The Role of Water during CO2 Adsorption by Ca-Based Sorbents at

Subscriber access provided by READING UNIV

Article

On the role of water during CO2 adsorption by Ca-based sorbents at high temperature Anna Lind, Knut Thorshaug, Kari Anne Andreassen, Richard Blom, and Bjørnar Arstad Ind. Eng. Chem. Res., Just Accepted Manuscript • DOI: 10.1021/acs.iecr.7b04052 • Publication Date (Web): 09 Feb 2018 Downloaded from http://pubs.acs.org on February 9, 2018

Just Accepted “Just Accepted” manuscripts have been peer-reviewed and accepted for publication. They are posted online prior to technical editing, formatting for publication and author proofing. The American Chemical Society provides “Just Accepted” as a service to the research community to expedite the dissemination of scientific material as soon as possible after acceptance. “Just Accepted” manuscripts appear in full in PDF format accompanied by an HTML abstract. “Just Accepted” manuscripts have been fully peer reviewed, but should not be considered the official version of record. They are citable by the Digital Object Identifier (DOI®). “Just Accepted” is an optional service offered to authors. Therefore, the “Just Accepted” Web site may not include all articles that will be published in the journal. After a manuscript is technically edited and formatted, it will be removed from the “Just Accepted” Web site and published as an ASAP article. Note that technical editing may introduce minor changes to the manuscript text and/or graphics which could affect content, and all legal disclaimers and ethical guidelines that apply to the journal pertain. ACS cannot be held responsible for errors or consequences arising from the use of information contained in these “Just Accepted” manuscripts.

Industrial & Engineering Chemistry Research is published by the American Chemical Society. 1155 Sixteenth Street N.W., Washington, DC 20036 Published by American Chemical Society. Copyright © American Chemical Society. However, no copyright claim is made to original U.S. Government works, or works produced by employees of any Commonwealth realm Crown government in the course of their duties.

Page 1 of 28 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

Industrial & Engineering Chemistry Research

On the role of water during CO2 adsorption by Ca-based sorbents at high temperature.

Anna Lind, Knut Thorshaug, Kari Anne Andreassen, Richard Blom, and Bjørnar Arstad*

SINTEF Industry Forskningsveien 1 0373 Oslo Norway E-mail: [email protected]

1 ACS Paragon Plus Environment

Industrial & Engineering Chemistry Research 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

Abstract Reactions of CaO, MgO and decarbonated dolomite (CaOMgO) with CO2 and added water has been studied with the goal of understanding fundamental issues related to these material's performance as CO2 sorbents. We used a fixed bed reactor, in-situ XRD, and DRIFTS to monitor the extent- and kinetics of carbonation, surface reactions and performance loss during repetitive adsorption-desorption cycles at industrial relevant conditions. From reactor and in-situ XRD experiments, we found that water is essential to reach high carbonation levels (solid conversion >40%) of CaO and CaOMgO which is in contrast to a situation where only a small fraction (< 10%) of the capacity is used. Water has a more pronounced effect when applying CaOMgO as sorbent compared to CaO, both when considering solid conversions and carbonation rates. DRIFTS shows that water together with CO2 do in fact react at the MgO surface into carbonates species, but without forming any bulk MgCO3. Furthermore, H2CO3 may be important for exploiting CaO and CaOMgO materials since hydrogen carbonate is observed as a surface species only during reactions with water.


Page 2 of 28

Page 3 of 28 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


Introduction With the growing concern of anthropogenic CO2 emissions and its consequences on the earth's climate, major efforts have been made to head for a low-emission and more sustainable future.1 In addition to increasing the deployment of technologies providing energy from renewable sources, it is accepted that a de-carbonization of power production, transport and industrial processes with subsequent underground storage of CO2 (carbon capture and storage, CCS) is important to meet the necessary and ambitious goals.2 At present, de-carbonization of large power sources by aqueous amine solvent based absorption processes have been developed the most and are likely the type of CO2 removal processes that will be the preferred one at short- to mid-term time scale.3 However, at present other technologies with potentially lower energy penalties and/or environmental impacts are being explored for de-carbonization of power production and/or chemical production. Membrane technologies are able to separate either H2 or CO2 from the gas streams4,5 and solid sorbents can be used to remove CO2 in-situ in the reactor in so-called Sorption-Enhanced Reactions (SER) to produce pure hydrogen6 in either steam reforming7 or water- gas shift processes.8 When CO2 is adsorbed by the solid sorbent, the reaction is driven towards the products as explained by the principle of Le'Chatelier. Another application aimed for CO2 capture by solid sorbents at elevated temperatures emerged almost 20 years ago: The socalled Ca-Looping (Ca-L) process, a post combustion CO2 capture technology.9,10 The Ca-L technology is a potentially low cost 2nd generation technology based on the use of CaO as a regenerable sorbent through carbonation/calcination cycles.11 Recently, dolomite (or more precisely its decarbonated form) has also been evaluated as a CO2 sorbent in the Ca-L technology and appears to be an advantageous alternative. The authors claim that ".... MgO grains in the decomposed dolomite are resistant to sintering under severe calcination conditions and segregate from CaO acting as a thermally stable support which mitigates the multicyclic loss of CaO conversion".12 In both cases (SER and Ca-L) the main reaction is that CaO reacts with CO2 to CaCO3 as described in Reaction 1: CaO(s) + CO2(g) CaCO3(s) ∆Hr,298K = -178 kJ/mol

(1)

Reaction conditions like PCO2, Ptot, and temperature will control the equilibrium. At 1 atm. pressure and 500-600 °C the carbonate is strongly favored, but above 900 °C Reaction 1 is driven 3 ACS Paragon Plus Environment


to the left even in a pure CO2 atmosphere.13 Due to the potential large scale uses of CaO and dolomite as described above, these materials have received much attention over the last decades in terms of stability and performance with prolonged use. Especially CaO derived from Limestone has been extensively tested for applications in the Ca-L process and mathematical relations for residual CO2 capacity after extended use have been derived.14 However, even with extensive work on CaO based materials for applications as sorbents, the influence and role of important co-reactants such as water is less understood.

In the present paper we report a study aimed to assess the effects of water when CO2 (w/wo cofed water) reacts with CaO, MgO, and CaOMgO (calcined dolomite). The basis for this study is a series of fixed bed reactor tests where CO2 (w/wo water) reacts with CaO and calcined dolomite at 600 °C at different process condition. To better understand these observations, we applied insitu X-Ray Diffraction (XRD) and Diffuse Reflectance Infrared Fourier Transform Spectroscopy (DRIFTS) to obtain information on bulk changes and surface chemistry of these materials. The conditions we have applied are similar to those used in hydrogen production via sorption enhanced processes and the Ca-L post-combustion CO2 capture technologies.


Page 4 of 28

Page 5 of 28 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


Experimental Section Materials: The materials studied in this work are MgO, CaO, and the natural mineral dolomite (CaMg(CO3)2) from Seljelid, Norway. The dolomite was calcined for 12 hours at 900°C in 10% H2 in order to obtain the oxide version of the mineral that can react with CO2. H2 was used to reduce S-compounds in the natural mineral. Calcination transforms dolomite into an oxide form that can be termed CaOMgO, or calcined dolomite (there is close to a 1:1 ratio of Mg and Ca). Particles were pressed and sieved to the 0.02-0.9 mm fraction.

Fixed bed reactor testing: In order to assess the CO2 uptake/desorption capacity and kinetics, the materials were in a fixed bed reactor for their adsorption capacities through 30 or 90 adsorptiondesorption cycles. A connected mass-spectrometer measured signals from relevant ions and integration for quantification were carried out by a routine developed in Microsoft Excel. The reactor test rig consisted of two parallel tubes, each with an inner diameter of 1 cm. The powder bed was 4 cm high and the amount sorbent was somewhat below 2.5 g. One of the reactor tubes was a dummy unit containing a 4 cm high quartz bed. Temperatures in the sorbent bed were measured at four points along the height. The gas flows were 50 ml/min through both reactors, one stream was pure N2 and the other was 10% CO2 in N2 that could be saturated with water at predetermined temperatures. We chose to apply saturation ca. 11°C, and ca. 35 °C i.e. app. 1.2% and 5.5% steam, respectively. An in-house developed Labview program controlled the whole sequence automatically. Before each multicycle experiment the sorbent was kept at 600 °C for ten minutes to decompose minor components of hydroxyls and carbonates before the adsorptiondesorption cycles commenced. After a predetermined adsorption time with the CO2 containing gas (10 or 120 minutes) an automatic valve switched pure N2 over the sorbent for regeneration and the bed was heated to 850 °C and kept at this temperature for an hour. After desorption for 1h at 850 °C and cooling of the bed to the sorption temperature, a cycle had passed. The choice of abovementioned reactions conditions were done based on experiments using different bed heights, reaction times, and gas flows to avoid major unfortunate transport limitations. The conversion of sorbent was evaluated and defined as moles CO2 adsorbed pr. mol CaO unit in the sample and is based on the following estimate of total capacity. Since the molecular weight of CaOMgO (calcined dolomite) is 96 g/mol CaO constitutes 100%* 56/96=71% of the weight. 1 gram CaOMgO sorbent contains therefore around 710 mg CaO, which corresponds to 12.68 5 ACS Paragon Plus Environment


mmol and is also the amount CO2 possible to adsorb if given enough time. Similarly, 1 gram CaO is able to react with 17.85 mmol CO2 and form CaCO3. X-Ray Diffraction (XRD): XRD measurements were performed using a PANalytical Empyrean diffractometer. The system is equipped with a PIXcel3D solid state detector. The measurements were carried out in reflection geometry using CuKα radiation (λ= 1.541874 Å) and a step size of 0.013 degrees. An Anton Paar XRK 900 in-situ high-temperature (25-900°C) and high pressure (1-10 bar) reactor cell coupled to an automated gas switching system was used for the in-situ XRD experiments. The pressure was 1 atm. and the feed gas employed was a 10% CO2 in N2 mix saturated with H2O at 10-11°C, i.e. app. 1% water, for a total flow 100 ml/min. Ca. 100-120 mg sample powder was used for each experiment. For analyses of the in-situ XRD data the reflection at 2θ=29.3° for CaCO3 and at 2θ=37.2° for CaO were used. The CaO crystallite size was calculated based on a line profile analysis performed in the PANalytical HighscorePlus software. The crystallite sizes are calculated using the (111), (200) and (220) reflections at 32.2°, 37.2° and 53.8° 2θ.

Diffuse Reflectance Infrared Fourier Transform Spectroscopy (DRIFTS): The spectra were collected on a Bruker Vertex 70 instrument equipped with a MCT detector, high temperature reaction chamber (HVC-DRP-3), Praying Mantis, and a temperature controller (ATC-024-2). The latter three parts were from Harrick Scientific Products. Each reported spectrum is the average of 128 scans collected at 2 cm-1 resolution. The sample spectrum just prior to gas dosing was used as background spectrum. CO2 (10% diluted in N2), and N2 (99.999% purity) were used as received from Yara Praxair. CO2, H2O or a mixture of CO2 in H2O (all balanced in N2) was dosed into the DRIFTS cell from a small pressurized (2-3 bar) steel container kept at room temperature. Before use, the water was deionized and degassed by boiling for 10 minutes followed by Arbubbling for another 10 minutes. Exact water content was difficult to control so the DRIFTS data are only applied for qualitatively analyses. In a typical run, the solid sample was loaded into the sample cup, the cell was mounted, and its position adjusted in the z-direction. The temperature was raised to 800 °C under N2 flow and kept at 800 °C under high vacuum (p < 5 × 10-3 mbar) for 30 min before adjusted to reaction conditions.


Page 6 of 28

Page 7 of 28 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


Results and Discussion CO2 capacity measurements in fixed bed. The basis for our discussion of the title reaction is data from measured CO2 sorption capacities of CaO and CaOMgO, and how these depend on process conditions. Repeated, multiple CO2 capacity measurements of these materials were measured using a fixed bed reactor, and two examples of data sets are presented in Figure 1.

Figure 1. Bar-graph that shows CO2 uptake in each sorption-desorption cycle (one bar for each cycle) in an experiment with long sorption times (120 minutes, (top)) and short sorption times (10 minutes, (bottom)) using CaOMgO as sorbent. The water content was 1.2 % in each experiment.



Page 8 of 28

The sorption time was either 10 (short) or 120 (long) minutes (w/wo water) at 600 °C followed by a regeneration period for 1 hour at 850 °C. A sorption step followed by a regeneration step is termed a cycle. The multi cycle experiments using short sorption times included 90 sorptiondesorption cycles while the experiment using long sorption cycles consisted of 30 sorptiondesorption cycles for a total of 15 and 60 hours sorption times respectively. The total accumulated CO2 uptakes in mmol/g sorbent through such a set of cycles at certain conditions is reported in Table 1. The numbers in Table 1 are calculated by adding up the measured CO2 uptake from each sorption-desorption cycle (each bar in the plots). From parallel runs, we estimate the uncertainties to be about ±3 mmol/g sorbent.

An immediate observation from the bar-graphs in Figure 1 is that for long adsorption times, i.e. relative high carbonation levels, the capacity declined for each cycle (except from #1 to #2), and when using shorter contact times, i.e. low carbonation levels, there is no sign of a declining trend in cyclic capacity. These trends are also observed for all other fixed bed experiments. Table 1. Accumulated CO2 uptake in mmol/g sorbent in the different multicycle reactor test runs. Short cycles are experiments with 10 minutes sorption (90 cycles for a total of 15 hours adsorption time) while long cycles are experiments with 120 minutes sorption time (30 cycles for a total of 60 hours absorption time).

CaO

CaOMgO

(Acc. mmol/g sorbent)

(Acc. mmol/g sorbent)

Short cycles wo. steam

80

81

Short cycles with 1.2% steam

78

80

Short cycles with 5.5% steam

79

82

Long cycles wo. steam

110

98

Long cycles,1.2% steam

113

160

Long cycles, 5.5 % steam

94

152

Experiment type

There are several interesting conclusions that can be extracted from Table 1. For the short sorption cycle experiments, the accumulated amount of sorbed CO2 is very similar for the two solids, and it is independent on water content in the reacting CO2 gas. In these experiments, the 8 ACS Paragon Plus Environment

Page 9 of 28 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


CaO conversion levels are below 10%. However, at long sorption time differences between the materials appears: Using dry CO2 we see a larger amount sorbed by CaO than CaOMgO but adding steam does not improve the accumulated CO2 sorption for CaO, it is actually a decline at the highest steam level. For CaOMgO however, addition of 1.2 % steam increases the accumulated sorbed CO2 by about 60 %, but further increase in steam does also in this case appear to be detrimental. Since the short sorption times, i.e. low carbonation levels, resemble conditions experienced in an industrial plant, it appears that from a capacity point of view there are small differences between CaO and CaOMgO. The data presented in Table 1 is in mmol CO2 sorbed pr. gram sorbent. If one considers that CaOMgO is on a mole basis approximately 50% MgO one can see that the exploitation of CaO is much higher for this material than for pure CaO, i.e. the extent of carbonation is higher using CaOMgO.

In-situ XRD. In order to approach explanations for the data reported in Table 1 we applied in-situ XRD during three sorption-desorption cycles. In these experiments sorption was performed at two temperatures, 600 °C or 650 °C, while desorption took place at 850 °C. Due to the configuration of the XRD instrument and the in-situ cell these experiments resemble the long cycle experiments with or without 1.2% steam. An illustration of the experimental protocol is shown in Figure 2.

Figure 2. An illustration of the reaction protocol for the in-situ XRD experiments.



Figure 3a shows an in-situ XRD data set of CaOMgO at 600 °C during sorption of dry CO2, and Figure 3b shows a zoom-in over the CaCO3 and CaO peaks that we integrate to follow the carbonation reaction. In-situ XRD data for all the experiments is provided in the Supporting Information. A specific comparison between the XRD data recorded at 850 °C and 600 °C just before the reactive gases were switched on is also included in the supporting information to show that there are no significant differences between these data sets. During sorption, we can follow how the intensity of the CaO peaks decline as the CaO reacts with CO2 to form CaCO3. Consequently, the intensity of the CaCO3 peaks increase during the sorption process. As expected, no changes were observed in the MgO peaks, indicating that MgO does not react to form MgCO3 at these conditions, nor become influenced enough at a sufficient scale to give measureable changes of long range effects. The same general trend is seen for all experiments, but there are variations in the extent and rate of the reaction of CaO, as well as whether the extent of carbonation is constant over the three cycles or not.


Page 10 of 28

Page 11 of 28 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


Figure 3. In-situ XRD data of dry CO2 sorption on CaOMgO at 600 °C. a) Data set in the 2 theta(θ) 25-55 degrees showing characteristic peaks for all three phases present in the sample; b) A zoom in over the main peaks of CaCO3 at 29.3 degrees and CaO at 37.2 degrees that are used to calculate the integrated peak area plotted in Figure 4. The symbols represent the following phases: * CaO, ● CaCO3 and ◊ MgO.

In Figure 4 the integrated peak areas of the reflections of CaCO3 and CaO at 29.3 and 37.2 degrees 2θ, respectively, are plotted in order to follow the reactions. The peak areas have been normalized against the area of the highest peak of unreacted CaO in each sample in order to be able to compare the extent of the carbonation reaction of the different samples. Filled symbols are for CaCO3 while open symbols are for CaO. Top row is for CaOMgO and left column is for the dry reactions. Blue symbols are for 650 °C. 11 ACS Paragon Plus Environment


Figure 4. Integrated peak area of the main reflections of CaCO3 and CaO at 29.3 and 37.2 degrees 2 theta, respectively, plotted for all experimental conditions in order to follow the CaO and CaCO3 evolution over the three sorption/desorption cycles.

A main feature is that the extent of carbonation is higher for CaOMgO (top row) compared to CaO (bottom row) on the applied time scale. The initial sorption rate can be seen to be qualitatively similar in all experiments up to about 0.15 of total peak area, with some experiments having an almost similar sorption rate up to maximum, e.g. top right for CaOMgO. Water addition promotes especially the experiments at 600 °C resulting in a larger extent of carbonation while the effect of water is less pronounced at 650 °C. This indicates that water promotes sorption kinetics even without being observed as part of some bulk species like Ca(OH)2, which is not observed in the XRD data. For all experiments, the CaCO3 build up curve is steeper and have a sharper bend at the highest temperature, which is a measure of the competition between surface carbonation and deep intra carbonation of the crystallites. For CaO the extent of reaction 12 ACS Paragon Plus Environment

Page 12 of 28

Page 13 of 28 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


appears to be constant in the three sorption/desorption cycles at 600 °C for both dry and wet CO2, while in the experiments at 650 °C the extent of carbonation declines for each cycle and indicates that deactivating mechanisms like e.g. sintering is more rapid at higher temperatures. For CaOMgO, it is only in the case of dry CO2 adsorption at 600 °C we see a slight decline in the extent of carbonation over the three cycles. There are previous reports in the literature on in-situ calcination of both dolomite15 and limestone.16 These studies show that during calcination in N2, which is similar to our first step before starting the sorption cycles, CaO derived from dolomite (MgCO3CaCO3) will have significantly smaller crystallites than the CaO derived from limestone (CaCO3). It is suggested that during calcination of limestone CaO crystallites grow through aggregation of newly formed nanocrystals due to attractive surface forces between them.17 In CaOMgO, however, these forces are weakened by the MgO grains that are interpositioned between the CaO nanocrystals. When calculating the crystallite size we see the same effect, the calcined dolomite samples have a smaller crystallite size than the calcined CaO samples. We also observe a difference in the initial crystallite size in the 600 °C experiments and the 650 °C experiments on CaO. This is due to the fact that the experiments were performed on two different batches of calcined CaO samples. However, after the three carbonation/decarbonation cycles the CaO crystallite size reaches very similar values, which is to be expected as they are re-calcined at the same conditions during the decarbonation cycles. After the three carbonation/decarbonation cycles, the calcined dolomite samples have a CaO crystallite size in the range 70-100 nm, while the crystallite size of the pure CaO samples lie in the range 140-160 nm. CaO crystallite size plots for all the samples are provided in the supporting information. Since a smaller crystallite size generally means a higher reactivity of the material, this would explain why we see a much higher CaO reactivity in the calcined dolomite samples compared to the CaO samples.

As mentioned, the data presented in Figure 4 resembles mostly the fixed bed tests using long cycles without and with 1.2 % steam. A qualitative analysis is possible. By taking into consideration, that CaOMgO has about half the amount CaO as pure CaO one can see that for the 600 °C series in the left column the extent of carbonation is about twice for CaOMgO compared to CaO. This translates into approximately the same total accumulated capacities as given in 13 ACS Paragon Plus Environment


Page 14 of 28

Table 1 for long cycles without added steam. A similar exercise on the right column in Figure 4 gives qualitatively results that are in agreement with the reactor tests.

For the short cycle reactor tests the comparisons are not so straightforward, but by considering that the carbonation level is less than 10 % in these experiments and that this corresponds to short reaction time in the XRD curves we can see that the initial formation rate of CaCO3 is quite similar up to about 15 % of the normalized area. After this extent of carbonation, some experimental 600 °C curves diverge from each other.

In-situ DRIFTS. To obtain complementary information on the chemistry of the title reactions we applied DRIFTS to study increased dosing of CO2 on the materials, including MgO. We only present selected spectra, but all spectra and a table of band assignments compiled from the literature are provided in the Supporting Information.18-23 The species we have identified are sketched below in Scheme 1. Scheme 1. Carbonate species observed by in-situ DRIFTS in this work O

O O M

O

O M

Unidentate

H

O

O

Bidentate

O M

OH

Hydrogen carbonate

Observed at low and high pressure

O M

O O

Formate

O M

O M

Bridging carbonate

Observed at high pressure only

In Figure 5, spectra collected at low CO2 loadings (p < 5 mbar) and 600 °C from the carbonate region is presented for all three materials together with curve fitting results.


Page 15 of 28 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


Figure 5. in-situ DRIFTS spectra collected at 600 °C during the exposure of CaO (upper panel), MgO (middle panel), and CaOMgO (lower panel) to CO2 (upper spectrum in each panel) and CO2 + H2O (lower spectrum in each panel), respectively.



During CO2 reactions at dry conditions we assign the observed major bands to be unidentate carbonate (1535 cm-1, 1369 cm-1, and 1041 cm-1), and the minor bands to bidentate carbonate (1621 cm-1, 1331 cm-1, and 1214 cm-1). The band intensities suggest that the concentration of unidentate carbonate is higher than the concentration of bidentate carbonate. Our observations and interpretations compares good with reported unidentate carbonate23,25 and bicarbonate22 formation when CaO is exposed to CO2. When CaO was exposed to a mixture of CO2 and H2O, the major specie was again unidentate carbonate (1551 cm-1 and 1383 cm-1), the same specie as observed under dry CO2. In contrast to the observations made during dry CO2, bidentate carbonate was not observed. Instead, weak bands assigned to hydrogen carbonate were detected (1482 cm-1 and 1230 cm-1), which show that the surface chemistry is influenced by the presence of H2O. DRIFTS spectra of MgO's reactions are shown in the middle panel in Figure 5. Exposure to dry CO2 does not indicate any detectable formation of IR-active surface species, however, we observed bands that can be assigned to bidentate carbonate (1613 cm-1 and 1302 cm-1) together with weaker bands assigned to hydrogen carbonate (1655 cm-1, 1481 cm-1, and 1240 cm-1 (very weak)), and unidentate carbonate (1373 cm-1 (weak)) when MgO was exposed to a mixture of CO2 and H2O. The unidentate specie should also show a weak band in the range 1560-1515 cm-1, but overlap with the strong band assigned to bidentate carbonate at 1613 cm-1. The band observed at 1655 cm-1 may also be partially due to H2O adsorbed on MgO, but H2O adsorption alone cannot explain the observed bands as no appreciable reaction was observed to occur between MgO and H2O in the absence of CO2, as shown in Figure 6 (vide infra). The presence of observable surface species at 600 °C in the system MgO + CO2 + H2O is highly interesting. XRD does not shows any formation of a crystalline bulk phase, and neither pure CO2 nor pure H2O interact with MgO to any significant degree. According to thermodynamics, MgO, CO2, and H2O does not react to form any new stable bulk phase at these conditions, yet we observe carbonate formation. It is important to note that the spectra are not representative for the bulk phase because DRIFTS is a surface sensitive technique. We speculate that the observed carbonates are due to metastable surface species and dynamic gas-solid interactions, and the data show that both H2O and CO2 must be included in any reaction mechanisms leading to carbonate formation on the MgO surface. It is known that a rather slow reaction takes place between CO2 and water to form 16 ACS Paragon Plus Environment

Page 16 of 28

Page 17 of 28 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


H2CO3. We speculate that the MgO surface might catalyze this reaction but the strong ionic bonds in the material are not easily broken up by CO2 at the chosen conditions to form a crystallographic phase (MgCO3). The spectra collected using CaOMgO are shown in the lower panel in Figure 5. At dry conditions a strong band at 1558 cm-1 corresponds to the band assigned to unidentate carbonate in CaO, and again we expect overlapping bands to be the reason for the lack of an observable band at ca 1370 cm-1. We did not detect any hydrogen carbonate. As seen, the bands assigned to unidentate carbonate (1550 cm-1 and 1360 cm-1) dominated the spectrum when the material was exposed to dry CO2. Under CO2 and H2O, the major bands assigned to unidentate carbonate (1558 cm-1 and 1333 cm-1) dominate the spectra, in consistency with CaO being the main interaction site – as observed under dry conditions. Additional bands assigned to hydrogen carbonate (1633 cm-1, 1333 cm-1, and 1308 cm-1) were detected, which may arise from interactions at either CaO or MgO sites. If both sites are involved, we should also observe two bands close to 1610-1620 cm-1, but we only detect the band at 1633 cm-1. Given that we observed hydrogen carbonate on CaO and not MgO, we find it reasonable to suggest that the formation of hydrogen carbonate occurs mainly at CaO sites in CaOMgO. Table 2 below is a summary of the observed species on the various solids during exposure to dry or wet CO2. Table 2. Carbonate species observed at low CO2 loadings. The structures are sketched in Scheme 1.

Gas phase CO2

Solid Amount CaO

MgO

CaOMgO

Major

Unidentate carbonate

-


Minor

Bidentate carbonate

-

-

Major


Bidentate carbonate

Minor

Hydrogen carbonate

CO2 + H2 O

Hydrogen carbonate Unidentate carbonate


Unidentate carbonate Hydrogen carbonate -


Under dry CO2 conditions, the interactions with CaO and CaOMgO mainly occur by unidentate carbonate coordination whereas no interaction with MgO was detected. Thus, a reasonable interpretation of the data is to suggest that the initial surface reactions between CaOMgO and dry CO2 occur by formation of unidentate carbonate on CaO sites. In the presence of both CO2 and H2O, we observed unidentate carbonate, bidentate, and unidentate carbonate species on CaO, MgO, and CaOMgO, respectively. Additionally, the initial surface reactions on CaOMgO also gave rise to hydrogen carbonate species, which also most likely occurs on CaO sites. Thus, when we compare the observed species that form in the absence and presence of H2O, it is clear that H2O significantly influences on the reactions that take place on the oxide surfaces. At higher gas loadings (5 mbar < p(tot) < 0.5 bar, shown in the Supporting Information) all spectra show a transition to mainly bridging and bidentate carbonate coordination.

Similar to the above we followed the OH band region, as shown in Figure 6. The background spectrum is the solid oxide just prior to gas dosing, thus bands associated with interacting/reacting surface OH groups point upwards. The spectra show sharp bands in the range 3800-3600 cm-1 and a broad featureless band in the range 3600-3200 cm-1. The sharp bands are associated with isolated OH groups, whereas the broad band is probably due to multi coordinated hydrogenbonded OH groups.24 As seen, the oxides OH-groups interact with CO2, H2O, and their mixture.


Page 18 of 28

Page 19 of 28 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


Figure 6. in-situ DRIFTS low pressure spectra from the OH region of the reaction between CaO (upper panel), MgO (middle panel), or calcined dolomite (lower panel) with CO2 (upper spectrum in each panel) and CO2 + H2O (lower spectrum in each panel).

The band maximum observed in CaO's spectra appears at 3698 cm-1 and in MgO the maximum appears at 3751 cm-1. The band shapes indicate that more than one OH-site is present on the surfaces. All CaOMgO spectra show maxima at 3740 cm-1 and 3700 cm-1, with the 3700 cm-1 band being the most intense. This implies that the OH sites on CaOMgO closely resembles those on CaO whereas the OH sites on MgO in the mixed oxide are perturbed compared to the pure MgO. We also note the presence of an additional maximum at 3672 cm-1 under dry CO2 which is 19 ACS Paragon Plus Environment


not observed for CaO nor MgO. Qualitatively one could argue that the isolated OH-groups (narrow bands) interacting with adsorbed gases are probably located at defect sites along corners and edges on outer surfaces of the crystallites whereas the broad band might arise from OHgroups located at crystallite interfaces and more confined spaces like pores. In these situations, the OH-groups possibly interact with O-atoms from the crystal and/or other OH-groups.

On the role of water. Based upon our presented data we are able to suggest possible roles of water in CO2's reactions with the studied materials, in particular CaO and CaOMgO. First, we note that from our reactor and XRD data we observe that water promotes CO2 uptake such that higher conversion levels of the sorbents can be reached, but that at low conversion it appears that CaCO3 formation rates are rather independent of water's presence. Another finding (DRIFTS) is that hydrogen carbonate as shown in Scheme 1 is only observed when water is present (for all three materials). From these findings CO2 probably has one or two main reaction routes depending on water's presence or not.

1

Reaction of CO2 to surface carbonates

2

Reaction of water and CO2 to H2CO3 with possible further reactions

Formation of H2CO3 from water and CO2 is a well-known reaction and has been studied repeatedly due to its importance, both experimentally and theoretically.25,26 By invoking H2CO3 as a central specie we are in a position where we now have two reactive species leading to various carbonates: CO2 and H2CO3/ HCO3-. Their different reactivities and diffusivities enable analyses with two or more very different adsorption sites. The linear CO2 molecule would probably need an electrophilic site to react to a unidentate or bidentate carbonate. This is in contrast to what can be envisaged for H2CO3 species as this is already a carbonate. Initially, reactive O-atoms on edges and kinks react with CO2 at a rate at least comparable to the formation of H2CO3, but in parallel, H2CO3 is formed to some extent. When most reactive O-atoms are consumed by CO2, less reactive sites must be exploited. In this situation, the H2CO3 species become important as they can attach to an O-vacancy instead. This might be a typical defect or an unsaturated Ca-ion. Such sites can accommodate one of the three O-atoms in the molecule/ion, and are likely not very reactive towards CO2. In addition, one could expect that the concentration 20 ACS Paragon Plus Environment

Page 20 of 28

Page 21 of 28 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


of sites ready for CO2, mainly various O-atoms, are more numerous compared to sites in need for an O-atom.

During adsorption hydrogen carbonate formation could involve surface OH group whereas the unidentate or bidentate carbonate formation occurs mainly via an O-atom part of the solid oxide surface itself. However, MgO + CO2 did not yield any detectable formation of carbonates in the absence of H2O (Figure 5), yet the surface OH-groups are shown to interact with CO2 in both the absence and presence of H2O (Figure 6). Consequently, the formation of hydrogen carbonate is not necessarily related to the involvement of the surface OH groups under our conditions, at least the influence is not significant.

Our observations are in line with Yang et al.'s study on the effects of water on CO2 uptake by various dolomites (different Mg:Ca ratios) and limestone.27 They found that water did not promote significant the kinetics during the fast adsorption step but led to higher conversion before the product layer stage was entered. From SEM analyses they suggested that water improves porosity providing more surface area for the fast kinetic reaction between CaO and CO2. In another work on this topic Yang and Xia found that co-fed steam with CO2 enhanced the carbonation rates and further suggested that the reasons should be attributed to steam catalysis and not formation of Ca(OH)2.27 These authors further suggested from their own, and others work, that adsorbed hydroxyl groups and bicarbonates (termed hydrogen carbonate in our work) could be important intermediates in the steam catalysis during CaO carbonation. This idea is to some extent supported in our work.

Variations in e.g. diffusion rates, sticking probability, surface coverage, translational diffusion, oxide surface dynamics, and rates of reactions will all depend on macroscopic process conditions.



Conclusions In this article, we have combined fixed bed reactor testing with in-situ XRD and DRIFTS analyses at elevated temperatures (600-850 ºC) to study reactions of MgO, CaO and calcined dolomite (CaOMgO) with CO2, in the absence and presence of water vapor. From the fixed bed reactor tests, we observed that at high sorbent conversion (>40%) water has an effect on degree of conversion and kinetics of the carbonation reaction, in contrast to a situation where only a small fraction (< 10%) of the capacity is in use. In the latter case, it appears that water has no effect on any of the materials studied. From both the fixed bed reactor tests and the in-situ XRD experiments it is clear that water has the largest effect on CaOMgO (calcined dolomite) reactions, both when considering sorbent conversion and carbonation rates, and it appears to be an optimal water concentration.

Even if there is no thermodynamically stable MgCO3 phase at our applied conditions, which also is evidenced from our in-situ XRD experiments showing no bulk MgCO3 formation, the DRIFTS data show that water together with CO2 do in fact react at the MgO surface into surface carbonates species. Such species most likely have implications for CO2 sorption in calcined dolomite (CaOMgO). It has been suggested that the MgO micro grains hinder sintering in calcined dolomite with repeated use and that this is a reason for the relative stable cyclic capacity.[12] However, one may also envisage that carbonate formation on MgO surfaces results in precursor carbonate species that rapidly reacts with CaO. This suggestion is supported by observation of faster carbonation of CaOMgO and higher conversion. Hence, formation of H2CO3 is important for exploitation of the full capacity of the sorbent. With this specie one is able to generate, in principle, a carbonate unit reactive towards both Ca and O ions with the effect of potential of using a larger capacity of the sorbent.


Page 22 of 28

Page 23 of 28 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


Supporting Information Plots of all in-situ XRD data (Fig. S1-S8), comparisons of in-situ XRD data at 600 and 850 °C before reactive gases are switched on (Fig. S9-S10), graphs showing all crystallite calculations (Fig.S11-S12), all DRIFTS spectra (Fig. S13-S21), table with relevant FTIR band assignments from literature (Table S1). Text describing DRIFTS data at high CO2 loading.



Acknowledgements This publication was produced with support from the BIGCCS Centre, performed under the Norwegian Research Program Centers for Environment-Friendly Energy Research (FME). The authors acknowledge the following partners for their contributions: Gassco, Shell, Statoil, TOTAL, ENGIE, and the Research Council of Norway (grant no. 193816). Some of the experimental work has been carried out in the iSEWGS project with support from the Research Council of Norway (grant no. 243736). Aud I. Spjelkavik at SINTEF is acknowledged for material preparations.


Page 24 of 28

Page 25 of 28 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


References (1) Schlömer S.; Bruckner, T.; Fulton, L.; Hertwich, E.; McKinnon, A.; Perczyk, D.; Roy, J.; Schaeffer, R.; Sims, R.; Smith, P.; Wiser, R. Annex III: Technology-specific cost and performance parameters. In: Climate Change 2014: Mitigation of Climate Change. Contribution of Working Group III to the Fifth Assessment Report of the Intergovernmental Panel on Climate Change; Edenhofer, O.; Pichs-Madruga, R.; Sokona, Y.; Farahani, E.; Kadner, S.; Seyboth, K.; Adler, A.; Baum, I.; Brunner, S.; Eickemeier, P.; Kriemann, B.; Savolainen, J.; Schlömer, S.; von Stechow, C.; Zwickel T.; Minx, J.C., Eds.; Cambridge University Press, Cambridge, United Kingdom and New York, NY, USA. 2014. (2)

Boot-Handford, M. E.; Abanades, J. C.; Anthony, E. J.; Blunt, M. J.; Brandani, S.; Mac

Dowell, N.; Fernandez, J. R.; Ferrari, M. C.; Gross, R.; Hallett, J. P.; Haszeldine, R. S.; Heptonstall, P.; Lyngfelt, A.; Makuch, Z.; Mangano, E.; Porter, R. T. J.; Pourkashanian, M.; Rochelle, G. T.; Shah, N.; Yao, J. G.; Fennell, P. S. Carbon capture and storage update. Energy Environ. Sci. 2014, 7, 130. (3) Rochelle, G. T. Amine Scrubbing for CO2 Capture. Science 2009, 325, 1652. (4) Peters, T.; Stange, M. S. S.; Sunding, M. F.; Bredesen, R. Stability investigation of microconfigured Pd-Ag membrane modules - Effect of operating temperature and pressure. Int. J. Hydrogen Energy 2015, 40, 3497. (5) He, X.; Lindbråthen, A.; Hägg, M.-B. Pilot testing on fixed-site-carrier membranes for CO2 capture from flue gas. Int. J. Greenhouse Gas Control 2017, 64, 323. (6) Harrison, D. P. Sorption-Enhanced Hydrogen Production: A Review. Ind. Eng. Chem. Res., 2008, 47, 6486. (7) Arstad, B.; Prostak, J.; Blom, R. Continuous hydrogen production by sorption enhanced steam methane reforming (SE-SMR) in a circulating fluidized bed reactor: Sorbent to catalyst ratio dependencies. Chem. Eng. J. 2012, 189-190, 413. (8) Jansen, D.; van Selow, E.; Cobden, P.; Manzolini, G.; Macchi, E.; Gazzani, M.; Blom, R.; Pakdel Henriksen, P.; Beavis, R.; Wright, A. SEWGS Technology is Now Ready for Scale-up! Energy Procedia 2013, 37, 2265. (9) Shimizu, T.; Hirama, T.; Hosoda, H.; Kitano, K.; Inagaki, M.; Tejima, K. A twin fluid-bed reactor for removal of CO2 from combustion processes. Trans IChemE 1999, 77, 62.



Page 26 of 28

(10) Abanades, J. C.; Anthony, E. J.; Wang, J.; Oakey, J. E. Fluidized bed combustion systems integrating CO2 capture with CaO. Environ. Sci. Technol. 2005, 39, 2861. (11) Perejón, A.; Romeo, L. M.; Lara, Y.; Lisbona, P.; Martínez, A.; Valverde, J. M. The Calcium-Looping technology for CO2 capture: On the important roles of energy integration and sorbent behavior. Applied Energy, 2016, 162, 787. (12) Valverde, J. M.; Sanchez-Jimenez, P. E.; Perez-Maqueda, L. A. X Ca-looping for postcombustion CO2 capture: A comparative analysis on the performances of dolomite and limestone. Applied Energy 2015, 138, 202. (13) Berstad, D.; Anantharaman, R.; Blom, R.; Jordal, K.; Arstad, B. NGCC post-combustion CO2 capture with Ca/carbonate looping: Efficiency dependency on sorbent properties, capture unit performance and process efficiency. Int. J. Greenhouse Gas Control 2014, 24, 43. (14) Rodriguez, N.; Alonso, M.; Grasa, G.; Abanades, J. C. Heat requirements in a calciner of CaCO3 integrated in a CO2 capture system using CaO. Chem. Eng. J. 2008, 138, 148. (15) Valverde, J. M.; Perejon, A.; Medina, S.; Perez-Maqueda, L. A. Thermal decomposition of dolomite under CO2: insights from TGA and in situ XRD analysis. Phys. Chem. Chem. Phys. 2015, 17, 30162. (16) Valverde, J. M.; Sanchez-Jimenez, P. E.; Perez-Maqueda, L. A. Limestone Calcination Nearby Equilibrium: Kinetics, CaO Crystal Structure, Sintering and Reactivity. J. Phys. Chem. C 2015, 119, 1623. (17) Valverde, J. M.; Medina, S. Crystallographic transformation of limestone during calcination under CO2. Phys. Chem. Chem. Phys. 2015, 17, 21912. (18) Rethwisch, D. G.; Dumesic, J. A. Effect of metal-oxygen bond strength on properties of oxides. 1. Infrared spectroscopy of adsorbed carbon monoxide and carbon dioxide. Langmuir 1986, 2, 73. (19)

Philipp,

R.;

Fujimoto,

K.

J.

FTIR

spectroscopic

study

of

carbon

dioxide

adsorption/desorption on magnesia/calcium oxide catalysts. Phys. Chem. 1992, 96, 9035. (20) Gruene, P.; Belova, A. G.; Yegulalp, T. M.; Castaldi, M. J.; Farrauto, R. J. Dispersed Calcium Oxide as a Reversible and Efficient CO2−Sorbent at Intermediate Temperatures. Ind. Eng. Chem. Res. 2011, 50, 4042. (21) Jensen, M. B.; Petterson, L. G. M.; Swang, O.; Olsbye, U. CO2 Sorption on MgO and CaO Surfaces: A Comparative Quantum Chemical Cluster Study. J. Phys. Chem. B 2005, 109, 16774. 26 ACS Paragon Plus Environment

Page 27 of 28 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


(22) Zaki, M. I.; Knözinger, H.; Tesche, B.; Mekhemer, G. A. H. Influence of phosphonation and phosphation on surface acid–base and morphological properties of CaO as investigated by in situ FTIR spectroscopy and electron microscopy. J. Colloid Interface Sci. 2006, 303, 9. (23) Cornu, D.; Guesmi, H.; Krafft, J.-M.; Lauron-Pernot, H. Lewis Acido-Basic Interactions between CO2 and MgO Surface: DFT and DRIFT Approaches. J. Phys. Chem. C. 2012, 116, 6645. (24) Fujimori, Y.; Zhao, X.; Shao, X.; Levchenko, S. V.; Nilius, N.; Sterrer, M.; Freund, H.-J. Interaction of Water with the CaO(001) Surface. J. Phys. Chem. C 2016, 120, 5565. (25) Terlouw, J.K.; Lebrilla, C.B.; Schwarz, H. Thermolysis of NH4HCO3—A Simple Route to the Formation of Free Carbonic Acid (H2CO3) in the Gas Phase. Angew. Chem. Int. Ed. Engl. 1987, 26, 354. (26) Lewis, M.; Glaser, R. Synergism of Catalysis and Reaction Center Rehybridization. A Novel Mode of Catalysis in the Hydrolysis of Carbon Dioxide. J. Phys. Chem A. 2003, 107, 6814. (27) Yang, X.; Zhao, L.; Yang, S.; Xia, Y. Investigation of natural CaO-MgO sorbent for CO2 capture. Asia-Pac. J. Chem. Eng. 2013, 8, 906. (28) Yang, S.; Xia, Y. Steam Catalysis in CaO Carbonation under Low Steam Partial Pressure. Ind. Eng. Chem. Res. 2008, 47, 4043.



Table of content graphic


Page 28 of 28

The Role of Water during CO2 Adsorption by Ca-Based Sorbents at

Recommend Documents