Introduction

Hydrogen sulfide (H2S) is a toxic malodorous gas originating from different natural and anthropogenic activities, including volcanic eruption, microbial degradation of vegetation, natural gas and oil industries, and sewage treatment facilities. H2S creates human discomfort even at a low concentration of 0.1 ppm. Long-term exposure to 50 ppm causes respiratory tract irritation, whereas a range of 500–1000 ppm could lead to death1, 2. H2S presence in various energy sources like natural gas, petroleum oil, and biogas is the major concern for industries as H2S corrodes transport pipelines and poison catalysts3. The combustion of H2S-rich fuel releases sulfur dioxide in the atmosphere, which eventually leads to the formation of aerosol sulfates and sulfuric acid by reacting with air components4. The presence of aerosol sulfates and sulfuric acid in the atmosphere adversely affects human health and the environment by acidifying oceans, lakes, and soil. Thus, mitigation strategies are required to remove H2S from energy gases.

Metal–organic frameworks (MOFs) are inorganic–organic hybrid materials formed by the self-assembly of metal ions and bridging organic ligands. Unique physicochemical properties like large surface area and pore volume, good structural stability, and ease of fabrication have popularized their use in various environmental applications5, 6. MOFs are being used as precursors for the fabrication of porous metal oxides with application-specific structural and functional characteristics7,8,9,10,11. MOFs transformation to metal oxides has provided an alternative pathway to develop metal oxides with newer morphologies, high surface area, and high porosity12.

Unlike oxides of Fe, Cu, and Zn, oxides and hydroxides of cobalt are less explored for room temperature desulfurization. Xue et al. screened several transition metal oxides for room temperature H2S removal. The study reported 6 mg g−1 of H2S uptake capacity for Co3O4, which increased to 134 mg g−1 for Zn3.5CoO4.9 mixed oxide13. Since the surface area of Co3O4 was undetermined, the low adsorption capacity was probably due to its low surface area. This inference was supported by the reported work of Pahalagedara et al. The study reported mesoporous Co3O4 with a surface area of 143 m2 g−1 and an adsorption capacity of 134 mg g−1 at room temperature14. Wang and coworkers integrated Co3O4 in three-dimensionally ordered macroporous silica for room temperature desulfurization process, where the adsorption capacity reached as high as 189 mg g−115. Besides Co3O4, CoOOH16 and Co(OH)217 integrated with graphite oxide has been successfully tested for the room temperature desulfurization process. Long et al. developed aerogels and xerogels of Na-MnOx and H-MnOx, which had the H2S adsorption capacity in the range of 17–680 mg g−118. Also. Some reports are available for the application of cobalt-based MOFs like ZIF-67 in the desulfurization process19. Though Co-MOFs and MOF-derived cobalt oxides have not been studied for H2S adsorptive/oxidative removal, Dong et al. have reported cataluminescence sensing of H2S using ZIF-67-derived porous Co3O4 dodecahedra.

The product of MOF pyrolysis/calcination depends on the temperature, metal ions, substrates, and rate of heating20, 21. While the calcination of monometallic MOFs yields single-phase oxides, the presence of a second metal in the MOF has unpredictable outcomes. Huang et al. reported the formation of CuCr2O4/CuO composite after the calcination of bimetallic Cr-embedded MOF-199 at 600 °C9. Yang et al. reported the formation of CoFe2O4 nanocrystals after the air calcination of bimetallic Co-Fe terephthalate MOF at 400 °C22. Lee and Kwak reported the formation of Mn-doped Fe2O3 after air calcination of the bimetallic FeMn-MOF21. Thus, the formation of metal oxides as composites, single-phase, or doped oxides is highly unpredictable.

In the literature, the role of sodium in MOFs is largely unexplored, with some reports on sodium metal–organic frameworks23. In the present study, the unexpected presence of sodium in Co-based MOFs has played a decisive role in the formation of MOF-derived oxides. The presence of Na in Co-MOF resulted in the formation of single-phase NaCoxOy as opposed to Co3O4/Na2O as the air calcination product. Apart from unique microsheets like morphology, the material showed exceptionally high H2S removal capacity in ambient conditions. The parameters associated with the column studies were optimized. Moreover, the H2S removal mechanism was studied in detail using different microscopic and spectroscopic analyses.

Materials and methods

Chemicals

Benzene-1,4-dicarboxylic acid (H2BDC), benzene-1,3,5-tricarboxylic acid (H3BTC), and cobalt(II) nitrate hexahydrate (Co(NO3)2·6H2O) were purchased from Sigma Aldrich. Ethanol, methanol, N,N-dimethyl formamide (DMF), and sodium hydroxide (NaOH) pellets were acquired from Samchun Pure Chemicals, Korea. H2S gas (500 ppm balanced with N2 gas) was procured from Union gas, Korea.

Synthesis of adsorbents

A 4.29 g of NaOH pellets were dissolved in 82.5 mL methanol, which served as the precipitating agent. The cobalt salt solution was prepared by dissolving 12.07 g of Co(NO3)2·6H2O in 75 mL of DMF. The NaOH solution was added to the cobalt salt solution to form hydroxide. Then, it was kept under ultrasonication, and H2BDC solution (6.64 g in 190 mL of DMF) was added to it. After 20 min of ultrasonication, the CoBDC MOF product was separated, washed with ethanol, and dried at 70 °C in a hot air oven. CoBTC MOF was synthesized using the same protocol with 8.50 g of H3BTC in 190 mL of DMF. Finely powdered MOF in an alumina crucible was calcined in air at 700 °C for 24 h. The CoBDC and CoBTC-derived oxide was labelled as NCO-D and NCO-T, respectively.

Material characterization

The surface morphology of materials was probed by field emission scanning electron microscopy (FE-SEM, Hitachi S-4300, Japan). A gold-platinum alloy was coated on dried samples using an E-1048 Hitachi ion sputter. The transmission electron microscopy (TEM) was conducted on a field emission TEM (FE-TEM, JEM-2010F, JEOL, Japan). Elemental mapping was done by energy-dispersive X-ray spectroscopy (EDAX, X-Maxn 80 T, Oxford, UK). N2 adsorption–desorption isotherms were recorded at − 196 °C over a Gemini series Micromeritics 2360 instrument and analyzed by the Brunauer–Emmett–Teller (BET) equation. Samples were pre-heated at 200 °C for 8 h for degassing purpose. XRD patterns were recorded on an X-ray diffractometer (Ultima IV Rigajku, Japan) with Cu Kα and a Ni filter. Fourier-transform infrared (FTIR) spectra were recorded on a Cary670 FTIR spectrometer after pelletization with KBr. For XPS analyses, a K-alpha XPS instrument (Thermo Scientific Inc., UK) with a monochromatic Al Kα X-ray source and 4.8 × 10−9 mbar of pressure was used. Spectra were charge-corrected to the main line of the carbon 1s spectrum (aromatic carbon) set to 284.7 eV. Spectra were analyzed using CasaXPS software (version 2.3.14).

Breakthrough studies

H2S breakthrough studies were carried out in a fixed bed micro-reactor at 25 °C. A known mass of an adsorbent packed between glass wool was supported on silica beads in a pyrex tube (height: 50 cm, diameter: 1 cm). The adsorbents were tested in dry and moist (by passing moist air through the adsorbent bed for 0.5 h with 0.3 L min−1 of flowrate) conditions. The concentration of the outgoing gas was measured by a multi-gas analyzer (GSR-310, Sensoronic, Korea) every 15 s until the effluent concentration reached 10 ppm (2% was the breakthrough condition). The adsorption capacity of an adsorbent (q, mg g−1) was calculated by integration of the area above the breakthrough curve.

$$q = \frac{{C_{0} Q}}{m}\int_{0}^{{t_{b} }} {\left( {1 - \frac{C}{{C_{0} }}} \right)dt}$$
(1)

where C0—initial concentration (500 ppm or 0.697 mg L−1), Q—flowrate, m—the mass of adsorbent (g), and tb—breakthrough time.

Results and discussion

Characterization of adsorbents

The SEM and TEM micrographs of Co-MOFs and derived oxides are shown in Fig. 1. CoBDC has a cluttered sheet-like morphology (Fig. 1a,e). CoBTC has distorted hexagonal microcrystals surrounded by nanothreads (Fig. 1b,f). The oxides formed by the calcination of MOFs have similar morphology. The NCO-D has smooth intercalated hexagonal sheets with some deposition of nanoparticles (Fig. 1c,g). The hexagonal microsheets and nanothreads in CoBTC were transformed to irregular hexagonal microsheets and polyhedral nanoparticles, respectively, in NCO-T (Fig. 1d,h). The TEM-EDS analyses of Co-MOFs and derived oxides are shown in Figs. S1 and S2. The EDS analysis confirmed the presence of Co, C, and O with an additional Na peak at ~ 1.0 eV for Co-MOFs (Fig. S1). For CoBTC, a metal-to-ligand ratio of 1:1 (opposed to the conventional 3:2) was adopted in the present study. The deficient Co ions for metal–ligand interactions in CoBTC was balanced by Na-ligand interactions24. The probable consumption of some of the cobalt hydroxide to form cobalt oxide was compensated by Na ions, which interacted with the carboxylate groups in a strong alkali medium25. The oxides derived from the calcination of Co-MOFs (NCO-D and NCO-T) (Fig. S2) have all peaks except for carbon. Based on EDS analysis, a compositional formula of NaCo0.7O2.4 and NaCo1.1O3.3 was assigned to NCO-D and NCO-T, respectively (Table S1). The excess oxygen could be from the hydroxyl density, adsorbed molecular oxygen, and mixed-valence states of cobalt ions.

Figure 1
figure 1

SEM images of (a) CoBDC; (b) CoBTC; (c) NCO-D; (d) NCO-T; TEM images of (e) CoBDC; (f) CoBTC; (g) NCO-D; (h) NCO-T.

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The surface and pore properties of MOFs and oxides were evaluated by N2 adsorption–desorption isotherms (Fig. 2a). The MOFs exhibited adsorption–desorption isotherms for mesoporous materials. The surface area of these MOFs (6.9–18.3 m2 g−1) was in agreement with the reported values25. The low surface area of MOFs was due to the formation of microparticles26. The MOF-derived binary metal oxides exhibited curves for macroporous or non-porous materials27. Metal oxides have a lower surface area than MOFs due to complete loss of mesoporosity after high-temperature calcination. The measured surface area of NCO-D and NCO-T was 1.15 and 1.90 m2 g−1, respectively (Table S2).

Figure 2
figure 2

(a) N2 adsorption–desorption isotherms (b) PXRD patterns; (c) FTIR spectra; (d) XPS surveys of Co-MOFs and derived oxides.

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The PXRD pattern of MOFs and derived oxides are shown in Fig. 2b. The PXRD pattern of CoBDC matched with the one reported by Ma et al. with a slight variation in the peak intensity at 28.5°28. Moreover, the peak at 9.0° has a minor split due to a slight distortion in the symmetry of MOF29. The PXRD pattern of CoBTC largely matched with the reported MOF by Nowacka et al. with an additional presence of diffraction peaks in the 5°–10° range30. Since MOFs were synthesized in a strong alkali medium, Na+ ions in the MOFs as the nodes had a strong impact on their PXRD pattern31. The calcination of these MOFs results in the formation of conventional oxides like Co3O432. On the other hand, mixed metal ions in a MOF yields binary metal oxides33. In the present study, (Na,Co)-based MOFs oxidized to yield single-phase NaCoxOy materials and not Co3O4/Na2O composite. The possible reason for single-phase NaCoxOy formation was a high calcination temperature of 700 °C and a long heating time of 24 h. In the literature, numerous reports are available on the fabrication of sodium cobalt oxides by the solid-state synthesis method with Na2O2 and Co3O4 as precursors34. During the calcination of (Na,Co)-MOFs, the oxide formation occurs through the initial growth of oxide nanoparticles on the MOF surface. These nanoparticles served as seeds for the development of microsheets35. In the case of the delocalized distribution of Na and Co in the MOF structure, both the metal cations took part in the formation of oxide to yield NaCoxOy type materials. On the contrary, localized distribution of Na and Co in the MOF may have formed Na2O/Na2O2 and Co3O4 nanoparticles, which served as the precursors for NaCoxOy type materials at a high temperature of 700 °C for 24 h. Thus, in both cases, NaCoxOy formation was possible. For this reason, the PXRD pattern of NaCoxOy matched with the pattern of NaCo2O436 with the absence of Na2O or Co3O4 (Fig. S3).

The FTIR spectra of MOFs and oxides are shown in Fig. 2c. The band at 3432 cm−1 was assigned to the stretching vibrations of O–H stretching vibrations of adsorbed water molecules. The high-intensity bands at 1579, 1387, and 1357 cm−1 were due to the asymmetric and symmetric O–C–O stretching of organic linkers. The band at 1502 cm−1 was for C=C stretchings of the aromatic skeleton. The mid-intensity bands at 825, 808, and 753 cm−1 were attributed to the C–H bending modes37, 38. The band at 457 and 510 cm−1 were due to the Co–O stretching39. For CoBTC, additional peak at 1668 and 1707 cm−1 were possibly due to the Na-bound carboxylate groups. In the FTIR spectra of NaCoxOy, the band at 1639 cm−1 was due to the bending mode of the adsorbed water molecules. The bands at 881 and 1442 cm−1 were ascribed to the asymmetric stretching Co–OH and Na–O vibrations, respectively. The band at 561 cm−1 was due to the Co–O stretching vibrations36, 40. The full scan XPS survey of MOFs and derived oxides confirmed the presence of Na in the materials along with C, O, and Co (Fig. 2d). The C peak in NaCoxOy materials was due to the adventitious carbon. Na 1s and Na KLL peak intensity increased in the NaCoxOy compared to the MOFs due to the loss of carbon after calcination.

The HRXPS spectra of CoBDC and CoBTC are shown in Fig. S4, and the curve-fitting parameters are listed in Tables S3–S5. The HRXPS C 1s spectrum of CoBDC has five contributions at 284.7, 285.8, 287.0, 288.6, and 290.6 eV for C=C/C–H, C–O, –COOCo, –COONa, and π-π* satellite, respectively (Fig. S4a)41. The HRXPS Co 2p spectrum of CoBDC has peaks at 782 and 798 eV for Co 2p3/2 and Co 2p1/2, respectively (Fig. S4b). The HRXPS Co 2p3/2 spectrum deconvoluted into two contributions at 781.1 and 782.7 eV for Co3+ (49.1%) and Co2+ (50.9%), respectively, with two satellites at 785.0 and 788.2 eV24. The HRXPS O 1s spectrum of CoBDC has four contributions at 531.5, 532.5, 533.7, and 535.4 eV for O–Co/O–Na, O–C=O, O–H, and H2O, respectively (Fig. S4c)42. For CoBTC, the HRXPS C 1s spectrum deconvoluted into four peaks at 284.7, 285.9, 288.3, and 290.0 eV for C=C/C–H, C–O, –COOCo, and –COONa, respectively (Fig. S4d). In the HRXPS Co 2p spectrum of CoBTC, the Co3+ and Co2+ peaks were observed at 781.0 and 782.5 eV with 48.7 and 51.3% contribution, respectively (Fig. S4e). The HRXPS O 1s spectrum of CoBTC has peaks at 531.5, 532.5, and 533.8 eV for O–Co/O–Na, O–C=O, and O–H, respectively (Fig. S4f).

The HRXPS spectra of NCO-D and NCO-T are shown in Fig. 3, and the curve-fitting parameters are listed in Tables S6 and S7. The HRXPS Na 1s spectra of NCO-D (Fig. 3a) and NCO-T (Fig. 3d) has a peak at 1070.6 and 1070.8 eV, respectively, for Na+ ions43. The HRXPS Co 2p spectrum of NCO-D has peaks at 779.8 and 794.9 eV with satellite peaks at 789.6 and 804.8 eV for Co 2p3/2 and Co 2p1/2, respectively (Fig. 3b). The Co 2p3/2 peak was deconvoluted into two contributions at 779.7 and 780.7 eV for Co3+ (36.4%) and Co2+ (63.6%) ions, respectively44. The HRXPS O 1s spectrum of NCO-D has three contributions at 530.1, 532.1, and 533.5 eV for O–Co/O–Na, O–H, and H2O, respectively (Fig. 3c)45. The HRXPS Co 2p spectrum of NCO-T has peaks at 779.7 and 781.1 eV for Co3+ (44.1%) and Co2+ (55.9%), respectively (Fig. 3e). The HRXPS O 1s spectrum of NCO-T has similar contributions as observed in NCO-D (Fig. 3f). Based on the analysis, NCO-D and NCO-T were NaCoIII0.25CoII0.45O1.33 and NaCoIII0.48CoII0.62O1.67, respectively.

Figure 3
figure 3

HRXPS (a) Na 1s; (b) Co 2p; (c) O 1s spectra of NCO-D; (d) Na 1s; (e) Co 2p; (f) O 1s spectra of NCO-T.

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H2S breakthrough studies

The breakthrough curves for MOFs and oxides in dry and moist conditions are shown in Fig. 4. CoBDC and CoBTC showed a low adsorption capacity of 1.6 and 5.7 mg g−1, respectively, in dry condition (Fig. 4a). The studies dealing with Co-based MOFs for H2S removal are absent in the literature. The closest study is the role of Co ions in UiO-67(bipy) for H2S removal. The post-synthetic inclusion of Co in the MOF could not substantially improve its H2S removal capacity. The Co-sites in a highly porous MOF with a surface area of ~ 2500 m2 g−1 failed to interact with H2S gas46. The low adsorption capacity of Co-MOFs in the present study was probably due to the low Co-density and poor accessibility of Co-sites in MOFs for H2S interaction. The NCO-D and NCO-T had an adsorption capacity of 133.9 and 134.6 mg g−1, respectively, which was a significant improvement compared to MOF precursors. In the moist condition, the adsorption capacity of MOFs and derived oxides decreased. In general, the presence of moisture plays a positive role in the H2S adsorption process by dissociating H2S molecules in the water film37, 47. The adsorption capacity could decrease due to the competitive behaviour of water molecules for the adsorption sites48. Moreover, the formation of sulfuric acid could lower the structural integrity of the adsorbent and decrease its adsorption capacity. In the present case, moisture alone can destroy the adsorbent structure due to the hygroscopic nature of sodiated transition metal oxides49. Nevertheless, the H2S adsorption capacity was satisfactorily preserved even in the presence of moisture. The higher adsorption capacity of NCO-T was due to its comparatively higher surface area than NCO-D. The surface area of nonporous adsorbent is highly relevant in the adsorption of gases. Zheng et al. reported an increased CO2 adsorption capacity (1.02–2.83 cm3 g−1) in KNbWO6·H2O pyrochlore with an increase in the surface area (1.82–2.90 m2 g−1) after Sn2+ substitution50. Thus, the surface area plays a decisive role in the gas adsorption capacity of nonporous materials. The N2 adsorption–desorption isotherms of spent NCO-D and NCO-T are available in Fig. S5. The surface area of spent NCO-D and NCO-T was 0.94 and 1.16 m2 g−1, respectively. The deposition of elemental sulfur and sulfate species was responsible for the decreased surface area. The decreased surface area after H2S exposure further supported the fact that a low surface area restricts the gas diffusion process and limits the adsorption capacity.

Figure 4
figure 4

H2S breakthrough curves of MOFs and derived metal oxides in (a) dry; (b) moist conditions. [Adsorbent] = 0.2 g, flowrate = 0.3 L min−1.

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The effect of dosage and flowrate on the adsorption capacity was studied for NCO-D and NCO-T (Fig. 5). The adsorption capacity was negatively impacted by the increasing adsorbent mass (Fig. 5a,b). The adsorption capacity of 154.6 and 168.2 mg g−1 with a 0.2 g dosage reached 117.6 and 118.9 mg g−1 with 0.4 g dosage for NCO-D and NCO-T, respectively. The decreased adsorption capacity was due to the formation of dead zones in the bed with the increasing bed loading, which remained unused during the initial phase of the adsorption process51. The adsorption capacity decreased with the increasing H2S flow rate (Fig. 5c,d). The adsorption capacity of 154.6 and 168.2 mg g−1 (0.1 L min−1) for NCO-D and NCO-T dropped to 134.8 and 133.9 mg g−1 (0.3 L min−1), respectively. The decrease in the adsorption capacity with the increasing flow rate was due to the insufficient contact time for adsorbate-adsorbent interactions at a higher flow rate. The impact was stronger due to the low surface area and porosity of the metal oxide adsorbents52.

Figure 5
figure 5

Effect of dosage on the adsorption capacity of (a) NCO-D; (b) NCO-T, flowrate = 0.1 L min−1; effect of flowrate on the adsorption capacity of (c) NCO-D; (d) NCO-T, [Adsorbent] = 0.2 g.

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The adsorption capacity of synthesized oxides was compared with the Co-based adsorbents reported in the literature (Table 1). The H2S uptake capacity of NCO-D and NCO-T were superior to many of the reported adsorbents. The higher adsorption capacity of the Co3O4–SiO2 composite was due to its high porosity. Only Zn-Co hydroxide had a higher capacity than NCO adsorbents. Nevertheless, the MOF-derived NaCoxOy adsorbents reported in these studies are unique and highly effective in removing H2S from effluent gases at room temperature.

Table 1 The adsorption capacity of cobalt-based adsorbents at room temperature.
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Adsorption mechanism

The distribution of sulfur in the spent NCO-D was probed through TEM-EDS analysis (Fig. 6a). The EDS map has peaks for Co, O, and Na with a new high-intensity peak for S at ~ 2.1 keV. The 2D elemental mapping confirmed the wide distribution of Na, Co, and O in the NCO-D adsorbent. Apart from the constituent elements, uniform distribution of S was observed in the spent adsorbent. The PXRD patterns of fresh and spent NCO-D are shown in Fig. 6b. The major peaks observed at 16.2°, 36.6°, 40.2°, 49.8°, 65.9°, and 68.6° for fresh NCO-D either disappeared or remained with decreased intensity after the H2S exposure. Wang et al. have studied the H2S oxidation process over Co3O4. Though the mechanism was probed by XPS analysis, the study suggested at least five products (CoS. CoSO4, CoSOH, CoOOH, and elemental sulfur)15. Park et al. reported the catalytic oxidation of H2S at 350 °C over Co3O4, where the XRD pattern confirmed the formation of CoSO453. Jun et al. reported the formation of Co9S8 after high-temperature oxidation of H2S over ZnCoTiO454. Pahalagedara et al. reported the formation of Co3S4 after the H2S desulfurization process over mesoporous Co3O4 at 200 °C14. The PXRD pattern of spent NCO-D has sharp peaks in the entire 15°–70° range. These peaks were assigned to Co3S4, CoSO4, Co3O4, and Co(OH)2 phases. The XPS surveys of fresh and spent NCO-D are shown in Fig. 6c. The XPS survey of spent NCO-D has peaks for S 2p and S 2s at ~ 165 and ~ 234 eV, respectively. The S present in the spent sample accounted for 14.2% of the total atomic composition.

Figure 6
figure 6

(a) TEM-EDS analysis; (b) PXRD patterns; (c) XPS surveys of spent NCO-D in dry condition.

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The HRXPS Na 1s peak at 1070.6 eV for fresh NCO-D shifted to 1071.1 eV for spent NCO-D (Fig. 7a). The shift in the binding energy of Na 1s peak by 0.3 eV was probably due to the redistribution of electron density after the formation of different sulfide and oxide phases of cobalt. In the HRXPS Co 2p spectrum of spent NCO-D, the peaks at 780.3 and 796.7 eV were assigned to the Co 2p3/2 and Co 2p1/2, respectively (Fig. 7b). The HRXPS Co 2p3/2 peak has two contributions at 779.7 and 781.2 eV for Co3+ (39.9%) and Co2+ (60.7%) sites, respectively. In the fresh NCO-D, the Co3+ and Co2+ contribution was 36.4 and 63.6%, respectively. The variation in the contributions of Co3+ and Co2+ sites in the NCO-D was due to the involvement of the Co3+/Co2+ catalytic cycle in the H2S oxidation process. The HRXPS O 1s spectrum of spent NCO-D has two contributions at 530.0 and 531.0 eV for O–Na/O–Co and O–S/O–H, respectively (Fig. 7c). The increase in the hydroxyl density over the adsorbent was most likely due to the interaction of H2S molecules with the surface lattice O2− to yield HS– and –OH groups55. Moreover, the absence of surface H2O in the spent sample hinted towards its full utilization to form sulfates and Co(OH)2. The HRXPS S 2p spectra of spent NCO-D is shown in Fig. 7d. The spectrum was fitted into six peaks for different sulfur species. The peaks at 162.0 and 163.2 eV were assigned to the S 2p3/2 and S 2p1/2 of sulfide species, respectively, which were bound to Co ions in Co3S415, 56. The sulfide species was primarily formed due to the reactive interaction of adsorbed H2S molecules with the lattice oxygen. The peaks at 164.1 and 165.3 eV were attributed to the S 2p3/2 and S 2p1/2 of elemental sulfur, respectively18. The peaks at 168.1 and 169.2 eV were assigned to the S 2p3/2 and S 2p1/2 of sulfate ions, respectively. These sulfate ions were considered as CoSO4, observed in the PXRD pattern as well18, 37. Thus, all three major sulfur species were conclusively detected in the XPS analysis. Moreover, the sulfide, sulfur, and sulfate contribution in the total sulfur content were 39.1, 11.1, and 49.8%, respectively. Based on the above discussion, the following reactions have been proposed for the H2S removal over NCO-D.

$$lattice\;(O^{2 - } ) + H_{2} S \to - OH + HS^{ - }$$
(2)
$$2Co^{3 + } + S^{2 - } \to 2Co^{2 + } + S$$
(3)
$$2Co^{2 + } + \frac{1}{2}O_{2} \to 2Co^{3 + } + O^{2 - }$$
(4)
$$S + \frac{1}{2}O_{2} + H_{2} O \to H_{2} SO_{4}$$
(5)
$$NaCo_{x} O_{y} + H_{2} O \to NaOH + Co_{3} O_{4} + CoO$$
(6)
$$Co_{3} O_{4} + 4H_{2} O \to 3Co(OH)_{2}$$
(7)
$$CoO + H_{2} SO_{4} \to CoSO_{4} + H_{2} O$$
(8)
Figure 7
figure 7

HRXPS (a) Na 1s; (b) Co 2p; (c) O 1s; (d) S 2p spectra of spent NCO-D in dry condition.

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Conclusion

We have reported a novel approach for the fabrication of NaCoxOy adsorbents by air calcination of (Na,Co)-organic frameworks. NaCoxOy were formed irrespective of the type of organic linkers used in the MOF precursor. Moreover, the oxides crystallized as microsheets of 100–200 nm thickness with the presence of some polyhedral nanocrystals. These macroporous oxides have a surface area in the range of 1.15–1.90 m2 g−1. X-ray photoelectron spectroscopy (XPS) analysis confirmed the near equal presence of Co2+ and Co3+ sites in MOFs, which were largely preserved in the NaCoxOy. The maximum adsorption capacity of 168.2 mg g−1 was recorded for NCO-T in dry conditions. The competitive nature of water molecules led to the decrease in adsorption capacity in moist condition. The adsorption capacity decreased with the increasing flow rate and bed loading due to the insufficient contact time for adsorbate-adsorbent interactions and the formation of dead zones, respectively. TEM-EDAX analysis confirmed abundant and uniform distribution of sulfur in the adsorbent. PXRD analysis of the spent sample suggested the formation of Co3S4, CoSO4, Co3O4, and Co(OH)2 after the H2S exposure. The products of the H2S adsorption-oxidation process were further confirmed by XPS analysis. Thus, we have reported highly efficient adsorbents for the adsorptive-oxidative removal of H2S gas in ambient conditions.