Use of Nanoporous FeOOH as a Catalytic Support for NaHCO3

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Use of Nanoporous FeOOH as a Catalytic Support for NaHCO3 Decomposition Aimed at Reduction of Energy Requirement of Na2CO3/NaHCO3 Based CO2 Separation Technology Bryce Dutcher,† Maohong Fan,*,† Brian Leonard,‡ M. Darby Dyar,§ Jinke Tang,|| Elly A. Speicher,§ Pan Liu,|| and Yulong Zhang^ Department of Chemical and Petroleum Engineering and ‡Department of Chemistry, University of Wyoming, Laramie, Wyoming 82071, United States § Department of Astronomy, Mount Holyoke College, South Hadley, Massachusetts 01075, United States Department of Physics and Astronomy, University of Wyoming, Laramie, Wyoming 82071, United States ^ Western Research Institute, Laramie, Wyoming 82072, United States

)

†

ABSTRACT: CO2 capture is typically a costly operation, usually due to the energy required for regeneration of the capture medium. Na2CO3 is one potential capture medium with the potential to decrease this energy requirement. Extensively researched as a potential sorbent for CO2, Na2CO3 is wellknown for its theoretically low energy requirement, due largely to its relatively low heat of reaction compared to other capture technologies. Its primary pitfalls, however, are its extremely low reaction rate during sorption and slow regeneration of Na2CO3. Before Na2CO3 can be used as a CO2 sorbent, it is critical to increase its reaction rate. In order to do so, this project studied nanoporous FeOOH as a potential supporting material for Na2CO3. Because regeneration of the sorbent is the most energy-intensive step when using Na2CO3 for CO2 sorption, this project focused on the decomposition of NaHCO3, which is equivalent to CO2 desorption. Using BrunauerEmmetTeller analysis, Fourier transform infrared spectroscopy, X-ray diffraction, X-ray photoelectron spectroscopy, scanning electron microscopy, transmission electron microscopy, magnetic susceptibility tests, and M€ossbauer spectroscopy, we show FeOOH to be thermally stable both with and without the presence of NaHCO3 at temperatures necessary for sorption and regeneration, up to about 200 C. More significantly, we observe that FeOOH not only increases the surface area of NaHCO3, but also has a catalytic effect on the decomposition of NaHCO3, reducing activation energy from 80 to 44 kJ/mol. This reduction in activation energy leads to a significant increase in the reaction rate by a factor of nearly 50, which could translate into a substantial decrease in the cost of using Na2CO3 for CO2 capture.

’ INTRODUCTION As the world’s demand for energy increases and because most renewable and “green” energy sources are not yet efficient enough for large-scale application, power suppliers are increasingly relying on fossil fuel combustion to meet the world’s energy needs. However, the combustion of fossil fuels causes large amounts of CO2 to be released into the atmosphere faster than natural processes can remove them. It has been noted over the past century that as the global concentration of CO2 has increased, so have average global temperatures. Known commonly as global warming, this phenomenon can lead to severe climate change, which can dramatically affect the planet and its life forms. To this end, it has become increasingly imperative to decrease levels of anthropogenic CO2. As with most separation technologies, it is easiest to separate a minor component when its concentrations are high. In the case of CO2, this is true at the point sources. The largest contributors to anthropogenic CO2 are fossil fuel power plants, particularly coal-fired plants. Concentrations of CO2 in the flue gas emitted r 2011 American Chemical Society

by these plants are typically around 1015%, making it an ideal gas stream for separating CO2 before its release into the atmosphere. While there is little technical challenge to capturing CO2, doing so cost-effectively is difficult. Current technologies include liquid solvent scrubbing, membrane separation, cryogenic separation, and solid-state sorption.16 However, all of these technologies require large amounts of energy, typically through regeneration of the capture medium, as with liquid scrubbing and sorption onto solids, or pressurizing the flue gas in advance of separation, as with membranes and at times with scrubbing and sorption as well. All of these methods exact large energy penalties, significantly increasing the cost of capturing CO2. The best technology currently available for CO2 capture is absorption using aqueous amines, that is, “amine scrubbing,” a method that is both effective and less costly than other available Received: May 25, 2011 Revised: June 28, 2011 Published: July 05, 2011 15532

dx.doi.org/10.1021/jp204899r | J. Phys. Chem. C 2011, 115, 15532–15544

The Journal of Physical Chemistry C

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methods. Amine scrubbing is also well-known, having been used for CO2 capture in the petroleum industry for over 60 years. Its primary advantage is the flexibility of the available solvents, because amines can be tailored to be easily regenerable while maintaining their sorption properties.13 But while amine scrubbing has become a benchmark for other technologies, it nonetheless faces problems, including poor thermal stability of the solvents,1,2,7 poisoning from SOx, NOx, and even O2,13 amine evaporation,7,8 and toxic and corrosive properties.1,2,7,8 Further, dilution of the amine in water increases regeneration costs while decreasing its overall capacity for CO2 sorption.13,7,8 Research has shown that use of solid sorbents can alleviate many of the issues of liquid scrubbing in general and amine scrubbing in particular. Generally, solid sorbents have a theoretically lower energy requirement for desorption due to higher loading, lower heats of reaction, and/or lower heat capacities.9 In addition, many materials used in aqueous scrubbing can be impregnated onto a solid support, with amines representing one of the primary materials studied for this technology.913 Impregnating a solid support with the amine eliminates the dilution effect of water, thereby increasing the apparent CO2 capacity and reducing the amount of energy needed to heat the material. Furthermore, due to the high water vapor content of flue gas, any reactions that require water can still proceed. While supported amines are currently used to capture CO2 from enclosed spaces such as submarines and space shuttles,13 the raw materials are currently too expensive to be applied in large-scale industrial settings.12 As an alternative to supported amine sorbents, alkali metal carbonates can be used as solid sorbents via the reaction M2 CO3 þ H2 O þ CO2 T 2MHCO3

ðR1Þ

where M represents the alkali metal (typically sodium or potassium). It has been shown that pure, solid Na2CO3 is capable of removing 90% of the CO2 from synthetic flue gas with little deterioration in sorption properties through five cycles.14,15 The primary disadvantage to using carbonate materials, however, is their very slow reaction rate.16 The reaction rate can be increased by impregnating the carbonate onto a porous matrix, but so far this has been observed only qualitatively.1719 It has been generally noted that potassium carbonate is superior to sodium carbonate in terms of both CO2 capacity and kinetics;2022 however, the primary advantage of using sodium carbonate over potassium carbonate is its lower price. When sodium carbonate is used, R1 can be written specifically as Na2 CO3 þ H2 O þ CO2 T 2NaHCO3

ðR2Þ

If the reaction rate can be increased sufficiently, Na2CO3 has several advantages. Foremost among these is cost and availability, as mentioned above, especially in Wyoming, which has the world’s largest trona (Na3H(CO3)2 3 2H2O) deposits and supplies 90% of the soda ash in the United States.23 In addition, due to the stoichiometry of their respective reactions, Na2CO3 can, in theory, have higher loadings than amines. Na2CO3 exhibits very high thermal stability, decomposing at temperatures above 800 C, whereas some amines decompose at temperatures as low as 120 C.1 Moreover, Na2CO3 is easier to handle because it is not corrosive. Na2CO3 has a negligible vapor pressure, so there is little loss due to evaporation. As with amines, some losses do occur from irreversible reactions with SO2 and NOx.14,24

However, due to its low price and ease of handling, Na2CO3 is more easily replaced. Sorption of CO2 onto Na2CO3 is spontaneous at flue gas conditions of roughly 5080 C. The most energy-intensive step, and therefore the most costly, involves regeneration of the spent material by heating the sorbent to temperatures high enough to shift the equilibrium of R2 toward the left, which typically occurs above 100 C. While the decomposition of pure NaHCO3 has been studied extensively,2528 there is little information on the effects of supporting materials in this method. A supporting material should not only provide a large surface area for the Na2CO3, but also have a catalytic effect on the decomposition of NaHCO3. The use of a catalyst reduces activation energy, and thus the total energy required for a reaction to proceed (in this case, the reverse reaction of R2). To the best knowledge of this research team, this effect has, to date, been ignored by other researchers studying Na2CO3. One potential support, FeOOH, is a widely available, nanoporous compound already studied as a catalyst.29,30 FeOOH decomposes via the reaction 2FeOOH T Fe2 O3 þ H2 O

ðR3Þ

However, the reaction does not proceed to the right except at temperatures above 200 C,3133 making FeOOH thermally stable for use in flue gas conditions. The aim of this research is to investigate the use of nanoporous FeOOH as a multifunctional support for Na2CO3-based CO2 separation. This paper focuses on the surface characteristics of the sorbent and the kinetics of desorption of CO2, or the decomposition of NaHCO3. It should also be noted that optimizing sorption capacity lies beyond the scope of this work.

’ EXPERIMENTAL SECTION Preparation of NaHCO3/FeOOH Samples. Pure NaHCO3 was obtained from BDF Chemical (purity >99.0%); FeOOH of an unknown polytype was provided by Siemens (purity >66.8%, the balance is mostly water). NaHCO3 was loaded onto FeOOH by mixing predetermined amounts of NaHCO3 and FeOOH in a volume of distilled water sufficient to dissolve the NaHCO3. The amounts of NaHCO3 and FeOOH were chosen to create the desired percentage by weight of the final NaHCO3/FeOOH product (hereafter called “NHF”). Target weight percentages were 20, 50, and 90 wt % of NaHCO3 in the sample. Specific steps for creating NHF included stirring the mixture for at least 5 h, removing all water under vacuum in a rotary evaporator at 60 C, and grinding and sieving the resulting solid to obtain particles of a diameter less than 300 μm. For clarification, any Na2CO3 resulting from the decomposition of NaHCO3 in NHF remains on the surface of the FeOOH, and the resulting composite material is referred to as “NF.” The wt% of Na2CO3 in a given NF sample refers to the weight percent of the Na2CO3 created from the decomposition of NaHCO3 in the corresponding NHF sample and can be calculated based on stoichiometry or the mass balances through R2. Characterization. BrunauerEmmetTeller analysis (BET) was used to study FeOOH before heating and for NaHCO3 and each NHF sample before and after NaHCO3 decomposition. BET analyses employed a Micromeritics TriStar 3000 V 6.04 A device. Fourier transform infrared spectroscopy (FTIR) spectra were collected for pure FeOOH before and after heating to 180 C and for pure NaHCO3, 20, 50, and 90 wt % NHF, both 15533

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Table 1. BET Results for Pure FeOOH, Pure NaHCO3, and 20, 50, and 90 wt % NHF before and after Decomposition at 180C surface area pore volume

pure

pore

sample

(m2/g)

(cm3/g)

size (nm)

as received

130.722

0.115

3.390

FeOOH pure

before decomposition

1.274

0.002

4.550

NaHCO3 after decomposition 20 wt % before decomposition

5.288 37.985

0.008 0.046

5.010 4.620

NHF

after decomposition

31.351

0.040

4.700

50 wt %


22.744

0.028

4.995

NHF

after decomposition

10.561

0.020

6.500

90 wt %


5.276

0.008

5.389

NHF

after decomposition

7.240

0.012

5.864

before and after decomposition at 180 C using a Thermo Nicolet Magna-IR 760 spectrometer. X-ray diffraction (XRD) data for pure FeOOH, NaHCO3, and Na2CO3 were collected using a SCINTAG XDS2000 automated powder diffraction system equipped with a θθ XRD goniometer and a solid-state X-ray detector. XRD data for all NHF samples before and after decomposition at 180 C were collected using a Philips X’Pert diffractometer with Cu KR radiation. X-ray photoelectron spectroscopy (XPS) was done on pure FeOOH and NaHCO3 before and after heat treatment at 180 C to check for the presence of impurities. XPS analyses were performed with a Physical Electronics PE5800. Further characterization was done for pure FeOOH before and after heating to 180 C and pure NaHCO3 and 20 wt % NHF samples before and after decomposition at 180 C using an FEI Quanta FEG 450 field-emission scanning electron microscope (SEM). Transmission electron microscopy (TEM) images were obtained for FeOOH before heating and 20, 50, and 90 wt % NHF before decomposition using a Hitachi (H-7000) TEM. Magnetic susceptibility curves were acquired for FeOOH before heating to 180 C and for 20, 50, and 90 wt % NHF both before and after decomposition at 180 C. Magnetic susceptibility curves were measured using a Physical Properties Measurement System (PPMS) from Quantum Design. M€ossbauer spectra at 295 and 4 K were acquired from pure FeOOH before heating and for 20 wt % NHF before and after heating to 180 C. A 50 mCi 57Co in Rh source on a WEB Research Co. model WT302 spectrometer (Mount Holyoke College) was used. The instrument is equipped with a Janus closed cycle He cooling system to allow measurements over the range from 4 to 295 K. M€ossbauer results were calibrated against a 25 μm R-Fe foil. CO2 Desorption or NaHCO3 Decomposition. The decomposition of NaHCO3 on the surface of NHF, which can be thought of as the desorption of CO2, was studied using a TA Instruments SDT Q600 thermogravimetric analyzer (TGA). Each test used 20100 mg of sample, depending on the density of the material, loaded in an alumina tray. All samples were heated to the test temperature at 20 C/min (the highest achievable heating rate), and then kept isothermally for 10 min. Argon flowing at 0.1 L/min was used as a carrier gas in all tests. As a reference, pure NaHCO3 with particles of less than 300 μm was decomposed at temperatures of 100, 120, 140, 150, and 160 C. CO2 desorption from 20, 50, and 90 wt % NHF

Figure 1. FTIR spectra of pure FeOOH before and after heating to 180 C, pure NaHCO3, and pure Na2CO3.

samples was tested in the temperature range of 100140 at 10 C increments. The 20 wt % samples were initially heated at 20 C/min to a temperature of 80 C, kept isothermally for 10 min, and then heated at 20 C/min to the test temperature. This step was required to remove water adsorbed to the surface of the sample. No appreciable water appeared on the 50 and 90 wt % NHF samples, based on material balance, so this step was bypassed. Each sample was tested at least three times at a given temperature, and the reported results are the average value.

’ RESULTS AND DISCUSSION Characteristics of Materials. BET results are shown in Table 1. FeOOH is initially a porous material with 3.4 nm pore diameter, 0.12 cm3/g pore volume, and 130 m2/g surface area. NaHCO3 is not porous, having only 0.0015 cm3/g of pore space. It becomes slightly more porous after decomposing into Na2CO3, upon which pore volume increases to 0.0075 cm3/g. The porosity of FeOOH is significantly reduced after loading NaHCO3 onto the surface with pore volume dropping by roughly a factor of 2.5 to 0.046 cm3/g in the 20 wt % NHF, and pore diameter increasing to 4.6 nm. Its surface area decreases more substantially from 130 to 38 m2/g. One would expect the pore volume of NHF to increase slightly as the NaHCO3 on its surface decomposes to Na2CO3, as with the pure components. In fact, pore volume decreases slightly, as does the surface area. However, pore diameter increases slightly, as expected. Absorption FTIR spectra are shown in Figures 1 and 2. Figure 1a,b shows pure FeOOH before and after heating to 180 C, respectively. This FTIR spectrum matches well with the spectrum of β-FeOOH, commonly known as akaganeite.34,35 The broad peak at 3400 cm1 and the peak at 1630 cm1 result from OH stretching and H2O bending motions, respectively, from surface-adsorbed water.3436 The peaks at 1490 and 1360 cm1 may be due to OH bending from surface-adsorbed water shifted uniquely for akaganeite due to its Hollandite crystal structure,35,37,38 but these peaks have also been attributed to organic contamination from the environment during preparation.39 The strong peak at 1100 cm1 is most likely due to an impurity, possibly sulfur in the form of sulfate,40,41 or SiO2, 15534



Figure 2. FTIR spectra of 20, 50, and 90 wt % NHF freshly prepared and after decomposition at 180.

naturally present as an impurity in many iron compounds.36 The weak peak at 980 cm1 may be due to FeOH bonds,34 with the shoulder at 700 cm1 possibly due to OH.35 Peaks at 600 and 440 cm1 are likely due to FeO bonds.3436 After heating the FeOOH to 180 C, the peaks at 1490 and 1360 cm1 appear to decrease in intensity, which is likely because either the surface adsorbed water or organic contaminants have been partially removed during heating. The decomposition of FeOOH via R3 is not likely the cause for the decrease in intensity because other characterization techniques, including M€ossbauer spectra and magnetic susceptibility curves, discussed below, indicate that there is no phase change. Figure 1c shows the FTIR spectrum of NaHCO3, and Figure 1d shows the spectrum of Na2CO3 derived from the decomposition of NaHCO3 at 180 C. Both spectra match well with values in the literature.4245 The noteworthy peaks at 1300 cm1 for NaHCO3 and 1400 cm1 for Na2CO3 are due to carbonate asymmetric stretching.45 Figure 2a,b shows the FTIR spectra of 20 wt % NHF before and after being heated to 180 C, respectively. In both curves, the characteristic peaks of FeOOH are prevalent. Two new peaks at 1500 and 1350 cm1 are most likely due to the formation of a slight amount of iron carbonate while drying the sample during preparation; these peaks match those of iron carbonate in the presence of water.46,47 A weak doublet appears at 860 cm1, which may be attributed to SiO2.36 The characteristic peaks of NaHCO3 and Na2CO3 do not appear, respectively, in Figure 2a,b, which is consistent with XRD results (discussed below) and likely due to the small amount of NaHCO3 loaded onto NHF. Figure 2c,d shows the FTIR spectra of 50 wt % NHF before and after decomposition. The peak at 1350 cm1 is likely due to a small amount of iron carbonate in the presence of trace amounts of water, as with the 20 wt % NHF. There are two potential explanations for the peaks at 1450 and 880 cm1, which are present both before and after decomposition. These may be due to Na2CO3, which forms during preparation of the sample; however, because of the conditions during preparation, this is unlikely. They may also be due to a small amount of iron carbonate in the absence of water.48,49 In either case, the amount formed can be assumed negligible because the expected material balance of the decomposition of

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Figure 3. XRD patterns of pure FeOOH before and after heating to 180 C, pure NaHCO3, and pure Na2CO3.

Figure 4. XRD patterns of 20, 50, and 90 wt % NHF freshly prepared and after decomposition at 180.

NaHCO3 is not significantly affected. Evidence that NaHCO3 decomposes to Na2CO3 when heated to 180 C is seen in the disappearance of the peak at 840 cm1, along with an increase in intensity in the peak at 880 cm1. The FTIR spectra of 90 wt % NHF, shown in Figure 2e,f, are very similar to the 50 wt % spectra. One notable distinction in the 90 wt % spectra is the disappearance of the peak at 600 cm1 after decomposition. This peak is assigned to the FeO bonds in FeOOH, and its disappearance is likely due to the heterogeneity of the sample rather than the decomposition of FeOOH. XRD results are given in Figures 3 and 4. Figure 3a,b shows the XRD patterns of pure FeOOH before and after heating to 180 C. Peaks appear at 2-θ values of 36 and 62. The peaks are broad, indicating that the crystallite size is small or the particles are amorphous.50,51 These peaks match well with many different nearly amorphous iron oxide and hydroxide compounds,5257 15535



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Figure 5. SEM images of pure FeOOH, pure NaHCO3, pure Na2CO3, and freshly prepared and decomposed 20 wt % NHF samples.

but other characterization techniques verify that the compound is in fact FeOOH. Because of the nearly amorphous nature of the material, it cannot be determined which polytype of FeOOH is present from XRD data. Figure 3c,d shows XRD patterns of

NaHCO3 and Na2CO3 derived from the decomposition of NaHCO3 at 180 C and matches well with data reported in the literature.58,59 Consistent with FTIR spectra, the XRD patterns of 20 wt % NHF in Figure 4a before CO2 desorption 15536



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Figure 6. TEM images of FeOOH and 20, 50, and 90 wt % NHF.

and Figure 4b after CO2 desorption do not show the characteristic peaks of NaHCO3 or Na2CO3. Neither do these peaks appear for the 50 wt % NHF before and after decomposition in Figure 4c,d; the peak due to FeOOH at 62 is not easily discernible for the 50 wt % NHF. The peaks for NaHCO3 and Na2CO3 do appear in the 90 wt % NHF before and after heating in Figure 4e,f. These findings are consistent with the work of Zhao and his colleagues,60 in which they studied the dispersion of Na2CO3 on γ-Al2O3. In each set of data for all NHF samples, both before and after heating, there are no unexpected peaks indicating the formation of a new crystalline material. This provides further evidence that NHF is stable for use as a support for NaHCO3 in CO2 desorption. The elemental analysis provided by XPS analysis matches well with the expected elemental composition of FeOOH, NaHCO3, and Na2CO3. FeOOH contains nearly 6% impurities by element. The primary elemental impurity in the FeOOH was sulfur, as high as 3.4% in the elemental analysis. Other impurities in the FeOOH include aluminum, calcium, magnesium, and silicon. The presence of sulfur and silicon supports both possible explanations for the peak at 1100 cm1 in the FTIR spectra of FeOOH. Trace impurities in NaHCO3 and Na2CO3 include titanium and chlorine, but always in amounts less that 0.5%. Due to these impurities, it should be noted that the use of “pure

NaHCO3” and “pure FeOOH” in this study implies each individual compound as received, not mixed with one another to form NHF. SEM images are provided in Figure 5. The porous nature of pure FeOOH can be seen in Figure 5a. Although the pores appear to decrease in size, as seen in Figure 5b, the heating process does not significantly affect the structure. Figurs 5c,d shows NaHCO3 before and after its decomposition to Na2CO3, respectively. While NaHCO3 has a very crystalline structure initially, consistent with BET results, the crystals rearrange and form a network of pores as the molecules thermally decompose and lose mass in the form of H2O and CO2. Figure 5e shows 20 wt % NHF. Much of the visibly porous nature of FeOOH is retained, and no clear sign of NaHCO3 is visible, indicating that the NaHCO3 is inside the pores. As shown in Figure 5f, there is no significant change after NaHCO3 on the surface of the NHF has decomposed to create 14 wt % NF. TEM images for pure FeOOH and each NHF sample are shown in Figure 6. The nanoporous nature of FeOOH can be seen on the surface of the particle in Figure 6a. Two locations are shown for the 20 wt % NHF in Figure 6b. As seen in Figure 6b1, a netlike structure forms after loading NaHCO3 onto the sample, which suggests self-assembly of the NHF material during preparation. However, large crystals of NaHCO3 are present in 15537



Figure 7. Magnetic susceptibility curves (χ-T curves) of FeOOH and 20, 50, and 90 wt % NHF before and after decomposition.

Figure 6b2. This netlike structure, as well as the crystals of NaHCO3, are visible in Figure 6c for 50 wt % NHF. In Figure 6d, it is apparent that the surface of the 90 wt % NHF is nearly all NaHCO3, which is expected due to the large wt % of NaHCO3. Magnetic susceptibility curves are shown in Figure 7. A Neel temperature of roughly 25 K is observed for the pure FeOOH as well as each of the NHF samples before decomposition. Initially, this suggests that the FeOOH is γ-FeOOH, or lepidocrocite, rather than akaganeite, as suggested by FTIR and M€ossbauer spectroscopy, discussed below, because the Neel temperature of akaganeite is 299 K, while that of lepidocrocite is 77 K.61 This is likely because the particle size of FeOOH is sufficiently small enough that it becomes superparamagnetic.62,63 The peak at 25 K is therefore likely to be the blocking temperature. After the NHF samples have been decomposed, the peak shifts slightly to about 45 K. This is attributed to decrease in pore size and increase in crystallinity of the FeOOH during heating, as seen in BET analyses of NHF samples and SEM images. Interpretation of the M€ossbauer results is based on the fundamental parameters of these spectra, which can be used to identify Fe-containing minerals that happen to have distinct peak locations (in this case, Fe oxides). If the material is paramagnetic, its M€ossbauer spectrum will be composed of doublets; if magnetic ordering is present, then the result is a sextet. M€ossbauer parameters include isomer shift (δ, or IS), which reflects the s-electronic charge density at the nucleus; quadrupole splitting (Δ, or QS), which arises from a distribution of surrounding charges with less than cubic symmetry; and the magnetic hyperfine field (Bhf) which indicates magnetic order or an externally applied field.64 The former two parameters are expressed in terms of velocity relative to the movement of the source toward and away from the sample. For example, for 57Fe, moving the source at a velocity of 1 mm/sec toward the sample, by eq (2) of Dyar et al.,64 increases the energy of the emitted photons by (14.413 keV)(v/c) = 4.808 108 eV, or about 10 natural linewidths. The hyperfine field magnitude is given in Teslas. M€ossbauer parameters vary with temperature, and magnetic ordering at extremely low temperatures (e.g., 4K) often produces sextets with diagnostic parameters characteristic of specific phases. The area of the peaks that make up each doublet or sextet is proportional to the

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percentage of Fe in that site. In this study, we are particularly looking for changes in the parameters of the starting material that would indicate a transformation to a different Fe oxide. M€ossbauer spectra at 4 K and at room temperature are shown in Figure 8 with individual distributions plotted along with the raw data. M€ossbauer parameters are given in Table 2. All spectra are alike within the acceptable errors,65 suggesting that the site in which the Fe atom resides is not changing. The 4 K parameters, which are most diagnostic, are IS = 0.45 and 0.48 mm/s, QS = 0.08 and 0.05 mm/s, and BHf = ∼45 and 49 T for each of the two distributions, respectively. On the basis of parameters of the five FeOOH polymorphs,66 this material is definitely not goethite (R-FeOOH) because QSgoethite = 0.25 mm/s. Lepidocrocite (γ-FeOOH) or feroxyhite (δ-FeOOH) are not a good match either, because QS ∼ 0.02 and 0.12 mm/s, respectively. The ε-FeOOH polymorph may be a match, though its 4 K spectra usually has only a single sextet with parameters of IS = 0.49 mm/s, QS = 0.12 mm/s, and BHf = 52.3 T. The best match for the material studied here is akaganeite (β-FeOOH), which commonly has multiple overlapping broad distributions with parameters close to those measured here. M€ossbauer spectra show that the FeOOH is not decomposing nor is it reacting in significant amounts with the addition of NaHCO3. NaHCO3 Decomposition or CO2 Desorption Kinetics. Rate Equation. The kinetics of desorption of CO2 or decomposition of NaHCO3 via R2 are studied using reported kinetic models.67 According to these models, the reaction rate is modeled by FðRÞ ¼ kt

ðE1Þ

where R is the mass fraction of decomposed NaHCO3, k is the kinetic rate constant of the reverse reaction of R2, and t is time. F(R) is a function that depends on the assumed mechanism of the reaction. It is necessary to determine the Rt curves in order to determine the F(R)t curves. The equation for R is w0 w ðE2Þ R¼ w0 where w0 is the initial weight of NaHCO3 and w is the weight of NaHCO3 at any t. The values of R for pure NaHCO3 and the NHF samples can be determined based on the stoichiometry of R2 and the corresponding weight loss measured by the TGA during reaction. The Rt curves of the three samples at 120 C are shown in Figure 9. The time scale in Figure 9 has been shifted so that only data collected under isothermal conditions is used. Seven common forms of F(R) are listed in Table 3.67,68 Appropriate forms of F(R) are determined by fitting F(R)t. On the basis of E1, a form of F(R) that gives a linear fit is an appropriate model for the decomposition reaction, where the slope of the line is equal to k. Furthermore, with those forms of F(R) involving reaction orders, the derived reaction order must vary within a narrow range. Of these seven forms of F(R), the best consistent fit was found to be the AvramiErofeyev model, Am, or FðRÞ ¼ ½ ðlnð1 RÞ1=m

ðE3Þ

where m is the reaction order, which agrees with literature findings for pure NaHCO3.25,27,28 This model implies that, for both pure NaHCO3 and NHF, random nucleation of Na2CO3 is the ratelimiting step of the reverse reaction of R2.67,68 Table 4 shows the linear regression coefficients of all seven models listed in Table 3 and the determined reaction orders according to E3. A plot of 15538



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Figure 8. M€ossbauer spectra at room temperature and at 4K of FeOOH and 20 wt % NHF before and after decomposition.

F(R)t is shown in Figure 10 for samples at 120 C. As the value of R approaches 1, however, the diffusion of the product gases (CO2 and H2O) through the reacted solid becomes the ratelimiting step. This stage can be observed visually as the linear section of the Rt curves for the NHF samples in Figure 9. Because of its slow reaction rate, pure NaHCO3 takes roughly 30 min to reach this stage. The Am model is not applicable in the diffusion-controlled region. In general, the rate equation for decomposition of NaHCO3 can be written as dR ¼ kð1 RÞm dt

ðE4Þ

Table 5 shows the values of m and k at each condition averaged over three experiments. The average value of m for pure NaHCO3 is found to be 1.1, consistent with values in the literature.2528 The value of m for all NHF samples ranges from

0.13 to 0.72, depending on temperature and the wt % of NaHCO3 present. This variability is not unusual, and has been observed in other studies using the Am model.67 The average value of m for NHF samples is 0.42, meanig that this reaction, when catalyzed by FeOOH, is roughly half order in regards to NaHCO3, whereas it is roughly first order when uncatalyzed. FeOOH not only reduces the activation energy of NaHCO3 decomposition, but also changes its reaction order since FeOOH changes the reaction pathways of the reaction. The decrease of m from 1.1 to 0.42 when FeOOH is used can lead to the increase in reaction rate, (dR)/(dt), since (1 R) is always equal to or smaller than 1 according to E4. It is apparent that the 50 wt % NHF has the lowest value of m, and determining if this is a significant trend will be part of a future study. The highest k values observed are for 90 wt % NHF, and the lowest k values for an NHF sample are for 20 wt % NHF. However, in all cases the value of k is increased by at least a factor of 8 compared to pure NaHCO3, and in some cases by almost a factor of 50. 15539



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Table 2. M€ ossbauer Spectroscopy Parameters 20 wt % NHF sample

pure β-FeOOH

spectrum temperature (K)

295


20 wt % NHF

after decomposition

295

pure β-FeOOH

295


after decomposition 4

4

4

δ (mm/s)

0.35

0.35

0.34

0.45

0.44

0.43

Δ (mm/s)

0.91

0.92

0.94

0.08

0.07

0.05

Γ (mm/s)

0.42

0.60

0.56

1.04

BHf (T) % Areaa

38

50

61

0.91

0.95

45.1

45.1

45.1

37

43

28

δ (mm/s)

0.35

0.35

0.35

0.48

0.48

0.48

Δ (mm/s) Γ (mm/s)

0.55 0.38

0.62 0.42

0.69 0.34

0.05 1.10

0.05 0.75

0.04 1.09

BHf (T) % Areaa χ2b |χ2|b

49.8

49.3

49.9

62

50

39

63

57

72

503

408

535

2533

1012

3300

0.98

0.79

1.03

4.94

1.97

6.44

a

Relative peak areas are tabulated but have little significance in this study, because FeOOH peaks are heavily overlapped and the sextets represent multiple overlapping distributions.65 b Values of |χ2| (normalized) and χ2 (for all 512 channels) are calculated using 1/(N n)∑N I=1[(YC(I) YD(I))/((YD(I))1/2)]2, where N is the number of points fit, n is the number of parameters fit, YC(I) is the calculated value at each channel, and YD(I) is the experimental data at each channel. The optimum value for χ2 is one times the number of degrees of freedom (N n).

Table 3. Seven Common Forms of F(R)a model type (based on mechanism)

a

Figure 9. Progression of R with time (heated at 20 C/min to 120 C, Ar flow = 0.1 L/min, isothermal at t = 0).

Arrhenius Form. The reaction rate coefficient is related to the activation energy through the Arrhenius form,69 ln k ¼

Ea þ ln A RT

ðE5Þ

where Ea is the activation energy, R is the molar gas constant, T is the absolute temperature, and A is the frequency factor. The activation energy and frequency factor can be determined for a sample by making a linear fit to the relationship between ln k and 1/T. Developed and shown in Figure 11, the Arrhenius plots for the three NHF samples as well as pure NaHCO3 are based on values of k reported in Table 5 and the temperatures at which

F(R)

symbol

1D diffusion 2D diffusion

D1 D2

R R + (1 R)ln(1R)

3D diffusion

D3

[1 (1R)1/3]2

GinstlingBrounshtein

D4

1 (2/3)R (1 R)2/3

ProutTompkins

Au

ln(R/(1 R)

contracting surface/volume

Rn

1 (1 R)1/n

AvramiErofeyev

Am

[ln(1 R)]1/m

2

The values of n and m are reaction orders.

they were measured. The regression coefficient of this fit and the resulting Arrhenius parameters are shown in Table 6. In all NHF samples, the value of the activation energy decreases relative to pure NaHCO3. Even though it does not have the highest k value, 50 wt % NHF has the highest reduction in activation energy, about 50% of the activation energy of pure NaHCO3. In Arrhenius theory, A represents the likelihood of molecules colliding, and by effectively diluting NaHCO3 with the FeOOH, the value of A decreases compared to pure NaHCO3, as expected. The 90 wt % NHF has a higher activation energy than 50 wt % NHF, which is expected because it approaches pure NaHCO3 in composition, and very little of the NaHCO3 is in contact with the FeOOH. This also makes sense when considering that the value of A for 90 wt % NHF is higher than for 50 wt % NHF. However, 20 wt % NHF also has a higher activation energy, which does not follow the expected trend. One would expect the activation energy of 20 wt % NHF to be low due to the abundant surface area provided by the presence of excess FeOOH. It is believed that agglomeration occurs when preparing 20 wt % NHF, creating large crystals of NaHCO3 that do not form a large contact area with FeOOH. The value of A for 20 wt % NHF, which is higher than 50 wt %, supports this hypothesis. Similarly, the reaction order follows the same trend as activation 15540



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Table 4. Typical Correlation Coefficients of Each Form of F(R)t listed in Table 3 (T = 120C, n = 3, TGA heated at 20C/min to listed temperature, 0.1 L/min flow Ar, 20 wt % NHF dehydrated at 80C for 10 min) sample

T (C)

D1

D2

D3

D4

Au

Rn

Am

m

pure NaHCO3

100

0.9815

0.9809

0.9803

0.9807

0.9434

0.9986

0.9992

0.9

120 140

0.9968 0.9942

0.9942 0.9851

0.9906 0.9695

0.9931 0.9804

0.9850 0.9728

0.9996 0.9994

1.0000 0.9991

1.0 1.2

150

0.9942

0.9851

0.9695

0.9804

0.9728

0.9994

0.9991

1.2

160

0.9986

0.9966

0.9867

0.9941

0.9929

0.9995

0.9999

1.2

100

0.9963

0.9952

0.9930

0.9945

0.9643

0.9943

0.9979

0.8

110

0.9855

0.9905

0.9948

0.9921

0.9650

0.9794

0.9995

0.4

120

0.9916

0.9945

0.9971

0.9955

0.9887

0.9916

0.9992

0.4

130

0.9913

0.9939

0.9964

0.9948

0.9900

0.9918

0.9993

0.3

140 100

0.9936 0.9973

0.9961 0.9984

0.9982 0.9986

0.9969 0.9986

0.9920 0.9753

0.9940 0.9921

0.9995 0.9989

0.4 0.6

110

0.9761

0.9799

0.9839

0.9813

0.9676

0.9737

0.9983

0.2

120

0.9858

0.9882

0.9907

0.9891

0.9814

0.9847

0.9993

0.2

130

0.9882

0.9896

0.9912

0.9902

0.9865

0.9880

0.9997

0.2

140

0.9969

0.9971

0.9973

0.9972

0.9969

0.9970

0.9988

0.1

100

0.9994

0.9996

0.9997

0.9997

0.9953

0.9982

0.9997

0.6

110

0.9983

0.9991

0.9996

0.9993

0.9929

0.9967

0.9997

0.5

120 130

0.9977 0.9959

0.9987 0.9977

0.9994 0.9991

0.9990 0.9983

0.9934 0.9936

0.9965 0.9957

0.9996 0.9999

0.5 0.5

140

0.9979

0.9990

0.9997

0.9994

0.9967

0.9979

0.9998

0.6

90 wt % NHF

50 wt % NHF

20 wt % NHF

Table 5. Average Reaction Order and Correlation Factor for Each Sample and Temperature (Am model used with respective reaction orders, TGA heated at 20C/min to temperature, 0.1 L/min flow Ar, 20 wt % NHF dehydrated at 80C for 10 min) sample

temperature (C)

m

k (min1)

pure NaHCO3

100

0.9

0.005

0.9992

120 140

1.0 1.2

0.02 0.06

1.0000 0.9991

150

1.2

0.13

0.9991

160

1.2

0.29

0.9999

100

0.7

0.19

0.9996

110

0.6

0.25

0.9994

120

0.4

0.49

0.9995

130

0.4

0.89

0.9990

140 100

0.3 0.6

1.32 0.20

0.9975 0.9989

110

0.4

0.32

0.9989

120

0.1

0.46

0.9994

130

0.1

0.59

0.9997

140

0.1

0.84

0.9995

100

0.5

0.06

0.9997

110

0.5

0.13

0.9998

120 130

0.5 0.5

0.23 0.35

0.9998 0.9998

140

0.5

0.50

0.9998

90 wt % NHF

50 wt % NHF

Figure 10. Progression of F(R) with time (heated at 20 C/min to 120 C, Ar flow = 0.1 L/min, isothermal at t = 0).

energy with the concentration of NaHCO3 in NHF, a further indication that the thermal decomposition characteristics of 20 and 90 wt % NHF samples are more like those of pure NaHCO3. This further suggests that large crystals of pure NaHCO3 are present in the 20 and 90 wt % NHF samples to a more significant extent than 50 wt % NHF. TEM provides evidence for such crystals, but they are not apparent in SEM. This phenomenon will be the subject of future investigation. Catalytic Mechanism. HCO3 has a resonance hybrid structure involving a central CdO bond, OH, and an O bonded to the C. Without the presence of FeOOH, decomposition of

20 wt % NHF

R2

NaHCO3 can be expressed as HCO3 T OH þ CO2 15541

ðR4Þ



ARTICLE

temperatures according to FeOOH þ HCO3 T FeOðOHÞ2 þ CO2

ðR6Þ

The inherently unstable intermediary, FeO(OH)2, can then readily release the OH when another HCO3 approaches to form H2O and CO32- via the reaction FeOðOHÞ2 þ HCO3 T FeOOH þ H2 O þ CO3 2 ðR7Þ This mechanism mirrors that of pure NaHCO3 decomposition. However, due to the presence of FeOOH, and therefore the intermediate FeO(OH)2, the bond between the OH and the CdO is more easily broken, improving the kinetics of R4. In a second proposed mechanism, HCO3 initially bonds with FeOOH in a manner similar to the sorption of aqueous arsenate and arsenite onto iron oxides and hydroxides7779 via FeOOH þ HCO3 T FeOðCO3 Þ þ H2 O Figure 11. Arrhenius plots of the decomposition reactions of pure NaHCO3 and 20, 50, and 90 wt % NHF.

ðR8Þ

In an additional step compared with the sorption of arsenate and arsenite, CO2 is released via the reaction FeOðCO3 Þ þ HCO3 T FeOOH þ CO3 2 þ CO2 ðR9Þ

Table 6. Arrhenius Equation Regression Correlation Coefficients, Frequency Factors, and Activation Energy of the CO2 Desorption of Pure NaHCO3 and NHF Samplesa sample

R2

A (min1)

Ea (kJ/mol)

pure NaHCO3

0.9988

9.66 109 ( 3.16 108

86 ( 2.5

90 wt % NHF

0.9529

2.65 108 ( 2.43 107

64 ( 5.8

50 wt % NHF

0.9493

4.86 105 ( 4.06 104

44 ( 3.5

20 wt % NHF

0.9899

4.02 108 ( 1.72 107

69 ( 2.8

a

(Am model used with reaction orders from Table 5, TGA heated at 20C/min to temperature, 0.1 L/min flow Ar, 20 wt % NHF dehydrated at 80C for 10 min).

Because of the high strength of the bond between OH and CdO, R4 is a very slow reaction.7072 The resultant OH reacts with another HCO3 and forms stable water and CO32 in the form of OH þ HCO3 T H2 O þ CO3 2

ðR5Þ

R5 is a fast reaction.7375 Obviously, R4 is the limiting step for the NaHCO3 decomposition when catalysts are not used. Therefore, catalysts should be used to avoid the occurrence of or increase the rate of R4 in the whole NaHCO3 decomposition process. This research shows that FeOOH can considerably reduce the activation energy of the reverse reaction of R2. Furthermore, characterization tests show that FeOOH is stable and that it does not react with NaHCO3 to form unexpected compounds in the NHF samples during decomposition of NaHCO3. It is therefore reasonable to assume that FeOOH plays a catalytic role in NaHCO3 decomposition. So far, the authors think that one of the three following potential mechanisms is likely involved with the catalytic decomposition of NaHCO3. In the first possible mechanism, the decomposition proceeds in a pathway similar to pure NaHCO3. Like other metal oxyhydroxides,76 FeOOH has a high affinity for OH and can form a complex with the OH loosely held in HCO3 at high

The third mechanism proposed in this work involves Lewis acidbase chemistry. As with many compounds that contain OH functional groups, FeOOH can both accept and donate a proton in acidbase reactions.80 In this case, HCO3 is more alkaline than FeOOH, meaning that the HCO3 can accept a proton from the hydroxide group of FeOOH via the reaction FeOOH þ HCO3 T FeOO þ H2 O þ CO2

ðR10Þ

The conjugate base of FeOOH, FeOO , is now more alkaline than HCO3, so it can readily take a proton from another HCO3 via FeOO þ HCO3 T FeOOH þ CO3 2

ðR11Þ

Of these three mechanisms, the one proposed in R10 and R11 is believed to be the most likely reaction path due to the simplicity involved in proton transfer. The least likely of these three is the second mechanism, proposed in R8 and R9. The intermediate in the second proposed mechanism, FeO(CO3), is relatively large and may sterically hinder R9. Each of these mechanisms is theoretical, and no experimental evidence exists to support them. As such, other mechanisms not discussed here may also be plausible. A detailed quantum chemical analysis to determine the reaction mechanism will be the subject of future work. Implications for CO2 Capture. As mentioned above, the sorption of CO2 onto Na2CO3 is spontaneous in flue gas conditions, and therefore the energy-intensive step of CO2 capture involves regeneration of the sorbent. For solid sorbents, the theoretical energy requirement is due to the heat capacity, heat of adsorption and/or the energy required for reaction. In the case of Na2CO3, the energy required for reaction comes from the heat of reaction and activation energy of the reverse reaction of R2. The heat of reaction is a constant for a given reaction at a given temperature; this energy requirement cannot be reduced without using a new material. However, as this study shows, the activation energy can be reduced. By reducing the activation energy through catalysis, the total energy requirement for regeneration of the sorbent is decreased. 15542



ARTICLE

Nonetheless, adsorption of CO2 onto the sorbent is also a critical issue that cannot be neglected. The primary pitfall of using Na2CO3 as a sorbent is its slow reaction rate. By reducing the activation energy of the reverse reaction of R2, the activation of the forward reaction of R2 should also be reduced according to the relation ΔH ¼ Ea, R2 Ea, R2

ðE6Þ

where ΔH is the heat of reaction, Ea,R2 is the activation energy of the forward reaction of R2, and Ea,-R2 is the activation energy of the reverse reaction of R2. Because ΔH is constant for a given temperature, a reduction in Ea,-R2 due to supporting the Na2CO3 on FeOOH means that Ea,R2 is also decreased. In theory, this reduction in the activation energy of adsorption can increase the rate of sorption, according to the Arrhenius form E5. Increasing the sorption rate reduces the necessary size of the adsorber, which means lower capital costs. Further, a smaller adsorber means a lower pressure drop and temperature change in the flue gas, reducing operating costs as well. In addition, supporting Na2CO3 on FeOOH has the potential to increase the total loading of CO2 on the sorbent, which can reduce the size of the adsorber even further.

’ CONCLUSIONS Nanoporous FeOOH was used as a supporting material for the decomposition of NaHCO3. The kinetic data of the decomposition of supported NaHCO3 were compared with those of pure NaHCO3. Depending on the amount of NaHCO3 present on the FeOOH surface, the Arrhenius form showed that activation energy decreases by 2050%, which implies a catalytic function of the FeOOH support material. Although supporting the NaHCO3 onto FeOOH significantly decreased the frequency factor of the Arrhenius form, the reduction in activation energy increased the kinetic rate constant by a factor of at least eight, and in certain cases by a factor of almost 50 relative to pure NaHCO3. Several material characterization techniques were used to verify the thermal stability of the FeOOH and the NHF samples. The only apparent changes due to heating the sample to 180 C can be attributed to a slight increase in particle size of the FeOOH. The use of nanoporous FeOOH as a support material for CO2 capture through adsorption with Na2CO3 has the potential to dramatically reduce the energy requirements and time needed for regeneration of the sorbent, which in turn can significantly reduce the overall cost of CO2 capture. ’ AUTHOR INFORMATION Corresponding Author

*Tel: +1 (307) 766-5633. Fax: +1 (307) 766-6777. E-mail: [email protected].

’ ACKNOWLEDGMENT This research was supported by the National Science Foundation and the University of Wyoming. The authors would like to thank Dr. Patrick McCurdy at Colorado State University for the FTIR analysis, Dr. Norbert Swoboda-Colberg at the University of Wyoming for the SEM images and XRD analysis of the pure FeOOH and NaHCO3, Dr. Erlei Jin for the TEM images, and Dr. Xiao Zeng and Yu Zhao for aid with the reaction mechanism,

as well as Ying Gao, Anthony Richard, Ming Chen Tang, Zhuoyan Sun, and Patrick Tcheunou for collecting TGA data.

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