Article pubs.acs.org/est
Room-Temperature Oxidation of Formaldehyde by Layered Manganese Oxide: Effect of Water Jinlong Wang,†,‡ Pengyi Zhang,*,†,‡ Jinge Li,† Chuanjia Jiang,§ Rizwangul Yunus,† and Jeonghyun Kim† †
State Key Joint Laboratory of Environment Simulation and Pollution Control, School of Environment, Tsinghua University, Beijing 100084, China ‡ Collaborative Innovation Center for Regional Environmental Quality, School of Environment, Tsinghua University, Beijing 100084, China § Department of Civil and Environmental Engineering, Duke University, Durham, North Carolina 90287, United States S Supporting Information *
ABSTRACT: Layered manganese oxide, i.e., birnessite was prepared via the reaction of potassium permanganate with ammonium oxalate. The water content in the birnessite was adjusted by drying/calcining the samples at various temperatures (30 °C, 100 °C, 200 °C, 300 °C, and 500 °C). Thermogravimetry-mass spectroscopy showed three types of water released from birnessite, which can be ascribed to physically adsorbed H2O, interlayer H2O and hydroxyl, respectively. The activity of birnessite for formaldehyde oxidation was positively associated with its water content, i.e., the higher the water content, the better activity it has. In-situ DRIFTS and step scanning XRD analysis indicate that adsorbed formaldehyde, which is promoted by bonded water via hydrogen bonding, is transformed into formate and carbonate with the consumption of hydroxyl and bonded water. Both bonded water and water in air can compensate the consumed hydroxyl groups to sustain the mineralization of formaldehyde at room temperature. In addition, water in air stimulates the desorption of carbonate via water competitive adsorption, and accordingly the birnessite recovers its activity. This investigation elucidated the role of water in oxidizing formaldehyde by layered manganese oxides at room temperature, which may be helpful for the development of more efficient materials.
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CeO2,25 Ag/SBA-15,26 Ag/MCM-41,27,28 Ag−SiO2,29 and Ag/CeO2.30 These supported noble metals were reported to oxidize HCHO into carbon dioxide (CO2) and water (H2O) at relatively low temperature. However, the wide application of supported noble metals is restricted by their high cost, and it is of great necessity and interest to develop a kind of cost-effective catalyst. Sekine et al. first compared the activity of commercial transitional metal oxides for HCHO conversion, and among them manganese oxides (MnOx) showed the best activity at 25 °C within 20 h.31 Since then, researchers have synthesized various manganese oxides with different morphologies and crystal structures to catalytically oxidize HCHO at relatively high temperature. For examples, the porous structure of KOMS-2 led to an increase in its HCHO conversion at 100 °C from 40% to 60%.32 β-MnO2 nanoparticles confined in mesoporous silica (SBA-15) enabled complete conversion of HCHO at 130 °C,33 while the complete HCHO conversion
INTRODUCTION Formaldehyde (HCHO) is one of the major indoor air pollutants, and the exposure to HCHO may threaten human health. For example, long-term exposure to HCHO at even as low as 0.03 ppm (ppm) can cause inflammation or nasopharyngeal carcinoma.1 Physical adsorption (e.g., by activated carbon), nonthermal plasma degradation, catalytic combustion, plant absorption, and photocatalysis are commonly used methods for HCHO removal. However, the application of these methods are restricted for various shortcomings, such as limited adsorption capacity, demands of high energy consumption and high reaction temperature, low efficiency, and formation of secondary pollutants, etc. Thus, it is still a challenge to remove indoor HCHO at room temperature.2 Room-temperature catalytic oxidation has been regarded as a promising strategy for the removal of indoor HCHO due to its mild reaction conditions and no demand for energy.3 Until now, most reported room-temperature catalysts for HCHO removal are supported noble metals, such as, Pt/TiO2,3−7 Pt/ MnOxCeO2,8 Pt/Fe2O3,9 Pt/MnO210 Pt/AlOOH,11 Pt/γAl2O3,12 Pd/TiO2,13,14 Au/CeO2,15−20 Au/Fe2O3,21 Au/ FeOx,22 Au/Co 3O4 CeO 2,23 Au-γAl 2O 3,24 Ag/MnO x © XXXX American Chemical Society
Received: April 25, 2015 Revised: September 30, 2015 Accepted: October 1, 2015
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DOI: 10.1021/acs.est.5b02085 Environ. Sci. Technol. XXXX, XXX, XXX−XXX
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Environmental Science & Technology was observed at 100 °C by mixed MnOx-CeO2 and at 75 °C by solid solution MnxCo3‑xO4,34,35 respectively. KxMnO2 hollow nanospheres also exhibited high HCHO conversion at 60 °C.36 Sidheswaran et al. synthesized a mixed MnOx consisting of nsutite (84%), manjiroite (13%), and cryptomelane (2%) for low-level HCHO (30 and 200 ppb) oxidation at room temperature.37 Chen et al. compared three tunnel-structure MnOx, i.e., pyrolusite (1 × 1), cryptomelane (2 × 2), and todorokite (3 × 3), and found that cryptomelane with the tunnel size similar to the dynamic diameter of HCHO showed the highest activity.38 Recently, Zhang et al. compared α-, β-, γ-, and δ-MnO2 and supposed that tunnel structure and active lattice oxygen species are the main factors that contribute to the best performance of δ-MnO2.39 As mentioned above, the effects of crystal structure and morphology of MnOx on HCHO conversion at temperatures higher than room temperature were extensively investigated. However, room-temperature oxidation of HCHO by MnOx was only minimally investigated, and the role of surface hydroxyl groups and water molecules contained in MnO2 has not been explored, although hydroxyl groups play a key role in noble-metal catalyzed HCHO oxidation. Here, we report the oxidation of HCHO at room temperature by a layered MnOx, (i.e., birnessite) without the need of any noble metals, as well as its mechanism and the role of water contained in birnessite. Birnessite, i.e., δ-MnOx, is built up from layers of edge-sharing MnO6 octahedra with a certain number of water molecules and foreign cations (e.g., Li+, Na+, K+, Ca2+) between the layers. We found that water molecules in birnessite promote HCHO adsorption, followed by oxidation by hydroxyl groups, and water molecules both in birnessite and in air can compensate the consumed hydroxyl groups to maintain the efficient oxidation of HCHO at room temperature. Mass transfer and catalyst lifetime are two key issues for the oxidation of indoor HCHO, due to its low concentration and continuous release, thus our findings are instructive for designing efficient catalysts and with sustained high activity during practical use.
recorded on a Bruker D8-Advance X-ray diffractometer (Germany) using Cu Kα radiation (λ = 0.1542 nm) operated at 40 kV and 40 mA. As for typical diffraction peak (2θ ≈ 12°), samples were collected at a scan step of 0.02° from 10° to 15°. The interlayer distance was calculated by Bragg’s law: 2d sin θ = λ
(1)
where d is the distance of interlayer, θ is the Bragg angle, and λ is the X-ray wavelength. Chemical states of surface elements were investigated by X-ray photoelectron spectroscopy (XPS, PHI-5300, ESCA) at a pass energy of 50 eV, using Al Kα as an exciting X-ray source. The spectra were calibrated with respect to the C 1s peak of adventitious carbon at 284.8 eV. The nitrogen adsorption−desorption curves of samples were recorded by using a Micromeritics ASAP 2020 nitrogen adsorption apparatus (U.S.A.). The Brunauer−Emmett−Teller (BET) specific surface area was determined by a multipoint method using adsorption data in the relative pressure (P/P0) range of 0.05−0.3. The degassed temperature is 100 °C for 10 h. The single-point pore volume was obtained from the nitrogen adsorption curve at the relative pressure of 0.97. Thermogravimetry−mass spectrometry (TG−MS) was used to determine the temperatures at which water molecules were lost as samples were heated. The analysis was performed on a Netzsch STA 409PG Luxx system, which was coupled to a Thermal Star mass spectrometer via a heated (200 °C) capillary transfer line. Samples (∼20 mg) were heated from 25 to 500 °C at 5 °C/min with the sample held under flowing dry N2 at 20 mL/min. In-situ diffuse reflectance infrared Fourier transform spectrometry (DRIFTS) was recorded on a Nicolet 6700 FTIR spectrometer (U.S.A.) with an in situ cell, which was to determine the intermediate species during HCHO oxidation process and to clarify the HCHO oxidation mechanism. The spectra were recorded with a resolution of 4 cm−1 and an accumulation of 32 scans. HCHO with concentration of ∼80 ppm was injected into the cell at a flow rate of 30 mL/min with the synthetic air as the balance gas. Evaluation of HCHO Removal Activity. To evaluate the activity of as-obtained samples for removing HCHO, a glass bottle of 3.5 L was used, at the bottom of which was placed a sealed weighing bottle with a diameter of 30 mm containing 50 mg of sample. The glass bottle was first vented with CO2-free synthetic air for 15 min to eliminate the interference of atmospheric CO2 on the measurement of CO2 formation during the reaction. Then 8 μL of HCHO solution was injected into the glass bottle, making the initial HCHO concentration at ∼200 ppm. After the HCHO solution volatilized completely and the concentration of HCHO in the reactor became stabilized, the cap of the weighing bottle was removed with a thread tied on the cap and thus the samples were exposed to formaldehyde. The effect of relative humidity (water vapor in air) on HCHO removal was investigated in a flow-through reactor. 50 mg catalyst (40−60 meshes) was used. The relative humidity (0%, 32%, 65%, and 92%) was adjusted by changing the ratio of dry air to humid air. HCHO was generated by vaporizing the paraformaldehyde, and its inlet concentration was set at ∼10 ppm. The total flow rate was 300 mL/min with the corresponding GHSV of 180 000 h−1. The reaction was conducted at room temperature. The concentration of formaldehyde in the bottle was determined by using MBTH method.1 The concentration of CO2 was determined by a gas chromatograph (GC) equipped
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EXPERIMENTAL SECTION Preparation of Birnessite Samples. All chemicals used were of analytical grade and used as received without further purification. Birnessite samples were prepared by a facile redox reaction between potassium permanganate (KMnO4) and ammonium oxalate ((NH4)2C2O4). The typical procedures are as follows: 1.0 g of KMnO4 and 0.4 g of (NH4)2C2O4·H2O were dissolved in 130 mL of deionized water in an Erlenmeyer flask, and the Erlenmeyer flasks were shaken in a water bath (90 °C) at a frequency of 150 rpm for 10 h. After the flask was cooled to room temperature, the precipitate was centrifuged and washed 3 times with deionized water to remove any possible residual reactants. After washing, the precipitate was kept in a desiccator for 24 h at ∼30 °C. Finally, the as-obtained powder was dried/calcined in air at various temperatures, i.e., 100 °C, 200 °C, 300 °C, or 500 °C for 3 h. The samples dried/ calcined at different temperatures hereafter are referred to as S30, S-100, S-200, S-300, and S-500, respectively. Characterization. Scanning electron microscopy (SEM) observations were carried out on a Hitachi 5500 field emission scanning electron microscope operated at 10 kV. All samples were sputter-coated with carbon before observation. Transmission electron microscopic (TEM) was recorded on JEOL 2011 operated at an accelerating voltage of 150 kV. X-ray diffraction (XRD) patterns of as-prepared samples were B
DOI: 10.1021/acs.est.5b02085 Environ. Sci. Technol. XXXX, XXX, XXX−XXX
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Environmental Science & Technology with a methanizer and a flame ionization detector (FID) (GC2014, Shimadzu, Japan). The removal ratio of HCHO was calculated as follows: HCHO convs(%) =
[HCHO]0 − [HCHO]t × 100% [HCHO]0 (2)
where [HCHO]0 is the initial HCHO concentration before the exposure of the sample to HCHO, and [HCHO]t is the HCHO concentration at different time after the exposure of sample to HCHO.
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RESULTS AND DISCUSSION Morphology and Phase Structures. When the drying temperature was no higher than 300 °C, all samples possessed nanosphere-like morphology (Figure 1). Moreover, when the
Figure 2. XRD patterns of as-synthesized samples dried/calcined at various temperatures.
high temperature and transformed into 2 × 2 tunnel structure (commonly called OMS-2). When the drying temperatures are below 500 °C, the layered structure of birnessite is prevented from collapsing by the interlayer cations (K+) and water molecules, which was demonstrated by TG−MS analysis. To learn the details of the layer structure of birnessite, the change of distance (2θ ≈ 12°) between interlayer was measured from 10° to 15° by XRD in the step scanning mode (Figure S2). The 2θ value changed from 11.9° to 12.5°, and the corresponding interlayer distance decreased from 7.44 to7.06 Å as the drying temperature increased from 30 °C to 300 °C, indicating that H2O was removed when dried under different temperatures.40 The sample was almost completely dehydrated at 500 °C, and the layer structure no longer remained, as evidenced by the disappearance of XRD reflections. BET Surface Area. Sample S-500 possessed the largest surface area (44 m2/g), followed by S-300 (21 m2/g), S-200 (15 m2/g), and S-100 (14 m2/g) (Table S1). The specific surface area of samples obtained at 30 °C was difficult to measure, due to interference by the large amount of physically absorbed H2O, which would be removed under high vacuum during the BET measurement. The trend in pore volume was consistent with that in the specific surface area: samples obtained at higher drying temperature possessed larger pore volume. These observations can be explained by the fact that enhanced desorption of water with increasing drying temperature (as detailed below) promotes the intercalation of N2 during measurement which in turn resulted in higher specific surface area as well as pore volume of the samples. Dedydration during Heating. The weight loss and simultaneous evolution of water molecules from the samples were investigated with TG/DTG−MS (Figure 3). As for sample S-30, the weight loss dropped dramatically when the temperature increased from 25 to 500 °C (Figure 3a). Three peaks (at 68 °C, 132 °C, and 323 °C) appeared on the differential thermograms (DTG), indicating the presence of three different sources of released water.41,42 An online mass spectrometer was used to simultaneously monitor the evolved gas, targeting for the ion with a mass/charge (m/z) ratio of 18, which is taken to correspond to H2O (Figure 3b). There is an obviously asymmetric H2O peaks from 25 to 500 °C. The curve was fit with the Lorenz equation and was deconvoluted into three peaks, which is consistent with the DTG result. The peak located at 68 °C (peak 1) corresponds to weakly physically adsorbed water, which can be easily removed below 100 °C or
Figure 1. FESEM images of as-synthesized samples dried at various temperatures: (a) 30 °C, (b) 100 °C, (c) 200 °C, (d) 300 °C, (e) 500 °C, and (f) low magnification of sample dried at 30 °C.
drying temperature increased from 30 to 100, 200, and 300 °C, the size of nanospheres showed little change and was generally around 500 nm. However, when the drying temperature further increased to 500 °C, the morphology dramatically changed and turned into a flower-like structure consisting of a cluster of nanorods. As with SEM characterization of sample morphologies, when the samples were dried below 500 °C, their intensity and location of the diffraction peaks did not show any obvious change (Figure 2). The intense diffraction peaks of 2θ located around 12.3°, 24.6°, 36.5°, and 65.5° can be assigned to the characteristic peaks of birnessite (JCPDS No. 80-1098), which is a type of manganese oxides with layer structures. TEM was also used to confirm the structure of birnessite (Figure S1). High resolution TEM showed that birnessite was poorly crystallized. The d-spacing of ∼0.71 nm is consistent with the spacing between the (001) planes. However, when the sample was dried at 500 °C, its XRD peaks greatly changed and can be assigned to cryptomelane structure (JCPDS No. 06-0547), which indicates that the layer structure of birnessite collapsed at C
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Evaluation of Catalytic Activity. The HCHO removal activity of different samples was tested with an initial HCHO concentration of 200 ppm at room temperature (Figure 4a).
Figure 4. (a) HCHO removal efficiency and generation of CO2 by different samples after reaction for 3 h. (b) HCHO decomposition and simultaneous formation of CO2 within 12 h by sample S-30.
Figure 3. (a)TG and DTG and (b) online MS spectrum of sample S30. Sample was heated from 25 °C to 500 °C at 5 °C/min with the sample held under flowing drying N2 at 20 mL/min. Mass chromatograms for the mass/charge (m/z) = 18 ions. Peaks 1, 2, and 3 represent three kinds of water during dehydration. The inset illustrates the three sources of water in birnessite.
The control experiment without a manganese oxide sample showed that the formaldehyde concentration remained stable during the test period. All samples prepared at different drying/ calcination temperatures showed the ability to remove HCHO at room temperature, and their HCHO removal efficiency decreased as the drying temperature increased from 30 °C to 300 °C. The HCHO conversion of sample S-30 was 84.7% after reaction for 3 h, while S-300 (with a much higher specific surface area: 21 m2/g) showed much lower activity (43.1%). Such a descending tendency becomes negligible with further increase in the calcination temperature up to 500 °C. This trend can be ascribed to the difference in the water contained in different samples; the HCHO removal efficiency was dependent on the state of water rather than the change of specific surface area. We also calculated the average oxidation state of Mn in different samples according to the energy level difference of Mn 3s peaks in XPS (Table S1). The Mn oxidation states in S-30, S-100, S-200, S-300, and S-500 are 3.81, 3.76, 3.85, 3.82, and 3.61, respectively. When the drying temperature is below 300 °C, the average oxidation state of surface Mn changes only little, which again implies that the state of bonded water mainly influence the activity of the samples. Simultaneous CO2 formation was detected during HCHO oxidation, while no CO was detected during reaction. The amount of formed CO2 changed a little for the samples S-30 and S-100. However, they showed significant difference in eliminating HCHO, which can be attributed to their difference in physisorbed water. Weakly bonded water may promote HCHO adsorption through hydrogen bonding; Xu et al. also
by vacuum, while the broad peak which reaches maximum at 132 °C (peak 2) is attributed to the loss of strongly adsorbed interlayer water, which maintains the stability of the layer structure. The third peak located at 323 °C (peak 3) is caused by the loss of structural hydroxyl. The state of three sources of water is schematically shown in the inset of Figure 3b. Samples dried/calcined at other temperatures were also analyzed with TG, and the desorbed water at different temperatures was quantitatively determined (Figure S3). All the samples showed three different kinds of water, the contents of weakly physisorbed water were 5.1%, 4.5%, 2.7%, 2.6%, and 2.6%, respectively, whereas those of the strongly adsorbed interlayer water were 9.8%, 8.2%, 6.6%, 5.1% and 1.5% as the drying/calcination temperature increased from 30 to 500 °C. The existence of physisorbed water in S-300 and S-500 indicates that desorbed physisorbed water can be partly recovered when samples are kept under ambient conditions, which is similar to the rehydroxylation of fried clay ceramics under humid environment.43 The strongly bonded water was recovered in lower amounts. In addition, it can be seen from Figure S3 that during the drying/calcination process, the amount of hydroxyl groups was relatively constant, with the exception of sample S-500. D
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Environmental Science & Technology reported a hydrogen-bonding interaction between positively charged H atoms of FeOH and negatively charged O atoms of HCHO molecules;44 however, it does not enhance the transformation of HCHO into CO2. Sample S-500 with low content of physisorbed and interlayer water can still oxidize HCHO into CO2, implying that structural hydroxyl groups play a role in HCHO oxidation, which was also reported in the literature.22,27 As shown in Figure S4, the degradation of HCHO followed pseudo-first order reaction kinetics, and the corresponding reaction rate constants (k) are summarized in Table S1. In addition, the complete mineralization of HCHO by S-30 within 12 h was observed (Figure 4b). The lagging of complete mineralization behind HCHO removal indicates that intermediates are formed during HCHO oxidation, which is consistent with that reported in the literatures.45,46 The effect of relative humidity (water vapor in air) on HCHO removal was also investigated. As can be seen from Figure S5, when the water vapor is very low (RH 0%), the S300 sample gradually deactivated with reaction time. While HCHO conversion remained stable at 85% when the relative humidity was increased to 33% or 65%. The presence of water vapor did not inhibit but rather enhanced HCHO oxidation at room temperature. When the relative humidity further increased to 92%, the HCHO conversion dropped to ∼65% due to the competitive adsorption of water with HCHO on the catalyst surface. The effect of relative humidity on sample S-30 (Figure S5b) was similar. However, the removal efficiency was much higher by S-30, reaching 93%, which further confirmed that the water contained in MnO2 played a key role for formaldehyde removal. Mechanism of Formaldehyde Oxidation. XPS was used to examine the surface of sample S-30 before and after exposure to HCHO for a short time. The high-resolution asymmetric C 1s (Figure S6) confirms the formation of intermediates during HCHO oxidation. The C 1s peak can be deconvoluted into three peaks at 284.8, 286.2, and 288.7 eV with XPSPEAK41 software, which are assigned to the sp3-hybridized C (adventitious carbon), C in CO bonds, and the carbonate C (CO32−), respectively.47 When the sample was exposed to 200 ppm of HCHO, the C 1s peak at 288.7 eV shifted to lower binding energy, which is deconvoluted into two peaks at 288.3 and 288.7 eV. The peak at 288.3 eV is due to the carboxylate C (OCO), as a result of formate formation during HCHO oxidation. Besides, the intensity of the carbonate peak (288.7 eV) significantly increased after reaction, further confirming the mineralization of formaldehyde at room temperature. To illustrate the oxidation mechanism of HCHO over birnessite, in situ observation of DRIFTS spectra of sample S30 exposed to a flow of 80 ppm of HCHO/O2 for 60 min at room temperature was performed (Figure 5). Absorption peaks at 1345 (νs(COO−)), 1380 (δ(CH)), 1569 (νas(COO−)) and 2827 cm−1 (ν(CH)) can be ascribed to formate species.48−50 The frequency difference between νs(COO−) and νas(COO−) is usually employed to judge the adsorption mode of formate species. Larger splits related to monodentate and smaller splits related to bidentate, while splits correlated to bridging configuration falls in between.51 The frequency difference of 224 cm−1 here is indicative of a bridging mode. Carbonate species over birnessite samples were also observed during HCHO oxidation process, the absorbance at 1503 (νas(COO−)) and 1324 cm−1 (νs(COO−)) can be ascribed to monodentate carbonate, while the peaks located at 1680 (ν(CO)) and 1220 cm−1 (νas(COO)) are assigned to
Figure 5. In-situ DRIFTS spectra of sample S-30 under the flow of 80 ppm of HCHO/O2.
bridging carbonate species, respectively.52 Thus, it can be concluded that formate and carbonate are formed and accumulated during HCHO oxidation over birnessite. In addition, there was a negative absorbance around 3584 cm−1, which is ascribed to structural hydroxyl groups (OH).4,53 The decrease of OH intensity suggests that the formation of surface formate and carbonate consumed OH groups. It has been reported that the surface OH plays an important role during HCHO oxidation at room temperature. Zhang et al. reported that HCHO oxidation over the promoted 2% Na-1% Pt/TiO2 catalyst involved the direct reaction between surface OH and formate species.4,14 Besides, Chen et al. suggested that HCHO adsorbed over supported gold catalysts (Au/CeO2 or Au/FeOx) could be oxidized by surface oxygen to formate, which could be further oxidized by surface hydroxyl groups.21,22 In addition, the XRD result (Figure S2) also indicates the loss of interlayer water during HCHO oxidation. The 2θ peak reflecting the interlayer distance shifted from 11.9° to 12.2° after reaction, i.e., the interlayer distance slightly reduced, which can be ascribed to the loss of interlayer water. These results demonstrate that both OH group and bonded water of birnessite are consumed for HCHO oxidation. To investigate the effect of water content of birnessite on HCHO oxidation, in situ DRIFTS spectra of different samples upon exposure to 80 ppm of HCHO/O2 at room temperature were examined (Figure S7). The shape and location of absorption peaks were similar for samples S-30, S-100, S-200, and S-300. The absorption intensity of formate species and carbonate species were highest in sample S-30, i.e., S-30 had the highest activity transforming HCHO into formate and carbonate. This result indicates that the activity of birnessite is closely associated with its water content, and the sample with the largest water content had the highest activity. To further elucidate the role of water in oxidizing HCHO, we compared the ability of the flow of dry O2 or wet O2 (O2/ H2O) to regenerate S-30 sample after it had been exposed to a flow of 80 ppm of HCHO/O2 for 60 min. In the case of dry O2 purge (Figure 6a), the intensity of formate absorption peak slightly decreased with further consumption of hydroxyl groups. Correspondingly, both peaks of bridging carbonate (1680 cm−1) and monodentate carbonate (1497 cm−1) increased a little. These results mean that under the flow of dry O2 surface hydroxyl groups cannot be compensated and carbonate further accumulate on the birnessite, i.e., the birnessite cannot be E
DOI: 10.1021/acs.est.5b02085 Environ. Sci. Technol. XXXX, XXX, XXX−XXX
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concentration of HCHO/O2 (∼80 ppm) reached stable, the gas was switched from the bypass to flow through S-30 sample. The concentration of HCHO dropped sharply; however, with a small amount of CO2 generation (stage 2). After that, dry O2 was used to flow through S-30 sample. No obvious changes occurred for CO2 generation (stage 3). However, when we use humid O2 to flow S-30 sample (stage 4), there was an obvious peak of CO2, which indicate water vapor favors desorption of carbonate species from catalyst. During the whole process, no formic acid species were flowed out by dry O2 or humid O2, which indicates that the disappearance of formate species in Figure 6b is caused by further oxidation rather than desorption from the surface of catalyst by humid O2. For long-term use, recycling tests were carried out. The used sample S-30 was regenerated with H2O/O2 purge for 1 h before the next use. And the recycling performance of sample S-30 without regeneration treatment was also tested for comparison. Sample S-30 after H2O/O2 treatment exhibited a high and stable HCHO removal efficiency (Figure S10a), further indicating the function of H2O. However, as for sample S-30 without H2O/O2 treatment, HCHO removal efficiency dropped gradually from 91% to 78% after 5 times test (Figure S10b). According to the above-mentioned Results and Discussion, we summarized the roles of water both in birnessite and in air, and proposed the mechanism of HCHO oxidation on birnessite at room temperature (Figure 7). First, adsorption of HCHO on
Figure 6. In-situ DRIFTS spectra of S-30 sample after exposed to a flow of 80 ppm of HCHO/O2 for 60 min followed by (a) O2 purging and (b) H2O/O2 purging at room temperature.
regenerated. However, when purged with H2O/O2 (Figure 6b), formate species disappeared within 10 min, and carbonate species did not accumulate with the decomposition of formate. Furthermore, the amount of hydroxyl groups did not reduce but increased. Chen et al. also reported the same phenomenon that water can react with surface oxygen ([O]s) to form surface hydroxyl group.21,22 The consumed surface OH can be regenerated by the reaction between surface active oxygen (O2−, O−, etc.) and water (water vapor or physisorbed/ interlayer water) through the reaction (O2−, O− + H2O → 2OH). Surface active chemisorbed oxygen (O2−, O−, etc.) can be formed by the complex migration between surface lattice oxygen (O2−) and oxygen vacancy with the continuous dissociation of molecular oxygen.54,55 These results indicate that water vapor can compensate the hydroxyl groups and be beneficial to the desorption of carbonate, which is probably due to their competitive adsorption on birnessite surface. To verify whether the used birnessite sample was regenerated by the flow of wet O2, again 80 ppm of HCHO/O2 flowed through the sample (Figure S8). Formate and carbonate were again formed, while surface hydroxyl groups again showed negative increase, which demonstrates that the birnessite can be regenerated by humid air. To further learn the function of water, online mass spectroscopy (MS) analysis was used to monitor HCHO, HCOOH, and CO2 under different conditions. S-30 was taken as the catalyst. As shown in Figure S9, after the inlet
Figure 7. Proposed mechanism of formaldehyde oxidation on birnessite at room temperature.
birnessite is promoted by bonded water via forming hydrogen bonding with HCHO. Second, the adsorbed HCHO is oxidized by structural hydroxyl groups to formate (bridging type) and carbonate (monodentate and bridging type). Third, the consumed surface OH can be regenerated by the reaction between surface active oxygen (O2−, O−, etc.) and water (gas phase or physisorbed/interlayer water) through this reaction (O2−, O− + H2O → 2-OH). The desorption of carbonate from the birnessite surface can be stimulated by competitive adsorption of water in air. In summary, a water-containing layered manganese oxide, i.e., birnessite, was synthesized and its water content was adjusted by changing the drying temperature. Birnessite materials show room-temperature activity for transforming HCHO into formate, carbonate and CO2. The activity is strongly associated with the water content of the birnessite materials, and the transformation of HCHO is accompanied by the consumption of hydroxyl group and bonded water. In addition, water vapor in air can compensate the consumed F
DOI: 10.1021/acs.est.5b02085 Environ. Sci. Technol. XXXX, XXX, XXX−XXX
Article
Environmental Science & Technology
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hydroxyl group and promote the desorption of carbonate. These findings shed light on the mechanism of HCHO oxidation by transitional metal oxides, which may be helpful for the development of new cost-effective materials for HCHO pollution control.
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ASSOCIATED CONTENT
S Supporting Information *
The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.est.5b02085. Table S1, Physical properties; Figure S1, S-30 with low and high magnification of TEM; Figure S2, Step scanning XRD; Figure S3, TG curves of different samples; Figure S4, HCHO decomposition kinetics of different samples; Figure S5, Effect of relative humidity on HCHO removal; Figure S6, XPS spectra of C1s; Figure S7, In-situ DRIFTS spectra of samples; Figure S8, In-situ DRIFTS spectra; Figure S9, On-line MS detection of HCHO; Figure S10, Change of the HCHO concentration; and additional references (PDF)
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AUTHOR INFORMATION
Corresponding Author
*Tel: +86 10 62773720; fax: +86 10 62796840-602; E-mail:
[email protected] (P.Y.Z.). Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS This work was funded by the National High Technology Research and Development Program of China (2012AA062701), National Nature Science Foundation of China (21221004, 21411140032) and Tsinghua University Initiative Scientific Research Program (20131089251).
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REFERENCES
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DOI: 10.1021/acs.est.5b02085 Environ. Sci. Technol. XXXX, XXX, XXX−XXX
