Introduction

Air contamination is a global issue which has been amplified by various anthropogenic activities in the last many decades. Among numerous air contaminants toxifying the air, hydrogen sulfide (H2S) and sulfur dioxide (SO2) has known to cause severe damage to human health and the environment. H2S is a toxic pungent-smelling gas released from decayed organic matter, oil industry, coal and natural gas-based thermal power plants, and sewage treatment facilities1,2. Acute exposure to H2S at levels of 200–500 ppm could paralyze the olfactory nerve and beyond 500 ppm could lead to sudden death2,3. Furthermore, H2S conversion to SO2 and its hydrolysis to form acid rain could acidify soil and water bodies, which could be disastrous to plants and marine life, respectively4,5. SO2 is a colourless toxic gas with a sharp odour, which could cause various respiratory ailments such as chronic bronchitis and infections of the respiratory tract6. While a low SO2 concentration of 1–5 ppm is enough for human discomfort, exposure above 100 ppm could be life-threatening7. The main sources of atmospheric SO2 are thermal power plants and vehicular emissions8. Thus, H2S and SO2 removal from point of origin should be prioritized to limit air contamination and prevent catastrophic events like smog formation and acid rain.

Chemical adsorption of these toxic gases over an adsorbent surface is one of the most simplistic and affordable methods to adsorb and mineralize H2S and SO2 gases to non-toxic by-products like sulfur and sulfates9. Moreover, chemisorption is highly efficient for flue gas desulfurization and natural gas purification applications, which are challenging, fundamentally and monetary-wise1,10. For this purpose, metal oxides have shown great potential due to the presence of weak basic sites (lattice oxygen) and basic OH– groups, which could react with acidic gases like H2S and SO2 (acting as electron donors)11,12. The surface reactivity of metal oxides for these gases could be amplified in the presence of water molecules. Firstly, the water layer on the metal oxide surface dissociatively reacts and improves the hydroxyl density. Secondly, the surface water film dissolves the gas molecules, which lowers the energy barrier for reactive interaction with the metal oxide surface and thus favours the overall chemisorption process13,14,15,16,17. Thus, it is worth exploring the positive effect of water during the adsorption of acidic gases over metal oxides, which is the focus of this research work. Also, it is equally important to explore adsorbent materials for the remediation of low H2S/SO2 concentrates to confirm the applicability of adsorbents in deep desulfurization and gas purification applications.

In this study, we have fabricated an affordable Mn-Zn-Fe metal oxide nanocomposite by a one-step coprecipitation method for room-temperature adsorptive removal of H2S and SO2 gases in wet conditions. The gas concentration of 500 and 100 ppm for H2S and SO2 was adopted for their industrial application and suitability in capturing these pollutants in the toxicity range for humans. The oxide showed better adsorption performance in wet conditions with complete mineralization to non-toxic by-products. Besides studying the factors affecting the adsorption process, the adsorption mechanism was studied in detail using various microscopic and spectroscopic techniques. The study confirmed that the oxide nanocomposite has the potential to eliminate and mineralize low concentrations of gaseous H2S and SO2 in dry–wet conditions.

Methods

Chemicals

Manganese(II) nitrate tetrahydrate (Mn(NO3)2·4H2O), zinc(II) nitrate hexahydrate (Zn(NO3)2·6H2O), iron(III) nitrate nonahydrate (Fe(NO3)3·9H2O), and 2.0 mol L−1 NaOH solution were procured from Samchum Pure Chemicals, Korea. H2S gas (0.05 Vol.%) and SO2 gas (0.01 Vol.%) balanced with pure N2 gas were procured from Union gas, Korea. All solutions were prepared in double-distilled water.

Synthesis of nanocomposite

In 50 mL of deionized water, 3.76 g of Mn(NO3)2·4H2O, 4.45 g of Zn(NO3)2·6H2O, and 4.04 g of Fe(NO3)3·9H2O were dissolved to give a final Mn2+:Zn2+:Fe3+ ratio of 3:3:2. A higher ratio of divalent to trivalent cations was adopted for the formation of MnO2 and ZnO in the nanocomposite (for a higher acidic gas adsorption capacity). Under vigorous stirring, 2.0 mol L−1 NaOH solution was added dropwise until the solution pH reached 12.5. This pH was sufficient for the formation of ternary oxide nanocomposites as reported earlier18. After stirring for 2 h, the precipitate was phase separated and dried at 393 K overnight in a hot air oven. The use of an excess of water for washing was avoided to reduce the overall impact on the environment during the material fabrication process.

Analytical instruments

The morphology in the surface and transmission mode was probed over field emission scanning electron microscopy (FE-SEM, Hitachi S-4300, Hitachi, Japan) and field emission TEM (FE-TEM, JEM-2010F, JEOL Ltd., Japan), respectively. SEM analysis was done on finely grounded dried samples after coating them with a gold-platinum alloy by ion-sputtering (E-1048 Hitachi ion sputter). The elemental analysis was done using energy-dispersive X-ray spectroscopy (EDAX, X-Maxn 80 T, Oxford Instruments, United Kingdom). The specific surface area and porosity were determined by analysing the standard N2 adsorption–desorption isotherm at 77 K using a Gemini 2360 series (Micromeritics, Norcross, United States) instrument after degassing at 423 K for 6 h with a mass of 0.324 g. The powder X-ray diffraction (PXRD) patterns were obtained at room temperature (2θ = 5–50°) on an Ultima IV (Rigaku, Japan) X-ray diffractometer with Cu Kα radiation (λ = 1.5406 Å) and a Ni filter. Fourier-transform infrared (FTIR) spectra of samples were recorded using KBr pellets over a Cary670 FTIR spectrometer (Agilent Technologies, United States). For X-ray photoelectron spectroscopy (XPS) analysis, a K-alpha XPS instrument (Thermo Fisher Scientific, United Kingdom) was used with a monochromatic Al Kα X-ray source. The pressure was fixed to 4.8 × 10−9 mbar. Spectra were charge corrected to the main line of the C 1 s (aromatic carbon) set to 284.7 eV. Spectra were analysed using CasaXPS software (version 2.3.14).

Breakthrough protocol

The gas adsorption experiments were performed by taking 0.5 g of the adsorbent in a Pyrex tube (height: 50 cm and diameter: 1 cm) at 298 K. The sample was fixed between the glass wool and supported on silica beads19. The H2S gas (0.05 Vol.%) or SO2 gas (0.01 Vol.%) was passed through it at a fixed flow rate. The outgoing gas was analyzed using an H2S gas analyzer (GSR-310, Sensoronic, Korea) or an SO2 gas analyzer (GASTIGER 6000, Wandi, Korea). The analyzer recorded the effluent gas concentration every minute in real-time until the breakthrough points of 20% (100 ppm for H2S and 20 ppm for SO2) were reached and after that the experiment was completed. The wet samples were prepared by passing water vapours (80% relative humidity) directly through the adsorbent bed before passing the gas through it. The gas adsorption capacity was measured using the following equation:

$$ q = \frac{{C_{0} Q}}{m}\mathop \int \limits_{0}^{{t_{b} }} \left( {1 - \frac{{C_{t} }}{{C_{0} }}} \right)dt $$
(1)

where C0-initial concentration (mg L−1), C-concentration at time ‘t’ (mg L−1), Q-flowrate (L min−1), m-the mass of adsorbent (g), and tb-breakthrough time (s).

Results and discussion

The SEM micrograph of oxide nanocomposite showed irregularly shaped nano-globules, which were uniformly distributed in the entire region (Fig. 1a). The controlled release of base (precipitation agent) assisted in regulating the nucleation and particle growth kinetics, which prevented the aggregation of metal oxides in the ternary nanocomposite. A more detailed investigation of morphology was conducted over high-resolution TEM, which confirmed that the nano-globules were constructed of polyhedral nanoparticles (Fig. 1b). The crystallite planes of nanoparticles were assigned by measuring the fringe width and correlating with the interplanar spacing (d) values from the XRD pattern. The fringe width of 0.308, 0.530, and 0.261 nm were assigned to the MnO2 (110)20, ZnO (0002)21, and MFe2O4 (311)22, respectively. The EDAX elemental analysis confirmed peaks for Mn, Zn, Fe, and O at respective energies having the atomic contribution of 13.80, 7.14, 3.57, and 75.48%, respectively (Fig. 1c). The 2D elemental mapping showed an abundant density of Mn and Zn with low-density regions in the ‘Fe’ map (respective high-density regions marked in ‘Mn’ and ‘Zn’ maps). The possible reason for such a distribution could be the formation of pure, binary, and ternary metal oxides in the nanocomposite (Fig. 1d).

Figure 1
figure 1

(a) SEM micrograph; (b) high-resolution TEM micrograph; (c) elemental distribution analysis; (d) 2D elemental mapping of MZFO.

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The PXRD pattern of MFZO nanocomposite has diffraction peaks for β-MnO2 (purple circle)23, ZnO (pink circle)24, ferrites (green square)25, and NaNO3 (blue square)26, which confirmed a poly-oxide nature of the composite (Fig. 2a). A similar report is available for the fabrication of Cu–Zn-Mn ternary oxide nanocomposite, where CuO, ZnO, and MnO2 nanoparticles were confirmed18. The absence of a washing step during the MFZO fabrication was responsible for the presence of NaNO3 in the sample. The N2 adsorption–desorption isotherm of MFZO nanocomposite exhibited a Type III behaviour, generally expected in macro-porous materials (Fig. 2b)23. The nanocomposite possessed a BET surface area of 21.03 m2 g−1 and a pore volume of 0.07 cm3 g−1. These values are higher than other nanocomposites like Mn2O3/Fe2O3 (6.18 m2 g−1, 0.12 cm3 g−1)27, CeO2/Mn2O3/Fe2O3 (15.64 m2 g−1, 0.09 cm3 g−1)28, and Fe2O3/Na2SO4 (2.89 m2 g−1, 0.01 cm3 g−1)29 used for the same applications. The spectrum has a broad band centred at 621 cm−1 for the metal–oxygen stretching vibrations30,31. The bands at 835 and 1385 cm−1 were attributed to the asymmetric stretching vibration (ν3-NO3−) and out-of-plane bending vibration (ν2-NO3−), respectively, of nitrate32. The bands at 3433 and 1635 cm−1 were assigned to the stretching and bending vibration modes of adsorbed water molecules, respectively (Fig. 2c)33. The XPS survey of MFZO confirmed peaks for Na 1 s, Zn 2p, Fe 2p, Mn 2p, O 1 s, and N 1 s at their respective binding energy. The Na 1 s and N 1 s peaks were associated with the presence of NaNO3 (Fig. 2d).

Figure 2
figure 2

(a) PXRD pattern; (b) N2 adsorption–desorption isotherm; (c) FTIR spectrum; (d) XPS survey of MZFO.

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The HRXPS Mn 2p3/2 signal of MFZO deconvoluted into three contributions at 640.7, 641.6, and 642.8 eV for Mn2+ (23.2%), Mn3+ (43.9%), and Mn4+ (32.9%) oxidation states of Mn ions, respectively34. The analyses showed that multivalent Mn ions were related to the formation of MnO2 and Mn-based ferrites (Fig. 3a, Table S1). The HRXPS Zn 2p spectrum has two peaks at 1021.4 and 1044.3 eV for 2p3/2 and 2p1/2 signals of Zn2+ ions, respectively (Fig. 3b, Table S2)35. In the HRXPS Fe 2p spectrum, the 2p3/2 signal was deconvoluted into two contributions at 710.8 and 712.9 eV for Fe2+ (70.4%) and Fe3+ ions (29.6%), respectively (Fig. 3c, Table S3)36. These two contributions were due to the formation of ferrites. The HRXPS O 1 s spectrum has three deconvoluted peaks at 530.0, 531.4, and 532.9 eV for lattice oxygen (56.9%), surface hydroxyl groups/nitrate ions (24.6%), and water molecules (18.4%), respectively (Fig. 3d, Table S4)37.

Figure 3
figure 3

High-resolution XPS (a) Mn 2p; (b) Zn 2p; (c) Fe 2p; (d) O 1 s spectra of MFZO.

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The synthesized nanocomposite was tested for H2S removal in breakthrough columns both in dry and wet conditions (Fig. 4a). The adsorption capacity of 0.73 mmol g−1 was achieved in the dry condition. However, in the wet condition, the capacity increased to 1.03 mmol g−1, showing the positive role of water in the adsorption process mediated by the dissolution and dissociation of H2S molecules in the water film over the oxide surface. The effect of parameters, i.e., gas flow rate (Fig. 4b) and adsorbent loading (Fig. 4c) on the adsorption capacity was studied in the wet conditions. The adsorption capacity decreased with the increasing flow rate, where the highest capacity of 1.21 mmol g−1 was achieved with a flow rate of 0.1 L min−1. Increasing flow rate disfavoured the adsorbate-adsorbent interaction due to a lowering in the gas retention time, which negatively impacted the capacity38. The adsorption capacity decreased with the increasing adsorbent loading and the maximum capacity of 1.31 mmol g−1 was achieved for 0.2 g of adsorbent and 0.2 L min−1 of flowrate. This behaviour could be associated with the presence of an unutilized mass of the oxide likely due to the cluttering of wet adsorbent particles in the adsorbent bed, which reduces the effective surface area for the reaction to occur31. However, no breakthrough experiment was conducted below 0.2 g. Since the material has a high density, loading adsorbent below 0.2 g led to a narrow bed length (since the tube diameter was 6 mm), which had a poor adsorbate-adsorbent interaction. The maximum adsorption capacity of 1.31 mmol g−1 achieved for the synthesized nanocomposite is similar to or higher than those reported for commercial ZnO (1.16 mmol g−1)39, CeO2-Mn2O3-Fe2O3 (0.83 mmol g−1)28, Mn2O3-Fe2O3 (0.35 mmol g−1)27, α-Fe2O3-Na2SO4 (1.06 mmol g−1)29, and ZnFe2O4 (0.05 mmol g−1)40 in similar experimental conditions.

Figure 4
figure 4

H2S breakthrough curves for (a) dry/wet adsorbents; (b) wet adsorbent at different flowrate; (c) adsorbent loading. Conditions: [Adsorbent] = 0.5 g, flowrate = 0.2 L min−1 (changed otherwise).

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The nanocomposite was also studied for SO2 adsorption in dry and wet conditions (Fig. 5a). The adsorption capacity of 0.22 mmol g−1 in the dry condition nearly doubled to 0.41 mmol g−1 in the wet condition. Such behaviour has been reported for oxidative SO2 adsorption over MnO216. SO2 adsorption in the presence of water molecules significantly accelerated the sulfate formation reaction, which favoured the overall adsorption process. The increasing gas flow rate negatively affected the adsorption process due to poor adsorbate-adsorbent interaction (Fig. 5b). The SO2 adsorption capacity of the composite significantly improved with the decreasing adsorbent loading, where the maximum adsorption capacity of 0.49 mmol g−1 was confirmed with 0.2 g of adsorbent and 0.2 L min−1 of flowrate. Here also, the negative role of increasing bed loading was according to the effect witnessed for H2S gas adsorption (Fig. 5c). Thus, in the optimized experimental conditions, the adsorbent could remove 0.49 mmol g−1 of SO2 gas. This value is highly significant and comparable to the reported values for ZnO (0.28 mmol g−1)41, MnO2 (0.48–1.23 mmol g−1)42, and NaMxOy (0.73 mmol g−1)7 in similar experimental conditions. The nanocomposite possessed a higher H2S adsorption capacity compared to the SO2 gas. The superior H2S adsorption was related to an easy dissociation of H2S molecules due to the much lower energy barrier and higher adsorption energy compared to SO2, which has been previously demonstrated for Zn–MoSe2 structure through computational calculations43.

Figure 5
figure 5

SO2 breakthrough curves for (a) dry/wet adsorbent; (b) wet adsorbent at different flowrate; (c) wet adsorbent at different adsorbent loading. Conditions: [Adsorbent] = 0.5 g, flowrate = 0.2 L min−1 (changed otherwise).

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The SEM and TEM micrographs post-H2S and SO2 adsorption showed no significant variation in the surface morphology (Figs. S1, S2) except for the transformation of a chunk of nanoparticles to cluttered nanorods. It could be due to the combined effect of moisture and gas acidity as the SEM micrographs of dry samples showed no such change in the surface morphology. The EDAX analysis of gas-exposed samples confirmed a new peak at ~ 2.3 keV for sulfur. The intensity of the S peak in the H2S-adsorbed sample is much higher than that of the SO2, which agreed with the experimental results (Fig. S3). The 2D elemental mapping of the H2S and SO2-adsorbed samples confirmed a high density of sulfur atoms over the oxide surface, which was uniformly distributed over the nanocomposite (Fig. 6).

Figure 6
figure 6

2D elemental mapping of wet MFZO after (a) H2S; (b) SO2 adsorption.

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The PXRD patterns of gas-exposed samples showed insignificant changes in the diffraction peaks, except for new peaks in the H2S-adsorbed sample (marked as purple stars). These two peaks were assigned to the presence of ZnSO344. The absence of additional new peaks in these samples could be related to the formation of oxidized sulfur species on the surface (Fig. 7a). The N2 adsorption–desorption isotherms are shown in Fig. 7b. For the H2S-adsorbed sample, the surface area and pore volume decreased by 26 and 17%, respectively (Table 1). However, for the SO2-exposed sample, a minimal drop in these values was observed. The drop in the surface area and porosity was linked to the deposition of oxidized sulfur species, which may have clogged the pores29. This clogging was expected more in the H2S-adsorbed sample due to a higher gas volume adsorption and subsequent mineralization onto the surface.

Figure 7
figure 7

(a) PXRD patterns; (b) N2 adsorption–desorption isotherms; High-resolution XPS (c) Mn 2p; (d) Zn 2p; (e) Fe 2p; (f) O 1 s spectra of wet MZFO after H2S and SO2 adsorption.

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Table 1 Surface area and pore characteristics of wet MFZO before and after gas adsorption.
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More detailed information on the adsorption mechanism was deduced for the XPS analysis of MFZO nanocomposite after the gas adsorption process. In the HRXPS Mn 2p spectrum of the H2S-adsorbed sample, all three contributions for Mn2+, Mn3+, and Mn4+ are present at a slightly lower binding energy with variation in the proportion of these oxidation species. The redshift in the binding energy could be associated with the partial sulfidation of the Mn oxides31. Moreover, the variation in the oxidation state proportions could be linked to the involvement of Mn2+/Mn3+/Mn4+ redox cycles during the chemisorption process17. However, for SO2-adsorbed samples, the only proportion of oxidation states varied with an insignificant shift in the position, which was related to the Mn redox behaviour responsible for the oxidation of SO2 (Fig. 7c)17,45. The HRXPS Zn 2p spectrum of H2S-adsorbed showed a minimal redshift in the peak position probably due to the formation of ZnSO3 species. However, no such shift was witnessed for the SO2-adsorbed sample, which further suggested the delocalized nature of the chemisorption process (Fig. 7d). In the HRXPS Fe 2p spectrum of H2S-adsorbed MFZO, the peak position shifted slightly but with a minimal change in the proportion of Fe2+ and Fe3+ species. DFT calculations have predicted that H2S dissociatively reacts better on the FeO (Fe2+ sites) than Fe2O3 (Fe3+ sites)46. Even in our previously reported work on the adsorption of H2S over Mn2O3/Fe2O3, bulk Fe2O3 phase did not take part in the oxidation process27. The slight variation in the peak position could be linked to the involvement of Fe2+ sites in the H2S adsorption process. For the SO2-adsorbed sample, the Fe2+ and Fe3+ peak positions red-shifted by 0.1 and 0.4 eV, respectively, with a significant drop in the Fe2+ proportion (70.4–62.3%). It has been proven that the SO2 molecules are much more reactive to the Fe2+ sites than the Fe3+ sites. Thus, a drop in the Fe2+ contribution suggested that the divalent Fe sites catered the oxidation of SO2 molecules (Fig. 7e)47. In the HRXPS O 1 s spectrum of H2S-adsorbed sample, the metal–oxygen bond contribution decreased, whereas the contribution at 531.6 eV for −OH/O–N increased due to the formation of metal-sulfide and sulfite (SO3−) species, respectively. For the SO2-adsorbed sample, the 531.5 eV peak improved even further due to the consumption of lattice oxygen and the formation of sulfate species (Fig. 7f).

The HRXPS S 2p spectrum of H2S-adsorbed MFZO was deconvoluted into three sets of doublets with their 2p3/2 peaks observed at 161.3, 163.6, and 167.9 eV for sulfide (36.1%), elemental sulfur (25.1%), and sulfite (38.8%)48. While the formation of metal-bound sulfide is initiated by the dissociated adsorption of H2S (into H+ and HS−) in the presence of water molecules. The formation of elemental sulfur and sulfide is mediated by the redox behaviour of transition metal ions, surface-adsorbed molecular oxygen, and water molecules31. Wang et al.have reported the dissociation of H2S over In2O3 thin film in the presence of moisture, where the reactive dissociation of H2S molecules with adsorbed water produced HS− and H+ species. The formed HS− and H+ ions reacted with the surface-chemisorbed oxygen species to yield sulfide and sulfite species49. The HRXPS S 2p spectrum of SO2-adsorbed MFZO has a set of doublets with a 2p3/2 peak at 168.4 eV, which was assigned to the sulfate species (Fig. 8, Table S5)48. The adsorption of SO2 over the oxide surface is generally driven by the reactive interaction of SO2 molecules with the lattice oxygen or surface hydroxyl groups to form sulfite/bisulfite, which further oxidized to sulfate via redox behaviour of transition metal oxide and gaseous oxygen molecules7,17. Moreover, just like H2S dissolution in the surface water, SO2 could be readily adsorbed and hydrolysed by surface water molecules, which makes the oxidation of SO2 molecules, energetically favourable17.

Figure 8
figure 8

High-resolution XPS S 2p spectra of MFZO before and after H2S and SO2 adsorption.

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Conclusion

In conclusion, we have fabricated an Mn-Zn-Fe metal oxide nanocomposite via a one-step coprecipitation reaction. The fabricated nanocomposite has MnO2, ZnO, and ferrites with a surface area and pore volume of 21.03 m2 g−1 and 0.07 cm3 g−1, respectively. The nanocomposite was tested for room-temperature adsorptive removal of H2S and SO2 in dry and wet conditions. The oxide exhibited better gas adsorption capacity in wet conditions owing to the dissolution and dissociation of gaseous molecules in the surface water film. The adsorbent showed a better adsorption capacity at a lower flow rate and adsorbent loading. In the optimized conditions, a maximum of 1.31 and 0.49 mmol g−1 of H2S and SO2 was removed by the nanocomposite, respectively. The in-depth spectroscopic analysis confirmed the mineralization of H2S gas into sulfide, sulfur, and sulfite, which was mediated by the Fe and Mn redox cycles in the presence of adsorbed water and molecular oxygen. Though Zn ions did not participate in the oxidation process, Zn2+ probably interacted with the sulfides and sulfites. The SO2 mineralization was associated with the formation of sulfates, driven by the redox behaviour of Fe and Mn in an oxidative environment. Thus, we have presented a novel adsorbent material for the successful mineralization of toxic sulfurous gases, which could be suitable for deep desulfurization applications.