was written with contributions from all authors. All authors have given approval to the final version of the manuscript.

Conflict of interest The authors declare no conflict of interest.

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AUTHOR BIOGRAPHIES

Dong-Dong Zhou was born in Jiangxi, China, in 1990. He received his PhD in 2016 under the supervision of Prof. J.-P. Zhang at Sun Yat-Sen University (SYSU). After serving as an associate researcher, he became an associate professor at SYSU in 2019. His current research interest focuses on the design, synthesis and application of PCPs/MOFs.

Xiao-Ming Chen obtained his BSc in 1983 and MSc under the supervision of Prof. H.-F. Fan in 1986 from SYSU, and his PhD under the supervision of Prof. T. C. W. Mak at the Chinese University of Hong Kong in 1992. He then joined SYSU and became a professor in 1995. He was elected to CAS in 2009 and TWAS in 2013. His current research interest is in synthesis and crystal engineering of functional metal complexes and coordination polymers.

Jie-Peng Zhang obtained his BSc in 2000 and PhD in 2005 under the supervision of Prof. X.-M. Chen at SYSU. He was a JSPS postdoc in Prof. S. Kitagawa's group at Kyoto University from 2005 to 2007. After that, he moved back and joined SYSU in 2007, and became a professor in 2011. His research focuses on adsorptive separation and electrocatalysis based on PCPs/MOFs.