Article pubs.acs.org/EF
Kinetics Analysis of CO2 Mineral Carbonation Using Byproduct Red Gypsum Omeid Rahmani,*,† Ali Kadkhodaie,‡ and James Highfield§ †
Department of Petroleum Engineering, Mahabad Branch, Islamic Azad University, Mahabad 59135-433, Iran Department of Petroleum Engineering, School of Chemical and Petroleum Engineering, Curtin University of Technology, Perth, Western Australia 6102, Australia § 560 Yishun Avenue 6 #08-25 Lilydale, Singapore 768966, Singapore ‡
ABSTRACT: In this study, a simplified model was proposed for CO2 mineral carbonation of red gypsum. The model was based on the PHREEQC-2.18 program, which is used for thermodynamic calculations. Test experiments in an aqueous carbonation reactor were used as reference (Rahmani, O.; Junin, R.; Tyrer, M.; Mohsin, R. Mineral carbonation of red gypsum for CO2 sequestration. Energy Fuels 2014, 28, 5953−5958 and Rahmani, O.; Tyrer, M.; Junin, R. Calcite precipitation from by-product red gypsum in aqueous carbonation process. RSC Adv. 2014, 4, 45548−45557) to verify the model, evaluate the possibility of implementing reactions, and predict the reaction progress over time. At the first step, the model thermodynamic constants and rate expressions were initially determined from experiments in an autoclave mini reactor. These parameters were then included in the model, and their quality was tested by comparing experimental and modeled data in the CO2 mineral carbonation of red gypsum. The evaluated model should prove valuable not only in applications of in situ or ex situ CO2 sequestration but more generally in computational geoscience.
1. INTRODUCTION One of the most important byproducts obtained in large amounts during extraction of titania (TiO2) from ilmenite mineral (nominally FeTiO3, containing 42−66% TiO2) is red gypsum (RG). TiO2 is extracted through stepwise processes, yielding sulfuric acid (H2SO4) as a byproduct. Neutralization of the latter with lime or limestone results in so-called RG, consisting of roughly 75% gypsum (CaSO4·2H2O) and 25% iron hydroxide (Fe2O3) by dry mass. Each year, Huntsman Tioxide produces over one million tons of RG worldwide. This material has been the focus of recent work in our laboratory, exploring its potential for ex situ CO2 mineralization.1,2 Increasing levels of atmospheric carbon dioxide (CO2) emissions are primarily responsible for the anthropogenic greenhouse effect or global warming. Despite the current attention given to physical disposal methods, by far the safest, verifiable, and permanent solution to the CO2 problem is by storing it as thermodynamically stable carbonate minerals.3 Naturally occurring and widely available minerals and certain industrial wastes rich in magnesium (Mg) and calcium (Ca) are obvious substrates to exploit because they are environmentally benign.4 However, industrial wastes offer several advantages over natural minerals in this respect.5 They are accessible “on-site”, hence cheaper, and are often geochemically unstable or more reactive. According to Prigiobbe et al.,6 mineral carbonation of industrial wastes rich in Ca is a thermodynamically favored reaction that mimics natural weathering. For example, during oil and gas production, a huge amount of saline wastewater containing dissolved Ca is produced annually and could also be used for the carbonation reaction.7 Many studies have attempted to model the kinetics of mineral carbonation on an industrial scale,8−12 but little attention has been given thus far to RG. Furthermore, a geochemical model © 2016 American Chemical Society
accurately predicting its reaction progress with CO2 is still lacking. This study aims to simulate and predict the processes involved in batch dissolution and carbonation experiments with RG, which produces calcium carbonate (CaCO3) as a stable carbonated mineral. To this end, the computer program PHREEQC-2.18 was used to simulate chemical reactions in the process of CO2 mineral carbonation. The geochemical model code PHREEQC is a well-established algorithm to estimate the dissolution and precipitation of various solid phases in technical carbonation applications.3,14−16 The value of the model was assessed by comparing the experimental data to predicted behavior.
2. METHODS 2.1. Experimental Procedure. According to Rahmani et al.,1 the process of CO2 mineral carbonation included two main steps. In the first (dissolution) step a Ca-ion-rich solution suitable for carbonation was targeted. A total of 10 g of dried RG samples with a defined grain size cut of less than 38 μm was dissolved under stirring in an excess (1.5 mol) of H2SO4 at a temperature of 60 °C in an open vessel. After 2 h, any residual (unwanted) solid was removed from the Ca-/Fe-rich solution (pH ∼ 2.5) by sedimentation and filtration. A total of 100 mL of 2.1 M NH4OH was then added to enhance the pH value to 8.6 and precipitate iron oxyhydroxides. The remaining solution was essentially rich in Ca2+ ion, as ideally represented by a single reaction (eq 1).
CaSO4 ·2H 2O(s) → Ca 2 +(aq) + SO4 2 −(aq) + 2H 2O(l)
(1)
The second (carbonation) stage was conducted in an autoclave mini reactor with 150 mL capacity. A digital set reactor controller including a magnetic stirrer, temperature sensor, and electrodes for electrical Received: February 1, 2016 Revised: August 14, 2016 Published: August 18, 2016 7460
DOI: 10.1021/acs.energyfuels.6b00246 Energy Fuels 2016, 30, 7460−7464
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Table 1. Concentration and Quantity of Material and Chemical Used in Dissolution and Carbonation Experiments for Modeling in the PHREEQC-2.18 Programa
a
material and chemical used
stage
experiment
CO2 (%)
n (rpm)
10 g of RG (d < 38 μm) 143 mL of H2SO4 (35 wt %) 100 mL of NH4OH (50 wt %) 3 mL of NH4OH (50 wt %) CO2 with a purity of 99.99%
dissolution
CaO dissolution CO2 dissolution
CaO CaO
0 10
400 100−400
carbonation
CaO carbonation CaO−RG
CaO CaO−CaSO4·2H2O
10 10
400 400
solid phase
T and P are 25 °C and 1 atm, respectively.
conductivity and monitoring pH was embedded to the mini reactor. At the starting point, the solution rich in Ca was bubbled with pure CO2 at a temperature of 25 °C for about 10 min,17 after which the solution became alkaline (pH ∼ 9.5) by the addition of 2.1 M NH4OH. The CO2saturated solution was loaded into the mini reactor, after which different partial pressures of CO2 (pCO2) ranging from 1 to 30% were introduced. The CO2 gas cylinder was equipped with a CO2 flow-meter regulator to monitor and calculate the flow rate of injected CO2 and the net volume (%) of inlet gas, respectively. Simultaneously, the mini reactor was heated to the set temperature between 25 and 150 °C at a stirring rate ranging from 100 to 400 rpm. To achieve an optimal level of dissolved CO2 throughout the Ca-rich solution, the stirring rates selected were necessarily in the low range. At higher speeds, a distinct case of “layering” was seen, in which CO2 evolution was agitated. The various experimental parameters are summarized in Table 1. 2.2. Model System. The PHREEQC-2.18 program13 was applied to thermodynamic calculations by providing the data from the Lawrence Livermore National Laboratory (LLNL). PHREEQC-2.18 was also used to determine the carbonation and dissolution degree of RG samples to predict the amount of CO2 sequestrated. Thermodynamic equilibrium constants for the reactions of CO2 mineral carbonation of RG were also incorporated from model databases. Reaction kinetics analysis was implemented using a BASIC interpreter (a computer language with statements on numbered lines) to predict the progress of reaction over time. The quality of predictions in the model system was evaluated by comparison to the determined rates (e.g., reaction rate “r” and rate constant “k”) under varying conditions [e.g., reaction temperature (T), time (t), stirring rate (n), and liquid/solid ratio (L/S)]. 2.3. Kinetics Analysis. Reactions and kinetics of CO2 sequestrated in RG suspension were determined experimentally1,2 on the basis of the process variables, such as T, t, n, and L/S, using a batch reaction system. Equation 2 was applied for calculation of dissolution and carbonation kinetics
r = − S[A e−Ea / RT (1 − DC)]
(4)
kt = 1 − 3(1 − XE)1/3
(5)
where k is the reaction rate constant and XE is the amount of Ca in solution at any time t. According to Dri et al.,19 a product layer diffusion model applies when the diffusion of the reaction products from the surface of the particle into the solution is rate-limiting. Alternatively, a film diffusion model applies for rate-limiting diffusion of the reactants from the solution to the surface of particles. Reaction control applies to rate-limiting chemistry occurring between the reactant and the particles at their surfaces. 2.4. CO2 Sequestration: Total Dissolved Inorganic Carbon (TDIC). The amount of CO2 sequestrated by conversion of the substrate to the carbonate form was calculated to differentiate vent level and supply of CO2 at fixed flow using an optical infrared (IR) sensor. n
CO2 sequestration (mol/g) =
∑ i
(pCO
2,in
− pCO
2,out
MTR
)i Q Δt (6)
where pCO2,in and pCO2,out denote to mean value of pCO2 (atm) at the outflow and inflow, respectively, Q and Δt are the flow rate (min−1) and time interval (min), respectively, R donates the gas constant (0.082 057 L atm mol−1 K−1), and M and T are the mass of RG (g) and temperature (K), respectively.
3. RESULTS AND DISCUSSION 3.1. Effects of Experimental Variables in the Dissolution Process. Two distinct mechanisms were noted in the dissolution process: internal diffusion of Ca to the surface of particles and dissolution of Ca from the exterior surface into the solution (see eq 1). The key variables in the dissolution process were the reaction temperature (T), reaction time (t), and grain size (d). The maximum amount of Ca in the solution (>99%) was achieved at a temperature of 60 °C using a particle size of less than 38 μm for 2 h.1,2 Huijgen et al.20 declared that leaching of Ca from the exterior surface of the steel slag sample into the solution rich in Ca may proceed more slowly at higher temperatures. This effect is apparently thermodynamic and consistent with reports that the aqueous solubility of CaSO4 passes through a maximum below 100 °C.21 A similar effect based on instability of the corresponding Ca2+/SO42−/H2O system may be expected in this work. The dissolution kinetics of Ca was measured in the aqueous phase at different pH values, as shown in Figure 1. The amount of dissolved Ca increased sharply in the first 5 min, reaching ∼0.018 mol/L and with an associated pH rise from 6.9 to 9.5. This trend continued for up to 10 min, with the Ca level increasing to 0.032 mol/L and the pH value reaching a maximum at around 12.5. 3.2. Effects of Experimental Variables in the Carbonation Process. The key variables in the carbonation process were T, t, n, and L/S. Two distinct procedure steps are distinguishable in the carbonation process: transformation of
(2)
where “r” is the reaction rate (mol/s), “S” is the mineral surface area (m2), “A” is the Arrhenius pre-exponential factor (mol m−2 s−1), “T” is the temperature (K), “Ea” is the activation energy (J/mol), and “R” and “DC” denote the gas constant and degree of conversion (%), respectively. Using this equation, the well-stirred mini reactor was simulated. Additionally, the minimum reaction time for equilibration and reaction route in a general rate were estimated. The kinetics analysis of the RG dissolution rate was estimated by applying “standard integral analysis”18 in Ca-rich solution using H2SO4 and NH4OH. According to Levenspiel,18 the integral method (after appropriate mathematical manipulation) shows a particular form of rate equation predicting a linear plot of concentration versus time. To evaluate the best fit to various heterogeneous reaction models (vide inf ra), a shrinking core model was assumed and multiple regression coefficients (R2) in the data were calculated.4 In similar (Ca-based) studies on steel slags, Dri et al.19 demonstrated that the reaction is initially limited to the outer surface of the virgin particles and reaction progress is subsequently regulated as a result of the predominance of a specific heterogeneous mechanism, such as product layer diffusion (eq 3), film diffusion (eq 4), and chemical reaction control (eq 5), as below
kt = 1 − 3(1 − XE)2/3 + 2(1 − XE)
kt = XE
(3) 7461
DOI: 10.1021/acs.energyfuels.6b00246 Energy Fuels 2016, 30, 7460−7464
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Figure 1. Ca concentration in solution based on the experimental variables in the dissolution process (T = 60 °C, n = 400 rpm, d < 38 μm, and pCO2 = 1 atm).
dissolved CO2 to CO32− (eq 7) and precipitation of CaCO3 (eq 8). CO2(g) + H 2O(l) → 2H+(aq) + CO32 −(aq)
(7)
Ca 2 +(aq) + CO32 −(aq) → CaCO3(s)
(8)
To investigate the reaction temperature and time in the carbonation process, the temperature was varied from 25 to 150 °C and t was limited up to 3 h. Preparing the experiment with the desired temperature (i.e., 25−150 °C) caused the maximum amount of Ca conversion (i.e., 98.8%) to CaCO3, which appeared at a temperature between 55 and 60 °C with a stirring rate of 400 rpm for 2 h.1,2 Temperatures higher than 60 °C caused a progressive reduction in the rate, probably linked to the lower aqueous solubility of CO2 at higher temperatures (up to 150 °C). An increasing reaction time up to 3 h did not improve the degree of carbonation. Furthermore, when “n” is increased above 400 rpm, the Ca conversion decreased. This could be due to nonlinearity in the efficiency of gas/liquid contact with the stirring rate, possibly linked to a change from turbulent to laminar flow. With regard to the effect of the L/S ratio on the process of CO2 mineral carbonation, the highest amount of Ca converted (∼99%) was achieved at the lowest L/S ratio (10 mL/g). When the L/S ratio increased to 30, it equaled the slight decrease of the Ca conversion to 91%. This trend continued up to L/S ≈ 100 mL/g, with Ca conversion falling to 75% relative, but with a further increase (to 200 mL/g), the efficiency of conversion recovered slightly to 82%. This variability in behavior may reflect the complexity of the sulfate/carbonate system, in which the known promoting effects of underlying nuclei of CaSO4 on CaCO3 (co)precipitation may depend upon access of CO32− (aqueous) through a more or less porous product layer.22 Park and Fan17 reported that the beneficial effect of a high L/S ratio during dehydration is useful for saving energy and reducing water treatment. It should be considered that, at a high L/S ratio, extra water is required to extract all Ca into the solution and enhance precipitation of CaCO3. A decrease in the L/S ratio led to a little increase in the rate of conversion, which is possibly due to a higher amount of Ca and, consequently, ionic strength in solution.23 3.3. Kinetics Analysis of Ca Dissolution by H2SO4. Three main kinetic models (see eqs 3−5) were used to evaluate the dominant mechanism(s) operative in the experimental dissolution data. Figure 2 illustrates the plot of the (a) product layer diffusion, (b) film diffusion, and (c) chemical reaction control at
Figure 2. Kinetics analysis of the Ca concentration corresponding to (a) product layer diffusion, (b) film diffusion, and (c) chemical reaction control models at reaction temperatures between 25 and 60 °C in 1.5 M H2SO4.
reaction temperatures between 25 and 60 °C in H2SO4. The product layer diffusion (Figure 2a) and chemical reaction control (Figure 2c) gave the best fit to the measured data, and their combination is evidently rate-limiting for Ca dissolution. The activation energy (Ea) for mineral dissolution was measured from a semi-log plot of the time-independent rate k versus 1/T. As shown in Figure 3, the qualities of fit for chemical reaction control (R2 = 0.9994) and product layer diffusion (R2 = 0.9992) were better compared to the film diffusion model (R2 = 0.9891). When these are taken as the mechanisms that control the dissolution of RG, the values of Ea were 14.78 and 16.52 kJ/mol for product layer diffusion and chemical reaction control, 7462
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Figure 3. Arrhenius plots for the concentration of Ca from byproduct RG using three chosen models of the chemical reaction control, product layer diffusion, and film diffusion.
respectively. Teir et al.24 suggested that chemical reaction control is rate-limiting at the outset but that product layer diffusion gradually predominates as the unreacted surface area decreases and impurities in the product layer accumulate. Gharabaghi et al.25 showed that a low value for the activation energy was associated uniquely with product layer diffusion. In this work, the similarity in Ea values indicates that the dissolution rate of RG is kinetically regulated by a combination of chemical reaction control and product layer diffusion. 3.4. Geochemical Modeling by PHREEQC-2.18. PHREEQC-2.18 was applied to perform equilibrium simulations with data from the LLNL. The saturation state of the CaCO3 and CO2−CaCO3 systems were modeled accordingly. The RG sample contains highly soluble constituents. Contact with water leads to a rapid release of Ca, SO42−, and OH− in the solution and a sharp increase in alkalinity (Figure 4a). The amount of Ca ion converted to CaCO3 is 0.01 and 0.032 mol/L after 5 and 20 min, respectively. The dissolution of Ca(OH)2 is coupled to simultaneous precipitation of calcite (pH > 6.9), which also reaches completion over this time frame. Moreover, the saturation state of CO2−CaCO3 was simulated in the presence of varying amounts of NH4OH. In the temperature range of 25−60 °C, CO2 gas sequentially made a reaction with the Ca-rich solution in the presence of NH4OH to neutralize any carryover of H2SO4 from the dissolution process. As revealed, the presence of H2SO4 played an essential role in the rapid enhancement of the concentration of Ca in the solution, although it has little influence on the saturation state of CO2− CaCO3. By the end of the experiment, the amount of calcite precipitation is increased to 0.05 mol/L, which is increasingly in line with the rate of CO2 uptake (Figure 4b). On the other hand, the high rate of CO2 uptake is qualified to the precipitation of calcite as the concentration of Ca is reduced. With regard to the effect of T on the overall reaction kinetics, the concentration of Ca was increased as a consequence of enhancing temperature up to 60 °C. The precipitation of calcite was kinetically hindered at temperatures less than 55 °C because of the high value of Ea for the desolvation of the strongly hydrated Ca ions. During low-temperature experiments, no evidence of stable calcite formation was found in any of the analyzed samples, indicating that this phase is possibly a metastable form of calcite; however, the kinetics analysis at this level reveals it as an acceptable first product of reaction under this condition. It should be noted that the solution stays undersaturated regarding
Figure 4. Simulation results performed by PHREEQC-2.18 for the saturation state of (a) CaCO3 and (b) CO2−CaCO3 in the solution.
calcite at temperatures less than 55 °C and the rate of CO2 uptake is negligible. This adds weight to our initial thoughts that the metastable form of calcite is the product of rapid precipitation prior to calcite (experimentally) and that, under equilibrium conditions (e.g., T between 55 and 60 °C), calcite would be the dominant calcium-bearing phase.
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CONCLUSION The experimental work demonstrates that the kinetics of sequestration of CO2 into RG are viable, with the slow step being dissolution of CaSO4 in H2SO4. To understand the reaction sequence of the dissolution and carbonation processes, the simulation was found to be a very useful tool because it simulates the exposure of all dissolved species and possible mineral phases at progressive time intervals. Furthermore, the modeling of mineral carbonation of RG delivered important hints for further studies, such as the extended availability of thermodynamic data of characteristic mineral phases contained in industrial wastes. Geochemical modeling was performed in relation to carbonation and dissolution experiments of RG. The results revealed that all Ca extracted in the dissolution of RG could be fully converted to CaCO3 in the carbonation process. Therefore, the geochemical model was found as an applicable method to present mineral phases at each stage. The amount of Ca concentration obtained from the dissolution experiment was successfully implemented in the model and was supplemented by the available amount of precipitated CaCO3. 7463
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AUTHOR INFORMATION
Corresponding Author
*Telephone: +98-914-442-2009. E-mail: omeidrahmani@gmail. com and/or
[email protected]. Author Contributions
Omeid Rahmani conceived and designed the experiments and their interpretations. Ali Kadkhodaie assisted with the manuscript draft. James Highfield contributed technical information and reworked the manuscript draft. Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS This study was supported by the Research & Technology Division in Mahabad Branch, Islamic Azad University, with Vote 35.3.608805. The authors extend their appreciation to Dr. Sahar Zarza for helpful comments.
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NOMENCLATURE d = particle size T = temperature (°C) t = reaction time (s) n = stirring rate pCO2 = partial pressure of carbon dioxide S = mineral surface area (m2) r = reaction rate (mol/s) k = rate constant L/S = liquid/solid ratio (mL/g) Ea = activation energy (J/mol) R = gas constant (0.082 057 L atm mol−1 K−1) DC = degree of conversion (%) XE = Ca concentration (mol/L) Q = flow rate (L/min) M = mass of red gypsum (g)
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REFERENCES
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