Simultaneous Removal of Ammonium and Nitrite in Aqueous Suspensions of Ferric Tannate Powder by Adsorption and Catalysis

Abstract

Ferric tannate (TA-Fe³⁺) has the potential to transform ammonium (NH₄⁺) and nitrite (NO₂⁻) into nitrogen gas (N₂) through adsorption and catalysis. Few reports have given detailed account of different adsorption behavior for NH₄⁺ and NO₂⁻, which is important to develop the potential of the material. TA-Fe³⁺ was synthesized here and its performance as an adsorbent/catalyst for the simultaneous removal of NH₄⁺ and NO₂⁻ from water was investigated. Results confirmed the adsorption and catalysis capability of TA-Fe³⁺ toward NH₄⁺ and NO₂⁻: (1) Following a 24 h adsorption and redox, the concentrations of NH₄⁺ and NO₂⁻ in the mixed solution decreased from initial values of 10.71 and 7.14 to 4.28 and 1.64 mmol/L, respectively, and 0.41 mmol/g of N₂ was produced with a maximal N₂ yield rate of 0.072 mmol/[g·h]. The N₂ yield was about 144 times that in the absence of TA-Fe³⁺; (2) intermediate products such as nitrous oxide (N₂O) and nitrate (NO₃⁻) were not detected; and (3) Raman spectrum analysis identified the site of Fe–O bond to be the center of adsorption and catalysis. Moreover, there were interesting findings: (1) TA-Fe³⁺ exhibited significantly distinct adsorption behavior toward NH₄⁺ and NO₂⁻. External diffusion and internal diffusion exerted a key influence on the adsorption toward NH₄⁺ and NO₂⁻, respectively; NH₄⁺ was adsorbed on TA-Fe³⁺ in the form of monolayer, and NO₂⁻ in the forms of both monolayer and multilayer; (2) TA-Fe³⁺ was easily regenerated with water; and (3) adsorption toward NO₂⁻ was the rate-determining step of the catalytic reaction. These findings will provide valuable enlightenment for the further work to reveal the adsorption and catalysis mechanisms of TA-Fe³⁺.

Introduction

Ammonium (NH₄⁺) and nitrite (NO₂⁻), two important pollutants in surface and ground water systems contaminated with nitrogenous compounds, can cause the eutrophication of receiving water systems and human health problems (Withers et al., 2014; Ye et al., 2018). Traditional biological (Daverey et al., 2015; Hasan et al., 2012; Kuypers et al., 2018) and physicochemical methods (Cheng et al., 2017; Du et al., 2017; Shin et al., 2014; Yao et al., 2019) for the removal of the two nitrogen species from wastewater require additional oxidizers (e.g., oxygen and chloride) or deoxidizers (e.g., formic acid, active metals, methanol), as well as energy sources (e.g., electrical, photochemical, thermal energy), which can lead to higher energy consumption and operating costs.

Anaerobic ammonium oxidation (Anammox) combined with partial nitrification, a novel biotechnology for nitrogen removal, can simultaneously remove NH₄⁺ and NO₂⁻ without additional oxidants and reductants (Liang and Liu, 2008). However, an extremely slow growth rate of anammox bacteria limits its wide practical application (Nozhevnikova et al., 2012).

New materials and technologies for nitrogen removal with high purification and energy efficiency have attracted more attention due to the current strain on energy sources and the stringent discharge regulations.

The redox reaction [Eq. (1)] can theoretically occur between NH₄⁺ and NO₂⁻ according to their standard redox potential (Table 1) (Fanning, 2000). Thus, ammonium and nitrite in solution can be simultaneously removed without any additional oxidant and reductant through the chemical reaction. However, the reaction requires high temperature and acidic condition for high reaction rate. Adsorption and catalysis are attractive technologies for wastewater treatment due to high removal efficiency (Klimenko et al., 2009; Rashid et al., 2021). If a material can not only adsorb NH₄⁺ and NO₂⁻, but also accelerate the reaction between them through catalysis under moderate conditions, this reaction would be developed as a promising alternative for nitrogen removal from wastewater. On one hand, NH₄⁺ and NO₂⁻ can be transferred from solution onto this material by adsorption, leading to the concentrated NH₄⁺ and NO₂⁻ on its surface. On the other hand, the adsorbed NH₄⁺ and NO₂⁻ can be converted to N₂ by catalysis, leading to the regeneration in situ of the adsorbent.

Table 1.

Standard Redox Potential

Couple	φ°_a (v)	φ°_b (v)
NO₂⁻/N₂	1.53	0.42
NH₄⁺/N₂	−0.27	0.74

φ°_a in acidic solution; φ°_b in basic solution.

NH₄⁺, ammonium; NO₂⁻, nitrite; N₂, nitrogen gas.

Furthermore, a novel nitrogen removal process can be formed through the combination of the catalytic redox reaction and biological partial nitrification: first, 50% ammonium in wastewater is converted to nitrite through biological partial nitrification; second, nitrite and the residual ammonium are catalyzed to produce N₂ through the Equation (1). Compared with the traditional aerobic nitrification and anoxic denitrification, the process can theoretically save 62.5% energy consumption and 100% supply of organic substrates. $N {H_{4}}^{+} + N {O_{2}}^{-} \to N_{2} + H_{2} O$ (1)

A titanium (IV) oxide photocatalyst was used to realize the simultaneous removal of nitrite and ammonia at room temperature (Hiroshi et al., 2014), but irradiation of UV light was needed. Ferric tannate (TA-Fe³⁺), a stable metal–polyphenol chelate, is water insoluble and cheaper than precious metals. This chelate has been found to exhibit a rough and porous surface with highly crosslinked network and to be capable of adsorbing both cation through its negative oxygen ions and anion through its ferric ions (Rahim et al., 2020; Xu et al., 2018). This material has been reported and applied in many fields, including antifouling, water distillation, and catalysis (Lin et al., 2020; Magro et al., 2020; Peres et al., 2017; Zhang et al., 2017).

However, research on the catalytic redox reaction in Equation (1) through TA-Fe³⁺ in wastewater treatment is limited. In this study, we prepared TA-Fe³⁺ using tannin and ferric chloride, and then evaluated the effectiveness of the TA-Fe³⁺ powder as an adsorbent/catalyst for the simultaneous removal of NH₄⁺ and NO₂⁻ under moderate temperature and neutral conditions. Moreover, by simultaneously monitoring nitrogen yield, the mechanism of adsorption and catalysis was elucidated.

The aim of this article is to aid in the development of new materials for the simultaneous removal of NH₄⁺ and NO₂⁻ from wastewater.

Materials and Methods

Materials

TA-Fe³⁺ raw material was prepared as follows: (1) 5% tannin solution (0.03 mol/L) reacted with ferric chloride solution (0.3 mol/L) for 1 h at room temperature (20°C), and the pH of the reaction mixture was adjusted to 7 through using 0.65 mol/L NaHCO₃ solution; (2) the black reaction product was washed with deionized water, separated centrifugally three times, and lyophilized in a freezer dryer at −50°C; and (3) the black TA-Fe³⁺ powder was stored in a vacuum desiccator.

After the adsorption and catalytic reaction, TA-Fe³⁺ was collected and dried in a freezer dryer at −50°C for characterization.

Stock solution of NH₄⁺ and NO₂⁻ was prepared with ammonium chloride (NH₄Cl) and sodium nitrite (NaNO₂), respectively.

Tannin, FeCl₃, NH₄Cl, and NaNO₂ were purchased from Sinopharm Chemical Reagent Co., China, with analytical reagent grade.

Experimental set-up

Adsorption tests

All adsorption and catalytic reaction experiments were performed as follows: TA-Fe³⁺ powder with predetermined weight was introduced into a 100-mL serum bottle. The aliquot deionized water, ammonium, and/or nitrite stock solution were added to obtain a mixture with a certain volume and initial concentration of NH₄⁺-N and NO₂⁻-N. To eliminate atmospheric nitrogen and oxygen in the bottle, the mixed solution was bubbled with highly purified helium for 5 min and the headspace was subsequently swept for 10 min before sealing with a rubber closure. The pH of the mixture was approximately equal to 7 and not adjusted with any acid or alkali. The serum bottle was then placed in a thermostatic shaker and agitated at 180 rpm and 30°C ± 1°C. The suspension was centrifuged and filtered with 0.45 μm membrane for chemical analysis.

The amount of NH₄⁺-N or NO₂⁻-N adsorbed by TA-Fe³⁺ powder at time t, q_t (mmol/g), was calculated as follows: $q_{t} = \frac{(C_{0} - C_{t}) V}{1, 000 m},$ (2)

where C₀ and C_t are the concentrations of the initial and final NH₄⁺-N or NO₂⁻-N in the solution at time t (mmol/L), respectively; V is the solution volume (mL); and m is the weight of the TA-Fe³⁺ powder (g).

The adsorption isotherms of NH₄⁺ and NO₂⁻ on TA-Fe³⁺ powder were determined using Langmuir [Eq. (3)] and Freundlich models [Eq. (4)], respectively. $q_{e} = \frac{q_{m} b C_{e}}{1 + b C_{e}},$ (3)

q_{e} = K C_{e}^{1 ∕ n},

(4)

where q_e is the equilibrium adsorption capacity (mmol N/g); q_m is the maximum adsorption capacity (mmol N/g); b (L/mmol N) is the equilibrium adsorption constant; and C_e is equilibrium liquid-phase concentration (mmol N/L); K is the equilibrium adsorption constant (L/g); and 1/n is the Freundlich intensity parameter.

Weber–Morris intraparticle diffusion model [Eq. (5)] was used to fit the kinetics data. $q_{t} = k_{p} t^{1 ∕ 2} + C,$ (5)

where q_t is the adsorption capacity (mmol N/g) at time t; k_p (mmol N/[g·h^1/2]) is the intraparticle diffusion rate constant; C is the particle diffusion equation constant.

Regeneration and reuse test

At the end of adsorption for 24 h, regeneration of adsorbent was finished as the following steps: first, the adsorbent (TA-Fe³⁺ powder) was centrifugally separated from suspension; and then the separated adsorbent was repeatedly washed and centrifugally filtered with 30 mL deionized water for three times; finally, washing water of total 90 mL in three times was fully mixed and then filtered with 0.45 μm membrane for chemical analysis.

Regeneration efficiency (RE) was calculated by using the Equation (6). $R E = \frac{C_{w} V_{w}}{(C_{0} - C_{t}) V} \times 100 %,$ (6)

where C₀, C_t, and V refer to the Equation (2), C_w is the concentration of NH₄⁺-N or NO₂⁻-N in the washing water (mmol/L), and V_w is the volume of the washing water (mL).

Four successive cycles of adsorption–desorption were conducted to investigate the reusability of TA-Fe³⁺ powder as the following steps: after adsorption toward NH₄⁺ or NO₂⁻ (the initial concentration of 14.29 mmol N/L) for 24 h (referring to Adsorption Tests section), TA-Fe³⁺ powder (1 g) was desorbed with water (referring to the above regeneration test) and then dried in a freezer dryer for the next cycle. The values of C_t/C₀ and q_t were used to evaluate the change of adsorption efficiency of TA-Fe³⁺.

Catalytic reaction

Experiment 1

The comparative tests (Table 2) were conducted to test the catalytic redox reaction and N₂ production under the coexistence of NH₄⁺ and NO₂⁻. Samples were taken at different time intervals to observe the variations of compounds in the gas and water with time during the catalytic reaction. To minimize the sampling error, water and gas from Runs 3 and 4 were sampled under the same reaction conditions, respectively. The bottles were sealed and 1 mL gas from the Run 4 bottle was immediately sampled and analyzed as a blank gas sample before the redox reaction.

Table 2.

Experiment Set-Up for Adsorption and/or Catalytic Nitrogen Removal

Run no.	Initial concentration (mmol N/L)		Catalyst (g)	V_total (mL)
Run no.	NH₄⁺	NO₂⁻	Catalyst (g)	V_total (mL)
1	10.71	0	0.43	40
2	0	7.14	0.43	40
3	10.71	7.14	0.43	40	Chemical analyses
4	10.71	7.14	0.43	40	Gas analyses

Experiment 2

To estimate the catalytic effectiveness of TA-Fe³⁺, a comparative test with/without TA-Fe³⁺ was performed under the same conditions. Following 24 h of adsorption and reaction, the catalyst was desorbed with deionized water. The total nitrogen (TN) was balanced according to chemical and gas analyses.

Experiment 3

Based on the different adsorption rate of TA-Fe³⁺ powder toward NO₂⁻ and NH₄⁺, effect of two-step adsorption on catalytic reaction rate was evaluated as follows: at the time of adsorption only toward NO₂⁻ for 12 h, NH₄⁺ was added in bottle. Hereafter, water and gas samples were taken out for analysis at 1, 4, 6, 8, and 12 h, respectively.

Analytic methods

TA-Fe³⁺ powder characterization

The surface physical morphology and major elements of TA-Fe³⁺ were observed using a scanning electron microscope (SEM), coupled with an energy-dispersive X-ray spectroscopy detector (S-3500N; Hitachi Co., Japan). BET surface area, pore diameter, and pore volume were determined by nitrogen adsorption/desorption, using an ASAP2000 surface analyzer (Micromeritics Co., USA). A Mastersizer 2000 (Malvern Co., USA) was employed to measure the particle size. Raman spectrum analysis was performed with a Raman analyzer (JY-HR800, France) under the following parameters: laser excitation line, 532 nm; laser output power, 3 mW; the laser power at the sample, 0.6 mW; and scan time, 100 s.

Chemical analyses

NH₄⁺, NO₂⁻, and NO₃⁻ concentrations were determined according to the standard methods issued by the Environmental Protection Agency (EPA) of China. Data were measured in millimole N per liter (mmol N/L). The pH value was determined using glass electrodes connected to a pH meter (320-S, Shanghai, China). N₂ and nitrous oxide (N₂O) analyses were performed through a Model HP6890 gas chromatograph with a thermal conductivity detector, equipped with an 8-ft molecule sieve column and a Porapak Q column, respectively. Helium was used as the carrier gas.

Results and Discussion

Characteristics of TA-Fe³⁺ powder

The prepared ferric tannate powder had a rough surface with a high specific surface area of 102.4 m²/g, a relatively homogeneous mesopore structure with average pore diameter of 56 Å and pore volume of 0.14 cm³/g, and an average particle size of 76 μm (Supplementary Fig. S1). These characteristics facilitated efficient adsorption. The raw ferric tannate mainly consisted of C, O, and Fe elements, and that the percentage content of C, O, and Fe were 38.84%, 43.73%, and 15.43%, respectively (Supplementary Fig. S2a).

Adsorption performance of TA-Fe³⁺ powder toward NH₄⁺ and NO₂⁻

Effects of contact time on the adsorption capability

Supplementary Figure S2b–d showed the appearance of nitrogen element and different percentage contents of N in different samples, suggesting the adsorption of NH₄⁺ or/and NO₂⁻ on the TA-Fe³⁺. The adsorption performances of ferric tannate powder toward NH₄⁺ and NO₂⁻ varied significantly with time (Fig. 1). In the Run 1, the NH₄⁺ concentration fell rapidly within half an hour from 14.29 to 3.88 mmol/L with removal efficiency of 72.8% and then gradually reached the minimum value of 2.68 mmol/L and maximum adsorption capacity of 0.58 mmol/g at 4 h. The NH₄⁺ concentration subsequently rose again and stabilized between 7.02 and 7.29 mmol/L after 12 h.

FIG. 1.

Adsorption rates of TA-Fe³⁺ toward NH₄⁺ and NO₂⁻. (volume of mixture = 50 mL; initial concentrations of both NH₄⁺ and NO₂⁻ are 14.29 mmol N/L; 1 g TA-Fe³⁺ powder; Run 1: only NH₄⁺ in solution; Run 2: only NO₂⁻ in solution; Run 3: NH₄⁺ and NO₂⁻ coexisting). NH₄⁺, ammonium; NO₂⁻, nitrite; TA-Fe³⁺, ferric tannate.

In comparison, the NO₂⁻ concentration in Run 2 exhibited a gradual reduction with time, with the adsorption capacity smoothly increasing to 0.5 mmol/g at t = 24 h. In Run 3, NH₄⁺ and NO₂⁻ were simultaneously adsorbed by TA-Fe³⁺, with adsorption trends following those of Runs 1 and 2, respectively. However, q(NH₄⁺) in Run 3 was less than that in Run 1 before 10 h, and subsequently enhanced. In contrast, the q(NO₂⁻) was consistently greater than that in Run 2.

To further characterize the adsorption process, data from Run 1 for NH₄⁺ and Run 2 for NO₂⁻ were fitted into the Weber–Morris intraparticle diffusion adsorption model (Arameh and Mousa, 2014; Weber and Morris, 1963). As shown in Supplementary Fig. S3 and Table 3, the intraparticle diffusion model showed a better fit to the experimental data, especially for NO₂⁻ with higher relation coefficient (R²) values. The deviation of the plots from the origin indicated that the intraparticle diffusion was not the only rate-controlling step for the adsorption of both NH₄⁺ and NO₂⁻. The adsorption processes for the two N species exhibited three stages: (1) Stage I before 0.5 h was the instantaneous adsorption, that is, external surface adsorption, which had the highest gradient (k_p) among the three stages (Table 3); (2) Stage II (0.5–12 h) was intraparticle diffusion process; and (3) Stage III (12–24 h) was the adsorption equilibrium stage.

Table 3.

Fitting Adsorption Parameters of Weber–Morris Intraparticle Diffusion Model

Adsorbate		NH₄⁺	NO₂⁻
Stage I	k_p (mmolN/[g·h^1/2])	0.1686	0.1977
	C	0.405	−0.0725
	R ²	0.9155	0.9985
Stage II	k_p (mmolN/[g·h^1/2])	0.0386	0.1200
	C	0.4977	−0.0466
	R ²	0.8871	0.9592
Stage III	k_p (mmolN/[g·h^1/2])	0.0116	0.0586
	C	0.3083	0.2000
	R ²	0.9272	0.9987

On the other hand, there were obvious differences between the two fitting curves: (1) As for NH₄⁺, almost highest adsorption capacity was acquired during Stage I among the three stages, inferring that adsorption toward NH₄⁺ was mainly owed to external diffusion.

As for NO₂⁻, the Stage II, that is, intraparticle diffusion, contributed more NO₂⁻ adsorbed on TA-Fe³⁺; (2) after a sharp decrease in q_t of NH₄⁺ during the later period of Stage II (4–12 h), there was a slight growth and adsorption equilibrium during the Stage III; and (3) the higher k_p and lower C for NO₂⁻ than those for NH₄⁺ (especially during the Stage II and Stage III), suggested that the intraparticle diffusion exerted more influence on the adsorption for NO₂⁻. The above observations implied that the adsorption process toward either NH₄⁺ or NO₂⁻ was controlled cooperatively by external particle diffusion and intraparticle diffusion, and that external diffusion and internal diffusion played an important role in the adsorption toward NH₄⁺ and NO₂⁻, respectively.

Regeneration and reusability of TA-Fe³⁺

As shown in Table 4, almost all adsorbed adsorbate was able to be released from the TA-Fe³⁺ powder either in Run 1 or Run 2, meaning that TA-Fe³⁺ could be easily regenerated with water. In contrast, after 24 h of adsorption in Run 3, the remaining adsorbates were obviously less than those in Run 1 and Run 2, indicating that more NH₄⁺ and NO₂⁻ were adsorbed from the solution onto the TA-Fe³⁺ powder than those in Runs 1 and 2 (Fig. 1). Moreover, the desorbed adsorbates from TA-Fe³⁺ were not equal to the adsorbed adsorbates, and thus the partial adsorbed adsorbates were not desorbed, leading to a lower RE than those in Runs 1 and 2. The total adsorption capacities increased upon the coexistence of NH₄⁺ and NO₂⁻, and not all of the adsorbed adsorbates were released from TA-Fe³⁺. These results suggested the occurrence of a reaction between the adsorbed NH₄⁺ and NO₂⁻ on the TA-Fe³⁺ powder.

Table 4.

Regeneration Results

Run no.	Initial (mmol N)		Final (mmol N)				Undesorbed adsorbate (mmol N)	RE (%)
	In solution		In solution		Desorbed adsorbate from TA-Fe³⁺
	NH₄⁺	NO₂⁻	NH₄⁺	NO₂⁻	NH₄⁺	NO₂⁻
1	0.714	0	0.360	0	0.336	0	0.018	95
2	0	0.714	0	0.216	0	0.501	−0.003	100
3	0.714	0.714	0.277	0.161	0.210	0.282	0.227 (NH₄⁺)0.271 (NO₂⁻)	48 (NH₄⁺)51 (NO₂⁻)

Volume of mixture = 50 mL; initial concentrations of both NH₄⁺ and NO₂⁻ are 14.28 mmol N/L; 1 g TA-Fe³⁺ powder; Run 1: only NH₄⁺; Run 2: only NO₂⁻; Run 3: NH₄⁺ and NO₂⁻ coexisting. Nitrate nitrogen could not be detected in all tests.

RE, regeneration efficiency.

Supplementary Figure S4 displayed the reusability of TA-Fe³⁺. As the number of regeneration cycles increased, there was a slight decrease in the adsorption capacity of TA-Fe³⁺ and a small increase in the concentration of adsorbates in solution. Despite this, adsorption capacity of TA-Fe³⁺ for NH₄⁺ and NO₂⁻ remained at about 90% and 89% of the original after the fourth regeneration, respectively, suggesting that the TA-Fe³⁺ powder can be reused to adsorb NH₄⁺ or NO₂⁻.

Adsorption isotherms of NH₄⁺ and NO₂⁻ on the TA-Fe³⁺ powder

The experimental data and theoretical data calculated by both Langmuir and Freundlich models are shown in Fig. 2, and the values of the adsorption isotherm parameters of the two models were listed in Table 5. It was found that the adsorption data of NH₄⁺ on the TA-Fe³⁺ powder were better fitted to the Langmuir model (R² = 0.9716) than to the Freundlich model (R² = 0.8388), and that the adsorption isotherms of NO₂⁻ demonstrated strong agreement with the two models with high correlation coefficients and n value more than 1 in the Freundlich model. These findings indicated that NH₄⁺ was adsorbed on the surface of TA-Fe³⁺ in the form of monolayer, and NO₂⁻ in the forms of both monolayer adsorption and multilayer (existence of a heterogeneous surface) (Sari et al., 2007).

FIG. 2.

Adsorption isotherms of NH₄⁺ and NO₂⁻ on TA-Fe³⁺ powder for 24 h contact time.

Table 5.

Adsorption Isotherm Parameters for NH₄ ⁺ and NO₂⁻ on TA-Fe³⁺

	Langmuir constants			Freundlich constants
	b (L/mmol)	q_m (mmol/g)	R ²	K (L/g)	n	R ²
NH₄⁺	0.137	0.732	0.9716	0.176	2.787	0.8388
NO₂⁻	0.075	1.110	0.9919	0.129	2.001	0.9315

Langmuir equilibrium adsorption constant b for NH₄⁺ exceeded that of NO₂⁻, and the adsorption equilibrium was achieved for C_e values close to 14 and 19 mmol/L for NH₄⁺ and NO₂⁻, respectively. More specifically, the adsorption equilibrium for NH₄⁺ could be realized with a lower initial concentration. These observations indicated the higher adsorption rate of TA-Fe³⁺ toward NH₄⁺ than toward to NO₂⁻. However, the theoretical maximum adsorption capacity (q_m) for NH₄⁺ was less than that for NO₂⁻ (Table 5). As mentioned in Effects of Contact Time on the Adsorption Capability section, despite the higher adsorption rate of NH₄⁺ than that of NO₂⁻ during the initial period of adsorption, the subsequent desorption of parts of adsorbed NH₄⁺ led to the decline in adsorption capacity of NH₄⁺, and the q_t of NO₂⁻ gradually rose with time and exceeded that of NH₄⁺ after 12 h (Fig. 1; Supplementary Fig. S3). Therefore, the q_m of NO₂⁻ surpassed that of NH₄⁺ through the 24-h adsorption equilibrium.

Adsorption capability of TA-Fe³⁺ powder toward NH₄⁺ and NO₂⁻

The aforementioned observations indicated that the ferric tannate adopted in this study could not adsorb anion (NO₂⁻) and cation (NH₄⁺) simultaneously, and was also desorbed by water with more than 95% desorption efficiency. According to the postulated structure of TA-Fe³⁺ by Ross and Francis (1978), three tannate ions react with one ferric ion to form a stable octahedral coordination compound (Supplementary Fig. S5a). In addition, as the tannin molecules are generally polymers of the basic flavonoid structure, each molecule can react with a number of ferric ions to form a network structure, where many phenolic hydroxyl groups combine with the ferric ions in the form of negative oxygen ions. In the network structure, those phenolic hydroxyl groups with negative oxygen ions on the external surface may serve as the negative function groups to absorb NH₄⁺, and Fe³⁺ in the interior as the positive function group to adsorb NO₂⁻ (Supplementary Fig. S5b), suggesting the electrostatic adsorption for NH₄⁺ and electrostatic attraction combined with coordination for NO₂⁻.

The network structure also implies that TA-Fe³⁺ is an amphoteric ion exchanger with a special structure, in which negative and positive function groups can attract each other, forming an inner salt band to neutralize part of the charges. This reduces the adsorption force between the groups and the ions with opposite charges in the solution, thus leading to its effective regeneration with water after adsorbing other ions in solution (Hatch et al., 1957).

However, TA-Fe³⁺ exhibited distinct behavior toward NH₄⁺ and NO₂⁻: (1) despite the initial rapid adsorption toward NH₄⁺, parts of the adsorbed NH₄⁺ were desorbed and released back to the solution, leading to a subsequent rise in the concentration and a fall in the adsorption capacity; and (2) despite the slow adsorption toward NO₂⁻, the adsorption capacity rose gradually.

The desorption of the partially adsorbed NH₄⁺ may be attributed to the following three aspects: (1) many NH₄⁺ rapidly were adsorbed on the external surface through electrostatic interaction, meaning the weak binding affinity and a short diffusion distance for the NH₄⁺ ion from the particle surface to solution; (2) the easy water regeneration property of TA-Fe³⁺; and (3) the NH₄⁺ concentration gradient between the solution and adsorbent. Actually, the adsorption and desorption occurred simultaneously, while their rates varied with the stage.

Initially, due to the greater NH₄⁺ concentration gradient between the exterior surface of particle and the solution, NH₄⁺ was quickly combined with the particle and adsorbed on the TA-Fe³⁺ through external diffusion, leading to a sharp rise in q_t and higher adsorption rate than desorption rate. At the interior diffusion stage, intraparticle diffusion started to slow down along with the decrease in the adsorbate concentration in the solution, as listed in the Table 3 (k_p of NH₄⁺ went down from 0.1686 to 0.0386). Moreover, the NH₄⁺ concentration on the TA-Fe³⁺ gradually exceeded that in the solution, reversing the NH₄⁺ concentration gradient between the solution and TA-Fe³⁺.

These factors induced the release of NH₄⁺ from the adsorbent to water and thereby the desorption rate surpassed the adsorption rate, resulting in the decline in the q_t of NH₄⁺. Upon gradually reaching the same point, the NH₄⁺ concentration in the solution and adsorption capacity were maintained at stable levels, resulting in the flattening of the curve during the adsorption equilibrium (Fig. 1 and Supplementary Fig. S3).

In comparison, ferric ion, standing in the interior of structure, was surrounded by the exterior network, indicating a longer diffusion distance for the NO₂⁻ ion from the solution to the adsorption site compared with the NH₄⁺ ion. Additionally, a larger ionic radius of NO₂⁻ (0.35 nm) compared with NH₄⁺ (0.148 nm) may lead to a slower diffusion rate of NO₂⁻ (Rowsell and Yaghi, 2004). Furthermore, three NH₄⁺ ions can be simultaneously adsorbed by three negative oxygen ions, while three NO₂⁻ ions can only be adsorbed by a single ferric ion (Supplementary Fig. S5b). These may be linked to the lower adsorption rate of the NO₂⁻ ion. Due to the slow adsorption rate toward NO₂⁻, the NO₂⁻ concentration in the solution exceeded that on the adsorbent during the whole adsorption process, and the concentration gradient diffused the NO₂⁻ from the solution to the exterior and then to the interior of particle, where it was adsorbed on the TA-Fe³⁺. Thus, the adsorption rate was higher than the desorption rate, bringing about a gradual increase in the q_t of NO₂⁻.

Catalytic redox capability of TA-Fe³⁺ powder toward NH₄⁺ and NO₂⁻

Catalytic redox reaction and N₂ production

The above results demonstrated that TA-Fe³⁺ was able to simultaneously adsorb NH₄⁺ and NO₂⁻, and that the partial NH₄⁺ and NO₂⁻ adsorbed on the TA-Fe³⁺ were not desorbed when they coexisted (Table 4). The comparative tests were performed to determine the existence of a reaction between NH₄⁺ and NO₂⁻ and N₂ yield (Table 2). N₂ production was only observed for the coexistence of NH₄⁺ and NO₂⁻ in Run 3 (Fig. 3).

FIG. 3.

Catalytic redox reaction and N₂ production. N₂, nitrogen gas.

N₂ production drastically increased from 0.03 to 0.27 mmol/g at 16 h, corresponding to an obvious decrease in C_t (NH₄⁺) and C_t (NO₂⁻) from 5.08 and 3.68 to 4.71 and 2.24 mmol/L in Run 3, respectively. Compared with Runs 1 and 2, q_t values in Run 3 increased at a greater extent, while the C_t values of either NH₄⁺ or NO₂⁻ were reduced, particularly for higher N₂ yields. When just NH₄⁺ or NO₂⁻ was present, almost all the adsorbed NH₄⁺ or NO₂⁻ could be desorbed from TA-Fe³⁺, and when NH₄⁺ and NO₂⁻ coexisted, part of the NH₄⁺ and NO₂⁻ disappeared from the solution and catalyst (Table 6). This implied that only adsorption occurred in Runs 1 and 2, while in Run 3, both adsorption and the reaction of the N₂ yield were present. Furthermore, the N₂ was produced from the adsorbates that had not desorbed. The N₂ production increased with the q_t of NH₄⁺ and NO₂⁻ and exhibited a marked growth at 16 h, indicating that significant production of N₂ took place when the q_t of NO₂⁻ gradually rose to the level close to the q_t of NH₄⁺. This demonstrated that redox reaction between NH₄⁺ and NO₂⁻ to form N₂ occurred when they coexisted with TA-Fe³⁺.

Table 6.

Results of Total Nitrogen Balance Tests

Run no.	Initial (mmol)			Final (mmol)						R-NH₄⁺ (mmol)	R-NO₂⁻ (mmol)
	In solution		TN	In solution		Desorption from TA-Fe³⁺		N₂	TN
	NH₄⁺	NO₂⁻	TN	NH₄⁺	NO₂⁻	NH₄⁺	NO₂⁻	N₂	TN
5	0.355	0.289	0.644	0.159	0.092	0.048	0.043	0.144	0.630	0.148	0.154
6	0.355	0.289	0.644	0.325	0.276			0.001	0.603	0.03	0.013

Volume of mixture of 40 mL; contact time of 24 h; initial concentrations of NH₄⁺ and NO₂⁻ are 8.88 and 7.24 mmol N/L, respectively; Run 5-with 0.43 g TA-Fe³⁺ powder; Run 6-without TA-Fe³⁺ powder. Analysis concentrations of N₂ in the Run 5 and Run 6 are 1.60 and 0.02 mmol/L, respectively, and mol mass of produced N₂ is from the product of the analysis concentration and volume in the headspace of bottle (90 mL). R-NH₄⁺ and R-NO₂⁻ represent removed NH₄⁺ and removed NO₂⁻, respectively.

TN, total nitrogen.

Table 6 compared the TN balance results between Run 5 with TA-Fe³⁺ and Run 6 without TA-Fe³⁺. In Run 5, 0.148 mmol NH₄⁺-N and 0.154 mmol NO₂⁻-N were removed from the initial 0.644 mmol TN, accompanied by a N₂ production of 0.144 mmol. Moreover, the ammonium and nitrite nitrogen removal and N₂ production yielded a mol ratio of 1:0.96:1.02, almost equal to the stoichiometric ratio given in Equation (1) (1:1:1). N₂O and nitrate nitrogen (NO₃⁻-N) were not detected, suggesting that all removed NH₄⁺ and NO₂⁻ was transformed to N₂. The 0.014 mmol TN from the initial 0.644 TN was attributed to the small amount of adsorbed NH₄⁺ that was not desorbed (Table 6).

In contrast, Run 6 yielded a minimal amount of N₂ (∼0.7% of that in Run 5) due to the limited removal of NH₄⁺ and NO₂⁻, that is, the N₂ yield in Run 5 was about 144 times that in Run 6. N₂O and NO₃⁻-N were also not observed in the products in Run 6. In the absence of the catalyst, high temperature or acidic condition is indispensable to the reaction in Equation (1) (Cox et al., 1994), and thus the N₂ formed under the moderate temperature and neutral conditions in Run 5 was certainly ascribed to TA-Fe³⁺.

Figure 4 presents the characteristic Raman bands in the 1,100–1,600 cm⁻¹ range corresponding to the phenolate ring vibrations of metal–phenolate complexes. The bands at 1,101, 1,351, and 1,436 cm⁻¹ were associated with C–O stretching and C–H bending modes, respectively, and those at 1,481 and 1,598 cm⁻¹ were attributed to the ring C–C stretching vibrations for the coordinated phenolate ligand (Chen et al., 2021; Li et al., 2019; Zou et al., 2013). The strong peaks at 587 cm⁻¹ in the spectra of Samples 2, 3, and 4 were obviously weaker than the spectrum of the raw TA-Fe³⁺ (Sample 1), while no other changes were observed.

FIG. 4.

Raman spectra in the 100–2,500 cm⁻¹ region. (1: raw TA-Fe³⁺; 2: TA-Fe³⁺ adsorbing NH₄⁺ only; 3: TA-Fe³⁺ adsorbing NO₂⁻ only; 4: TA-Fe³⁺ adsorbing both NH₄⁺ and NO₂⁻ to complete redox reaction.)

Studies of ferric complexes with phenol have identified that such complexes exhibit a characteristic strong absorption band in the 400–600 cm⁻¹ region due to a phenolate oxygen to Fe (III) charge transfer transition, and that the band at 587 cm⁻¹ can be assignable to ν (Fe–O) stretching vibration frequencies (Thakur et al., 2017; Varshney and Kumar, 2013). Despite the weakening of the peak intensity, the peak site of the Fe–O bond was not shifted, indicating that the bond was not destroyed, rather the adsorption toward NH₄⁺ or/and NO₂⁻ reduced the Raman absorptivity toward the Fe–O bond.

No differences were detected among the Samples 2, 3, and 4, and the N₂ peak at 2,330 cm⁻¹ did not appear in Sample 4 spectrum. These observations revealed the following: (1) either NH₄⁺ or NO₂⁻ attacked the same adsorption/reaction site in TA-Fe³⁺ in all waters, and as a result, the same characteristics and changes in the Fe–O bond site were presented in the three samples; and (2) for Sample 4, the produced N₂ quickly broke away from the TA-Fe³⁺, rather than accumulating on it. This agreed with the results in Table 6, whereby the produced N₂ was almost equal to the removed NH₄⁺ or NO₂⁻. Moreover, no other compounds were produced, thus leading to the same spectrum as those in Samples 2 and 3. These results revealed the Fe–O bond site to be the center of adsorption and catalysis, and while TA-Fe³⁺ acted as an adsorbent and catalyst in the system. The catalytic redox capability of the transition metal complex was realized through the cycle Mⁿ⁺↔M⁽ⁿ⁻¹⁾⁺ (Haber and Weiss, 1934; Rahmania et al., 2021; Tamami and Ghasemi, 2015). As for TA-Fe³⁺, the cycle of Fe³⁺↔Fe²⁺ was estimated to help the electron transfer between NH₄⁺ and NO₂⁻.

Serving as an electron transfer agent, ferric tannate transported the electrons between NH₄⁺ and NO₂⁻ to form a stable redox system with NH₄⁺ (electron donor) and NO₂⁻ (electron acceptor). Three N₂ production pathways were postulated in this cycle: (1) intermediates A₁ and B₁ from NH₄⁺ and NO₂⁻, respectively, interreacted to form N₂, (2) NH₄⁺ was gradually oxidized into nitrogen through the loss of electrons and (3) NO₂⁻ was gradually reduced to nitrogen by obtaining electrons. The cyclic transformation of Fe³⁺↔Fe²⁺ will be further characterized in the next work to reveal the catalytic mechanism.

Above observations suggested that: (1) in the absence of TA-Fe³⁺, the N₂ yield was negligible, even with the coexistence of NH₄⁺ and NO₂⁻; (2) with only NH₄⁺ or NO₂⁻ in the solution, adsorption occurred through just TA-Fe³⁺; and (3) when TA-Fe³⁺, NH₄⁺ and NO₂⁻ coexisted in the system, adsorption and redox reactions simultaneously occurred, and the reaction occurred on the TA-Fe³⁺, not in the solution.

Based on this, it could be concluded that: (1) The coexistence of NH₄⁺, NO₂⁻, and TA-Fe³⁺ was necessary for N₂ yield through the redox reaction between NH₄⁺ and NO₂⁻, where the three components functioned as the reducer, oxidizer, and adsorbent/catalyst, respectively; and (2) TA-Fe³⁺ initially adsorbed NH₄⁺ and NO₂⁻ from the solution and subsequently catalyzed the redox reaction between the adsorbed NH₄⁺ and NO₂⁻ to form N₂. The adsorbed NH₄⁺ and NO₂⁻ mutually reacted to produce N₂ and were subsequently removed from the solid phase (Supplementary Fig. S5c). Thus, more empty adsorption sites were available to continuously absorb more NH₄⁺ and NO₂⁻ ions, and hence the adsorption capacities toward either NH₄⁺ or NO₂⁻ exceeded those under the conditions of only NH₄⁺ (Run 1) or NO₂⁻ (Run 2).

Effect of two-step adsorption on the N₂ production

According to the two-step adsorption test (Fig. 5), TA-Fe³⁺ adsorbed NO₂⁻ in advance for 12 h, and q (NO₂⁻) was 0.33 mmol/g at 12 h, when NH₄⁺ was added in solution. One hour later, q (NH₄⁺) and q (NO₂⁻) were 0.41 and 0.35 mmol/g, respectively, and a small quantity of N₂ of 0.02 mmol/g was formed. At t = 18 h, that is, 6 h later after adding NH₄⁺, N₂ production largely increased to 0.15 mmol/g, which was advanced for 10 h compared with that in Fig. 3. The test indicated that, only when q (NO₂⁻) increased gradually to the level close to q (NH₄⁺), did the N₂ production rise remarkably.

FIG. 5.

Effect of two-step adsorption on reaction rate. (volume of mixture = 40 mL; initial concentrations of NO₂⁻ and NH₄⁺ are 7.14 and 10.71 mmol/L, respectively; 0.43 g TA-Fe³⁺ powder.)

Based on Fig. 3, the fit curve of N₂ production rate was acquired (Fig. 6a, b). The N₂ production with time was characteristic of “S”-type growth curve, which appeared to illustrate three N₂ yield stages: preparation period before 13 h, stepped-up growing period between 13 and 20 h, and succedent slowed-up growing period, leading to the corresponding rate changes and the maximal rate of 0.073 mmol N₂/(g·h) at 16 h in the second period. The change trend of N₂ production with q (NO₂⁻) was almost same as that with time (Fig. 6c, d), implying that N₂ production depended on q (NO₂⁻). Additionally, the value of produced N₂ gradually approached that of q (NO₂⁻) with time, and the differentials (ΔN) between q (NO₂⁻) and produced N₂ also presented three different periods in response to those in Fig. 6a: gradual increase from 0 to 0.29, decrease from 0.29 to 0.06, and relative stabilization in the range of 0.06–0.09. q (NO₂⁻) represented the amount of adsorbed NO₂⁻ from solution on TA-Fe³⁺, including two parts: left on TA-Fe³⁺ (N1) and changed into N₂ (N2), and therefore, ΔN was namely N1.

FIG. 6.

Effects of time and q (NO₂⁻) on N₂ production. The production rate (dc/dt) was from the derivative of the fit curve at a time. (a, b) change trend of N₂ production with time; (c, d) change trend of N₂ production with q (NO₂⁻).

In the first period, more adsorbed NO₂⁻ was accumulated on adsorbent, in the second period, more adsorbed NO₂⁻ was transformed into N₂ with time [about 88% of q (NO₂⁻) at 20 h], leading to less N1. More adsorbed NO₂⁻ was transformed, more NO₂⁻ should be transferred from solution on TA-Fe³⁺ synchronously, however, the fact was that N1 gradually reduced, implying that adsorption rate was less than N₂ yield rate.

Above results demonstrated that the three N₂ development stages were directly related to q (NO₂⁻): (1) preparation period was used to accumulate enough adsorbed NO₂⁻ for the redox reaction, (2) more adsorbed NO₂⁻ on TA-Fe³⁺ ensured higher reaction rate and more N₂ production in the second stage, and (3) finally, reaction slowed down with less and less NO₂⁻ left on TA-Fe³⁺. Although adsorption toward NO₂⁻ in advance shortened the preparation period as shown in Fig. 5, how to improve the adsorption toward NO₂⁻ at the same time of N₂ formation required other measures.

Conclusions

Under moderate temperature and neutral conditions, the redox reaction between NH₄⁺ and NO₂⁻ was realized to produce N₂ by adsorption and catalysis through TA-Fe³⁺. Results identified this redox reaction to be a promising chemical denitrification technology. Facile synthesis, easy regeneration, reusability, and inexpensiveness determined TA-Fe³⁺ to have the potential for practical application in nitrogen removal from water. This work also revealed the relatively lower adsorption capacities of TA-Fe³⁺ for both NH₄⁺ and NO₂⁻, and adsorption toward NO₂⁻ to be the rate-determining step of the catalytic reaction. To develop the engineering application of the material, future work will mainly focus on three aspects: (1) to improve the understanding of the adsorption and catalysis mechanisms; (2) to identify key parameters impacting adsorption and redox reaction; and (3) to find an effective solution to the problem of solid–liquid separation in this system.

Footnotes

Authors' Contributions

Z.L. was responsible for the conception of the research and did the experiments. Z.L. and M.Z. wrote the article. Z.X. and Y.Y. participated in the discussion and article revision.

Author Disclosure Statement

No competing financial interests exist.

Funding Information

This work was supported by Sichuan Science and Technology Program [Grant No. 2023JDR0277], Yibin Science and Technology Program [Grant No. 2022NY029], Research Foundation of Yibin University for Talented Scholars [Grant No. 2021QH024], and the Key Research and Development Program of Sichuan Province, China [Grant No. 2021YFS0289].

Supplementary Material

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Supplementary Material

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Simultaneous Removal of Ammonium and Nitrite in Aqueous Suspensions of Ferric Tannate Powder by Adsorption and Catalysis

Abstract

Introduction

Materials and Methods

Materials

Experimental set-up

Adsorption tests

Regeneration and reuse test

Catalytic reaction

Experiment 1

Experiment 2

Experiment 3

Analytic methods

TA-Fe3+ powder characterization

Chemical analyses

Results and Discussion

Characteristics of TA-Fe3+ powder

Adsorption performance of TA-Fe3+ powder toward NH4+ and NO2−

Effects of contact time on the adsorption capability

Regeneration and reusability of TA-Fe3+

Adsorption isotherms of NH4+ and NO2− on the TA-Fe3+ powder

Adsorption capability of TA-Fe3+ powder toward NH4+ and NO2−

Catalytic redox capability of TA-Fe3+ powder toward NH4+ and NO2−

Catalytic redox reaction and N2 production

Effect of two-step adsorption on the N2 production

Conclusions

Footnotes

Authors' Contributions

Author Disclosure Statement

Funding Information

Supplementary Material

References

Supplementary Material

TA-Fe³⁺ powder characterization

Characteristics of TA-Fe³⁺ powder

Adsorption performance of TA-Fe³⁺ powder toward NH₄⁺ and NO₂⁻

Regeneration and reusability of TA-Fe³⁺

Adsorption isotherms of NH₄⁺ and NO₂⁻ on the TA-Fe³⁺ powder

Adsorption capability of TA-Fe³⁺ powder toward NH₄⁺ and NO₂⁻

Catalytic redox capability of TA-Fe³⁺ powder toward NH₄⁺ and NO₂⁻

Catalytic redox reaction and N₂ production

Effect of two-step adsorption on the N₂ production