Promoted MgO (A= Na, K, Rb and Cs)

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C: Surfaces, Interfaces, Porous Materials, and Catalysis 2

Interfacial Interactions Governs the Mechanisms of CO Absorption and Desorption on ACO-Promoted MgO (A= Na, K, Rb and Cs) Absorbents 2

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Jin-Su Kwak, Kang-Yeong Kim, Ji Woong Yoon, Kyung-Ryul Oh, and Young-Uk Kwon J. Phys. Chem. C, Just Accepted Manuscript • DOI: 10.1021/acs.jpcc.8b04895 • Publication Date (Web): 13 Aug 2018 Downloaded from http://pubs.acs.org on August 13, 2018

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The Journal of Physical Chemistry

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Interfacial Interactions Governs the Mechanisms

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of CO2 Absorption and Desorption on A2CO3-

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Promoted MgO (A= Na, K, Rb and Cs) Absorbents Jin-Su Kwaka, Kang-Yeong Kima, Ji Woong Yoonb, Kyung-Ryul Oha and Young-Uk Kwona*

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a

Department of Chemistry, Sungkyunkwan University, Suwon 16419, Korea

b

Research Center for Nanocatalysts and Molecular Simulations, Korea Research Institute of Chemical Technology (KRICT), Daejeon 305-600, South Korea

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*

E-mail: [email protected]

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ACS Paragon Plus Environment

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Abstract

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In this study, we investigated the detailed mechanisms of CO2 absorption and desorption on

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A2CO3-promoted MgO absorbents (A = Na, K, Rb, and Cs). We analyzed the materials

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formed at various stages of the absorbents during CO2 absorption and desorption by thermal

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decomposition-gas chromatography/mass spectroscopy, in-situ IR spectroscopy, solid-state

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MAS NMR spectroscopy on

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spectroscopy, and in-situ X-ray diffraction as well as the conventional thermogravimetric

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analysis, ex-situ X-ray diffraction, and elemental distribution mapping by energy dispersive

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X-ray analysis using scanning electron microscopy. The absorption of CO2 of the absorbents

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occurs in two steps. The first step is a fast process involving basic sites on the MgO surface

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formed by the interaction between A2CO3 and MgO. Because the basicity of these sites

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depends on the size of A ion, the kinetics and capacity of CO2 absorption and the desorption

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properties of this process are strongly dependent on the nature of A. Since basic sites are

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formed at the interface between A2CO3 and MgO, the observation of this first step depends on

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the method of sample preparation among other factors, explaining the failure of observing

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this step in the previous studies on similar absorbents. The second step is a slower process

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during which the double carbonate phase between Mg and A is formed. The diffusion of solid

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state materials required to form the double carbonate phases explains the slow kinetics of this

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step. The phase stability of the double carbonates also influences the kinetics and capacity of

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CO2 absorption in the second step. We also find that, against the conventional belief, a

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physical mixture between A2CO3 and MgO undergoes chemical changes in addition to

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dehydration and dehydroxylation during the pre-treatment step of thermogravimetric analysis,

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which, in fact, is a necessary step for the first step process to occur.

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C,

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Na, 25Mg and

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K nuclei, diffuse reflectance UV/Vis

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1. Introduction

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Curtailing the emission of greenhouse gas, especially CO2, is an issue of utmost urgency for

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mitigating the on-going global warming problem.1 Ultimately, the present level of heavy

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reliance on fossil fuels of human activity will have to be lifted by replacing fossil fuels with

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more sustainable energy sources, which, however, will take a long time.2 In the meantime,

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reducing CO2 emission while maintaining the present level of human activity is also

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important, for which the idea of carbon capture and sequestration (CCS) technologies have

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emerged and pursued.3-5

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There are a number of scientific and technological issues to be addressed for CCS

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technology to be realized, however. Focusing on the CO2 capture in CCS, for instance, in the

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post-combustion process which is considered to be the least demanding for implementation

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and, thus, is expected to find applications the earliest among the proposed process to capture

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CO2, reliable agents to capture CO2 is an important issue. Energy and cost-efficient operation

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of the process is also important,6-9 for which Park et al. proposed the multiple-stage CO2

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capture technology to be applied to the post-combustion process. In this technology, solid

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CO2 absorbents operating at three different temperature ranges are to be used so that the

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waste heat from a higher temperature process can be used in the next lower temperature

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one.10 In this concept, the amine-impregnated absorbents can be used at low temperatures

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below 200 °C.11-15 For the high temperature (500 °C and higher) absorbents CaO-based,

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lithium zirconate, and lithium silicate materials are considered.11, 13, 16-20

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The candidate materials for the mid-temperature (200-500 °C) absorbents are mostly

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centered around MgO because MgCO3 decomposes to MgO and CO2 at below 500 °C, while

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all of the alkali carbonates and all of the other alkaline earth carbonates decompose at 3



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temperatures higher than 500 °C, which is due to the largest charge density of Mg2+.21

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However, the very large lattice energy of MgO originating from the high charge density of

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Mg2+ renders it and its derivatives extremely sluggish reaction kinetics for both absorption

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and desorption, whose resolution is a prerequisite for any MgO-based absorbent ever to be

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used in real applications.22 In order to improve the kinetics of MgO-based absorbents, alkali

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metal nitrates22-26 and nitrate-carbonate mixture27-30 have been introduced as a promoter.

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However, although they significantly enhance CO2 absorption capacity of MgO that the

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nitrate promoters are in their molten state limits their applicability. Certain types of

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operations of CCS, including the aforementioned multiple-stage CO2 absorption process,

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require all solid state absorbents.10

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In the literature, there are a few examples of all solid state absorbents of MgO, among which

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alkali carbonate (A2CO3) promoters appear to be promising with the highest absorption

84

capacity and kinetics. Studies using Na2CO3, K2CO3, and Cs2CO3 as a promoter for MgO

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have been reported. They explained the CO2 absorption as due to the formation of double

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carbonates such as Na2Mg(CO3)2 and K2Mg(CO3)2.31-33 However, many of the details of the

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absorption and desorption mechanism of these absorbents remain unexplained.

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In this study, we prepared A2CO3 (A = Na, K, Rb, and Cs)-promoted MgO and investigated

89

their CO2 absorption mechanisms in details. We found that the CO2 absorption occurs in two

90

steps, not in the single step reported in the previous works.31-33 Moreover, we found that the

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two steps involve different chemical reactions, both of which involve compounds of A in

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different ways. Based on these results, we propose the CO2 absorption and desorption

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mechanism of A2CO3-promoted MgO absorbents. The details are given in the followings.

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2. Experiment

97 98

Materials

99 100

Mg(OH)2 (Alfa Aesar, 95-100.5%), Mg5(CO3)4(OH)2.4H2O (Alfa Aesar, reagent grade,

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MgO = 40 ~ 43.5%, -325 mesh), Na2CO3 (Sigma-Aldrich, 99.5+%), K2CO3 (Sigma-Aldrich,

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99+%), Rb2CO3 (Alfa Aesar, 99.8%, metal basis), Cs2CO3 (Alfa Aesar, 99.9%, metal basis),

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and Mylar film (Chemplex Industry Inc.; thickness = 6.0 µm, major impurities (in ppm

104

level) = Ca, P, Sb, Fe, Zn).

105 106

Preparation of absorbents

107 108

A2CO3-promoted MgO absorbents (A = Na, K, Rb, and Cs) with the A2CO3/MgO molar

109

ratio 0.10 were prepared by physically mixing A2CO3 and MgO in an agate mortar-and-pestle.

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The reagents were dried at 120 °C to remove absorbed water before mixing. MgO was

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obtained by heating Mg(OH)2 at 450 °C for 4 h under an air flow, which produced pure MgO

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by X-ray diffraction. Hereafter, these absorbents will be denoted as A2Mg; when the identity

113

of A matters it will be indicated in the denotation as in Na2Mg and absorbents with molar

114

ratios other than 0.10 will be denoted with the ratios indicated as in 0.3Na2Mg. 0.01A2Mg

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samples were prepared for infra-red (IR) spectroscopic studies. For the study on the elemental

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distribution in the absorbents with A = Na and K during the CO2 absorption and desorption,

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0.4Na2Mg and 0.3K2Mg absorbents with higher A2CO3 contents were prepared to see the

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locations of A atoms clearly. In such cases, ball-milling technique was used instead of mixing

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by a mortar-and-pestle. A slurry composed of Mg5(CO3)4(OH)2⋅4H2O and A2CO3 (A = Na 5



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and K) in the desired A/Mg molar ratio and distilled water was placed in a plastic bottle

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(Nalgene, 250 mL) with zirconia balls (Grinding Media, φ = 10 mm), and the bottle was

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rotated at 60 rpm for 24 h, followed by drying at 120 °C and calcination at 450 °C for 4 h

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under an air flow. The same ball-mill method was used for the preparation of Na2Mg(CO3)2

124

(mixed with Na2CO3) and K2Mg(CO3)2 to be used as a reference for 23Na, 25Mg and 39K solid-

125

state NMR studies. The Na/Mg molar ratio in Na2Mg(CO3)2 was 2.4, 20% excess in Na2CO3

126

than the stoichiometric ratio of 2.0 and the K/Mg molar ratio in K2Mg(CO3)2 was 2. In the

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case of K2Mg(CO3)2, it was found that drying at 50 °C before drying at 120 °C was a

128

necessary step to increase the reproducibility of synthesis. Calcination was not required since

129

the desired A2Mg(CO3)2 (A = Na and K) phase forms upon drying. However, they were heat-

130

treated at 450 °C for 4 h in air before TG analysis. After synthesis, all samples were kept in

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an oven at 120 °C to keep the amount of absorbed water minimized while not inducing any

132

reaction within them.

133 134

Characterization

135 136

Thermogravimetric (TG) analyzer (Hitachi Thermal Analysis, STA 7200) was used to

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assess the CO2 absorption and desorption properties of the A2Mg absorbents. Typically, the

138

environment was 100% CO2 condition (flow rate = 100 mL.min-1) for both absorption and

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desorption; N2 was used to dilute CO2 when lower concentrations of CO2 were required. Each

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TG operation was preceded by heating the sample up to 400 °C in a 100% N2 flow (pre-

141

treatment). For 0.4Na2Mg and 0.3K2Mg absorbents, the pre-treatment temperature was 350

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°C. CO2 absorption capacity on the A2Mg absorbents will be expressed by wt%, which means

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the percentage of the mass of absorbed CO2 divided by the mass of absorbent. 6


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In order to investigate the changes in Na2Mg absorbent during the pre-treatment, thermal

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decomposition-gas chromatography/mass spectroscopy (TD-GC/MS; Agilent TDSA2) was

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used. Na2Mg absorbent was heated from room temperature to 400 °C in 1 atm He to collect

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the emitted gas and the gas was analyzed by a GC/MS. A φ= 250 µm column filled with a

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0.25 µm thick film composed of 5% biphenyl and 95% polydimethylsiloxane was used for

149

the GC. MAS NMR spectroscopy (Rigaku, 400 MHz, in Korea Basic Science Institute) on

150

13

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derivatives obtained by heating at 200 °C, 400 °C and 800 °C in N2 flow. The change of

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sample state during heating up to 500 °C was studied by monitoring the carbonate ions by IR

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spectroscopy (Thermo ScientificTM, Nicolet is50 FT IR spectrometer) with 100 °C

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temperature intervals. 15 mg of a sample without any additive such as KBr was pelletized

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into a disc (φ = 16 mm) for the measurement. The sample disc was placed in the sample

156

holder, heated to a designated temperature in a vacuum (1.0 x 10-6 torr) for 60 min, and

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cooled to room temperature before recording the IR spectrum. Since the intensity was too

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strong on A2Mg absorbents, 0.01A2Mg samples with ten times lower A2CO3 contents were

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prepared for the IR studies. Diffuse reflectance UV/Vis spectrometer (DR-UV/Vis, Shimadzu

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UV-3600) was used to record the absorption spectra of A2Mg absorbents in the wavelength

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range of λ = 200-500 nm in the diffuse reflectance mode. The background spectrum of DR-

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UV/Vis was calibrated by pelletized BaSO4 powder (Samchun, 97.5%). UV/Vis spectra of

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A2Mg were obtained on samples placed on the top of pelletized BaSO4. In-situ XRD (Rigaku

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SmartLab, CuKα (λ = 1.5406 Å), in UNIST Next-generation Catalysis Center) was used to

165

monitor the phase change of absorbents after pre-treated states and CO2 absorption state of

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the absorbents at 330 °C for 180 min in 100% CO2. The samples were heated by an infra-red

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lamp (200 V, 1800 W) under a CO2 flow (flow rate = 100 mL.min-1). The heating rate was

C,

23

Na, 25Mg and 39K nuclei were performed on Na2Mg and K2Mg absorbents and their

7



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10 °C.min-1. The phase identification of the absorbents after cycling test was done by powder

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X-ray diffraction (XRD, Rigaku D/Max 2200 Ultra, Cu Kα). Mylar film was used when the

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sample was to be kept from contacting the moisture in the environment during the XRD

171

measurement. Elemental distribution of Mg and A atoms in samples was investigated by an

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energy dispersive X-ray spectroscopy (EDS) analyzer equipped on a high resolution scanning

173

electron microscope (FESEM, JEOL JSM-7100F).

174 175

3. Results & Discussion

176 177

3.1. CO2 absorption and desorption properties on A2Mg absorbents

178 179

In a preliminary study on two series of A2CO3-promoted MgO (A = Na, K, Rb, and Cs)

180

absorbents, A2Mg and 0.2A2Mg with the A2CO3/MgO molar ratios of 0.10 (Figure 1) and

181

0.20 (Figure S1), respectively, we found that the former series is better than the latter one in

182

terms of the overall CO2 absorption and desorption abilities across the series. In their report

183

on using a triple eutectic (Li,Na,K)2CO3 to promote the CO2 absorption by MgO, Dagle et al.

184

discussed that carbonate in excess to a certain amount could block the access to MgO by CO2

185

and decrease the absorption ability, which may be the reason for the decrease of absorption

186

capacity of 0.2A2Mg absorbents in the present study.34 Based on this observation, we

187

investigated the series of A2Mg. However, it is noted that for the case of A = Na, 0.2Na2Mg

188

shows a larger absorption capacity than Na2Mg.

189

Figure 1(a) shows the weight changes of A2Mg absorbents with the variation of temperature

190

under 100% CO2 along with that of MgO as a reference. The reference MgO shows a small

191

weight gain of 2 wt% with a broad peak centered at ~180 °C, which can be attributed to the 8


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192

adsorption of CO2. This adsorbed CO2 is completely desorbed at temperatures higher than

193

360 °C. On the contrary, the A2Mg absorbents show peaks centered at 330-360 °C, higher

194

than the desorption temperature of MgO. The amount of weight gain varies from sample to

195

sample, 5 wt% for Na2Mg, 7 wt% for K2Mg, 13 wt% for Rb2Mg, and 6 wt% for Cs2Mg.

196

Because these weight gains are much larger and their peaks are at higher temperatures than in

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the reference MgO, they are attributed to absorption of CO2 rather than adsorption. In all of

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the absorbents, desorption is completed at 480 °C or higher. We also tested the Li-analogue,

199

but it does not show any sign of enhancing the absorption of CO2 and thus is not included in

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this study (Figure S2). These observations indicate that all of alkali carbonates except

201

Li2CO3 can enhance the CO2 absorption of MgO which by itself has no reactivity to CO2 at

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all. This is in agreement with the literature on the enhancing effects of A2CO3 (A = Na, K and

203

Cs) 31-33; the effect of Rb2CO3 in this study has not been reported before, making our data the

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first report on it.

205 206

Figure 1. (a) Weight changes of A2Mg (A=Na, K, Rb, and Cs) absorbents with the temperature 9



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variation from 60 to 600 °C under 100% CO2 at a heating rate of 10 °C.min-1. The data on the

208

reference MgO is also shown. (b) Weight changes of A2Mg absorbents with time under a 100% CO2

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condition at 330 and 460 °C. Red marks in (b) indicate inflection points between the first and second

210

steps.

211 212

Figure 1(b) shows the time variations of CO2 absorption on A2Mg absorbents at a constant

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temperature of 330 °C, followed by desorption at 460 °C, both under 100% CO2. The analysis

214

results of these data are summarized in Table 1. In all of the curves, absorption occurs in two

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steps. The proportion of the first step in the total absorption increases as the size of A atom

216

increases. Therefore, while the second step (5.6 wt%) is mainly responsible for the total

217

absorption in Na2Mg (7.3 wt%), it is the first step in Cs2Mg (5.8 wt%) that accounts for most

218

of its overall absorption (7.6 wt%). In the two intermediate members of the series, K2Mg and

219

Rb2Mg, the two steps are comparable in the amount of absorption. K2Mg shows the largest

220

capacity of 14.4 wt% among the four absorbents. In all of the absorbents, the first step is

221

faster than the second one in absorption rate. The onset time of the second step also shows a

222

trend with the size of A. In Na2Mg, the first step lasts for a short period of time and is

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superseded by the second step after 7 min. The onset of the second step is delayed to 30 min

224

in K2Mg and 50 min in Rb2Mg, but in Cs2Mg it occurs at 20 min.

225 226

Table 1. Summary of CO2 absorption rate and capacity of A2Mg absorbents in Figure 1(b). A

a

K

Rb

Cs

Maximum absorption

1st step

1.10

2.50

2.34

2.30

rate (wt%.min-1)a

2nd step

0.35

0.24

0.07

0.11

1st step

1.70

7.80

7.40

5.80

2nd step

5.60

6.60

4.10

1.80

Total

7.30

14.4

11.5

7.60

Absorption capacity (wt%)

227

Na

Maximum rate was obtained by taking the peak in the first derivative plot of the weight vs. time plot 10


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228


in Figure 1(b).

229 230

After CO2 absorption is saturated, the absorbents were heated at 460 °C under 100% CO2 to

231

induce desorption. Unlike absorption, desorption appears to occur in one step in all of the

232

absorbents. In all of the absorbents, the final weights after desorption show some changes

233

from the initial weights. The weight changes show an interesting variation with the size of A.

234

While Na2Mg and K2Mg show small weight losses of -1.0 − -1.3 wt%, Rb2Mg and Cs2Mg

235

show larger weight gains of 2.6 − 2.8 wt %. Because the only way to account for the weight

236

changes in these materials is the amount of CO2 uptaken or released, a weight gain means the

237

formation of stable carbonate ions and a weight loss means that a part of carbonate in the

238

starting absorbent is lost by decomposition after desorption. If one is to assume that the CO2

239

absorption/desorption is solely attributed to the formation/decomposition of MgCO3 and/or

240

A2Mg(CO3)2 double salts as most of the literature works discuss,31-33 then there cannot be any

241

change of weight after desorption let alone the variation of the final weight depending on the

242

alkali metal ion. Therefore, one can conclude that there are formation and decomposition of

243

A2CO3 involved in addition to those of Mg-compounds during the absorption and desorption

244

in Figure 1(b). That A2CO3 can be decomposed may be surprising because their

245

decomposition temperatures are reported to be higher than 800 °C.35 The details of this point

246

will be discussed later.

247

The CO2 absorption data measured at 330 °C while varying the CO2 concentration from 3%

248

to 100% in steps are shown in Figure 2. The threshold CO2 concentration for absorption to

249

take place changes with A. In Cs2Mg, the CO2 absorption occurs when the CO2 concentration

250

is as low as 3%, while Na2Mg requires 75% CO2 for the absorption to occur. In the case of

251

Rb2Mg, absorption starts at 3% CO2 at a lower rate than Cs2Mg and a second step absorption

252

occurs when the CO2 concentration is increased to 50%. These onset CO2 pressures of 11



253

absorbents appear to be related with the first step absorption seen in Figure 1(b).

254 255

Figure 2. Weight changes of A2Mg absorbents with the variation of CO2 concentration from 3% to

256

100% at 330 °C. Flow rates of N2 and CO2 were controlled to get the designated CO2 concentrations

257

while keeping the total flow rate at 100 mL.min-1.

258 259

The absorbents were subjected to cycling tests to measure their capacities in repeated

260

absorption and desorption operation (Figure 3). In all of the absorbents, the absorption

261

capacity is the maximum in the first cycle and is decreased in the following cycles. In Na2Mg,

262

K2Mg and Rb2Mg, repeatable capacity is attained from the second cycle onward. In Cs2Mg,

263

the loss of capacity is continued until the fourth cycle after which a stable capacity is attained.

264

The reasons for the capacity loss also appear to be dependent on the size of A. In Na2Mg and

265

K2Mg, the peak weight is decreased on moving from the first cycle to the second cycle,

266

accounting for the major part of the capacity loss. In fact, the baseline is lowered gradually as

267

the cycle number increases partly compensating the loss of capacity due to the decreased

268

peak weight although its contribution to the total capacity is small. In Rb2Mg and Cs2Mg, the

269

baseline is raised from that of the first cycle. In Rb2Mg, the base-line is constant after the

270

second cycle but, in Cs2Mg, it continues to rise until the end of the third cycle. This baseline 12


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271

rising adds to the decrease of the peak weight in reducing the capacity in the two absorbents.

272

The change of the baseline depending on A appears to be related with the final weights vs.

273

initial weight seen in Figure 1(b) which, we believe, is related with the formation and

274

decomposition of A2CO3.

275

It must be noted that the varying behavior of the baseline is induced because the CO2

276

desorption of Figure 1 and Figure 3 was conducted under a 100% CO2 condition. If the

277

desorption is conducted under 100% N2 (and at a lower temperature of 400 °C than the 460

278

°C in the Figures 3), the base-line is unchanged with cycling (Figure S3).

279 280

Figure 3. Cyclic CO2 absorption and desorption properties of A2Mg absorbents under 100% CO2. The

281

temperature was varied between 330 (for absorption of CO2) and 460 °C (for desorption). The top line

282

is the temperature profile.

283 284

3.2. Mechanism studies on CO2 absorption of A2Mg absorbents

285

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286

3.2.1. State of A2Mg absorbent before CO2 absorption

287 288

In order to understand the promoting effects of A2CO3 and their different effects as seen

289

above, we performed detailed studies on the mechanism. In this section, we will examine the

290

state of A2Mg absorbents. As will be shown below, the state of absorbents bears important

291

consequences in their CO2 absorption properties. The TG data on A2Mg absorbents discussed

292

thus far were obtained after heating physical mixtures of A2CO3 and MgO at 400 °C in N2 as

293

a pre-treatment. We thought that this pre-treatment could affect the CO2 absorption

294

characteristics of the absorbents.

295

The TG data of the physical mixtures of A2CO3 and MgO without pre-treatment from room

296

temperature to 400 °C in N2 show two-step weight losses (Figure S4). The weight loss at

297

below 150 °C can be ascribed to the loss of water molecules taken up by the hygroscopic

298

A2CO3 during the sample storage and handling. On the other hand, the nature of the second

299

step of weight loss occurring at 250-400 °C is unclear. Typically, weight losses in this

300

temperature range are ascribed to the loss of hydroxyl groups as water. However, it is also

301

possible that a part of carbonate ions of A2CO3 is decomposed by losing CO2. In order to

302

clarify this uncertainty, we performed a few different analyses. First, we analyzed the emitted

303

gas from a physical mixture between Na2CO3 and MgO during the heat treatment at 400 °C

304

by GC-MS. The spectrum shows a strong peak with m/z = 44, which can be explained as CO2,

305

along with other peaks such as O2 (Figure S5). Second, we used IR spectroscopy to monitor

306

the variation of the absorption peaks of carbonate groups as a function of heat-treatment

307

temperature from room temperature to 500 °C. We prepared 0.01A2Mg samples for this

308

purpose because the peaks of A2Mg samples were too strong to analyze in details (Figure S6).

309

The only difference between 0.01A2Mg and A2Mg samples is that the former has ten-times 14


Page 14 of 38

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310

less A2CO3 than the latter because both of them are physical mixtures between A2CO3 and

311

MgO. The reference MgO shows peaks of carbonate group at 865 cm-1 (out of plane bending),

312

1071 cm-1 (bidentate, symmetric stretching), 1343 cm-1 (unidentate, symmetric stretching),

313

1411 cm-1 (asymmetric stretching), 1512 cm-1 (unidentate, asymmetric stretching) and 1611

314

cm-1 (bidentate, asymmetric stretching) (Figure S6(a)). These peaks decrease in intensity

315

slightly at 200 °C and disappear almost completely at 300 °C, consistent with the TG data in

316

Figure 1(a). These peaks are originated from adsorbed CO2. The spectra on 0.01A2Mg

317

samples in Figure 4 also show an abrupt decrease of intensity at 300 °C, indicating the loss

318

of carbonate groups from A2CO3. Different from the MgO case, however, some peaks remain

319

up to 500 °C, which can be identified as those of corresponding A2CO3.

320 321

Figure 4. IR spectra of 0.01A2Mg samples (A=Na, K, Rb and Cs) treated at different temperatures

322

under vacuum condition.

323 324

From the TG, TD-GC/MS and IR spectra, it is evident that a part of A2CO3 in A2Mg (and

325

0.01A2Mg) samples is decomposed upon heating. Our TG data on pure A2CO3 show a very

326

small amount of decomposition of carbonate ions (less than 1 wt%) up to 600 °C (Figure S7). 15



Page 16 of 38

327

Therefore, we were led to conclude that the decomposition of A2CO3 in A2Mg samples was

328

induced by MgO.

329

We used solid-state MAS NMR on

13

C,

23

Na,

25

Mg and

39

39

K nuclei to probe the state of

330

samples during the sample preparation and pre-heating.

331

prepared by mixing K2CO3 and MgO. According to Papenguth et al. who studied 13C MAS

332

NMR on various carbonates including Na2CO3 and MgCO3 among many others, the chemical

333

shifts of carbonates ions are not strongly affected by the cation species and crystal structure

334

of carbonate samples. Instead, the peak width is noticeably changed by the crystallinity

335

although no quantitative analysis seems possible.36 In our case, the physical mixture between

336

Na2CO3 and MgO, the peak position and peak width are different from pure Na2CO3 (Figure

337

S8). Evidently, the chemical environment of carbonate ions in Na2CO3 is changed by mixing

338

it with MgO.

K NMR was used for samples

339

In order to probe the state of Na atoms during sample preparation and pre-treatment, we

340

analyzed Na2Mg sample and its heat-treated derivatives at 200 and 400 °C in N2 with solid-

341

state

342

Na2Mg(CO3)2/Na2CO3 was prepared by mixing Na2CO3 and Mg5(CO3)4(OH)2⋅4H2O (with

343

Na2CO3 in excess by 20%) in water and drying the slurry at 120 °C, whose XRD pattern

344

showed Na2Mg(CO3)2 as the major phase and Na2CO3 as the minor phase (Figure S9(a)).

345

Commercial Na2CO3 and MgO were also used for the 23Na and 25Mg references, respectively.

346

Comparing with the spectra of the references, the peaks in the 23Na NMR spectra of Na2Mg

347

and its heat-treatment derivatives can be explained as those of Na2CO3 (at 5.1, -6, and -11.5

348

ppm) and Na2Mg(CO3)2 (at -0.3 ppm). It is interesting that the peak of Na2Mg(CO3)2 is

349

present in the physical mixture state. Unlike the Na2Mg(CO3)2/Na2CO3 reference, it was

350

prepared by mixing Na2CO3 and MgO without the aid of water. Moreover, the mixing was

23

Na

and

25

Mg

MAS

NMR

spectroscopy

16


(Figure

5).

The

reference

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


351

done in air where the CO2 concentration was very low. This peak increases in intensity, when

352

the mixture is heated at 200 °C, but is almost disappeared upon heating at 400 °C. The 25Mg

353

NMR spectra support these observations. While MgO (at 26 ppm) predominates in all of the

354

samples, a small peak at -9 ppm due to Na2Mg(CO3)3 appears at 200 °C and disappears at

355

400 °C.

356 357

Figure 5. (a) 23Na and (b) 25Mg solid-state MAS NMR spectra of a physical mixture between Na2CO3

358

and MgO (molar ratio Na2CO3/MgO = 0.1) and samples obtained by heating it at 200 and 400 °C in

359

N2.

360 361

We also tried to probe the states of K and Mg atoms in K2Mg absorbent during pre39

25

362

treatment by solid-state

363

because of the low magnetogyric ratios of

364

spectral quality of 39K NMR is poorer than that of 23Na NMR.37 Nevertheless, it is clear that

365

the local environment of K ions in K2Mg sample and its heat-treated derivatives is much

366

different from that of pure K2CO3 but rather close to that of K2Mg(CO3)2 (Figure S10).

367

K and

Mg MAS NMR spectroscopy. Unfortunately, however, 39

K and

41

K of spin -3/2 quadrupolar nuclei the

The formation of A2Mg(CO3)2 in the physical mixing can be expressed as in the following

368

equation:

369

0.1A2CO3 + MgO 17



Page 18 of 38

370

→ x A2Mg(CO3)2 + (1-x) MgO + x A2O + (0.1-2x) A2CO3 (1)

371

Including the formation of CO2 during the pre-treatment, the equation is changed to the

372

following one:

373

0.1A2CO3 + MgO

374

→ x A2Mg(CO3)2 + (1-x) MgO + (x+y)A2O + (0.1-2x-y) A2CO3 + yCO2(g)

375

In order to form A2Mg(CO3)2 (A = Na and K) during the physical mixing and to produce CO2

376

during the pre-treatment, these equations indicate that A2O must be formed as one of the

377

products. Since pure A2CO3 is stable up to much higher temperatures, A2O in these equations

378

must be stabilized through interaction with other materials. Because there is MgO in large

379

excess, it is likely that A2O is reacted with it forming A-doped MgO.

(2)

380

Doping of alkali metal ions in MgO has been extensively investigated in relation with the

381

catalysis of MgO. According to the related literature, bulk doping of alkali metal ions in the

382

MgO lattice is unfavorable because of the large difference of ionic size between alkali metal

383

ions and Mg2+ ion and the charge imbalance between them. Even for Li whose ionic size is

384

the closest to that of Mg among the alkali metal, its solubility in MgO lattice is reported to be

385

very small. On the other hand, alkali metal ions can induce defects on the surface of MgO,

386

namely surface doping.38-40

387

We, therefore, sought for evidences for the formation of A2O by monitoring the MgO in the

388

pre-treated A2Mg absorbents. The XRD patterns of the pre-treated A2Mg samples show the

389

peaks of MgO and A2CO3 (Figure S11(a)). The MgO peaks of A2Mg absorbents are slightly

390

shifted to higher angles from those of MgO reference, indicating shrinkage of MgO lattice.

391

The lattice parameters calculated by the least-squares fits are smaller than those of MgO

392

reference and are gradually decreased as the size of A increases (Figure S11(b)). These

393

changes of the lattice parameter of MgO corroborate with what is expected if A-doped MgO 18


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


394

is to form. The magnitude of lattice shrinkage may be used as a measure of the concentration

395

of oxygen defect. Therefore the trend of the lattice parameter can be taken to be suggesting

396

the increasing efficiency of doping and, hence, the extent of reaction (2) with the size of A.

397

Additional support for surface doping was obtained from the DR-UV/Vis spectroscopic

398

data on pre-treated A2Mg absorbents shown in Figure 6. The spectrum of MgO shows a weak

399

absorption at ~290 nm in addition to the major absorption at wavelengths shorter than 260

400

nm. In the literature, these are attributed to the exciton excitations at threefold and fourfold

401

coordinated surface O2- ions, respectively.41 The presence of these surface defects of MgO

402

can be understood by noting that the MgO in this study is obtained by decomposing Mg(OH)2.

403

In addition to these peaks of MgO, all of A2Mg absorbents show a peak at 375 nm.

404

In a study on MgO doped with various amounts of Li, Zavyalova et al. also observed this

405

additional peak. Their data show that the intensity of this peak increases with the amount of

406

Li reacted to prepare the samples, indicating that this peak is induced by the action of Li.38

407

However because their samples were heat-treated at 800 °C, most of Li was lost by

408

evaporation. Therefore, although the detailed origin is not identified, this peak cannot be

409

associated with the Li-content but is attributed to the surface defect of MgO induced by Li in

410

the process of sample preparation. That the same peak is observed with alkali metal atoms

411

other than Li in our study fortifies this interpretation. Furthermore, our data show that the

412

intensity of this absorption increases with the size of A, indicating that the efficiency to create

413

the surface defect responsible for this peak increases with the size of A. This trend of peak

414

intensity corroborates with the trend of the lattice parameters mentioned above.

415

The results presented above have led us to conclude that (1) pre-treatment of A2Mg

416

absorbents at 400 °C under N2 induces decomposition of A2CO3 into A2O and CO2, (2) the

417

A2O so-formed exist as a dopant on the MgO surface, (3) the doping creates defects on the 19



418

surface of MgO, and (4) the concentration of such defect increases with the size of A making

419

Cs the most efficient surface dopant.

420 421

Figure 6. Diffuse reflectance UV/Vis spectra of pre-treated A2Mg absorbents at 400 °C in N2.

422 423

3.2.2. Mechanism of the first step absorption of A2Mg absorbents

424 425

In the literature on surface doping on MgO, it is well established that variously introduced

426

alkali metal can induce defect sites on MgO and the basicity of such sites is in the order of Cs

427

> K > Na > Li.42 Although the Rb case is not reported in the literature, it is reasonable to

428

locate it between Cs and K in this order. The same trend was observed in the CO2 absorption

429

rates of the first step absorption curves in Figure 1(b). Therefore, it seems that the different

430

CO2 absorption properties of A2Mg absorbents seen in Figure 1(b) are related with the

431

different extents of surface doping on MgO depending on A. That is, the reaction responsible

432

for the first step of CO2 absorption of A2Mg absorbents is the carbonation of A2O on the

433

surface defect sites of MgO: 20


Page 20 of 38

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434

A2O (on MgO defect) + CO2––––→ A2CO3

435

The weight gains of Rb2Mg and Cs2Mg absorbents after desorbing CO2 may be explained by

436

the large extent of the A2O formation in the pre-treatment step. Probably, the A2CO3 particles

437

formed by this reaction grow large in size and are detached from MgO particles making the

438

product A2CO3 free of the effect of MgO, explaining the irreversible weight gains after the

439

desorption of CO2 seen in Figure 1(b) and Figure 3. On the contrary, K and especially Na

440

have weaker propensity to form surface defects and there are fewer surface defects formed in

441

the pre-treated samples to begin with, explaining the absence of net weight gains in the K2Mg

442

and Na2Mg absorbents after the desorption of CO2. In addition, because the desorption

443

temperature, 460 °C, is higher than that used for pre-treatment, 400 °C, it is possible to

444

induce additional decomposition of A2CO3, explaining the weight losses in these samples.

(3)

445

These findings of chemical changes of absorbents upon heat-treatment below 400 °C give

446

especially important insight in understanding the CO2 absorption data in the literature. Most

447

of CO2 absorption properties reported in the literature are based on measurements on TG, in

448

which pre-treating the sample under an inert gas flow before submitting it to a CO2

449

environment is a conventional practice. Since there has been no report on the reactions during

450

pre-treatment, it seems that no one has suspected any reaction during the pre-treatment. In

451

this regard, our results show the need to pay attention to the pre-treatment conditions.

452

Depending on the pre-treatment condition, the state of the absorbents just before the CO2

453

absorption reaction in TG can be different from the state of sample analyzed and reported.

454 455

3.2.3. Mechanism of the second step absorption of A2Mg absorbents

456 457

The CO2 absorption behavior of A2Mg absorbents is governed by the nature of A not only 21



458

through the surface doping on MgO as discussed in the previous section but also through the

459

phases to be formed during the CO2 absorption. In the papers reporting on MgO absorbents

460

promoted by A2CO3 (A = Na31 and K32), the phases responsible for the CO2 absorption are

461

reported to be A2Mg(CO3)2 double carbonates. However, the double carbonates show varying

462

phase stability depending on A. In the case of A = Na, Na2Mg(CO3)2 is known to be a stable

463

phase under an ambient condition.43 We confirmed this by showing that it could be formed

464

even when Na2CO3 and MgO were ground together and that this phase could be obtained

465

quantitatively by mixing Na2CO3 and Mg5(CO3)4(OH)2⋅4H2O in water followed by drying.

466

On the contrary, it appears that the analogous double carbonates are increasingly unstable as

467

the size of A increases. Although we demonstrate that K2Mg(CO3)2 can be synthesized in a

468

similar way as Na2Mg(CO3)2, its formation is reported to require a higher temperature and a

469

higher CO2 pressure condition.44 In the case of Rb, there is no report on its formation at all. In

470

the case of Cs, Liu et al. observed a new phase by in-situ XRD during the CO2 absorption of

471

a Cs2CO3-promoted MgO absorbent at 300 °C. Although they did not determine the

472

composition or structure of this phase besides that it was a high-temperature phase, it is likely

473

to be a double carbonate between Cs and Mg.33

474

On the contrary, in the literature, the hydrated double salts, A2Mg(CO3)2⋅4H2O, for A = K,

475

Rb, and Cs are readily formed at room temperature, with the K and Rb compounds in the

476

same crystal structure type45, 46 and the Cs one in the other,47 while there is no report for A =

477

Na. Apparently, the size matching between A and Mg ions dictates the stability of the double

478

carbonates both anhydrous and hydrated forms. Therefore, although K2Mg(CO3)2 is reported

479

to form during CO2 absorption,32 its formation may not be as strongly favorable as

480

Na2Mg(CO3)2.

481

The XRD patterns of A2Mg absorbents taken at various stages of CO2 absorption and 22


Page 22 of 38

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482

desorption are shown in Figure 7. In the case of Na2Mg absorbent, the pattern of the pre-

483

treated state before absorption (black line in Figure 7(a)) shows peaks of MgO and Na2CO3.

484

The Na2Mg(CO3)2 seen in the NMR data in Figure 5 is not observed in this pattern probably

485

because of its small quantity and low crystallinity. Upon CO2 absorption (green line), peaks

486

of Na2Mg(CO3)2 appear and those of Na2CO3 decrease in intensity, indicating that the

487

formation of Na2Mg(CO3)2 is responsible for the CO2 absorption of Na2Mg absorbent:

488

MgO + Na2CO3 + CO2––––→ Na2Mg(CO3)2

489

In this pattern, the peak positions of Na2CO3 are shifted from those in the pre-treated state

490

probably because of the unusually large thermal expansion properties of Na2CO3.48 There is

491

no peak that can be assigned to MgCO3 in the XRD of the absorption product, eliminating the

492

possibility of the direct reaction between MgO and CO2 without Na2CO3 involved. In the

493

CO2 desorbed state, after 10 absorption and desorption cycles, the pattern shows peaks of

494

Na2CO3 and MgO (orange line), similar to the pre-treated state although the Na2CO3 peaks

495

are broader, indicating lower crystallinity.

(4)

496

The cases of A = K and Rb show different variations of XRD patterns from the Na2Mg

497

case. Upon CO2 absorption, the peaks of A2CO3 in the starting pre-treated A2Mg absorbents

498

disappear and new peaks appeared (green lines in Figure 7(b) and 7(c)). These peaks could

499

not be identified with any of known related phases, suggestive of new phases of double

500

carbonate. These patterns change into yet different and unidentified patterns upon CO2

501

desorption after 10 cycles of absorption and desorption (orange lines). We found that the last

502

patterns change upon keeping in an ambient condition for as short as 20 min and the changed

503

patterns can be indexed as the hydrated double carbonates A2Mg(CO3)2⋅4H2O (A= K and Rb)

504

(purple lines). Covering the samples with a Mylar film to keep moisture from contacting

505

them prevents the formation of the hydrated double carbonates (Figure S12). Apparently, the 23



506

moisture in the environment induced the formation of the hydrated double salt phases. Since

507

this transformation is done at room temperature for a short period of time, it must be a low

508

energy process which is devoid of any high energy processes such as diffusion of ions. In

509

other words, the formation of hydrated double carbonates in these samples strongly suggests

510

that their pre-forms are present in the CO2 absorbed states enabling their ready transformation.

511

Therefore, the unidentified phases in the XRD patterns of the CO2 absorbed states of these

512

absorbents are likely to be anhydrous double carbonates.

513 514

Figure 7. XRD patterns of A2Mg absorbents at different states of sample treatment. (Color coding)

515

Patterns of absorbents pre-treated at 400 °C in N2 before CO2 absorption (black), absorbents after CO2

516

absorbed at 330 °C for 180 min under 100% CO2 (taken at 330 °C) (green), absorbents after CO2

517

absorbed immediately after 10 cycles of CO2 absorption at 330 °C and desorption at 460 °C under 100% 24


Page 24 of 38

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518

CO2 (orange), and absorbents, which showed the patterns in orange, after exposed to atmosphere for

519

20 min (purple) are shown in different colors. (a) Na2Mg, (b) K2Mg, (c) Rb2Mg and (d) Cs2Mg.

520 521

The series of XRD patterns of Cs2Mg are in good agreement with the report by Liu et al.

522

(Figure 7(d)). In the CO2 absorbed Cs2Mg absorbent (green line), the pattern shows a peak of

523

MgO at 37° and the peaks which Liu et al. attributed to a Cs-containing phase33 and we

524

believe to be a double carbonate of Cs and Mg. It is interesting that the patterns in the pre-

525

treated (black line) and the CO2 desorbed state after 10 cycles (orange line) are featureless.

526

The pattern of the pre-treated state with a broad unassigned peak at ~28° and a small peak at

527

~37° for MgO is again similar with that by Liu et al. Probably, the surface doping occurring

528

during the pre-treatment induces the decomposition of Cs2CO3 extensively to see that there is

529

no Cs2CO3 in this pattern. Considering that Cs2Mg absorbent gains weight after desorption,

530

the CO2 desorbed state after 10 cycles must be composed of Cs2CO3 and MgO.

531

EDS elemental distribution maps taken at various stages of CO2 absorption and desorption

532

on each absorbent give additional insights into the mechanism of CO2 absorption. In order to

533

clearly see light A atoms such as Na and K in the images, we prepared 0.4Na2Mg and

534

0.3K2Mg absorbents with increased A contents (Figure 8). In the cases of heavy A atoms

535

such as Rb and Cs, increasing the A content was not necessary since the distribution of both A

536

and Mg atoms are clearly seen (Figure S13).

537

In the as-prepared 0.4Na2Mg sample, the Na and Mg atoms are evenly distributed all over

538

the sample. The EDS map taken on the CO2 absorbed sample, however, shows that Na atoms

539

are accumulated at some locations while the distribution of Mg atoms is not changed from the

540

as-prepared sample. When CO2 is desorbed from the absorbent, the Na atoms are re-

541

distributed to the entire area. These series of images clearly indicate that Na atoms diffuse in

542

the process of absorption and desorption. The accumulated Na atoms in the CO2 absorbed 25



543

sample can be understood as the Na atoms in Na2Mg(CO3)2 crystals. On the contrary, the

544

EDS maps on 0.3K2Mg absorbent do not show any change in the distribution of K with

545

respect to that of Mg, suggesting retarded diffusion of K atoms during the CO2 absorption

546

and desorption. The series of images on Rb2Mg and Cs2Mg are similar to those of 0.3K2Mg

547

(Figure S13), indicating the lack of diffusion of Rb or Cs during absorption and desorption of

548

CO2 on them.

549

There are two factors that make the absorption behavior of Na2Mg different from the

550

others. Because Na is lighter than the other A elements, its diffusion is expected to be faster

551

than the other A elements. Therefore, if there is any process which involves diffusion of A

552

atoms, Na2Mg is the most likely material that can show such an effect. The other factor to

553

consider is the different phase stabilities of the double carbonates that are formed as a result

554

of CO2 absorption. As seen before, the formation of Na2Mg(CO3)2 is a thermodynamically

555

favorable process and, thus, it is likely to grow into large crystals under the right conditions.

556

Soon after the crystals start to grow, the regions near the crystals will be depleted of Na2CO3

557

necessitating diffusion of Na2CO3 from regions away from them, resulting in the

558

accumulation of Na atoms in some locations. In the case of other A2Mg absorbents, the

559

formation of anhydrous A2Mg(CO3)2 phases is not as favorable as Na2Mg(CO3)2. Even if they

560

are formed during the CO2 absorption, they may not grow into large crystals because the

561

driving force to form them is not as strong as that of Na2Mg(CO3)2. These two factors make

562

the formation of the double carbonate phase increasingly less favorable as the size of A

563

increases. Probably, this is the reason for the gradually decreasing CO2 absorption capacity in

564

the second step with the size of A.

26


Page 26 of 38

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565 566

Figure 8. SEM-EDS elemental mapping images of (a) 0.4Na2Mg and (b) 0.3K2Mg absorbents. As-

567

prepared samples are ball-milled physical mixtures between Mg5(CO3)4(OH)2⋅4H2O and A2CO3

568

heated at 450 °C for 4 h in air, samples for after carbonation are prepared by heating the as-prepared

569

samples at 330 °C under 100% CO2 for 60 min and samples after decomposition are obtained after 10

570

cycles of absorption (330 °C, 100% CO2) and desorption (400 °C, 100% N2).

571 572

3.3. Proposed mechanism

573 574

Based on the results discussed above, we propose the CO2 absorption and desorption

575

mechanisms on A2Mg absorbents as in the followings (Figure 9). During the mixing of

576

A2CO3 and MgO to form A2Mg absorbents, a part of A2CO3 is decomposed into A2O and

577

CO2. A2O in A2Mg absorbent exists as a surface dopant on MgO forming a basic site. The

578

formation of double carbonates detected in the 23Na and 39K NMR spectra also suggests that

579

inter-diffusion of A and Mg occurs between A2CO3 and MgO. The extent of this reaction

580

increases as the size of A increases probably because of the increasing basicity of A2O.

581

Therefore, the amount of A2O present in the pre-treated A2Mg samples increases as the size

582

of A increases. Pre-treatment of heating at < 400 °C under N2 facilitates this process inducing 27



583

more basic sites.

584

The first step CO2 absorption takes place on these basic sites forming A2CO3 (reaction (3)).

585

The extent of this reaction is governed by the amount of the basic sites generated during the

586

pre-treatment step, explaining the growing proportion of the first step absorption to the total

587

CO2 absorption as the size of A increases. A part of A2CO3 formed in this step is segregated

588

from the MgO surface and becomes stable A2CO3 that does not decompose at temperatures

589

below 500 °C, explaining the incomplete desorption in Rb2Mg and Cs2Mg (Figure 1 and 3).

590

That Na2Mg shows a much smaller first step and Li2CO3 does not function as a promoter can

591

be explained by the tendency of generating the basic sites across the series of alkali metal

592

ions.

593

After the CO2 absorption on the basic sites is saturated, the second step reaction ensues. In

594

this step, CO2 absorption occurs through the formation of double carbonate between A and

595

Mg ions. A2CO3 formed in the first step absorption or present unchanged from the reagent is

596

reacted with MgO and CO2. Because this reaction involves the diffusion of A2CO3 on the

597

MgO surface, it is bound to be slow and the rates reflect the mass of A, explaining the

598

observed order of absorption rate of Na2Mg > K2Mg > Rb2Mg. Cs2Mg stands out from this

599

trend probably because the Cs2CO3 formed during the first step reaction is more active than

600

that carried over from the reagent state.

28


Page 28 of 38

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601 602

Figure 9. The proposed mechanism of CO2 absorption and desorption on A2Mg absorbents.

603 604

As mentioned in the introduction, there are precedent works on the present system.

605

However, none of them report on the existence of the two-step mechanism as in the present

606

work. A probable explanation is that the A2CO3/MgO ratio and the method of sample

607

preparation are important to differentiate the two steps.

608

In the work on K2CO3-promoted MgO absorbents by Xiao et al., the smallest K/Mg ratio is

609

0.45, much larger than 0.20 of the present work. Moreover, they synthesized samples by

610

impregnating K2CO3 on MgO in water.32 while our samples were synthesized by physical

611

mixing. In light that the first step CO2 absorption is due to the basic sites, it seems that our

612

physical mixing method induces more intimate mixing of the reagents than impregnation

613

method. Also the large amount of K2CO3 hinders the observation of the first step making the

614

differentiation between the first and second steps difficult. Our data on the time variation of

615

the weight of 0.2A2Mg absorbents show that the differentiation between the first and second

616

steps is not as sharp as in A2Mg absorbent, supporting this possibility (Figure S1). On the 29



617

contrary, the time variation of sample weight during CO2 absorption on Cs2CO3-promoted

618

MgO absorbent in the paper by Liu et al.33 looks similar to our data on Cs2Mg in Figure 1(b)

619

although they did not comment on it probably because the distinction between the first and

620

second steps in their plot is not as clear as ours. In addition, their sample contains more

621

Cs2CO3 (15 mol%) than ours (10 mol%) and their sample was prepared by the impregnation

622

method in water, both of which may hinder the observation of the two-step mechanism.

623 624

4. Conclusion

625 626

In this study, we have elucidated the detailed mechanism of CO2 absorption of A2CO3-

627

promoted MgO absorbents. All alkali carbonates except for Li2CO3 can promote the CO2

628

absorption of MgO at the intermediate temperature range of 300-500 °C. In all of the cases,

629

CO2 absorption occurs in two steps. The first step is a fast process involving the carbonation

630

of the highly basic sites generated by the action of A2CO3 on the MgO surface. The second

631

step involves the formation of double carbonate between A and Mg atoms. It is a slow

632

process because the formation of double carbonate requires diffusion of A2CO3. Because both

633

steps are influenced by the size of A, the four A2CO3-promoted MgO absorbents show

634

systematic variations in their CO2 absorption properties. The absorption rate and capacity of

635

the first step are mainly governed by the basicity of the base sites and, thus, increase as the

636

size of A increases. Rb2CO3- and Cs2CO3-promoted absorbents show very high absorption

637

rates and the ability to absorb CO2 even when CO2 concentration is as low as 3%. However,

638

their carbonation products are so stable making the CO2 desorption from them hard. For

639

cyclic operations by temperature swing under a 100% CO2 condition, these absorbents are not

640

suitable. Na2CO3- and K2CO3-promoted absorbents can undergo stable cycling by 30


Page 30 of 38

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


641

temperature swing under 100% CO2 after losing parts of the capacity of the first cycle.

642

However, the cycling properties of these two absorbents have different origins. In Na2CO3-

643

promoted absorbent, it is mainly from the second step. Its first step is slow and lasts for a

644

short period of time. Because the diffusion of Na or Na2CO3 is involved in the second step,

645

the absorption kinetics of this absorbent during cycling is bound to be slow. On the other

646

hand, the cycling properties of K2CO3-promoted absorbent are originated from the first step,

647

showing a relatively fast rate. Collectively, our results indicate that K2CO3-promoted MgO

648

absorbent appears to be the most promising for practical applications. There are a number of

649

issues to be resolved before its real application, however. Since the absorption properties of

650

the first step arise from the interface between K2CO3 and MgO, methods to increase the

651

contact area and to stabilize such interfaces need to be developed.

652 653

Supporting Information

654 655 656

Brief statement in non-sentence format listing the contents of the material supplied as supporting information.

657 658

Acknowledgments

659 660

This work was supported by the Korea CCS R&D Center (Korea CCS 2020 project) grant

661

funded by the Korea government (Ministry of Science, ICT & Future Planning) in 2017

662

(KCRC-2014M1A8A1049257).

663 664

5. Reference

665 31



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Yue-qing,

Z.;

Arnold,

A.,

Synthesis

and

Crystal

781 782 783 784 785 786 787 788 789 790 791 792 793 794 795 796 797 798 799 800 36


Structure

of

Dicesium

Trans-

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

801


TOC Graphic

802

37



224x137mm (150 x 150 DPI)


Page 38 of 38

Promoted MgO (A= Na, K, Rb and Cs)

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