Synergic Adsorption of H2S Using High Surface Area Iron Oxide

Article Cite This: Energy Fuels XXXX, XXX, XXX−XXX

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Synergic Adsorption of H2S Using High Surface Area Iron Oxide− Carbon Composites at Room Temperature Kexin Ling,† Varun Shenoy Gangoli,† and Andrew R. Barron*,†,§,‡ †

Department of Chemistry and §Department of Materials Science and Nanoengineering, Rice University, Houston, Texas 77005, United States ‡ Energy Safety Research Institute (ESRI), Swansea University, Bay Campus, Swansea SA1 8EN, U.K.

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S Supporting Information *

ABSTRACT: Developing an efficient hydrogen sulfide (H2S) sorbent is of great importance to natural gas industries, biomedical applications, and environmental conservation. Activated carbon, metal oxides, and their composite materials show potential for desulfurization. This work explores the synergic effects in composites of iron oxide (Fe2O3) and oxygenated porous carbon (OPC) for the removal of H2S at room temperature. Two types of Fe2O3-OPC composite samples were prepared: physically mixed (PM) and chemically mixed (CM). The two types of composites were tested for H2S uptake performance at ambient conditions, and a systematic study of the synergic effects of Fe2O3 and OPC was performed. Thorough characterization and analysis were used to reveal detailed structural and compositional properties of these samples. The CM sample with the best uptake capacity was also tested further for the desulfurization rate and the mechanism of action. The PM samples showed a lower H2S uptake capacity within 24 h compared to the theoretical value for the Fe2O3 and OPC working independently, indicating a negative synergic effect. The CM samples reached a maximum uptake capacity higher than the components working independently and importantly an increased rate of H2S uptake, which indicates positive synergy, showing potential in applications where rapid adsorption is required.

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100 °C; however, these metal oxides generally lack the high surface area offered by activated carbon.

INTRODUCTION Hydrogen sulfide (H2S) is a flammable, colorless, and extremely poisonous gas with a characteristic odor reminiscent of rotten eggs. It is found naturally in oil and natural gas reservoirs1 and is produced by the degradation of human/ animal wastes2 and organic compounds.3 A meager amount of H2S (a few parts per million) is sufficient to cause issues ranging from pipeline and tunnel corrosion4,5 to severe environmental pollution from acid rain.6 H2S removal methods are needed in industries where the gas is either a common byproduct or simply present along with other chemicals of interest, including purification of natural gas, waste gas desulfurization for steel/coal combustion industries, waste water treatment, and specific biomedical needs including ostomy filters, especially those that can be applied under ambient conditions.7−10 One of the approaches used today for the removal of H2S is via gas adsorption as it is a cost-effective approach and has been demonstrated to achieve a relatively high level of purification,11 including room temperature and pressure conditions for added economic feasibility. The H2S sorbents that show good performance at these mild conditions include activated carbon,12−17 mesoporous silica,8 zeolites,18,19 and metal oxides.7−9,20 Among these sorbents with porous structures, activated carbon has a high surface area and large pore volume, which allows for significant adsorption capacities in general. Metal oxides, on the other hand, adsorb H2S via chemical reaction to form metal sulfides, for example, eq 1.21 For most metal oxides, including zinc oxide,7 copper oxide,9 and iron oxide,8,20−22 this reaction takes place rapidly at relatively low reaction temperatures ranging from ∼25 °C to © XXXX American Chemical Society

MOx + x H 2S → MSx + x H 2O

(1)

Desiring both a large adsorption capacity and a high uptake rate has resulted in recent studies into composite sorbents for better H2S uptake performance.23,24 In particular, there are reports showing that composites of metal oxides and carbon exhibit enhanced H2S adsorption.25−27 Identifying the role of both carbon and metal oxide components in said composites and also the active compounds produced during H2S adsorption process is crucial to optimizing the chemical composition of such sorbents. Zhang and co-workers studied the H2S removal capacity of MgO-loaded mesoporous carbon millimeter spheres at 30 °C28 and found that a load amount less than 20 wt % enhanced the H2S uptake performance. This was attributed to the relatively large pore size and 3D pore structure, all of which allowed easier diffusion of reactants inward and products outward. Fauteux-Lefebvre et al. employed iron-functionalized carbon nanofilaments as an H2S sorbent at 100 °C for a test case of 500 ppm H2S in helium gas and demonstrated that carbon worked as both the sulfur adsorbent and a support to help uniformly distributed Fe.21 In a similar manner, Hernández et al. reported that 10% ZnO-functionalized activated carbon had a higher adsorption capacity compared to the unloaded commercial activated carbon for simulated biogas desulfurization (200 ppm H2S in Received: April 2, 2019 Revised: June 20, 2019 Published: June 24, 2019 A

DOI: 10.1021/acs.energyfuels.9b01012 Energy Fuels XXXX, XXX, XXX−XXX

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Energy & Fuels N2) at ambient conditions (28 °C, 1 atm).29 Finally, ArcibarOrozco et al. used iron (hydr)oxide/graphite oxide composites at 25 °C to detect sulfur “spices” including elemental sulfur and ferric and ferrous sulfates, which helps attribute to a redox reactive adsorption mechanism.30 There is thus evidence to support synergic effects in composite materials, in particular metal oxides and porous carbon for this application. We have previously shown that nanoscale iron oxide (Fe2O3) is suitable for removal of H2S from natural gas streams based on a scalable synthesis and high uptake reaction speed at low temperature and pressure.20 More recently, we have reported the synthesis of oxygenated porous carbon (OPC) from a polymer precursor with not only extremely high surface area and pore volume but also a narrow pore size distribution that enable great CO 2 adsorption and excellent CH 4 /CO 2 selectivity.31,32 This laid the background for the work published herein where we aim to investigate whether composite materials of iron oxide particles and OPC show a positive synergic effect that can help increase H2S sorption at ambient conditions. This work demonstrates two different ways of generating the iron oxide/carbon composites to study H2S uptake performance, and all the composites were characterized extensively to identify and analyze the sulfur species formed during the reaction in order to better understand the reaction process for further optimization.

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Synthesis of Iron Oxide (Fe2O3) Particles via the Hydrothermal Method. Fe2O3 particles were synthesized following a modified process described in the literature.35 A typical synthesis involved the dissolution of PVP (60 mmol) in 35 mL of deionized water followed by the addition of FeCl3 (3.0 mmol). The solution was stirred for 10 min and then transferred to a 50 mL volume hydrothermal cell followed by heating in an oven to 150 °C for 16 h. The resulting product was Fe2O3 particles, which were separated by centrifugation (2000 × g for 5 min at 25 °C in a Sorvall ST 16 centrifuge), washed with deionized water (5 × 25 mL) and then ethanol (25 mL), and finally dried at 70 °C overnight before use. Synthesis of Iron Oxide (Fe2O3) Particles Via the Precipitation Method. Fe2O3 particles were synthesized using a modification of a protocol previously reported.36 In a typical synthesis, NaOH (135 mmol) was dissolved in 25 mL of deionized water. FeCl3 (50 mmol) was separately dissolved in 25 mL of deionized water, and the solution was heated to 100 °C. Then, the NaOH solution was added to the FeCl3 solution drop-by-drop within 5 min of the latter being heated and under continuous stirring. The mixed solution was stirred for another 5 min following which dark red Fe(OH)3 precipitated out of the solution gradually. The suspension was transferred to a hydrothermal cell and heated in an oven set to 100 °C for 48 h. The resulting Fe2O3 particles were separated by centrifugation (2000 × g for 5 min at 25 °C in a Sorvall ST 16 centrifuge), washed with deionized water (5 × 25 mL) and ethanol (25 mL), and then dried at 70 °C overnight prior to use. Synthesis of Oxygenated Porous Carbon (OPC). OPC was generated using polyanisyl alcohol (PAA) with modification of the literature procedure.32 Concentrated H2SO4 (6 mL) was added dropwise into anisyl alcohol (10 g) in a 150 mL glass beaker with continuous stirring using a glass rod. The beaker should be kept in a water/ice bath to avoid excessive heating. The resulting dark purple product mixture was washed with deionized water (4 × 50 mL) and then acetone (200 mL). The brown-colored product (PAA) obtained after washing was crushed into a powder, quickly washed with acetone (100 mL) in a beaker, and dried at 25 °C for 12 h. PAA (500 mg) was then thoroughly mixed with KOH (1.5 g, crushed to powder) in a mortar using a pestle for 10 min. This mixture was then transferred to a quartz tube (1″ diameter), which was placed inside a tube furnace, dried under Ar (600 sccm, standard cubic centimeters per minute) for 20 min at 25 °C, and then heated to 750 °C for 30 min. The activated carbon product was washed with HCl (100 mL, 1.4 M) and then deionized water till the filtrate reached a neutral pH of 7. The product was then heated to 70 °C under vacuum for 12 h to dry before use. Synthesis of Physically Mixed (PM) Fe2O3-OPC Samples. A range of Fe2O3/OPC composites was created by simply mixing known quantities of the two in an agate mortar for 10 min. The weight percent of iron in each sample is shown in Table 1, calculated theoretically based on the amounts of each used in the mixture.

EXPERIMENTAL SECTION

Materials. FeCl3 (97% purity), poly(vinylpyrrolidone) (PVP, MW = 10,000 g/mol), NaOH (97% purity), anisyl alcohol (98 vol %), KOH (≥85% purity, pellets), HCl (37 vol %), acetone, ethanol, and H2S (all >99.5% purity) were obtained from Sigma-Aldrich. H2SO4 (95−97%) was purchased from Merck Millipore. Argon (99.999% purity) was supplied by Airgas. Deionized water was supplied via an industrial-grade RO/DI system capable of providing water with a resistivity >15 MΩ cm. All chemicals were used as received. Characterization. Scanning electron microscopy (SEM) measurements were conducted on an FEI Quanta 400 high-resolution field emission microscope at an acceleration voltage of 20 kV. SEM− energy-dispersive X-ray spectroscopy (SEM-EDS) was performed a minimum of thrice per sample for statistical accuracy. Transmission electron microscopy (TEM) images were acquired on a JOEL 2100 field emission gun microscope under an operating voltage of 200 kV. Wide-angle powder X-ray diffraction (XRD) measurements were performed on a Rigaku D/Max Ultima II powder XRD with Cu Kα radiation (40 kV, 40 mA). The crystallite size was calculated via the Halder−Wagner method.33 Nitrogen adsorption−desorption isotherms were recorded on a Quantachrome Autosorb-iQ-MP/Kr BET (Brunauer−Emmett−Teller) surface analyzer at −196 °C (77 K). The samples were degassed at 140 °C for 12 h under vacuum before measurements were taken. The surface area, total pore volume, and pore size distributions of various composites were determined via non-local density functional theory (NLDFT). X-ray photoelectron spectroscopy (XPS) measurements were carried out on a PHI Quantera XPS spectrometer. The weight percent of each element was determined by XPS survey scans and converted from atomic percent, with a passing energy of 140 eV. Each measurement was repeated at least three times for statistical accuracy. High-resolution elemental analysis was performed with a passing energy of 26 eV via multiple scans again. Each spectrum was corrected using 284.8 eV as a reference binding energy for the C 1s peak. XPS deconvolutions followed Gaussian−Lorentzian functions.34 Raman spectral measurements were performed with the samples on glass microscope slides inside a Renishaw Raman microscope at 25 °C using a 532 nm laser, and multiple acquisitions were collected for statistical accuracy. Fourier transform infrared (FT-IR) spectra were recorded on a Nicolet FTIR infrared microscope equipped with an ATR (attenuated total reflection) support.

Table 1. Iron Concentration in Each Physically Mixed (PM) Fe2O3-OPC Composite Sample As Calculated Based on the Amounts of Each Species Used in the Mixture sample

Fe2O3 (mg)

OPC (mg)

Fe wt % by calculation

PM1 PM2 PM3 PM4 PM5

2.2 6.7 11.0 22.2 43.4

22.3 21.2 20.5 19.0 16.6

6.29 16.81 24.44 37.72 50.63

Synthesis of Chemically Mixed (CM) Fe2O3-OPC Samples. This follows a similar synthesis process as OPC alone except that PAA (250 mg) was thoroughly mixed with KOH (0.75 g, crushed to powder) and also Fe2O3 particles (amount varied between different samples) simultaneously. This mixture was then activated in the same manner as OPC described previously. The weight percent of iron in each sample before activation is calculated and seen in Table 2. B


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reveals that the pure OPC material had a measured uptake of 33.62 wt % sulfur. Each band in the high-resolution XPS spectra (Figure 2) was deconvoluted based on possible functional groups that could be part of the sample. In the OPC C 1s spectrum (Figure 2a), peaks were assigned for the following functional groups: sp2 hybridized CC (284.6 eV), sp3 hybridized C−C (286.2 eV), C−O−C (288.3 eV), and CO (289.5 eV).31,32 The resolved peaks under O 1s spectra (Figure 2b) were attributed to two primary functional groups: C−O−C (533.2 eV) and CO (531.9 eV). Both C 1s and O 1s spectra analyses correspond to previously reported results.32 The C 1s and O 1s spectra of OPC after H2S exposure (Figure 2c,d, respectively) showed near-identical peak assignments. The sulfur related peaks (Figure 2e) could be fitted with two doublets (each comprising S 2p3/2 and S 2p1/2). The first doublet is at 164.1 and 165.4 eV, and these positions could either be showing the C−S bond or merely elemental sulfur (S0) due to the similar electronegativity of S and C. We propose that this doublet is from S0 2p3/2 and S0 2p1/2, respectively, based on previous reports of activated carbon behaving as surface-catalyzed oxidants to form elemental sulfur, and even sulfate groups.38,39 The second doublet is at 168.9 and 170.6 eV and corresponds to SO42− 2p3/2 and 2p1/2, respectively. The presence of SO42− also confirms that some of the H2S was catalytically oxidized by OPC during the test. Characterization and H2S Uptake Test of Fe2O3 Particles. As noted in the Experimental Section, two synthesis methods were adopted for preparing pure Fe2O3 particles. The hydrothermal method is commonly used for growing goodquality crystals40 since poly(vinylpyrrolidone) used as the capping agent helps keep Fe2O3 microspheres monodispersed and away from each other.41 The SEM image in Figure 3a shows that the synthesized Fe2O3 particles from this method are of a spherical morphology and have a uniform size of ∼1 μm in diameter. A higher-magnification image (Figure 3a, inset) shows that these particles have a rough surface, indicating that they are formed as aggregates of smaller particles. The wide-angle XRD pattern in Figure S2a (see the Supporting Information) shows diffraction peaks corresponding to hematite (α-Fe2O3), thus confirming successful synthesis of Fe2O3, with a calculated crystallite size of 20.8 nm consistent with the features in the SEM images (Figure 3a, inset). Figure 3b shows the Fe2O3 particles obtained via the precipitation method, which exhibit a pseudo-cubic morphology (as seen in the inset of Figure 3b) with a uniform size distribution around 450 nm. The XRD pattern (Figure S2b, Supporting Information) was also assigned to hematite, with a crystallite size of 14.8 nm. SEM images (Figures 3c,d) showed that the uniform morphology of both samples was lost after the H2S uptake test and no further uniform particles could be observed. The hydrothermal Fe2O3 was found to have a sulfur uptake of 9.19 wt %, while the precipitated Fe2O3 had an equivalent sulfur uptake of 13.11 wt %. Since the precipitated Fe2O3 not only had a smaller average particle size but also had a smaller crystallite size consistent with synthesis at a lower temperature, H2S uptake was thus found to be inversely proportional to the Fe2O3 particle size and crystallite size. This shows that H2S can likely only react with the surface of Fe2O3 particles. It has been previously reported that the interaction between solid metal oxides and adsorbed H2S is related to lattice diffusion.8,42 This kind of ion migration within the

Table 2. Calculated Iron Concentration in each Chemically Mixed (CM) Fe2O3-OPC Composite Sample before Activation sample

Fe2O3 (mg)

PAA (mg)

Fe wt % before activation

CM1 CM2 CM3 CM4 CM5 CM6

25.6 53.1 101.1 152.3 200.6 299.6

250.2 250.8 249.4 254.1 252.5 253.8

6.50 12.23 20.19 26.23 30.99 37.90

H2S Uptake. The H2S uptake tests for pure OPC, Fe2O3 particles, PM samples, and CM samples were all set up using a 250 mL roundbottom borosilicate glass flask with a vacuum line connector placed inside a fume hood at room temperature (held steady at 25 °C). The samples under study (20 mg OPC/Fe2O3/PM/CM in separate experiments) were put inside the flask, which was then evacuated under vacuum, and 99.5% H2S gas was allowed in the flask using a syringe through a corker stopper on the flask. Another syringe allowed for displacement of the flask environment such that H2S was used to fill the flask three times before the syringes were removed, thus allowing for a saturated H2S environment inside the flask. The contents in the flask were let to stand for 24 h, thus behaving as a batch process for the study of gas uptake by the various samples of interest. All uptake experiments were performed under the same operating conditions to ensure valid comparison. For adsorption rate experiments, 5 mg of the specific sample (OPC, Fe2O3, or CM5) was placed inside a series of six 25 mL flask, and the solutions were treated with 99.5% H2S as described above. At each time period (1, 2, 4, 8, 12, and 24 h), one of the flasks was opened, and the sample was collected for XPS characterization. Selected specific samples (OPC, Fe2O3, PM5, and CM5) were also tested for H2S uptake for longer durations (48 and 72 h) following the same operation procedures.

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RESULTS AND DISCUSSION Characterization and H2S Uptake Test for OPC. The structural and textural morphology of the synthesized OPC sample, as obtained by SEM, is shown in Figure 1. The entire

Figure 1. Scanning electron microscopy image of the typical morphology of oxygenated porous carbon (OPC).

sample is uniform in morphology, which is an indication of consistent synthesis from the protocol employed. H2S uptake via OPC was tested as described in the Experimental Section, and the H2S uptake amount represented by wt % S on the sample was measured by XPS. All tested samples showed that more than 95% H2S was adsorbed within the first 24 h (Figure S1 and Table S1, see the Supporting Information), which can be considered equivalent to the breakpoint (95−99%) in a fixed bed setup.37 Thus, the H2S adsorption amount at 24 h can be regarded as the adsorption capacity. The detailed information on chemical states of C, O, and S before and after the H2S uptake test was acquired by high-resolution XPS elemental scans shown in Figure 2. Analysis of the raw data C


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Figure 2. XPS elemental scanning for OPC of (a) C 1s and (b) O 1s before the H2S test and (c) C 1s, (d) O 1s, and (e) S 2p after exposure to H2S.

4) to show four main species formed: Fe1.95O3, lepidocrocite (γ-FeO(OH)), pyrite (FeS2), and greigite (Fe3S4).

Figure 3. SEM images for as-synthesized Fe2O3 by the (a) hydrothermal method and (b) precipitation method. Also seen here are images of (c) hydrothermal Fe2O3 after the H2S uptake test and (d) precipitated Fe2O3 after the H2S uptake test. The insets in panels (a) and (b) are the corresponding images at higher magnification to show the structural morphology.

lattice of iron oxide is very slow at the relatively lower temperatures used in our test conditions. As such, it is hard for the bulk iron oxide to react with H2S, and iron sulfides only end up forming on the surface of iron oxide particles, as described in the literature before.43,44 A smaller particle size and crystallite size on the precipitated Fe2O3 result in a more accessible surface to react with, in turn leading to higher uptake. Thus, precipitated Fe2O3 was used as the iron source in the synthesis of Fe2O3-OPC composites. The expected reaction between Fe2O3 and H2S is shown in eq 2; however, chemical speciation of the precipitated sample after the H2S test was determined by wide-angle XRD (Figure

Figure 4. X-ray diffraction pattern for precipitated Fe2O3 particles after H2S exposure.

Fe2O3 + H 2S → Fe2S3 + H 2O

(2)

This product distribution may be rationalized by the initial formation of Fe1.95O3; FeOOH, which was originally present on the surface of Fe2O345,46 and more of which was formed due to partial hydrolysis of Fe2O3,47,48 could then also react with H2S (eq 3). D


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from 27.12 to 15.80 wt % sulfur (see Table 3). We note that while XPS is a surface-sensitive technique, in terms of a quantitative analysis, it has been reported to be satisfactory to evaluate relative intensities of elements.50,51 Also, in our case, we only use it to compare the trend of sulfur content, which is more of a surface phenomenon than a bulk phenomenon. Nevertheless, EDS can detect up to 1 μm of sample depth, and the Fe2O3 particles are around 450 nm in diameter; thus, the surface characterization would be expected to be representative of the entire material. As may be seen from Figure S3 (see Supporting Information), EDS analysis follows the same trend as XPS (Figure 6a). The H2S uptake thus has an inversely proportional behavior with iron concentration in the physically mixed composites. In order to best interpret this result, we propose a theoretical, “no synergy” scenario where there is no interaction, synergic or otherwise, between Fe2O3 and OPC in these composites as they were physically mixed instead of undergoing a chemical reaction. Thus, an ideal H2S uptake of the PM samples can then split up into two exclusive parts: H2S adsorbed by Fe2O3 and H2S adsorbed by OPC. The total uptake will thus be the addition of uptakes from the Fe2O3 portion and OPC portion as a function of their relative concentration. The Fe2O3 portion line (Figure 6b, blue dashed line) is connected by two points: zero uptake at 0% iron concentration and 13.11% uptake at 70% iron concentration, indicating zero uptake with no Fe2O3 and uptake equaling that of pure Fe2O3 with 100% Fe2O3, respectively. In the same manner, the OPC portion line (Figure 6b, green dashed line) has boundary conditions of 33.62% uptake at zero iron concentration and zero uptake at 70% iron concentration, which represents pure OPC uptake without Fe2O3 and zero uptake made by OPC with pure Fe2O3, respectively. A total uptake line (Figure 6, red line) would be the addition of the Fe2O3 portion and OPC portion if the two species acted independently. Comparing the experimental data and the “no synergy” curve clearly shows a lower actual uptake for the PM samples that would be expected. SEM imaging discussed previously shows that Fe2O3 covers large sections of the OPC surface in the as-made PM sample. This, in turn, blocks part of the open pores, leading to a decrease in the number of accessible pores and available surface area of the composite accordingly (Table 3 and Figure S4, Supporting Information). The relationship between Fe concentration and surface area for the PM samples is shown in Figure S5. Since the capability for gas uptake is highly related

(3)

Fe2S3 quickly and spontaneously decomposes to form Fe3S4 and FeS2 due to its thermodynamic instability (eq 4).49 2Fe2S3 → FeS2 + Fe3S4

(4)

One of the products of decomposition, FeS2, can easily become oxidized in the presence of water, including ambient moisture (Eq. 5).8 This also explains why the FeS2 signal in the XRD pattern was weaker than that of Fe3S4. 4FeS2 + 15O2 + 2H 2O → 2Fe2(SO4 )3 + 2H 2SO4

(5)

Characterization and H2S Uptake Test of Physically Mixed Fe2O3-OPC Composites. The physically mixed (PM) sample series (PM1−PM5) had an iron concentration (Fe wt %) varying from 6.29 to 50.63% by calculation and from 4.55 to 45.60% by XPS survey scan as seen in Table 3. The around 5% weight difference was contributed by the loss of Fe2O3 during the mixing and transferring process. Table 3. Calculated Iron Concentrations Compared to Values Obtained from XPS, Physical Properties, and H2S Uptake in Physically Mixed (PM) Fe2O3-OPC Composite Sample sample

Fe calculated (wt %)

Fe analysis XPS (wt %)

Surface area (m2/g)

Pore volume (cm3/g)

S analysis XPS (wt %)

PM1 PM2 PM3 PM4 PM5

6.29 16.81 24.44 37.72 50.63

4.55 12.56 20.32 33.24 45.60

902 759 713 467 256

0.45 0.37 0.34 0.23 0.13

27.12 24.53 22.49 19.14 15.80

A series of SEM images (Figure 5a−e) obtained from these PM samples with different Fe wt % values indicate that Fe2O3 particles were predominantly located on the surface of OPC. The Fe2O3 particles were also distributed arbitrarily on the OPC surface at lower Fe concentration, indicating a relatively good dispersion. As the Fe concentration reached around 45 wt %, the surface of OPC was seen to be entirely covered by Fe2O3 particles. The relationship between iron concentration and H2S uptake is shown in Figure 6a below. As the iron concentration increased from 4.55 to 45.60 wt%, the H2S uptake dropped

Figure 5. SEM images for samples PM1−PM5, (a−e) as-synthesized and (f−j) after H2S test. The overall Fe concentration dictates the distribution of the Fe2O3 particles on the surface of OPC. E


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Figure 6. Relationship between (a) iron concentration (as wt %) and H2S uptake for the PM sample series and (b) theoretical H2S uptake of PM samples determined as a simple addition of the individual uptakes of the Fe2O3 and OPC components to show the no synergy scenario.

to the surface area and pore volume of OPC,52 the addition of Fe2O3 may cause lower uptake of physically mixed Fe2O3-OPC composites relative to even pure OPC by itself. At the same time, SEM images (Figure 5f−j) of these PM samples after the H2S test show that a small portion of Fe2O3 particles remains unreacted, especially for those samples with a higher iron concentration. The XRD pattern of sample PM5 after reacting with H2S also showed that the main iron species was unreacted iron oxide (Figure S6, Supporting Information). These results collectively indicate that Fe2O3 not only likely blocks the pores of OPC but that the existence of OPC also decreases the accessibility of the surface of Fe2O3. This tells us that physically mixing Fe2O3 and OPC to prepare composites results in negative synergy and other methods of mixing these two materials should be taken into consideration to achieve positive synergy instead. Characterization of Chemically Mixed (CM) Fe2O3OPC Composites. Samples of Fe2O3 particles with iron concentration varying from 6.50 to 37.90% (as a function of overall reagent) were added into PAA to form a series of chemically mixed samples (CM1−CM6, see Table 2). The iron concentration in the final CM samples was expected to increase slightly due to the carbon mass loss during the conversion of PAA to OPC. The final iron concentration was in the range of 11.04 to 49.57%, as seen in Table 4, which is in a similar range as compared to the PM samples from before. In contrast to the PM samples, no obvious Fe2O3 particles were seen on the surface of OPC (Figure 7). Sample CM1 shows a similar morphology to OPC and as the iron

Figure 7. SEM images for as-synthesized samples (a−f) CM1−CM6, respectively.

concentration increased, the samples began to exhibit a more “cataclastic” texture. The crystalline phase for one of the CM composites, CM5, was studied by X-ray diffraction (Figure 8).

Table 4. Calculated Iron Concentrations Compared to Values Obtained from XPS, Physical Properties, and H2S Uptake in Chemically Mixed (CM) Fe2O3-OPC Composite Sample sample name

Fe before activation (wt %)

Fe analysis XPS (wt %)

Surface area (m2/g)

Pore volume (cm3/g)

S analysis XPS (wt %)

Figure 8. Powder X-ray diffraction pattern for as-synthesized CM5.

OPC CM1 CM2 CM3 CM4 CM5 CM6

0 6.50 12.23 20.19 26.23 30.99 37.90

0 11.04 20.74 31.57 35.48 45.18 49.57

968 1131 1229 876 863 907 651

0.43 0.55 0.55 0.43 0.47 0.54 0.44

33.62 23.61 22.71 22.40 26.26 33.20 27.99

Fe2O3 originally added was reduced to elemental iron by OPC during the activation process, which occurred at a high temperature (750 °C) under an inert atmosphere. Fe3O4 was also detected as either the partially reduced product or as a native oxide layer surrounding the iron in the composite. Particles can be seen from TEM imaging (Figure S7, Supporting Information), with an average diameter of 50 nm. F


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Figure 9. Nitrogen adsorption−desorption isotherms at 77 K for (a) CM samples and (b) OPC. This indicates a mesopore- and micropore-type structure for CM composites.

Figure 10. Pore size distribution curves for OPC and CM samples at (a) 0.5−3.0 nm and (b) 3.0−5.0 nm.

range of methods.53 We measured the total volume of pores with diameter < 1.25 nm (V1), which are included in the first primary peak in the pore size distribution, and the total volume of pores with the diameter between 1.25 and 5 nm (V2) where the other pores are located within the said distribution. The V1/V2 metric for OPC and CM samples is shown in Table 5,

The salient structural properties, including surface area, total pore volume, and pore size distribution, were determined by nitrogen adsorption−desorption isotherms for the CM samples and compared to OPC by itself. The total pore volumes of the CM samples are close to that of OPC, as seen in Table 4, having a gradually decreasing trend with their respective surface areas (Figure S5, Supporting Information). Within the CM samples, CM6 has the smallest surface area (651 m2/g) as a result of the largest amount of Fe2O3 added, which in turn has a larger density and significantly smaller surface area compared to OPC. Nitrogen adsorption−desorption isotherms of all the obtained CM samples (Figure 9a) show typical type I curves, indicating that these CM samples predominantly consist of micropores on the inside. Compared to the isotherm for pure OPC (Figure 9b), a type H4 hysteresis loop was observed in each isotherm for the CM samples, with the characteristic step down at a partial pressure of P/P0 ≈ 0.45. The hysteresis loop was likely formed due to capillary condensation related to narrow fissures in these CM composite samples. This also indicates that these materials have both mesopores and micropores. Pore size distribution curves in Figure 10a derived from the adsorption branch show that the CM samples mainly consist of micropores in the range of 1−2 nm. Compared to the pore size distribution plot of OPC, a peak at around 4 nm was observed for the CM samples (Figure 10b). We hypothesize that this peak is attributed to the addition of Fe2O3 owing to previous studies on similar characterization of materials prepared by a

Table 5. Ratio of Pore Volumes for Pores under 1.25 nm and Pores between 1.25and 5 nm sample

V1/V2 ratio

OPC CM1 CM2 CM3 CM4 CM5 CM6

3.05 2.30 2.40 2.23 1.66 1.15 1.16

with the value for OPC being 3.05. This V1/V2 value drops gradually from 2.30 for sample CM1 to 1.16 for sample CM6, in order of increasing iron concentration. This indicates that the addition of Fe2O3 promotes the CM composite materials to form larger pores, which in turn result in an increased surface area for H2S intake. Samples CM3−CM6 were also characterized for detailed compositional and structural information. Fourier transform infrared (FTIR) spectroscopy was employed to understand better the functional groups present in these samples. The IR G


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Figure 11. (a) FT-IR and (b) Raman spectra for as-synthesized OPC and samples CM3−CM6.

which matches findings from Raman spectroscopy in that there are more sp3 carbons in the CM samples relative to those in OPC. The O 1s spectrum (Figure 12b) can be resolved into three main peaks that can be assigned to the following functional groups: C−O−C group (533.2 eV) and CO group (531.8 eV) from the carbon frame in the sample and lattice oxygen in Fe3O4 at 530.4 eV. The Fe 2p spectrum (Figure 12c) can be similarly fitted with the following peaks: Fe−O 2p3/2 at 711.2 eV can be deconvoluted into two peaks corresponding to Fe2+ (710.9 eV) and Fe3+ (713.2 eV). Fe−O 2p1/2 at 724.6 eV can also be fitted with Fe2+ (724.3 eV) and Fe3+ (726.2 eV).57 An additional shoulder between these two main peaks is attributed to the Fe−O 2p3/2 satellite peak (719.3 eV).57 Elemental iron was not detected by XPS as it is normally covered by a native oxide layer, the thickness of which exceeds the reliable measurement depth of XPS. H2S Uptake Test of Chemically Mixed Fe2O3-OPC Composites. In order to identify the chemically mixed (CM) composite with the highest H2S uptake capacity, we measured the wt % of sulfur for each CM sample and applied a similar ideal addition plot as a reference, similar to analysis of the PM composites from earlier. As seen in Figure 13, the addition of Fe2O3 into OPC results in a drop in H2S uptake for sample CM1 relative to OPC by itself (0 wt % Fe), and this trend continues with samples CM2 and CM3 having decreased uptake of H2S relative to the previous samples. As the iron concentration further increases, however, in samples CM4 and CM5, the H2S uptake increases and even goes beyond that of the simple addition reference line. It is noteworthy that sample CM5 with 45.18 wt % Fe has H2S uptake that reaches 33.20%, which is pretty much at the same level of pure OPC. In contrast, sample CM6 with the highest iron concentration has the H2S uptake drop to 27.99% while still being significantly higher than the simple addition reference would suggest. The same trend was also observed from EDS analysis (Figure S3, Supporting Information). The measured adsorption capacity of H2S for the CM samples is in agreement with reported values for similar materials in the literature, as shown in Table 7. We propose that the H2S uptake capacity for CM samples is likely controlled by the following factors: blocking of the accessible surface between OPC and the active iron species, pore size distribution, and total surface area. When the iron concentration in the CM composite is extremely low, the samples have a similar microstructure to OPC itself, including pore size and surface area. Addition of Fe2O3 causes blockage of accessible surface area for OPC, which in turn results in the

spectra for CM samples (Figure 11a) show vibrational stretches that are close to that for the OPC sample by itself, and these bands can be attributed to the CC symmetrical and asymmetrical stretching vibrations in the conjugated system (1538 and 1575 cm−1, respectively) and C−O−C asymmetrical stretching vibration (1043 cm−1). The peak at 787 cm−1 is relatively strong in the IR spectrum for OPC, but its intensity keeps decreasing in the CM samples as the iron concentration increases. This peak can be assigned to the outof-plane C−H or CC bend. The peak at 548 cm−1 is one of the characteristic bands for the iron species53 and shows a higher intensity as the iron concentration increases. The structural information obtained thus far was further enriched by Raman spectroscopy, with the collected spectra for the OPC and CM composites samples presented in Figure 11b. A series of peaks between 200 and 700 cm−1 were observed in the CM samples, corresponding to the characteristic peaks for the iron species.54 The relative intensity of this batch of peaks gradually increased as we go from CM3 to CM6 wherein the iron concentration also increases. The other two major bands, G band at 1590 cm−1 and D band at 1340 cm−1, can be attributed to aromatic sp2 carbon and amorphous sp3 carbon, respectively.55 The intensity ratio of these two bands (ID/IG) for the CM samples varies from 0.91 to 1.06, as seen in Table 6, which is larger than that for OPC (0.86). The trend in the Table 6. Intensity Ratio of D Band and G Band (ID/IG) for OPC and CM Samples before and after H2S Uptake Test sample name

as-synthesized

after treatment with H2S

OPC CM3 CM4 CM5 CM6

0.86 0.91 0.97 1.07 1.06

0.91 1.10 0.97 1.14 2.04

ID/IG value is an increasing one with increasing iron concentration for CM samples. This result indicates that by adding Fe2O3, some of the sp2 carbons were removed and more sp3 carbons were formed for the composite materials.56 The elemental composition and chemical states for one of the CM composites, CM5, were also studied by XPS (Figure 12). The C 1s spectrum (Figure 12a) shares nearly the same peaks as OPC (cf. Figure 2a), which can be assigned to sp2hybridized CC (284.8 eV), sp3-hybridized C−C (286.2 eV), C−O−C (287.1 eV), and CO (288.9 eV).31,32 In contrast to OPC, CM5 exhibited a much higher C−C peak intensity, H


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Energy & Fuels

Figure 12. XPS elemental scanning of as-synthesized chemically mixed sample CM5 in the vicinity of (a) C 1s, (b) O 1s, and (c) Fe 2p.

been shown before that Fe3O4 has a better H2S uptake performance relative to Fe2O3,59 and iron has also been proposed to be used as a successful H2S sorbent.60,61 No iron signal was seen in the XRD pattern of sample CM5 after H2S uptake, which also points toward the high reactivity of iron formed in the composite. We see this positive synergy, where the net composite effect is higher than the sum of the individual effects, happen in a relatively short range of iron content on OPC, which happens to be where samples CM4 and CM5 lie. The subsequent drop of H2S uptake for sample CM6 is due to the much lower effective surface area available as surface area is dominant in this gas adsorption performance phase. FTIR spectroscopy was employed to further study the structure change of CM samples after being treated with H2S (Figure S8a, Supporting Information). Carbon-related vibrations are much weaker compared to the as-synthesized materials based on the FTIR spectra for CM samples after being treated with H2S. The Raman ID/IG value for both OPC and CM samples increased after being treated with H2S (Table 6 and Figure S8b), with the value varying from 1.10 to 2.04 for the CM composites. The value change is particularly significant for sample CM6, which had a much smaller surface area relative to the other samples. This also shows that a significant amount of addition of iron oxide reduces the OPC ordered

Figure 13. Relationship between iron concentration and H2S uptake for the CM sample series.

drop of H2S uptake performance (Figure 13). A further increase in iron oxide content results in a shift in the pore size distribution (see above), resulting in an increasing fraction of larger micropores than before in the composites as the iron concentration increases. These larger pores will increase the potential for H2S molecules to get inside OPC via the larger bulk volume available, even compared to pure OPC, and also have the iron species react with H2S simultaneously. It has

Table 7. Comparison of Adsorption Properties of a Range of H2S Sorbents SBET (m2 g−1) capacity (g H2S/g sorbent)

sorbent

H2S conc

NaOH/activated carbon 20% Fe/C nanofilaments 1% iron oxide/activated carbon 1% manganese oxide/activated carbon 5% ferrihydrite/graphene oxide 10% ferrihydrite/graphene oxide 20% ferrihydrite/graphene oxide 5% magnetite/graphene oxide 10% magnetite/graphene oxide 20% magnetite/graphene oxide Fe2O3/SiO2 CM1 CM2 CM3 CM4 CM5 CM6

H2S in N2 (1000 ppm) H2S in helium (500 ppm) H2S in N2 (3000 ppm, O2/H2S = 1:1) H2S in N2 (3000 ppm, O2/H2S = 1:1) H2S in air (1000 ppm) H2S in air (1000 ppm) H2S in air (1000 ppm) H2S in air (1000 ppm) H2S in air (1000 ppm) H2S in air (1000 ppm) H2S in N2 (500 mg/m3) 99.5% H2S 99.5% H2S 99.5% H2S 99.5% H2S 99.5% H2S 99.5% H2S I

815 a 886 905 345 359 282 111 174 176 113 1131 1229 876 864 907 651

1.05 0.007 0.0521 0.142 0.0404 0.0286 0.0367 0.0102 0.0099 0.0137 0.727 0.329 0.312 0.307 0.379 0.528 0.413

temperature (°C) 30 100 180 180 25 25 25 25 25 25 80 25 25 25 25 25 25

reference 25 21 58 58 30 30 30 30 30 30 8 this this this this this this

work work work work work work


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Figure 14. XPS elemental scanning of sample CM5 after the H2S test in the vicinity of (a) C 1s, (b) O 1s, (c) Fe 2p, and (d) S 2p.

2p3/2 satellite peak at 720.5 eV, and two doublets attributed to Fe−O and Fe−S. The Fe−O doublet is located at the same location as in the CM sample before the H2S test. The Fe−S doublet is 714.9 and 726.8 eV, corresponding to Fe−S 2p3/2 and 2p1/2, respectively. Comparisons of the PM and the CM Samples. The chemical and structural properties as well as the H2S uptake performance of the PM samples and the CM samples are summarized and compared in Table 8. The PM samples

structure. The characteristic peaks for iron species between 200 and 700 cm−1 also had a decreased intensity after H2S uptake, which further confirms the reaction between iron species and H2S. Sample CM5 after H2S uptake was characterized by XPS, which helps confirm the composition change of the sample as seen in Figure 14. The C 1s spectrum shows nearly the same peak assignments as described above. However, for the O 1s spectrum (Figure 14b), the peak assignments are the same as for as-synthesized sample CM5, except with a much lower Fe− O intensity. This also points to the reaction between iron oxide and H2S during the uptake test and agrees also with the other data. Several sulfur-containing species appear in the S 2p spectrum (Figure 14d), which can be resolved into the following peaks: FeS2 at 161.8 eV and two doublets located at the same position compared with pure OPC after the H2S test. The first doublet at 164.0 and 165.2 eV corresponds to S0 2p3/2 and S0 2p1/2, respectively; see above. No detection of bands related to C−S bonds in both FT-IR and Raman spectra (typically observed at 570−710 cm−1 for C−S) also confirms the hypothesis that this set of peaks is from elemental sulfur instead of any C−S bond in the sample. X-ray diffraction pattern of sample CM5 after the H2S uptake test also shows that elemental sulfur is the predominant product (Figure S9). The second doublet at 168.9 and 170.2 eV (Figure 14d) corresponds to SO42− (S 2p3/2 and 2p1/2, respectively). The intensity of the second doublet is much higher than that of the same doublet in OPC after H2S uptake since SO42− groups are also attributed to the oxidization of iron−sulfur species, as has been discussed previously. The reaction between iron species and H2S can be also reflected from the Fe 2p spectrum. The Fe 2p spectrum can be fitted to several peaks that can be assigned to the following functional groups: FeS2 at 707.7 eV, the Fe−O

Table 8. Comparisons of the PM and the CM Samples characteristic iron species morphology surface area pore structure

PM samples

CM samples

Fe2O3 Fe2O3 on the surface of OPC decreases with the amount of Fe2O3 added same as OPC

Fe3O4, Fe iron species uniformly distributed onto and into OPC remains the same for a range of iron concentration (lower than 45 wt %) larger micropores as more Fe2O3 is added good with carefully tuned iron amount positive

H2S uptake poor performance synergic effect negative

contain Fe2O3 on the surface of OPC, causing the blockage of each other’s active sites, while the CM samples have the more reactive iron species Fe3O4 and Fe, which are uniformly distributed on OPC. Pore structures remain the same for all the PM samples but with a gradual decrease in the surface area. Contrary to this, for the CM samples, the surface area generally remains the same as that of OPC, while larger micropores form to help enhance gas diffusion. J


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CONCLUSIONS We employed two different methods to synthesize a new composite of oxygenated porous carbon (OPC) and Fe2O3 particles for H2S adsorption. In the first approach, we physically mixed as-synthesized Fe2O3 and OPC. The H2S adsorption capacity for physically mixed samples with different iron concentration was much lower than an ideal scenario employing simple additions of the individual uptake from the two species, which can be attributed to the blockage of pores and accessible surface for both OPC and Fe2O3 when they were physically mixed. This results in effective negative synergy with physical mixing and is not recommended for use. In the second method, we created a series of chemically mixed samples by adding Fe2O3 into the polymer precursor of OPC and activating them afterward to synthesize the composites. Fe2O3 turns into more reactive iron and Fe3O4 after the activation process. The addition of Fe2O3 did not influence the surface area of CM samples compared to pure OPC but gradually increased the pore size of the composites. As the iron concentration in CM samples increased, the H2S total uptake dropped at first due to the blockage of the accessible surface and then started to increase even past the simple addition scenario as the larger pore size enhanced the gas uptake. This result agrees with previously reported work that large pores in the carbon framework allow easier diffusion of H2S.28 The chemically mixed sample with an iron concentration of 45.18% showed the best performance, and following this, the uptake dropped again with further increased iron concentration as a result of the additional amount of Fe2O3 disrupting the carbon frame of OPC. This caused a sharp decrease in surface area, which is the dominant factor in this operating regime. The sample with the highest H2S uptake capacity was also found to have a much higher initial uptake rate compared to OPC as well as a high adsorption capacity relative to Fe2O3, taking advantage of the two species collectively in a positive synergic effect for the composite. Faster uptake rates were not explored in previously published works as a benefit of synthesizing activated carbon loaded with metal oxides for H2S adsorption.28,30 Further optimization of the process will enable a replicable volcano-style graph of gas adsorption activity to the composition of the adsorbent, allowing users to fine tune the composite as needed. In addition, the removal of H2S from a multigas atmosphere containing moisture, CO2, and other gases can be further tested to explore the selective desulfurization performance of the composite materials. Regeneration strategies can also be further explored by treating the spent materials with steam or controlled oxidation to effectively reuse the sorbents for multiple desulfurization cycles.62−64 This work shows potential for the chemically mixed composites to be use in practical real-life situations, especially those that require fast and consistent adsorption, such as biomedical filters or waste water treatment.

Based on the differences, the H2S uptake performance for the PM samples is highly dependent on the surface area and pore volume (Figure S10, Supporting Information). In other words, the influence of pore blockage is predominant and thus shows negative synergic effects. For the CM samples, although the influence of pore blockage still exists, there is no obvious trend between surface area and H2S uptake (Figure S11a, Supporting Information). Instead, there are two positive synergic effects. First, the chemically mixed samples end up having larger micropores as more Fe2O3 is added, and the H2S uptake amount has a general trend of inverse proportionality with the pore diameter (Figure S11b). Second, the chemical compositions of iron in the CM materials are more reactive than the original Fe2O3. With the advantages of the more reactive iron species, larger pore diameter, and better diffusion of reactant and products, the CM samples have better potential to be used as a H2S sorbent owing to the positive synergic effect of iron species and OPC. H2S Adsorption Rate for OPC, Fe2O3, and Sample CM5. The results thus far showed that the sample CM5 composite has the highest H2S uptake among all the CM samples in our tests. Aside from total uptake, the instantaneous adsorption rate is also crucial for a sorbent if we want to extend its use to a realistic case. In order to further compare the uptake rate of the chemically mixed composites, the H2S uptake for OPC, precipitated Fe2O3, and sample CM5 for the 24 h duration of the tests was plotted in Figure 15. Compared

Figure 15. Cumulative uptake curve of H2S adsorbed by OPC, precipitated Fe2O3, and sample CM5.

to pure OPC, Fe2O3 shows a much faster rate at the beginning (9.03 wt % S uptake within 1 h), and the uptake increased at a slower rate after that. This faster rate is due to the chemical reaction between Fe2O3 and H2S initiating immediately after exposure. OPC had only a 5.46% uptake in the first 4 h but then took a giant leap forward for the next 4 h, and the uptake rate slowed down afterward. The relatively slower but consistent uptake ensures a long time use for OPC as a sorbent. Sample CM5, by comparison, not only had a high instant uptake (9.38% within 1 h) similar to Fe2O3 but also continuous uptake over the test duration of 24 h as with OPC. This is proof of the positive synergy achieved by chemically mixing iron oxide and ultraporous carbon to create composite sorbents for H2S uptake and differs strongly from the negative synergy we saw with the physically mixed composites.

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ASSOCIATED CONTENT

* Supporting Information S

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.energyfuels.9b01012. XRD patterns, N2 adsorption−desorption isotherms, pore size distribution, plots of surface area versus wt % K


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(10) Xue, M.; Chitrakar, R.; Sakane, K.; Ooi, K. Screening of adsorbents for removal of H2S at room temperature. Green Chem. 2003, 5, 529−534. (11) Stirling, D. The Sulfur Problem: Cleaning Up Industrial Feedstocks, 1st Ed.; Royal Society of Chemistry: Cambridge; 2007. (12) Adib, F.; Bagreev, A.; Bandosz, T. J. Effect of pH and surface chemistry on the mechanism of H2S removal by activated carbons. J. Colloid Interface Sci. 1999, 216, 360−369. (13) Bandosz, T. J. On the Adsorption/Oxidation of hydrogen sulfide on activated carbons at ambient temperatures. J. Colloid Interface Sci. 2002, 246, 1−20. (14) Bagreev, A.; Adib, F.; Bandosz, T. J. pH of activated carbon surface as an indication of its suitability for H2S removal from moist air streams. Carbon 2001, 39, 1897−1905. (15) Adib, F.; Bagreev, A.; Bandosz, T. J. Analysis of the relationship between H2S removal capacity and surface properties of unimpregnated activated carbons. Environ. Sci. Technol. 2000, 34, 686−692. (16) Habeeb, O. A.; Ramesh, K.; Ali, G. A. M.; Yunus, R. M.; Thanusha, T. K.; Olalere, O. A. Modeling and optimization for H2S adsorption from wastewater using coconut shell based activated carbon. Aust. J. Basic Appl. Sci. 2016, 10, 136−147. (17) Liang, M.; Zhang, C.; Zheng, H. The removal of H2S derived from livestock farm on activated carbon modified by combinatory method of high-pressure hydrothermal method and impregnation method. Adsorption 2014, 20, 525−531. (18) Förster, H.; Schuldt, M. Infrared spectroscopic study of the adsorption of hydrogen sulfide on zeolites NaA and NaCaA. J. Colloid Interface Sci. 1975, 52, 380−385. (19) Karge, H. G.; Raskó, J. Hydrogen sulfide adsorption on faujasite-type zeolites with systematically varied Si-Al ratios. J. Colloid Interface Sci. 1978, 64, 522−532. (20) Barron, A. R.; Coker, C. E.; Loscutova, J. R. Method to remove sulfur or sulfur-containing species from a source. US Patent US 7,569,199 B1, 2007. (21) Fauteux-Lefebvre, C.; Abatzoglou, N.; Blais, S.; Braidy, N.; Hu, Y. Iron oxide-functionalized carbon nanofilaments for hydrogen sulfide adsorption: The multiple roles of carbon. Carbon 2015, 95, 794−801. (22) Lin, C.; Qin, W.; Dong, C. H2S adsorption and decomposition on the gradually reduced α-Fe2O3 (001) surface: A DFT study. Appl. Surf. Sci. 2016, 387, 720−731. (23) Florent, M.; Wallace, R.; Bandosz, T. J. Removal of hydrogen sulfide at ambient conditions on cadmium/GO-based composite adsorbents. J. Colloid Interface Sci. 2015, 448, 573−581. (24) Giannakoudakis, D. A.; Bandosz, T. J. Zinc (hydr)oxide/ graphite oxide/AuNPs composites: role of surface features in H2S reactive adsorption. J. Colloid Interface Sci. 2014, 436, 296−305. (25) Castrillon, M. C.; Moura, K. O.; Alves, C. A.; Bastos-Neto, M.; Azevedo, D. C. S.; Hofmann, J.; Möllmer, J.; Einicke, W. D.; Gläser, R. CO2 and H2S removal from CH4-rich streams by adsorption on activated carbons modified with K2CO3, NaOH, or Fe2O3. Energy Fuels 2016, 30, 9596−9604. (26) de Falco, G.; Montagnaro, F.; Balsamo, M.; Erto, A.; Deorsola, F. A.; Lisi, L.; Cimino, S. Synergic effect of Zn and Cu oxides dispersed on activated carbon during reactive adsorption of H2S at room temperature. Microporous Mesoporous Mater. 2018, 257, 135− 146. (27) Bagreev, A.; Bandosz, T. J. On the mechanism of hydrogen sulfide removal from moist air on catalytic carbonaceous adsorbents. Ind. Eng. Chem. Res. 2005, 44, 530−538. (28) Zhang, Z.; Wang, J.; Li, W.; Wang, M.; Qiao, W.; Long, D.; Ling, L. Millimeter-sized mesoporous carbon spheres for highly efficient catalytic oxidation of hydrogen sulfide at room temperature. Carbon 2016, 96, 608−615. (29) Hernández, S. P.; Chiappero, M.; Russo, N.; Fino, D. A novel ZnO-based adsorbent for biogas purification in H2 production systems. Chem. Eng. J. 2011, 176-177, 272−279. (30) Arcibar-Orozco, J. A.; Wallace, R.; Mitchell, J. K.; Bandosz, T. J. Role of surface chemistry and morphology in the reactive adsorption

Fe, H2S uptake versus time, wt % S versus wt % Fe, FTIR spectra, Raman spectra, and TEM image (PDF)

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected] ORCID

Kexin Ling: 0000-0003-1112-2428 Andrew R. Barron: 0000-0002-2018-8288 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS Financial support was provided by the Robert A. Welch Foundation (C-0002) and the Reducing Industrial Carbon Emissions (RICE) research operations funded by the Welsh European Funding Office (WEFO) through the Welsh Government. The authors would also like to thank Dr. Saunab Ghosh and Dr. Bruce E. Brinson for useful discussions and help in collecting TEM images of a sample.

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NOMENCLATURE ATR = attenuated total reflection BET = Brunauer−Emmett−Teller CM = chemically mixed EDS = energy-dispersive X-ray spectroscopy FT-IR = Fourier transform infrared NLDFT = non-local density functional theory OPC = oxygenated porous carbon PAA = polyanisyl alcohol PM = physically mixed PVP = poly(vinylpyrrolidone) SEM = scanning electron microscopy TEM = transmission electron microscopy XPS = X-ray photoelectron spectroscopy XRD = X-ray diffraction

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REFERENCES

(1) Fan, H.; Li, C.; Xie, K. Sorbents for high-temperature removal of hydrogen sulfide from coal-derived fuel gas. J. Nat. Gas Chem. 2001, 10, 256−270. (2) Gostelow, P.; Parsons, S. A.; Stuetz, R. M. Odour measurements for sewage treatment works. Water Res. 2001, 35, 579−597. (3) Pluth, M. D.; Bailey, T. S.; Hammers, M. D.; Hartle, M. D.; Henthorn, H. A.; Steiger, A. K. Natural products containing hydrogen sulfide releasing moieties. Synlett 2015, 26, 2633−2643. (4) Nam, S.; Hur, K.-B.; Lee, N.-H. Effects of hydrogen sulfide and siloxane on landfill gas utility facilities. Environ. Eng., Res. 2011, 16, 159−164. (5) Vollertsen, J.; Nielsen, A. H.; Jensen, H. S.; Wium-Andersen, T.; Hvitved-Jacobsen, T. Corrosion of concrete sewersThe kinetics of hydrogen sulfide oxidation. Sci. Total Environ. 2008, 394, 162−170. (6) O’Neill, P. Environmental chemistry, 3rd Ed; Boca Raton: CRC Press, 2017. (7) Garces, H. F.; Galindo, H. M.; Garces, L. J.; Hunt, J.; Morey, A.; Suib, S. L. Low temperature H2S dry-desulfurization with zinc oxide. Microporous Mesoporous Mater. 2010, 127, 190−197. (8) Huang, G.; He, E.; Wang, Z.; Fan, H.; Shangguan, J.; Croiset, E.; Chen, Z. Synthesis and characterization of γ-Fe2O3 for H2S removal at low temperature. Ind. Eng. Chem. Res. 2015, 54, 8469−8478. (9) Wang, J.; Wang, L.; Fan, H.; Wang, H.; Hu, Y.; Wang, Z. Highly porous copper oxide sorbent for H2S capture at ambient temperature. Fuel 2017, 209, 329−338. L


Article

Energy & Fuels of H2S on iron (hydr)oxide/graphite oxide composites. Langmuir 2015, 31, 2730−2742. (31) Ghosh, S.; Sevilla, M.; Fuertes, A. B.; Andreoli, E.; Ho, J.; Barron, A. R. Defining a performance map of porous carbon sorbents for high-pressure carbon dioxide uptake and carbon dioxide-methane selectivity. J. Mater. Chem. A 2016, 4, 14739−14751. (32) Ghosh, S.; Barron, A. R. Optimizing carbon dioxide uptake and carbon dioxide-methane selectivity of oxygen-doped porous carbon prepared from oxygen containing polymer precursors. ChemistrySelect 2017, 2, 11959−11968. (33) Halder, N. C.; Wagner, C. N. J. Separation of particle size and lattice strain in integral breadth measurements. Acta Crystallogr. 1966, 20, 312−313. (34) Jain, V.; Biesinger, M. C.; Linford, M. R. The GaussianLorentzian sum, product, and convolution (Voigt) functions in the context of peak fitting X-ray photoelectron spectroscopy (XPS) narrow scans. Appl. Surf. Sci. 2018, 447, 548−553. (35) Wang, D.; Song, C.; Zhao, Y.; Yang, M. Synthesis and characterization of monodisperse iron oxides microspheres. J. Phys. Chem. C 2008, 112, 12710−12715. (36) Sugimoto, T.; Sakata, K.; Muramatsu, A. Formation mechanism of monodisperse pseudocubic α-Fe2O3 particles from condensed ferric hydroxide gel. J. Colloid Interface Sci. 1993, 159, 372−382. (37) Seader, J. D.; Henley, E. J.; Roper, D. K. Separation Process Principles: Chemical and Biochemical Operations; 3rd Ed, John Wiley & Sons, 2011; 601−603. (38) Bashkova, S.; Baker, F. S.; Wu, X.; Armstrong, T. R.; Schwartz, V. Activated carbon catalyst for selective oxidation of hydrogen sulphide: on the influence of pore structure, surface characteristics, and catalytically-active nitrogen. Carbon 2007, 45, 1354−1363. (39) Xiao, Y.; Wang, S.; Wu, D.; Yuan, Q. Catalytic oxidation of hydrogen sulfide over unmodified and impregnated activated carbon. Sep. Purif. Technol. 2008, 59, 326−332. (40) O’Donoghue, M. A Guide to Man-made Gemstones; 1st Ed., Van Nostrand Reinhold Company: New York, 1983. (41) Arndt, D.; Zielasek, V.; Dreher, W.; Bäumer, M. Ethylene diamine-assisted synthesis of iron oxide nanoparticles in high-boiling polyolys. J. Colloid Interface Sci. 2014, 417, 188−198. (42) Carnes, C. L.; Klabunde, K. J. Unique chemical reactivities of nanocrystalline metal oxides toward hydrogen sulfide. Chem. Mater. 2002, 14, 1806−1811. (43) Davidson, J. M.; Lawrie, C. H.; Sohail, K. Kinetics of the absorption of hydrogen sulfide by high purity and doped high surface area zinc oxide. Ind. Eng. Chem. Res. 1995, 34, 2981−2989. (44) Davidson, J. M.; Sohail, K. A drifts study of the surface and bulk reactions of hydrogen sulfide with high surface area zinc oxide. Ind. Eng. Chem. Res. 1995, 34, 3675−3677. (45) Gendler, T. S.; Shcherbakov, V. P.; Dekkers, M. J.; Gapeev, A. K.; Gribov, S. K.; McClelland, E. The lepidocrocite−maghemite− haematite reaction chainI. Acquisition of chemical remanent magnetization by maghemite, its magnetic properties and thermal stability. Geophys. J. Int. 2005, 160, 815−832. (46) Baltrusaitis, J.; Cwiertny, D. M.; Grassian, V. H. Adsorption of sulfur dioxide on hematite and goethite particle surfaces. Phys.Chem. Chem. Phys. 2007, 9, 5542−5554. (47) Galicia, P.; Batina, N.; González, I. The Relationship between the surface composition and electrical properties of corrosion films formed on carbon steel in alkaline sour medium: an XPS and EIS Study. J. Phys. Chem. B 2006, 110, 14398−14405. (48) Barr, T. L. An ESCA Study of the Termination of the Passivation of Elemental Metals. J. Phys. Chem. 1978, 82, 1801−1810. (49) Davydov, A.; Chuang, K. T.; Sanger, A. R. Mechanism of H2S oxidation by ferric oxide and hydroxide surfaces. J. Phys. Chem. B 1998, 102, 4745−4752. (50) Carter, W. J.; Schweitzer, G. K.; Carlson, T. A. Experimental evaluation of a simple model for quantitative analysis in X-ray photoelectron spectroscopy (No. CONF-740429-1). Syracuse Univ., NY USA; Oak Ridge National Lab., Tenn. USA. 1974.

(51) Powell, C. J.; Seah, M. P. Precision, accuracy, and uncertainty in quantitative surface analyses by Auger-electron spectroscopy and Xray photoelectron spectroscopy. J. Vac. Sci. Technol., A 1990, 8, 735− 763. (52) Huffman, W. P. The importance of active surface area in the heterogeneous reactions of carbon. Carbon 1991, 29, 769−776. (53) Blatt, O.; Helmich, M.; Steuten, B.; Hardt, S.; Bathen, D.; Wiggers, H. Iron oxide/polymer-based nanocomposite Material for hydrogen sulfide adsorption applications. Chem. Eng. Technol. 2014, 37, 1938−1944. (54) Namduri, H.; Nasrazadani, S. Quantitative analysis of iron oxides using Fourier transform infrared spectrophotometry. Corros. Sci. 2008, 50, 2493−2497. (55) De Faria, D. L. A.; Venâncio Silvia, S.; de Oliveira, M. T. Raman microspectroscopy of some iron oxides and oxyhydroxides. J. Raman Spectrosc. 1997, 28, 873−878. (56) Dresselhaus, M. S.; Dresselhaus, G.; Saito, R.; Jorio, A. Raman spectroscopy of carbon nanotubes. Phys. Rep. 2005, 409, 47−99. (57) Ferrari, A. C.; Robertson, J. Interpretation of Raman spectra of disordered and amorphous carbon. Phys. Rev. B 2000, 61, 14095. (58) Wright, K. D.; Barron, A. R. Catalyst residue and oxygen species inhibition of the formation of hexahapto-metal complexes of group 6 metals on single-walled carbon nanotubes. C 2017, 3, 17. (59) Fang, H.; Zhao, J.; Fang, Y.; Huang, J.; Wang, Y. Selective oxidation of hydrogen sulfide to sulfur over activated carbonsupported metal oxides. Fuel 2013, 108, 143−148. (60) Janetaisong, P.; Lailuck, V.; Supasitmongkol, S. Pelletization of iron oxide based sorbents for hydrogen sulfide removal. Key Eng. Mater. 2017, 751, 449−454. (61) Tian, H.; Wu, J.; Zhang, W.; Yang, S.; Li, F.; Qi, Y.; Zhou, R.; Qi, X.; Zhao, L.; Wang, X. High performance of Fe nanoparticles/ carbon aerogel sorbents for H2S removal. Chem. Eng. J. 2017, 313, 1051−1060. (62) Zeng, B.; Yue, H.; Liu, C.; Huang, T.; Li, J.; Zhao, B.; Zhang, M.; Liang, B. Desulfurization behavior of Fe-Mn-based regenerable sorbents for high-temperature H2S removal. Energy Fuels 2015, 29, 1860−1867. (63) Cheah, S.; Carpenter, D. L.; Magrini-Bair, K. A. Review of midto high-temperature sulfur sorbents for desulfurization of biomassand coal-derived syngas. Energy Fuels 2009, 23, 5291−5307. (64) Cal, M. P.; Strickler, B. W.; Lizzio, A. A.; Gangwal, S. K. High temperature hydrogen sulfide adsorption on activated carbon: II. Effects of gas temperature, gas pressure and sorbent regeneration. Carbon 2000, 38, 1767−1774.

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Synergic Adsorption of H2S Using High Surface Area Iron Oxide

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