Absorption of SO₂ from flue gas by aqueous fulvic acid solution

Abstract

A new regenerable flue gas desulfurization process was proposed, in which fulvic acid derived from biomass residues was used as an absorbent to absorb SO₂ from flue gas, based on acid–base buffering capacity. Experiments have been carried out to examine the absorption, desorption and reabsorption performance of fulvic acid solution in a lab-scale reactor. The results show fulvic acid solution (0.04 g mL⁻¹, pH 5.5) could excellently absorb SO₂ with a maximum absorption efficiency of 97.5% (298 K, 2200 ppm SO₂, 5% O₂, 0.14 m³ h⁻¹). In the process of SO₂ absorption, chemical absorption is the predominant mechanism. The SO₂-loaded solution is readily desorbed and regenerated under ambient pressure by heating at 343 K, and the regenerated fulvic acid solution still exhibits good absorption performance after seven absorption/desorption cycles. Trace metal ions binding to fulvic acid play a decisive role in the absorption process. Fulvic acid samples before and after absorbing SO₂ were well characterized by Fourier transform infrared spectroscopy, near-edge X-ray absorption fine structure and X-ray photoelectron spectroscopy. These results demonstrate that no chemical change is found, except that carboxylate groups are protonated to carboxylic groups, indicating that fulvic acid is stable as a regenerable absorbent.

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Introduction

The emissions of SO₂ from the burning of fossil fuels have caused significant environmental and human health effects.¹ To control SO₂ emissions, the once-through non-regenerative wet flue gas desulfurization (WFGD) processes, based mainly on limestone scrubbing, are frequently used.² Although these processes have a number of advantages, they also have some serious disadvantages such as the high capital and operating expenses, the large volume of water required, poor quality of the byproduct, and even causing secondary pollution.³ Hence, an environmentally friendly and cost-effective technology for removing SO₂ from flue gas needs to be developed.

Recently, researchers studied SO₂ absorption in room-temperature ionic liquids (ILs), such as 1,1,3,3-tetramethylguanidinium lactate ([TMG]L),¹ TMG-based ILs,⁴ hydroxyl ammonium ILs,⁵ and imidazole-based ILs.⁶ In 2010, Heldebrant et al. reported an emerging technology that can capture SO₂ and regenerate it for reuse by the reversible zwitterionic SO₂-binding organic liquid produced from the reaction of SO₂ with N,N-dibutylundecanolamine.⁷ Although these technologies can effectively absorb SO₂, the high production costs would limit their industrial applications in large-scale flue gas desulfurization. Humic substances have been widely utilized as environmentally friendly organic fertilizers because of their molecular structure.⁸ With the aim of a resourceful type of WFGD technology, Hu et al. proposed humate for the removal of SO₂ from flue gas and the production of compound humic acid fertilizer,⁹ investigated the absorption of SO₂ with sodium humate solution and elucidated the forming mechanism of desulfurization byproduction.¹⁰ Fulvic acid (FA) is the fraction of humic substances that is soluble under all pH conditions. One of the most important properties of FA is the high buffering capacity in a wide pH range, which essentially arises from the dissociation of acidic functional groups (e.g. carboxylic and phenolic groups, Fig. 1).¹¹ FA can be extracted from sources such as peat, soil, water, lignite, agricultural and forest biomass residues etc.¹² Now, FA extracted from biomass residues has been industrialized, owing to accessible raw materials and low cost. Although FA has good buffering capacity, to our knowledge there have been no reports related to removing SO₂ from flue gas by its absorption. Thus we propose a novel regenerable flue gas desulfurization process in which SO₂ is first absorbed by aqueous FA solution in a bubbling reactor and then desorbed from the loaded solution by heating; the regenerated FA solution is recovered and reused for absorbing SO₂ from flue gas. Meanwhile, the desorbed SO₂ can be recovered and used to produce sulfuric acid or liquid sulfur dioxide. Compared with conventional limestone wet scrubbing, the regenerative desulfurization process would neither require a large quantity of water nor cause the problem of disposal of large amounts of waste such as wastewater, CaSO₄ and waste carbon dioxide.

Fig. 1 Representative structure of fulvic acid as proposed by Schnitzer and Khan (1972)^11d.

The aim of the present study is to investigate experimentally the feasibility and durability of FA as a regenerative flue gas desulfurization absorbent by the performances of absorption/desorption cycles in a lab-scale reactor; to obtain a more detailed understanding of the absorption mechanisms of SO₂ with Fourier transform infrared spectroscopy (FTIR), near-edge X-ray absorption fine structure (NEXAFS) and X-ray photoelectron spectroscopy (XPS), etc., and demonstrate that FA is stable as an absorbent during SO₂ absorption. Our findings signify the possibility that conventional limestone wet scrubbing could be replaced by regenerative absorption in FA solution.

Experimental

Sample preparation and treatment

In this work, powdered FA (≥95 wt%), extracted from crop straw, was provided by Shanghai Tongwei Biological and Technology Co., Ltd. (Shanghai, China). The FA sample was further purified using a DAX-8 resin (Supelco, Bellefonte, USA) column technique described by Thurman and Malcolm.¹³ The purified FA was converted to the protonated form by quickly passing it through a column of Amberlite IR-120 (H⁺ form) cation-exchange resin (Acros Organics, USA), and then freeze-dried for desulfurization experiment and chemical analysis. Deionized water (18 MΩ, Milli-Q, Millipore) was used for all experimental procedures.

After the SO₂ absorption test, the SO₂-loaded solution was dried for approximately 72 h at 313 K under vacuum, and then the dried FA desulfurization products were obtained and provided for characterization analysis.

Absorption and desorption experiments

A schematic diagram of the experimental apparatus is shown in Fig. S1 (see ESI†). The experiments for SO₂ absorption were conducted in a lab-scale bubbling reactor at ambient pressure and room temperature. A simulated flue gas consisted of 2200 ppm SO₂, 5 vol% O₂, and the balance was N₂. The SO₂ (99.95%), O₂ (99.99%) and N₂ (99.99%) gases were supplied from cylinders. The total flow rate of the simulated flue gas was controlled by rotameters. The SO₂ absorption efficiency was defined in our previous literature.¹⁴ The FA solution (0.04 g mL⁻¹, initial pH of 5.5) transformed into loaded solution after absorbing SO₂ for some time; the SO₂-loaded solution was also called FA desulfurization liquid. The desorption of the SO₂-loaded solution was conducted in nitrogen gas by heating at 343 K under ambient pressure and FA solution was regenerated, from which the escaping steam–SO₂ mixture was passed through a condenser for condensing water vapour and recovering the water into the regenerated FA solution that would be used to reabsorb SO₂ gas. The desorption process was continued until the pH of the solution remained constant (pH 4.7). When the desorption was done, the first absorption/desorption cycle (Cycle 1) finished. The next reabsorption of SO₂ in the regenerated FA solution was carried out, and so on. In the present study, we investigated eight absorption/desorption cycles. Furthermore, the two absorption experiments in purified FA solution and deionized water were also performed, which would be helpful to discussion of the role of trace metal ions binding to FA in the SO₂ absorption. All absorption experiments were performed under the same conditions. In the absorption and desorption processes, the pH of the solution was monitored using a digital pH meter (PHB-5, Shanghai Weiye Instrument Plant, China).

Analysis methods

The changes in SO₂ concentration at the inlet and outlet of the reactor were monitored by a flue gas analyzer (Testo-350XL, Germany). The S-containing ion (SO₄²⁻ and SO₃²⁻) was measured by modular ion chromatograph (Metrohm, Switzerland). Both FA and FA desulfurization products were analyzed for their elemental compositions. The ultimate analysis was performed on a CNS-Analyzer (Vario EL III, Elementar, Germany). The inorganic elements concentrations were analyzed by inductively coupled plasma atomic emission spectrometry (ICP-AES, IRIS Advantage 1000, THERMO, U.S.).

The functional groups of FA and FA desulfurization products were identified by a Fourier transform infrared spectrometer (Equinox 55, Germany Bruker), in the range 4000–400 cm⁻¹ with a resolution better than 1 cm⁻¹. To record FTIR spectra, the KBr pressed disk technique was used with 1 mg of sample dispersed in 200 mg KBr. The spectra were processed by the OPUS 5.5 software.

XPS was performed with a Kratos Axis Ultra DLD X-ray photoelectron spectrometer, employing a monochromated Al-Kα X-ray source (hν = 1486.6 eV) at 150 W (15 kV, 10 mA), hybrid (magnetic/electrostatic) optics, a hemispherical analyzer, a multichannel plate and delay line detector (DLD). The powdered FA and FA desulfurization products were pressed onto a stainless-steel sample-holder using double-sided scotch tape. Spectrometer pass energies of 160 eV for the survey spectra and 40 eV for high resolution spectra were used for all elemental spectral regions. As charge effects were expected in the study of these poorly conducting samples, the binding energy of C1s from contamination at 284.7 eV was used as an internal reference to calibrate the spectra. The CasaXPS software (version 2.3.14) was used for all quantitative analysis and peak-fitting routines. The C1s, O1s and S2p envelopes were fitted using mixed Gaussian–Lorentzian component profiles after subtraction of a Shirley background.

The oxygen K-edge NEXAFS spectroscopy characterizations were carried out at the soft X-ray beamline BL08U of Shanghai Synchrotron Radiation Facility (SSRF) with a 150 MeV linac, a 3.5 GeV booster, a 3.5 GeV storage ring, and seven beamlines. The FA solution and FA desulfurization liquid were uniformly deposited on Au-coated silicon wafers, and then dried by gentle heating with a hair dryer to shape thin films of samples. The film samples were mounted onto a sample bar, and then put into the beamline chamber. The chamber was evacuated to a pressure of less than 10⁻⁷ mbar. The spectra were recorded through adjusting the incident photon energy across the oxygen K-edge (520–560 eV) with a step of 0.5 eV. The spectra were acquired in the total electron yield mode (TEY). The TEY signals were normalized by the concurrent signal from the reference grid coated with clean gold.

Results and discussion

Absorption efficiency and reabsorption efficiency of SO₂

To verify the desulfurization performance of FA solution and understand the mechanisms of SO₂ absorption, we performed absorption experiments in FA solution, regenerated FA solution, purified FA solution and deionized water. Fig. 2 shows the SO₂ absorption efficiency with time. It can be seen that a high SO₂ absorption efficiency (95–97.5%) is maintained in the first 20.4 min for the FA solution (Fig. 2a). Then the SO₂ absorption efficiency starts to decline, when the pH of the FA solution is 3.1 (Fig. S2, see ESI†). It is finally reduced to 75% at 29.4 min, when the pH is 2.7 (Fig. S2, see ESI†) and the concentrations of SO₄²⁻ and SO₃²⁻ are 772.8 and 69.3 mg L⁻¹ (Table 1), respectively. Besides the concentrations of SO₄²⁻ and SO₃²⁻, Table 1 also presents the total SO₂ absorption and the molar ratio of S-containing ions to total SO₂ absorption. These results indicate that the FA solution exhibits excellent absorption performance and confirms the feasibility of FA as a SO₂ absorbent. It should be pointed out that our FA sample contains small amounts of SO₄²⁻ (98 mg L⁻¹). Furthermore, Table 2 shows the ratios of both sulfur and oxygen increase from 4.482 to 6.706% and from 49.386 to 51.957%, respectively, which indicate that FA has interacted with SO₂ in the absorption process.

Fig. 2 The SO₂ absorption efficiency with time for the FA solution of Cycle 1 (a), regenerated FA solution of Cycle 8 (b), (c) and purified FA solution (d). Experimental conditions: 100 mL solution, gas flow of 0.14 m³ h⁻¹, inlet SO₂ of 2200 ppm, inlet O₂ of 5 vol%, at 298 K and ambient pressure.

Table 1 The total SO₂ absorption and products of removing SO₂^a

Sample	Maximum absorption efficiency of SO2 (%)	DTHE (min)	Total SO2 absorption^b (mmol)	S-containing ion concentration after absorption (mg L^−1)		Molar ratio of S-containing ion to total SO2 absorption (%)
				SO3^2− ion	SO4^2− ion
a FA solution (100 mL, 0.04 g mL^−1, the initial pH 5.5); gas flow of 0.14 m^3 h^−1; inlet SO2 of 2200 ppm; 5 vol% O2; 298 K; ambient pressure; the same SO2 absorption time of 29.4 min for Cycles 1–8. b Total SO2 absorption is calculated by multiplying the integral of the absorption efficiency curve by inlet SO2 concentration and by gas flow.
Cycle 1	97.5	20.4	6.06	69.3	772.8	14.7
Cycle 2	97.3	16.2	6.05	75.2	876.3	16.6
Cycle 3	96.8	15.2	5.91	87.9	926.0	18.2
Cycle 4	97.5	16.4	5.90	104.2	948.5	19.0
Cycle 5	96.6	14.3	5.77	109.8	952.3	19.6
Cycle 6	96.6	12.7	5.51	124.6	974.8	21.3
Cycle 7	96.2	12.8	5.38	167.9	977.5	22.8
Cycle 8	96.7	11.6	5.02	223.1	1059.0	27.5

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Table 2 The elemental compositions of FA before and after absorption of SO₂ and purified FA (wt%)

Element	FA	FA desulfurization products^a	Purified FA
a Gas flow, 0.14 m^3 h^−1; inlet SO2, 2200 ppm; O2, 5 vol%; 298 K; ambient pressure; FA absorption solution of Cycle 1, 100 mL, 0.04 g mL^−1. b nd: not determined.
Ultimate analysis
Carbon	41.935	38.020	ndb^b
Nitrogen	0.1215	0.102	ndb^b
Sulfur	4.482	6.706	ndb^b
Oxygen	49.386	51.957	ndb^b
ICP–AES analysis
Calcium	3.657	ndb^b	0.110
Sodium	3.486	ndb^b	0.098
Magnesium	0.202	ndb^b	0.015

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By heating, 97% of the absorbed SO₂ was driven off from the loaded solution, and a regenerated FA solution was obtained. Fig. 2 shows the reabsorption performance of SO₂ in the regenerated FA solution of Cycle 8. For absorption curves of Cycles 2–7 see Fig. S3, ESI.† Although the FA solution has absorbed SO₂ gas seven times, the regenerated FA solution still exhibits good absorption performance with a Duration Time of High Efficiency (DTHE, the time of the SO₂ absorption efficiency above 95%) of 11.6 min and a maximum absorption efficiency of 96.7% (Table 1). The total SO₂ absorption of this regenerated solution is 83% of the FA solution of Cycle 1. The results of eight successive absorption/desorption cycles indicate that the interaction between FA and SO₂ is neither too strong nor too weak, and absorption and desorption can occur. Therefore, SO₂ absorption in FA solution is a viable method for flue gas desulfurization.

As can be seen from Table 1, the highest SO₂ absorption by 4 g FA is 6.06 mmol among Cycles 1–8, which translates to 0.10 g SO₂ g⁻¹ FA. This value is lower than results obtained by Wu et al. (0.31 g SO₂ g⁻¹ [TMG]L)¹ and Heldebrant et al. (0.54 g SO₂ g⁻¹ DBUA),⁷ but higher than Huang et al. (0.02 g SO₂ g⁻¹ [TMG][BF₄]]).^4a It should be noted that the absorption of 6.06 mmol is not the maximum SO₂ absorption capacity of FA. Because too low SO₂ absorption efficiency has no engineering value, we terminate the SO₂ absorption of Cycle 1 when the efficiency decreases to 75%. However, the FA absorption liquid still has SO₂ absorption capacity. So FA is an absorbent with high desulfurization capacity. Interestingly, the pH of all regenerated FA solutions is almost 2.7—the same as that of the FA solution at the end of SO₂ absorption. When the eighth absorption finishes, the concentrations of SO₄²⁻ and SO₃²⁻ are 1059.0 and 223.1 mg L⁻¹ (Table 1), respectively. As shown in Fig. 2, this regenerated FA solution of Cycle 8 has a shorter DTHE than the FA solution of Cycle 1. In other words, the SO₂ absorption capacity of the regenerated FA solution is weaker than the FA. This is because a certain amount of sulfate radical generated during SO₂ absorption causes an important effect on SO₂ reabsorption.

When SO₂ is dissolved in FA buffer solution (eqn (1)), a SO₂–FA–H₂O system forms and the following chemical equilibrium should be considered:

SO₂(g) ↔ SO₂(aq) (1)

SO₂(aq) + H₂O(aq) ↔ H₂SO₃(aq) (2)

H₂SO₃(aq) ↔ HSO₃⁻(aq) + H⁺(aq) (3)

HSO₃⁻(aq) ↔ SO₃²⁻(aq) + H⁺(aq) (4)

FA⁻(aq) + H⁺(aq) ↔ FA(aq) (5)

where FA⁻ represents the FA anion. These are instantaneous and reversible reactions. The FA solution can be thermally regenerated by reversing the reactions (1)–(5). An important feature of the SO₂ absorption is that it dissociates instantaneously and reversibly to form H⁺ and HSO₃⁻ in an aqueous phase, and this dissociation reaction makes the absorption rate much higher than the physical absorption rate.¹⁵ However, reactions (3) and (4) are limited by their respective equilibria. Importantly, binding H⁺ can move the two proton-transfer reactions to the right and increase the solubility of SO₂ gas in solution (eqn (1)). This increase is preliminarily considered to result from the buffering effect of FA⁻ (eqn (5)), because FA can provide some proton-binding sites. It is inevitable that small amounts of metal ions are introduced into the industrialized production of FA. Our FA sample contains trace metal ions (Na⁺, Ca²⁺ and Mg²⁺) (Table 2). These cations can bind to FA and occupy proton-binding sites of FA, thus metal fulvates occur in the FA sample. In water, the occupied proton-binding sites can be discharged (eqn (6)).

nFA–M(aq) ↔ nFA⁻(aq) + Mⁿ⁺(aq) (6)

where FA–M and Mⁿ⁺ represent metal fulvates and metal cations (Na⁺, Ca²⁺ and Mg²⁺), respectively. As stated by Tipping, the most abundant cation-binding groups in FA are carboxylate, and next come the phenolate groups.¹⁶ The binding of metal cations is strongly influenced by pH, the first owing to the competition between protons and metal cations for fulvic binding sites, and the second due to cation hydrolysis.¹⁶ For our FA sample (pH 5.5), it possesses protonated phenolic groups (–OH) rather than deprotonated phenolate groups (–O⁻), because phenolic groups of FA can be completely protonated at pH values lower than 7.5.¹⁷ Consequently, our attention is focused on carboxylate cation-binding groups. When the concentration of H⁺ increases as a result of SO₂ absorption, H⁺ ions can compete effectively with metal ions for anionic binding sites on FA. Thus FA⁻ reacts with H⁺ to form FA, and the equilibrium of eqn (3), (4) and (5) shift to the right. Then the mass-transfer resistance of the liquid phase is lessened, and the SO₂ gas mass transfer is improved. Therefore, the SO₂ absorption capacity of water is increased. As oxygen is present in simulated flue gas, sulfate can form (eqn (7)), while the SO₄²⁻ formation rate is very slow in contrast with the hydrolysis of SO₂ itself (eqn (2) and (3)).¹⁸ So the ability of FA solution to absorb SO₂ does not depend on sulfate formation as a significant step in the absorption.

2SO₃²⁻(aq) + O₂(g) → 2SO₄²⁻(aq) (7)

By comparing the (c) and (d) curves in Fig. 2, we find that deionized water and purified FA solution (0.04 g mL⁻¹) have a similar poor SO₂ absorption performance with no DTHE, in which the absorption can be considered as a physical absorption process controlled by molecular diffusion. The findings suggest that trace metal ions binding to FA play a decisive role in the SO₂ absorption. In other words, SO₂ absorption depends on carboxylate cation-binding groups of FA.

Chemical analysis of FA absorbent before and after absorbing SO₂

To obtain some information about the chemical interaction between SO₂ and FA, we characterized the FA and FA desulfurization products by FTIR, XPS and NEXAFS. Fig. 3 shows the FTIR spectra of the FA, FA desulfurization products and purified FA. The interpretations of the spectra are based on numerous studies, which are summarized in Table S1 (see ESI†). In general, the spectra of the three samples are very similar and show the same pattern of bands. In spite of the similarities, the differences among the spectra are still obvious. The first important difference between the spectra of FA (Fig. 3a) and FA desulfurization products (Fig. 3b) is that the absorption band at 3405 cm⁻¹ broadens slightly in the latter. This finding suggests an increase in the content of –OH groups for FA after absorbing SO₂. A second difference is that the carboxylic C=O stretching (υ_C=O(COOH)) band at 1705 cm⁻¹ shifts to a slightly higher frequency (1717 cm⁻¹), and the absorption intensity of the υ_C=O(COOH) increases markedly in the latter (Fig. 3b). The modifications signify an increase in the content of carboxyl group. A third pronounced difference is that the asymmetric stretch υ_as(COO)⁻ band of the carboxylate functional groups overlapped by aromatic C=C (Table S1, ESI†) is shifted towards the υ_C=O(COOH) from 1600 to 1617 cm⁻¹ owing to SO₂ absorption; additionally, as compared to the FA spectra, the υ_as(COO)⁻ stretch of FA desulfurization products at 1617 cm⁻¹ becomes sharper with a slight intensity decrease, as well as a decrease in the intensity of the symmetric stretch υ_s(COO)⁻ at 1426 cm⁻¹, which suggests a decrease of the content of COO⁻ groups. This result is ascribed to the partial protonation of COO⁻ groups of the FA.

Fig. 3 FTIR spectra of FA (a), FA desulfurization products (b) and purified FA (c).

On the basis of the above findings, we conclude that the carboxylate groups of FA are protonated to carboxylic groups after SO₂ absorption. Because of one important acidic binding site of carboxylic hydroxyl groups, FA is capable of interacting with trace metal ions to form metal fulvates. However, at low pH, carboxylate groups (COO⁻) can be protonated to carboxylic groups (COOH). Importantly, with SO₂ absorption in FA solution, a large number of hydrogen ions are supplied by the dissociation of sulfurous acid (eqn (3) and (4)), which facilitates the protonation of COO⁻. To further corroborate the conclusion, we also recorded the spectra of the purified FA (Fig. 3c). Interestingly, a similar frequency shift of υ_C=O(COOH) for the purified FA is observed at 1710 cm⁻¹, where the absorption intensity also increases compared to the FA. Moreover, we also notice a simultaneous decrease and shift of υ_as(COO)⁻ at 1608 cm⁻¹ (Fig. 3c). This is because COO⁻ was transformed into COOH by proton-saturated resin in the purification of FA, which also triggered a strong loss of cations (see ICP-AES results in Table 2). Hence, the explanation for spectral changes of carboxylate groups and protonated carboxylic groups is well supported by the spectra of purified FA.

Finally, a fourth noticeable difference is that the 603 cm⁻¹ band of FA desulfurization products is more prominent than FA, and simultaneously the 1118 cm⁻¹ band broadens. A reasonable explanation is that more sulfates are generated, since bands at 603 and 1118 cm⁻¹ can be assigned to sulfate (SO₄²⁻).¹⁹ Note that the generation of the sulfate radical prevented a few carboxylate groups from being regenerated during desorption, which would explain why the regenerated FA solution has shorter DTHE than the FA through successive absorption/desorption cycles, as shown in Fig. 2. Unfortunately, we can not find evidence for the HSO₃⁻ and SO₃²⁻ ions, whose intense absorptions are centered at 1025 and 950 cm⁻¹,²⁰ respectively. To detect the HSO₃⁻, we adopt the new method of rhodamine-based fluorescent probe selective for the bisulfite anion in aqueous ethanol media, which is proposed by Yang et al.²¹ Fig. S4 (ESI†) shows changes in the absorption spectra of a rhodamine-based fluorescent probe upon mixing with FA solution and SO₂-loaded FA solution. We observe that the solution of the rhodamine-based probe with FA exhibits almost no absorption peak in the visible region (400–680 nm), as shown in Fig. S4 (see ESI†). However, an obvious peak at around 565 nm appears when mixing with the SO₂-loaded FA solution, especially with a SO₂-saturated FA solution. This confirms that HSO₃⁻ ions are generated during SO₂ absorption. For detailed results, see ESI.† With respect to the SO₃²⁻ ions, ion chromatography has provided evidence for their existence (Table 1).

As an ideal tool for identifying differences in surface chemistry, XPS was used to characterize the surface composition and chemical environment of FA before and after absorbing SO₂. Fig. S5 (ESI†) shows the survey spectra of FA and FA desulfurization products. The most prominent photoelectron peaks at 286 eV and 533 eV are assigned to C1s and O1s. Other peaks include the Na1s peak at 1072 eV, as well as Ca2s, Ca2p, S2s, S2p, Na2s and O2s. In addition, auger lines of O and Na are present. For the two samples, we find no noticeable difference in the relative intensity of peaks (Fig. S5, ESI†). However, Table S2 (ESI†) presents several changes in relative atomic concentrations, which are calculated using the intensity of an appropriate line and XPS cross sections. After SO₂ absorption, the surface atomic concentrations of both S and O increase from 2.0 to 2.6% and from 33.0 to 36.3%, respectively, which is consistent with the results of the elemental analysis (Table 2). This suggests that an interaction of FA and SO₂ occurs. In spite of detailed FTIR analysis, we do not completely understand the speciation transformation of sulfur from SO₂ and the changes of functional groups on FA, which would determine the nature of the interaction between SO₂ and FA. In addition to the compositional information, XPS also provides information on the chemical state of each element. Thus, we can investigate surface chemical changes that occur on the FA surface exposed to SO₂ by qualitative and quantitative analysis of the C1s, O1s and S2p XPS spectra.

Fig. 4 shows the deconvolution of high-resolution XPS spectra of C1s, O1s and S2p along with fit peak parameters and assignments, of both the FA and FA desulfurization products. In general, the two samples are very heterogeneous. The C1s spectra can be deconvoluted into five different component peaks: C–C/C–H (284.7 eV), C–O (286.3 eV), COO⁻ (288.2 eV), COOH (288.6 eV) and π–π* shake-up transitions in aromatic rings (291.7 or 292.0 eV), as shown in Fig. 4a and b. We assign COO⁻ (deprotonated carboxyl group) and COOH (protonated carboxyl group) on the basis of the results of Yoshida et al.²² In fact, the FTIR results discussed above also confirm both COO⁻ and COOH in the two samples. In Fig. 4a and b, the peak profiles representing carboxyl groups at 288–290 eV are obviously different, which means the component changes in carboxyl groups. Inspection of Fig. 4a and b also shows that the amount of deprotonated carboxyl carbon decreases (from 3.2 to 1.2%), while protonated carboxyl carbon increases (from 2.9 to 5.1%) after SO₂ absorption, indicating that carboxyl groups bonded to metal ions were cleaved and free carboxylic acid groups were formed. As expected, with the decrease of pH, there is a transition from COO⁻ to COOH on FA. However, the ratio of peak areas of C–C/C–H and C–O for the C1s peaks remains essentially constant. This reveals that these bonds are not directly involved in the interaction of FA with SO₂, which is in good agreement with the FTIR result that no changes in the C–H band (2940 and 2844 cm⁻¹), the C–O band (1330 cm⁻¹) and the C–C band (532 cm⁻¹) are observed in Fig. 3.

Fig. 4 High-resolution XPS spectra of the C1s, O1s and S2p regions for FA (a, c and e) and FA desulfurization products (b, d and f) with the corresponding fits. The fit parameters, assignments and relative areas of the peaks are given in the insert Tables. The S2p doublets were constrained to a separation energy of 1.18 eV with an area ratio of 2 : 1 and equivalent full-width-at-half-maximum (FWHM) for both components (S2p_3/2 and S2p_1/2).

The O1s spectra show the presence of three oxygen components: oxygen atoms from C=O in carboxyl or carbonyl groups at 531.7 eV, the one related to C–O in carboxyl or phenolic groups at 533.1 eV, oxygen atoms in water or chemisorbed oxygen species at 535.7 eV. Note that binding energies for O=S and O–S (sulfonic groups, sulfate and sulfite), appear in the same range as those for O=C and O–C, respectively.²³ From the insert tables, it can be observed the ratio of peak areas at 531.7 eV decreases significantly after SO₂ absorption, while the ratio at 533.1 eV increases obviously. Since oxygen atoms in the carboxyl group have both single and double bonds, we can hardly determine the contribution of COOH according to the variation of area ratio. However, we could tentatively attribute it to the introduction of sulfate and sulfite, which would be verified in the following O K-edge NEXAFS and S2p analysis.

Oxygen K-edge NEXAFS spectra are highly sensitive to the chemical nature around the oxygen atom and are capable of detecting some significant differences in the spectra. For comparison, the O K-edge spectra of FA and FA desulfurization products are shown in Fig. 5. The spectrum of FA exhibits a prominent pre-edge peak at 532 eV, which is assigned to the 1s → π∗(C=O) transition from a carbonyl oxygen atom of the carboxylic group,²⁴ a strong main-edge peak at 535.6 eV, which corresponds to the 1s → σ*(C–O) transition from the carboxylate (COO⁻) group,^24a and a weak post-edge peak at 537.4 eV, which is attributed to the transition from the sulfate oxygen O(1s) → S(3p,4s)–O(2p) anti-bonding orbitals.²⁵ After SO₂ absorption, a new weak shoulder arises at 530.7 eV; this small resonance can be assigned to the 1s → π∗ transitions of S=O bonds from the absorbed SO₂ molecules.^24b Importantly, we observe that the absorption feature of 1s → σ*(C–O) at 535.6 eV disappears and a new absorption feature of σ*(C=O, OH) at 538.2 eV appears,^24b which result from the fact that the carboxylate group was protonated to the carboxyl group. Meanwhile, the intensity of 1s → S(3p, 4s)–O(2p) at 537.4 eV increases, indicating more sulfates in the FA desulfurization products. Note that sulfite species might also contribute to the increase. Interestingly, the spectra of FA desulfurization products show both a red shift of 0.24 eV and a slight broadening. The results agree with the findings of Messer et al.^24a As previously discussed, the carboxylate group was protonated at low pH, resulting in the breaking of the degeneracy of the π*(COO) orbital, lifting the degeneracy of the two C=O bonds and creating a C–O bond.

Fig. 5 O K-edge NEXAFS of FA (a) and FA desulfurization products (b).

To determine the chemical state of the sulfur atom and to gain an insight into the nature of sulfur-containing reaction products following the interaction of SO₂ and FA, we carefully deconvoluted the S2p spectra and referred to the many sulfur standards published.^23a,26 The S2p_3/2 binding energies are used to elucidate the chemical state of sulfur. Fig. 4e shows two doublets of S2p, and there are two types of sulfur species in FA: one at low binding energy (BE) (168.0 eV), attributed to the sulfonic acid group (–SO₃H),^23a and another at a higher BE (169.1 eV), corresponding to sulfate.²⁶ Sulfate salts exist as inorganic impurities in the FA sample, which has been confirmed by our previous analysis by ion chromatography and FTIR. Likewise, the sulfonic acid group is detected at two bands (1043 and 654 cm⁻¹) in Fig. 3. A new doublet of S2p_3/2 (167.0 eV) and S2p_1/2 (168.2 eV) appears in Fig. 4f, which implies sulfite formation.²⁶ Apart from the aforementioned three types of sulfur species, we cannot detect any other sulfur species in the process of fitting the S2p spectra, as shown in Fig. 4f. A comparison of quantitative results in Fig. 4e and f clearly shows that the contribution of sulfite is much smaller than sulfate for FA desulfurization products, and the ratio of sulfate increases dramatically from 8.6 to 53.4%, which likely results from the partial oxidation of sulfite prior to analysis.²⁷ The results demonstrate that the absorbed SO₂ is transformed into sulfite and sulfate. Although the ratio of sulfonic acid groups dramatically reduces from 91.4 to 42.9%, this does not imply that some sulfonic acid groups are transformed into other forms of sulfur-containing groups.

Conclusions

Based on the above analysis, we draw a conclusion that the buffer of fulvate plays a decisive role in the process of SO₂ absorption. Because of the lower acidity of FA (pK_a1 = 3.8, carboxyl groups;²⁸ pK_a2 = 8.6, phenolic groups²⁹) compared to that of SO₂ (pK_a1 = 1.8, H₂SO₃³⁰), deprotonated carboxylate in FA can indirectly interact with SO₂. The absorption of SO₂ results in an increase in the H⁺ concentration, which triggers the protonation of COO⁻ to COOH. Namely, the Na⁺, Ca²⁺ and Mg²⁺ ions binding to the COO⁻ groups are replaced by H⁺. No other chemical changes are found for FA after SO₂ absorption in spite of detailed chemical characterization, which demonstrates that FA is stable as a regenerable desulfurization absorbent. The pH of the FA solution is below 5.5, so the bisulfite ions are the principal reactive sulfite species in solution,³¹ and the absorption mechanism of SO₂ can be described as follows:

R–COONa + SO₂ + H₂O → R–COOH + NaHSO₃ (8)

R–(COO)₂Ca + 2SO₂ + 2H₂O → R–(COOH)₂ + Ca(HSO₃)₂ (9)

where R–COOH, R–COONa and R–(COO)₂Ca are the structural formulae of FA, sodium fulvate and calcium fulvate, respectively. For the very low content of Mg²⁺ in FA, we ignore magnesium fulvate in this work. The buffering effect of FA⁻ can be thermally regenerated by reversing the reactions (8)–(9). Admittedly, the partial dissociation of bisulfite to sulfite (eqn (4)) and the oxidation of sulfite to sulfate (eqn (7)) also take place during SO₂ absorption, but both are relatively minor (see Table 1).

In this paper, we focused on the SO₂ absorption mechanisms by FA solution without considering the influences of other components in the flue gas, such as water vapour, CO₂, NO and the ash on the absorption, desorption and reabsorption processes. In further research work, we will investigate their influence in detail. The experimental conditions in this work are appropriate for a batch reactor, but the results can also be applied to a continuous reactor where the FA absorption solution can be circulated between the absorption tower and a desorber with the pumps.

Overall, the FA solution can efficiently absorb SO₂, the absorbed SO₂ can be reversibly desorbed by heating, and the FA solution can be reused. This novel desulfurization method with low costs and benign environmental effects is feasible and potentially an alternative to the WFGD processes based on limestone scrubbing.

Acknowledgements

We gratefully acknowledge financial support by the National Science Foundation of China (No. 50876062, No. 51176113) and the Ministry of Science and Technology of China (No. 2007AA05Z313). We thank the Instrumental Analysis Center of SJTU for FT-IR, ICP-AES and ion chromatography measurements. We also thank Shanghai Synchrotron Radiation Facility (SSRF) for providing beam time at the BL08U beamline.

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Footnote

† Electronic supplementary information (ESI) available: Experimental apparatus, pH curve of the FA solution and regenerated FA solution, SO₂ absorption efficiency curve of Cycles 1–8, and characterization data. See DOI: 10.1039/c2ra21536e

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