Physico–Chemical Processes Limiting CO2 Uptake in Concrete during

Mar 9, 2013 - ABSTRACT: Accelerated curing of fresh concrete using CO2 is a possible ... of CO2 uptake during accelerated carbonation curing of concre...
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Physico−Chemical Processes Limiting CO2 Uptake in Concrete during Accelerated Carbonation Curing Sormeh Kashef-Haghighi and Subhasis Ghoshal* Department of Civil Engineering, McGill University, Montreal, QC, Canada S Supporting Information *

ABSTRACT: Accelerated curing of fresh concrete using CO2 is a possible approach for value-added, high-volume usage products from waste CO2 emitted from stationary sources. The extent of CO2 uptake and the spatial distribution of the CaCO3(s) precipitates formed during accelerated carbonation curing of compacted, 4-h hydrated cement mortar (fresh concrete mixture with fine aggregates) samples were investigated in this study. The maximum carbonation efficiency achieved was 20% of the theoretical uptake. Microprobe imaging was used to analyze the composition of the compacted cement mortar microstructure and showed extensive filling of pores of diameters 4 μm and smaller, with CaCO3(s). The carbonation efficiency, however, reached 67% when an aqueous suspension of cement was carbonated in a completely mixed reactor, where interparticle pores do not exist and a higher surface area of cement particles is exposed to dissolved CO2. The theoretical efficiency was not achieved because all reactive cement surfaces were saturated with carbonation products, as indicated by equilibrium concentrations of dissolved calcium, silica, inorganic carbon, and pH. This study shows that both deposition of CaCO3(s), on reactive surfaces, and pore filling may regulate the extent of CO2 uptake during accelerated carbonation curing of concrete.

1. INTRODUCTION Carbon dioxide capture and storage (CCS) technologies are being developed for stabilization of atmospheric CO 2 concentrations. Potential CO2 storage methods include storage of CO2 in geological formations1 such as oil and gas fields, unmineable coal beds, deep saline formations,2,3 on the deep seafloor,4−6 or on land as inorganic carbonates through chemical reactions with abundant metal oxide minerals such as olivine, serpentine, or wollastonite7−11 or industrial waste products such as cement kiln dust (CKD)12 and recycled concrete. Surface absorption of CO2 by recycled concrete, mainly composed of calcium hydroxide and calcium silicate hydrate (C−S−H) gel, showed significant CO2 uptake.13 With the exception of CO2 injection into oil fields for enhanced oil recovery or into coal beds for enhanced methane extraction,14,15 the above CO2 storage procedures do not lead to value-added products. Accelerated concrete curing using CO2 is an alternative process to the conventional accelerated steam curing of concrete, and has the potential for creation of value-added commercial concrete products such as masonry units, while providing mineral sequestration of CO2. The main concrete products which have the potential to be CO2 cured with this technology are nonreinforced concrete masonry units, concrete blocks and bricks, or concrete products with nonmetallic reinforcements. Compared to conventional accelerated curing methods using steam, accelerated carbonation curing is based on CO2 reactions with cement minerals and results in rapid hardening of concrete and a stronger and more durable concrete when used without steel reinforcement.16−18 Cement has potential for mineral sequestration of CO2 because it is rich in calcium, mainly tricalcium silicate (3CaO·SiO2 or C3S) and dicalcium silicate (2CaO·SiO2 or C2S), which can react with CO2 under ambient conditions to © 2013 American Chemical Society

form thermodynamically stable calcium carbonate. Carbonation of pure calcium silicate mortars (C3S and C2S) with CO2 gas and moisture results in exothermic carbonation reactions leading to the formation of CaCO3(s) and SiO2(s).19,20 Accelerated carbonation of calcium silicates occurs in a sequence of steps,21 and the overall carbonation reactions of C3S and C2S, the dominant cement phases, are shown in eqs 1 and 2. A schematic of carbonation of different cement mineral phases is presented in Figure 1.21,22 3CaO·SiO2 (s) + 3CO2 (aq) → SiO2 (s) + 3CaCO3(s) (1)

2CaO·SiO2 (s) + 2CO2 (aq) → SiO2 (s) + 2CaCO3(s) (2)

Cement can store CO2 up to 50% of its weight based on reactions with the oxides present in its composition according to the Steinour formula,23 assuming all the metal oxides in cement are available for carbonation reaction. An average carbonation efficiency of 17.7% ± 1.5 was obtained in our previous study24 for accelerated CO2 curing of compacted cement mortar samples in a reactor with CO2 flow-through, and was comparable with the carbonation efficiency of 17−21% by compacted specimens measured by Shao et al.16 in CO2 pressure chambers with the same cement type and mix ratios. As with concrete, cement mortars are mixtures of cement, aggregates, and water, but for a cement mortar the aggregates used are smaller than 9.5 mm in diameter. There are no systematic studies identifying the exact causes for limited CO2 uptake during accelerated CO2 curing of Received: Revised: Accepted: Published: 5529

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20% high purity CO2 in nitrogen balance (Praxair Inc.) to simulate as-captured flue gas from the cement industry. A 40% CO2 gas mixture was also used to study the effect of CO2 partial pressure on carbonation efficiency. 2.2. Compacted Cement Mortar Experiment. The specimens were prepared with a water/cement ratio of either 0.32 or 0.26 (w/w) and an aggregate (sand)/cement ratio of 4 (w/w). Each sample contained 151 g of cement, 607 g of sand as fine aggregate, and 39 or 48 mL of water. The mixture was compacted in a steel mold (125 mm diameter × 40 mm height) by 8 MPa pressure, corresponding to a compaction force of 100 kN, to gain a concrete density of 2103 kg/m3 acceptable for precast concrete products.28 An hour after casting, the samples were mounted in a PVC shell (125 mm internal diameter × 35 mm height) and sealed using 5-min epoxy (McMaster-Carr). Prior to carbonation, the mounted samples were stored in a sealed 100% humidity chamber at room temperature for 3 h hydration, to gain an initial strength for handling in the carbonation curing reactor. The compaction and subsequent hydration produces a compacted, solid matrix from the paste like mixture of cement, water, and sand. The compacted, fresh cement mortar in the PVC shell was carbonated in a 1-D flow-through, stainless steel reactor at constant temperature, inlet gas pressure, gas flow-rate, CO2 partial pressure, and relative humidity.24 The gas flow-rate was 1.17 Lpm. The effluent CO2 concentration was monitored inline during carbonation using a Quantek model 906 NDIR gas sensor. The CO2 uptake and carbonation efficiency was calculated based on the inline measurement of effluent CO2 concentration and the known influent concentration. The carbonation efficiency was also calculated based on measurements of the total carbon content of carbonated specimens carried out with an Eltra CS-800 carbon combustion infrared analyzer. 2.3. Aqueous Suspension of Cement. Cement suspension samples were prepared with 3.8 g of Type 10 Portland cement and water/cement ratio of 30. The aqueous cement suspension was carbonated in a three-port completely mixed flow-through (CMFT) reactor with the same gas mixture as used for the compacted specimens. The reactor was made of glass with an approximate volume of 250 mL. The gas flow-rate was 0.5 Lpm. The reactor was kept on a shaker operating at 300 rpm during carbonation to maintain a homogeneous suspension. The effluent CO2 concentrations were measured inline as described above. The overall experimental set up (Figure 2) was designed to accommodate either reactor, depending on whether the experiment involved carbonating a compacted cement mortar specimen or an aqueous suspension of cement and includes the analytical instrumentation, control systems, and gas supply. 2.4. Analytical Methods. The concentrations of total dissolved Ca and Si in the carbonated and noncarbonated aqueous cement suspensions were measured using a PerkinElmer AAnalyst 100 atomic absorption spectrometer. The samples were prepared by filtering 10 mL of the solution through a 0.45 μm filter. Total dissolved inorganic carbon (DIC) was also measured by a carbon analyzer (Folio DC-80 Instruments Inc.). A Mettler DL25 autotitrator was used to measure the alkalinity in the sample with two solutions of HCl and NaOH both at 0.01 N concentrations. Alkalinity was measured to the end point pH of 4.3. Raw cement particles, hydrated cement, and carbonated cement solids were characterized by X-ray diffraction (XRD). The XRD facility

Figure 1. Schematic showing reaction products formed during accelerated CO2 curing from various reactive cement phases.

concrete. Possible reasons for limited CO2 uptake could be that the dissolution of the metal oxides and/or CO2 transport is hindered by the formation of CaCO3(s) precipitates25,26 or that reaction conditions are altered over time leading to limited availability of dissolved metals and carbonates.10 The objective of this study was to investigate the spatial distribution of precipitated CaCO3(s) in the compacted cement mortar and its relationship to the pore structure to gain insight on matrix effects that may limit CO2 uptake. To maintain the homogeneity of the matrix, only aggregates smaller than 2.5 mm were used in the mixture. The distribution of precipitates in the microstructure of compacted cement mortar subjected to accelerated CO2 curing was examined by a scanning electron microprobe imaging technique. An associated objective was to determine the carbonation efficiency of cement in aqueous media, in the absence of interparticle pores and to characterize the changes in aqueous chemistry during the carbonation process. This was achieved by the carbonation of cement and water in a completely mixed, flow-through reactor, which was periodically sampled to determine the carbonation efficiency, and relevant aqueous chemistry parameters, namely, pH and concentrations of dissolved Ca, Si, and inorganic carbon.

2. MATERIALS AND METHODS 2.1. Materials. The compacted cement mortar specimens and aqueous suspension of cement were prepared with Type 10 Portland cement (St. Lawrence Cement), kiln dried sand (Bomix, diameter range 0.08−2.5 mm), and tap water as used in standard concrete manufacturing.27 The composition of the water is reported in Supporting Information, Table S.1. The raw cement was characterized for its elemental composition using X-ray fluorescence (XRF) spectroscopy to determine the relative abundance of different cement phases. Particle size distribution of cement was measured by laser diffraction of cement suspended in isopropyl alcohol. The gas mixture was 5530

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intensities higher than the threshold value. For instance, calcium carbonate was defined as the pixels containing both calcium and carbon with intensities more than the threshold values. The digitized images were postprocessed to reduce noise and to clearly define the boundaries of different solid cement phases. Image analysis was done by MATLAB to define the materials in each pixel and to reduce the noise.

3. RESULTS AND DISCUSSION 3.1. Characterization of Cement before Carbonation. The chemical composition of the dry cement as characterized using X-ray fluorescence (XRF) is shown in Table 1. The mass percentages of the major solid phase cement minerals were estimated using the modified Bogue equations.32 Table 1. Composition and Particle Size of Cementa

Figure 2. (A) Carbonation experiment set up which includes either of the following carbonation reactors: (B) completely mixed flowthrough (CMFT) carbonation reactor and (C) flow-through reactor with vertical (1-dimensional) gas flow.

used in this study was a Rigaku D/MAX 2400 (12 kW) with a rotating anode diffractometer. In all situations, the solution was dried in the oven, and the powder was compacted in the specific mold. 2.5. Microstructure Imaging and Analysis. The electron probe microanalysis instrument (JXA JEOL-8900L) was operated with an accelerating voltage of 15 kV and the current of 20 nA for high resolution digital X-ray mapping with both electron dispersive spectroscopy (EDS) and wavelength dispersive spectroscopy (WDS). Solid core samples with dimensions of 27 mm diameter and approximately 7 mm depth were taken from compacted specimens subjected to carbonation as described above or to hydration only. To image the noncarbonated (hydrated) matrix, the compacted cement mortar was hydrated for 24 h to gain an initial strength sufficient for coring. For both samples, pores were filled with an epoxy resin and the surface was sawed and polished using the procedure reported by Stutzman.29 The surface was coated with chrome to provide a conductive surface for electron microscopy. The imaging technique was similar to Stutzman29 and Scrivener30 who applied scanning electron microscopy (SEM) with X-ray microanalysis for studying the composition of cement and clinker. The backscattered electron images as well as the X-ray images for Ca, C, Mg, Na, K, Si, Fe, and Al were collected at the same location. Knowledge of the different elements at the same location allowed determination of the mineral phase at each pixel. Calibration was performed on standard materials such as wollastonite and calcite as listed in the Supporting Information, Table S.2. The analysis was based on segmentation.29 In this method, threshold values of X-ray intensity for each element were assigned based on the frequency of intensity levels.31 The solid cement phases were identified based on the elements at each pixel which have

major oxides (percent of dry mass)

cement phasesb (percent of dry mass)

CaO MgO Na2O K2O SO3 SiO2 Al2O3 Fe2O3 total

C3S C2S C3A C4AF

67.8 8.7 14.3 2.6

0.8−3.3 μm 3.9−7.8 μm 9.3−19 μm 22−44 μm 53−105 μm 125−210 μm >250 μm

30.2 27.6 21.2 15.1 5.6 0.4 0.0

total

93.4

total mean diameter μm median diameter μm

100.1 15.3 7.4

63.1 2.0 0.0 0.0 3.8 19.8 5.0 2.0 95.7

particle size distribution (percentage)

a

The cement phases are reported in standard cement science nomenclature: C = CaO, S = SiO2, A = Al2O3, F = Fe2O3, H = H2O. bCement phase composition estimated with modified Bogue equations.32

During 4 h of hydration, cement phases such as C3S and C2S undergo hydration reactions and form solid products. The hydration products formed during the first 4 h are calcium silicate hydrate gel (3CaO·2SiO 2 ·3H 2 O(s)), ettringite (3CaO·Al2O3·3CaSO4·32H2O(s)), and calcium hydroxide (Ca(OH)2(s)) as a result of the following reactions:33 2(3CaO· SiO2 (s)) + 6H 2O(l) → 3CaO·2SiO2 ·3H 2O(s) + 3Ca(OH)2 (s)

(3)

2(2CaO· SiO2 (s)) + 4H 2O(l) → 3CaO·2SiO2 ·3H 2O(s) + Ca(OH)2 (s)

(4)

(3CaO· Al 2O3(s)) + 3(CaSO4 · 2H 2O(s)) + 26H 2O(l) → 3CaO·Al 2O3 · 3CaSO4 ·32H 2O(s)

(5)

C3S and C2S remained the most abundant cement phases after the 4-h hydration period, as confirmed by the XRD analysis of the raw and 4-h hydrated cement (Supporting Information, Figure S.1). The amount of hydration products produced in the 4-h hydrated sample is approximately only 3% of cement weight.33 3.2. Extent of CO2 Uptake in Compacted Cement Mortar. The mass of sequestered CO2 in carbonation experiments was calculated based on its concentration in the effluent gas which was measured and recorded every 5 seconds 5531

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and the inlet and outlet gas flow-rates. The carbonation efficiency (ξ) is defined as the ratio of actual mass of CO2 stored within the sample to the Steinour-derived theoretical mass uptake.24 The theoretical CO2 uptake, XCO2,Tot (wt %), is calculated on a stoichiometric basis from the relative mass of oxides present in raw cement (Xi) listed in Table 1. The theoretical uptake in cement was estimated to be 49.62 wt % by the Steinour formula (eq 6).23

3.3. Changes in the Cement Mortar during Carbonation and Hydration. To determine the spatial distribution of the solid carbonation products in the porous matrix, core samples were taken from noncarbonated and carbonated compacted specimens used in CO2 flow-through experiments (with water/cement = 0.26 and flow-rate = 1.17 Lpm). Scanning electron microscopy and X-ray imaging analysis of representative small areas of the noncarbonated and carbonated specimens were used to determine how CaCO3(s) formed during carbonation altered the pore spaces and solid matrix. In Figure 4, the backscattered electron (BSE) images for the

XCO2,Tot = 0.785(XCaO − 0.560XCaCO3 − 0.700X SO3) + 1.091XMgO + 1.420X Na 2O + 0.935X K 2O

(6)

The CO2 uptake efficiency achieved within an hour of carbonation, corrected for the pre-existing CO2 content of 0.7 wt % contained within the unreacted cement as carbonate minerals, was 17.7% ± 1.5 using 20% CO2 and flow-rate of 1.17 Lpm.24 Using the same inlet gas pressure but a higher CO2 percentage of 40% in the mixture, the maximum CO2 uptake did not improve and the 1-h carbonation efficiency remained at 17.8% ± 1.3 (Table 2). Changing the water/cement ratio from Table 2. Carbonation Efficiency in Compacted Cement Mortar with 20% and 40% CO2 in Nitrogen Balance and Water/Cement Ratios of 0.26 and 0.32

a

CO2

w/c

carbonation efficiency %

20% 20% 40%

0.26 0.32 0.26

17.70 ± 1.50a 14.50 ± 1.00 17.80 ± 1.30

Derived from the data presented in our previous publication.24

0.26 to 0.32, that is, water saturation from 0.43 to 0.53 (v/v), decreased the 1-h uptake efficiency of carbonation to 14.5 ± 1% which is likely because higher water content retards the CO2 transport through the pores. Goodbrake et al.19 explained that the reduction in carbonation rate at higher water content is due to slower CO2 diffusion in the higher volume of water. Figure 3 shows rapid uptake of CO2 by compacted specimens in the first 10 min where almost all CO2 was consumed. Following the rapid CO2 uptake, the CO2 concentration in the outlet gradually approached to 93% of the inlet concentration.

Figure 4. Scanning electron microprobe imaging with X-ray microanalysis of 24-h hydrated and noncarbonated compacted cement mortar. (A) Backscattered Electron (BSE) image; (B) reconstructed image of pores (black), aggregate (gray), calcium carbonate (green), and remaining solid phase (white). A1 and A2 are two different locations from the same sample.

noncarbonated samples are shown in two different locations (A1 and A2). The image shows the cement paste between two silica aggregates and a large pore (in A1). CO2 gas advection would occur through the pores and connected voids and dissolve in the pore waters (primarily as carbonate ions) and would react with the dissolved Ca to form CaCO3(s). The distribution of CaCO3(s) after image processing of the X-ray maps, which involves assigning pixels containing both Ca and C above threshold values, is shown in B1 and B2. The pore spaces are shown in black, aggregates in gray, CaCO3(s) in green, and other solid phases (C3S, C2S, etc) in white to distinguish the distribution of CaCO3 in the matrix. The pore space map from the BSE image was overlapped with the solid cement phase map. The narrow pores in noncarbonated specimens, Figures 4B1 and B2, had pore diameters smaller than 4 μm. The purpose of this image reconstruction was to clearly show the presence of a network of narrow pores in the cement matrix, as

Figure 3. Measured ratio of outlet to inlet CO2 gas concentration and carbonation efficiency in 1-D flow-through reactor with compacted cement mortar prepared with water/cement ratio of 0.32 and carbonated with 20% CO2. 5532

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Figure 5. Scanning electron microprobe imaging with X-ray microanalysis of compacted cement mortar subjected to accelerated CO2 curing. (A) Backscattered Electron (BSE) image; (B) reconstructed images of pores (black), aggregate (gray), calcium carbonate (green), and remaining solid phase (white); (C) reconstructed color map of Ca and distribution of calcium carbonate (white). A1 and A2 are different locations on the same sample.

pores smaller than 4 μm in diameter accounted for 27% of the total pore volume before carbonation,24 and thus removal of this pore size range would decrease the porosity from 0.30 to 0.22. Some crystals were selected in the BSE images for mineral analysis. The weight percentage of different oxides was measured and the mineral was defined as listed in the Supporting Information, Table S.3. These minerals were used to verify the results from image analysis of X-ray maps. In Figure 5, the panels C1 and C2 show the maps of calcium overlapped with the CaCO3(s) derived in carbonated specimens. The colors in the map from blue to pink indicate increasing concentration of Ca. The CaCO3(s) are shown in white. These images confirmed the existence of calcium in solid phases other than CaCO3(s) indicating that all calcium was not reacted and converted to CaCO3. The calcium-containing minerals in most cases are surrounded by the precipitated CaCO3(s) in narrow pores which would inhibit mass transport of Ca2+ and CO32‑ ions. 3.4. CO2 Uptake by Aqueous Cement Suspension. To study the other factors that limit the CO2 uptake by cement, the effect of interparticle pore structure was eliminated by carbonation of cement in an aqueous suspension form. The dynamics of CO2 uptake is depicted in Figure 6A. The CO2 concentration in the outlet gas increased rapidly during the first 20 min and CO2 uptake ceased around 68% efficiency. The error bars represent the standard deviation of CO2 measurements in three experiments conducted under identical conditions. Although the maximum theoretical CO2 uptake potential was not achieved with the aqueous slurries, the carbonation efficiency compares well with those reported for cement in

well as the presence of some background CaCO3(s) which is the unreacted limestone in cement clinker. The BSE images for carbonated specimens are shown in Figures 5A1 and A2. The reconstructed images for carbonated cement mortar and CaCO3(s) distribution are shown in Figures 5B1 and B2. These figures show that CaCO3(s) deposited in the abundant narrow pores that were visible in Figure 4, as well as in the form of larger irregular deposits particularly close to the large pores. Deposition of CaCO3(s) was also observed as a layer along the walls of the large pores or voids (>10 μm). Some of these larger deposits may also be unreacted limestone in the cement clinker. The other precipitate generated during carbonation, SiO2(s), was randomly scattered in the carbonated specimens (data not shown) and was not systematically distributed in the proximity of large pores or as large deposits as was the case for CaCO3. In contrast to the noncarbonated samples shown in Figure 4, the carbonated samples shown in Figure 5 have virtually no remaining unfilled narrow pores that are visible at the image resolution of 0.65 μm per pixel, clearly indicating that the pores of 4 μm become clogged with CaCO3(s) during the 1-h carbonation period or filled to below 0.65 μm in diameter. The decrease in pore volume because of CaCO3(s) deposition is easily observed from the images. For a quantitative analysis, porosity was calculated based on the number of pixels containing pores and the total number of pixels in each image. The porosity in noncarbonated specimens based on images at 7 locations of the sample was calculated as 0.30 ± 0.03 decreasing to 0.18 ± 0.04 after carbonation which correlates with the reduction in porosity determined by mercury intrusion porosimetry (MIP) reported in our previous study.24 Mercury intrusion porosimetry analyses showed that 5533

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significant changes during the carbonation of the aqueous cement suspension. The concentrations of dissolved Ca, C as well as the pH of the aqueous cement suspension were monitored during carbonation (Figure 6B). The Si concentrations over time are included in the Supporting Information. The concentration of Ca and pH in the 4-h hydrated aqueous cement suspension, prior to carbonation, were measured as 28.5 ± 0.3 mmol/L and 13.2 ± 0.1, respectively, which compared well with the equilibrium model for aqueous cement suspension developed using PHREEQC as explained later (32.0 mmol/L and 12.5 for Ca and pH, respectively) and the measurements reported by Rothstein et al.35 in Portland cement pore waters (18.4 mmol/L of Ca and pH of 12.9). The CO2 uptake by cement ceased within 30 min, and the concentrations of Ca, dissolved inorganic carbon (DIC), and pH in the aqueous cement suspension reached steady-state values. The steady-state values for pH, DIC, Ca, and Si were 6.84, 26.4 mmol/L, 9.5 mmol/L, and 2.9 mmol/L, respectively. The alkalinity of cement solution decreased from an initial value of 64.96 meq/L to 21.57 meq/L after 30 min of carbonation. We tested the hypotheses that if all reactive surfaces are covered with carbonation products (CaCO3(s) and SiO2(s)), the aqueous concentration of various ions will be equal to that predicted from the chemical equilibrium of a system containing only those products and CO2. A chemical equilibrium analysis was performed using an ion association model, PHREEQC,36 with the Pitzer ion interaction model for the calculation of ion activity coefficients for solutions with high solute concentrations.37 The equilibrium constants and ion interaction data for reactions at 25 °C were obtained from the literature and shown in the Supporting Information, Tables S.4 and S.5.38−40 The system modeled for chemical equilibrium composition was a pure aqueous phase in contact with excess calcite and silica, the two carbonation products given by eqs 1 and 2, in the presence of CO2 gas with partial pressure of 0.21 atm. The model results show that at equilibrium, the solution pH is 6.5 and the concentration of DIC, Ca, and Si are 24.2 mmol/L, 5.4 mmol/L, and 2.0 mmol/L, respectively. Thus there was good agreement between the predicted and measured equilibrium concentrations. The equilibrium model showed that in the presence of reactive phases, such as C3S or Ca(OH)2(s), the equilibrium pH is 8.4 and 11.6, respectively, and much higher than the measured value of 6.8. Therefore, in the absence of exposed C3S or Ca(OH)2(s), the aqueous cement suspension reached a chemical equilibrium with the CaCO3(s) and SiO2(s) product layer. However, it should be noted that there are inadequacies in the model developed by PHREEQC, because of the chemical complexity of the cement/water system, which may explain the relatively minor disparities between the model and experimental measurements.37 Some sources of these disparities are the possible formation of some magnesium calcite along with calcite which will affect Ca concentration at equilibrium, and the possible formation of some calcium carboaluminate because of interaction of C3A and calcite which may affect Ca and DIC equilibrium concentrations. These products were not included in the model because the solubility data are not known. The changes in concentrations of Ca, DIC, and pH during the 90-min carbonation period, shown in Figure 6B, can be explained on the basis of the various carbonation reactions. C3S dissolution starts immediately upon addition of water to cement.41 The pH of a C3S solution increases very rapidly

Figure 6. (A) Ratio of outlet to inlet CO2 gas concentration in CMFT reactor over time during carbonation of aqueous cement suspensions, calculated carbonation efficiency from inline CO2 measurements, and carbonation efficiency from direct measurement of carbonate content by a combustion IR analyzer; (B) measured pH and concentrations of dissolved C and Ca in CMFT reactor during carbonation of aqueous cement suspensions.

noncompacted matrixes. The carbonation efficiency of cement in CKD was approximately 80%.12 The carbonation efficiencies in various calcium-rich minerals such as steel slag, wetted powders of C3S and C2S and Ca(OH)2 have also been reported in the range of 40%−80%.7,19,32,34 The ultimate carbonation efficiency achieved with pure CaO with CO2 gas was 70− 80%.25 All of those studies were conducted in reaction chambers where contact between CO2 and the cement mineral phase surfaces was not limited as with the pore structure of the compacted specimens. The difference between the theoretical CO2 uptake (i.e., 100% efficiency) and the experimentally determined carbonation efficiencies are generally attributed to the formation of a CaCO3(s) product layer that limits mass transport of Ca2+ ions from the cement phase and of carbonate ions from the bulk phase. In Figure 6A, there is a period of rapid, steady uptake of CO2, during the first 8 min (Stage 1), followed by a declining rate of uptake for the next 10 min (Stage 2). It is likely that the diffusion of ions through the deposited product layer becomes more rate controlling during this time, and after 30 min the rate of carbonation becomes too slow to be detected. An average thickness of 50 nm for the CaCO3(s) layer was determined by Alvarez and Abanades25 as the start of the slow reaction period. 3.5. Changes in Chemistry of Aqueous Cement Suspension and Relationship to CO2 Uptake. As shown in Figure 6B, the chemistry of the aqueous phase showed 5534

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because of the formation of OH− from dissolution of C3S (eq 7). The solution pH increases to approximately 12.5, and the concentration of Ca increases to 10−20 mmol/L in the first 30 s. Dicalcium silicate (C2S), the second most abundant phase in cement, also dissolves in water but more slowly compared to C3S. The Ca(OH)2(s), formed during the hydration period, also dissolves to Ca2+ and OH− ions. C3S + 3H 2O → 3Ca 2 + + H 2SiO24 − + 4OH−

during carbonation, it is clear that with the formation of a product layer on the cement surfaces, additional CO2 uptake is no longer feasible.

4. CONCLUSIONS The CO2 uptake by compacted cement mortar (concrete mixture containing fine aggregates) was rapid and a carbonation efficiency of 14−19% was achieved within an hour of carbonation with a gas flow-rate of 1.17 Lpm with 20% or 40% CO2 in the influent gas mixture. On the basis of an average carbonation efficiency of 18% with 20% CO2, and the estimated 16.4 Mt of cement used in concrete products in the United States in 2004,45 approximately 1.5 Mt of CO2 can be sequestered by this technology.24 In this study, the spatial distribution of CaCO3(s) in a compacted cement mortar matrix subjected to accelerated CO2 curing was characterized by microprobe imaging. The analysis revealed extensive filling of narrow pores by CaCO3(s). We suggest that such pore filling leads to limited mass transport of CO2 to reactive cement materials and is a critical phenomenon contributing to lower than stoichiometric amounts of CO2 uptake in concrete during accelerated CO2 curing. Previous studies have reported that formation of CaCO3(s) during carbonation of calcium oxide-rich particles occurs directly on these reactive particles during carbonation in noncompacted matrixes.7,46,47 The filling of a large portion of interparticle pores smaller than 4 μm by CaCO3(s) in compacted cement mortar imposes an additional barrier to achieving high carbonation efficiency. We demonstrated that carbonation efficiency increased when the mass transport limitations imposed by interparticle pore filling were removed by carbonating cement in an aqueous suspension. The CO2 uptake by cement in aqueous suspensions was approximately 68% of the theoretical (stoichiometric) uptake. The implication of this finding is that efforts to increase CO2 uptake by altering the concrete mixes with cements of higher reactive mineral content or cements with smaller particle sizes and hence larger reactive surface areas may show only limited increases in carbonation efficiencies. In fact, a recent study comparing the CO2 uptake by cement mortar prepared with Type 30 cement (Blaine fineness or specific surface area of 469 m2/kg) and Type 10 cement (373 m2/kg) showed only 3.8% improvement in carbonation efficiency.26 One possible way for further improving the CO2 uptake of concrete to be subjected to accelerated CO2 curing would be to carbonate CO2 absorbing waste materials such as steel slag, ladle slag,48 or electric arc furnace (EAF) slag,49 and use them to make suitable aggregates for concrete and mortar mixes prior to accelerated curing.

(7)

During Stages 1 and 2 of the carbonation experiment, the pH decreased over time with the decrease in Ca2+ associated with CaCO3(s) formation, as shown in Figure 6B. A detailed explanation of these trends is presented in the Supporting Information. In Stage 3, pH dropped to 6.8 and reached a plateau. At this pH (close to pK1), DIC is primarily present as HCO3− and undissociated carbonic acid (H2CO3), and thus there is no uptake. CO2 concentration measured in the outlet gas reached the inlet value. In Stage 3, some decalcification of the CaCO3(s) and formation of water-soluble calcium bicarbonate, Ca(HCO3)2, is likely to occur at these pH levels.42 Dissolved Ca increased from 2.9 mmol/L to 7.5 mmol/L and DIC concentration started to increase from 1.3 mmol/L to 13.6 mmol/L. The presence of calcite, in conjunction with the absence of reactive mineral phases such as C3S and C2S, in the carbonated cement suspension as the main carbonation product was confirmed by XRD analysis on 1-h carbonated cement in the CMFT reactor (Figure 7). Similar observations were reported

Figure 7. XRD results of (A) 4-h hydrated and 1-h carbonated Portland cement and (B) 4-h hydrated Portland cement; (Δ) C3S (3CaO·SiO2 (s)), (×) C2S (2CaO·SiO2 (s)), (■) calcite (CaCO3(s)), (●) C3A (3CaO·Al2O3 (s)), (+) quartz (SiO2(s)), and (−) magnesite (MgCO3(s)).

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ASSOCIATED CONTENT

S Supporting Information *

by Goodbrake et al.19 and Young et al.20 Main peaks of calcite (2θ = 29.4°, 39.4°, 43.1° and 47.5°) are dominant in Figure 7A, while main peaks of C3S (2θ =29.3°, 34.3° and 51.7°) are dominant in Figure 7B, although there is an overlap in the calcite peak at 29.4° and the C3S peak at 29.3°. Amorphous phases with low degree of crystallinity, such as C−S−H, show up as broad diffraction peaks as compared to narrow peaks for well-defined crystals. No clear indication of amorphous phases such as C−S−H were observed in the XRD image of the 4-h hydrated cement in this study compared to XRD images of amorphous cement phases in previous studies.43,44 Thus, from the analyses of the aqueous chemistry of the cement suspension

(1) Explanation regarding changes in the chemistry of aqueous cement suspension during carbonation; (2) XRD results of 4-h hydrated Portland cement and raw Portland cement; (3) measured concentrations of dissolved Si in CMFT reactor during carbonation of aqueous cement suspension; (4) composition of water used in aqueous cement suspension; (5) calibration standards and detection limits for each element analyzed with electron microprobe; (6) results of electron microprobe analysis for selected crystals in Figures 4 and 5; (7) equilibrium constants for major carbonation reactions in cement at 25 °C; and (8) values of electrolyte interaction 5535

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parameters at 25 °C. This material is available free of charge via the Internet at http://pubs.acs.org.

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(13) Haselbach, L. M.; Ma, S. Potential for Carbon Adsorption on Concrete: Surface XPS Analyses. Environ. Sci. Technol. 2008, 42 (14), 5329−5334. (14) van Bergen, F.; Pagnier, H.; Krzystolik, P. Field Experiment of Enhanced Coalbed Methane-CO2 in the Upper Silesian Basin in Poland. Environ. Geosci. 2006, 13 (3), 201−224. (15) Gentzis, T. Subsurface Sequestration of Carbon Dioxide–an Overview from an Alberta (Canada) Perspective. Int. J. Coal Geol. 2000, 43 (1−4), 287−305. (16) Shao, Y.; Mirza, M. S.; Wu, X. CO2 Sequestration Using Calcium-Silicate Concrete. Can. J. Civ. Eng. 2006, 33 (6), 776−784. (17) Shi, C.; Wu, Y. Studies on Some Factors Affecting CO2 Curing of Lightweight Concrete Products. Resour. Conserv. Recycl. 2008, 52 (8−9), 1087−1092. (18) Watanabe, K.; Yokozeki, K.; Ashizawa, R.; Sakata, N.; Morioka, M.; Sakai, E.; Daimon, M. High Durability Cementitious Material with Mineral Admixtures and Carbonation Curing. Waste Manage. 2006, 26 (7), 752. (19) Goodbrake, C. J.; Young, J. F.; Berger, R. L. Reaction of Hydraulic Calcium Silicates with Carbon Dioxide and Water. J. Am. Ceram. Soc. 1979, 62 (9−10), 488−491. (20) Young, J. F.; Berger, R. L.; Breese, J. Accelerated Curing of Compacted Calcium Silicate Mortars on Exposure to CO2. J. Am. Ceram. Soc. 1974, 57 (9), 394−397. (21) Fernandez Bertos, M.; Simons, S. J. R.; Hills, C. D.; Carey, P. J. A Review of Accelerated Carbonation Technology in the Treatment of Cement-Based Materials and Sequestration of CO2. J. Hazard. Mater. 2004, 112 (3), 193. (22) Dilnesa, B. Z.; Lothenbach, B.; Le Saout, G.; Renaudin, G.; Mesbah, A.; Filinchuk, Y.; Wichser, A.; Wieland, E. Iron in Carbonate Containing AFm Phases. Cem. Concr. Res. 2011, 41 (3), 311−323. (23) Steinour, H. H. Some Effects of Carbon Dioxide on Mortars and Concrete-Discussion. J. Am. Concr. Inst. 1959, 30, 905−907. (24) Kashef-Haghighi, S.; Ghoshal, S. CO2 Sequestration in Concrete through Accelerated Carbonation Curing in a Flow-through Reactor. Ind. Eng. Chem. Res. 2010, 49, 1143−1149. (25) Alvarez, D.; Abanades, J. C. Determination of the Critical Product Layer Thickness in the Reaction of CaO with CO2. Ind. Eng. Chem. Res. 2005, 44, 5608−5615. (26) Monkman, S.; Shao, Y. Assessing the Carbonation Behavior of Cementitious Materials. J. Mater. Civ. Eng. 2006, 18 (6), 768. (27) Lamond;, J.; Pielert, J. Significance of Tests and Properties of Concrete and Concrete-Making Materials; ASTM International: Philadelphia, PA, 2006. (28) Warszawski, A. Industrialized and Automated Building Systems 2nd ed.; Taylor & Francis: Abingdon, U.K., 1999. (29) Stutzman, P. Scanning Electron Microscopy Imaging of Hydraulic Cement Microstructure. Cem. Concr. Compos. 2004, 26 (8), 957. (30) Scrivener, K. L. Backscattered Electron Imaging of Cementitious Microstructures: Understanding and Quantification. Cem. Concr. Compos. 2004, 26 (8), 935. (31) Bentz, D. P.; Stutzman, P. E. SEM/X-Ray Imaging of CementBased Materials. In Proceedings of the 7th Euroseminar on Microscopy Applied to Building Materials, 1999. (32) Taylor, H. F. Cement Chemistry, 2nd ed.; Telford: London, U.K., 1997. (33) Lea, F. M. The Chemistry of Cement and Concrete; Edward Arnold: London, U.K., 1970. (34) Shih, S. M.; Ho, C. S.; Song, Y. S.; Lin, J. P. Kinetics of the Reaction of Ca(OH)2 with CO2 at Low Temperature. Ind. Eng. Chem. Res. 1999, 38 (4), 1316−1322. (35) Rothstein, D.; Thomas, J. J.; Christensen, B. J.; Jennings, H. M. Solubility Behavior of Ca-, S-, Al-, and Si-Bearing Solid Phases in Portland Cement Pore Solutions as a Function of Hydration Time. Cem. Concr. Res. 2002, 32 (10), 1663−1671. (36) Parkhurst, D. L.; Appelo, C. A. J. User’s Guide to PHREEQC (Version 2) a Computer Program for Speciation, Batch-Reaction, OneDimensional Transport, and Inverse Geochemical Calculations; U.S.

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Phone: +1 514 398 6867. Fax: +1 514 398 7361. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The authors would like to acknowledge the Natural Sciences and Engineering Research Council of Canada (NSERC), St. Lawrence Cement, CJS Technology, for providing funding for this project, the Global Environmental and Climate Change Centre (GEC3) at McGill University and le Fonds québécois de la recherche sur la nature et les technologies (FQRNT) for providing fellowships to S.K., and Robert Niven (Civil Engineering, McGill) for the XRF analyses.

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REFERENCES

(1) Peters, C. A. Accessibilities of Reactive Minerals in Consolidated Sedimentary Rock: An Imaging Study of Three Sandstones. Chem. Geol. 2009, 265 (1−2), 198−208. (2) White, C. M.; Strazisar, B. R.; Granite, E. J.; Hoffman, J. S.; Pennline, H. W. Separation and Capture of CO2 from Large Stationary Sources and Sequestration in Geological Formations- Coalbeds and Deep Saline Aquifers (Critical Review). J. Air Waste Manage. Assoc. 2003, 53, 645−715. (3) Wilson, E. J.; Morgan, M. G.; Apt, J.; Bonner, M.; Bunting, C.; Gode, J.; Haszeldine, R. S.; Jaeger, C. C.; Keith, D. W.; McCoy, S. T.; Pollak, M. F.; Reiner, D. M.; Rubin, E. S.; Torvanger, A.; Ulardic, C.; Vajjhala, S. P.; Victor, D. G.; Wright, I. W. Regulating the Geological Sequestration of CO2. Environ. Sci. Technol. 2008, 42 (8), 2718−2722. (4) Hoffert, M. I.; Caldeira, K.; Benford, G.; Criswell, D. R.; Green, C.; Herzog, H.; Jain, A. K.; Kheshgi, H. S.; Lackner, K. S.; Lewis, J. S.; Lightfoot, H. D.; Manheimer, W.; Mankins, J. C.; Mauel, M. E.; Perkins, L. J.; Schlesinger, M. E.; Volk, T.; Wigley, T. M. L. Advanced Technology Paths to Global Climate Stability: Energy for a Greenhouse Planet. Science 2002, 298 (5595), 981−987. (5) Tsouris, C.; Brewer, P.; Peltzer, E.; Walz, P.; Riestenberg, D.; Liang, L.; West, O. R. Hydrate Composite Particles for Ocean Carbon Sequestration: Field Verification. Environ. Sci. Technol. 2004, 38 (8), 2470−2475. (6) IPCC Special Report on Carbon Dioxide Capture and Storage; Cambridge University Press: New York, 2005. (7) Huijgen, W. J. J.; Witkamp, G. J.; Comans, R. N. J. Mineral CO2 Sequestration by Steel Slag Carbonation. Environ. Sci. Technol. 2005, 39 (24), 9676−9682. (8) Lackner, K. S.; Wendt, C. H.; Butt, D. P.; Joyce, E. L.; Sharp, D. H. Carbon Dioxide Disposal in Carbonate Minerals. Energy 1995, 20, 1153. (9) Jarvis, K.; Carpenter, R. W.; Windman, T.; Kim, Y.; Nunez, R.; Alawneh, F. Reaction Mechanisms for Enhancing Mineral Sequestration of CO2. Environ. Sci. Technol. 2009, 43 (16), 6314−6319. (10) Huijgen, W. J. J.; Witkamp, G. J.; Comans, R. N. J. Mechanisms of Aqueous Wollastonite Carbonation as a Possible CO2 Sequestration Process. Chem. Eng. Sci. 2006, 61 (13), 4242−4251. (11) Maroto-Valer, M. M.; Fauth, D. J.; Kuchta, M. E.; Zhang, Y.; Andresen, J. M. Activation of Magnesium Rich Minerals as Carbonation Feedstock Materials for CO2 Sequestration. Fuel Process. Technol. 2005, 86 (14−15), 1627. (12) Huntzinger, D. N.; Gierke, J. S.; Kawatra, S. K.; Eisele, T. C.; Sutter, L. L. Carbon Dioxide Sequestration in Cement Kiln Dust through Mineral Carbonation. Environ. Sci. Technol. 2009, 43 (6), 1986−1992. 5536

dx.doi.org/10.1021/ie303275e | Ind. Eng. Chem. Res. 2013, 52, 5529−5537

Industrial & Engineering Chemistry Research

Article

Department of the Interior, U.S. Geological Survey: Denver, CO, 1999. (37) Reardon, E. J.; Dewaele, P. Chemical Model for the Carbonation of a Grout/Water Slurry. J. Am. Ceram. Soc. 1990, 73 (6), 1681−1690. (38) Reardon, E. J. Problems and Approaches to the Prediction of the Chemical Composition in Cement/Water Systems. Waste Manage. 1992, 12, 221−239. (39) Bullard, J. W.; Enjolras, E.; George, W. L.; Satterfield, S. G.; Terrill, J. E. A Parallel Reaction-Transport Model Applied to Cement Hydration and Microstructure Development. Model. Simul. Mater. Sci. Eng. 2010, 18, 711−738. (40) Reardon, E. J. An Ion Interaction Model for the Determination of Chemical Equilibria in Cement/Water Systems. Cem. Concr. Res. 1990, 20 (2), 175−192. (41) Bullard, J. W. A Determination of Hydration Mechanisms for Tricalcium Silicate Using a Kinetic Cellular Automaton Model. J. Am. Ceram. Soc. 2008, 91 (7), 2088−2097. (42) Kutchko, B. G.; Strazisar, B. R.; Dzombak, D. A.; Lowry, G. V.; Thaulow, N. Degradation of Well Cement by CO2 under Geologic Sequestration Conditions. Environ. Sci. Technol. 2007, 41 (13), 4787− 4792. (43) Chen, J. J.; Thomas, J. J.; Taylor, H. F. W.; Jennings, H. M. Solubility and Structure of Calcium Silicate Hydrate. Cem. Concr. Res. 2004, 34 (9), 1499−1519. (44) Kim, J. J.; Foley, E. M.; Reda Taha, M. M. Nano-Mechanical Characterization of Synthetic Calcium-Silicate-Hydrate (C-S-H) with Varying CaO/SiO2 Mixture Ratios. Cem. Concr. Compos. 2013, 36 (0), 65−70. (45) EIA Emissions of Greenhouse Gases in the United States 2006; Energy Information Administration, U.S. Department of Energy: Washington, DC., 2007. (46) Bhatia, S. K.; Perlmutter, D. D. Effect of the Product Layer on the Kinetics of the CO2-Lime Reaction. AIChE J. 1983, 29 (1), 79−86. (47) Mess, D.; Sarofim, A. F.; Longwell, J. P. Product Layer Diffusion During the Reaction of Calcium Oxide with Carbon Dioxide. Energy Fuels 1999, 13 (5), 999−1005. (48) Monkman, S.; Shao, Y.; Shi, C. Carbonated Ladle Slag Fines for Carbon Uptake and Sand Substitute. J. Mater. Civ. Eng. 2009, 21 (11), 657−665. (49) Yokoyama, S.; Arisawa, R.; Hisyamudin, M. N. N.; Murakami, K.; Maegawa, A.; Izaki, M. Applicability of Carbonated Electric Arc Furnace Slag to Mortar. J. Phys.: Conf. Ser. 352 (1), 012049.

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