Reaction Mechanism for the Removal of NO_x by Wet Scrubbing Using Urea Solution: Determination of Main and Side Reaction Paths

Abstract

Secondary problems, such as the occurrence of side reactions and the accumulation of by-products, are a major challenge in the application of wet denitrification technology through urea solution. We revealed the formation mechanism of urea nitrate and clarified the main and side reaction paths and key intermediates of denitrification. Urea nitrate would be separated from urea absorption solution only when the concentration product of [urea], [H⁺] and [NO₃⁻] was greater than 0.87~1.22 mol³/L³. The effects of the urea concentration (5–20%) and reaction temperature (30–70 °C) on the denitrification efficiency could be ignored. Improving the oxidation degree of the flue gas promoted the removal of nitrogen oxides. The alkaline condition was beneficial to the dissolution process, while the acidic condition was beneficial to the reaction process. As a whole, the alkaline condition was the preferred process parameter. The research results could guide the optimization of process conditions in theory, improve the operation efficiency of the denitrification reactor and avoid the occurrence of side reactions.

1. Introduction

Denitration at low temperatures is a common demand area for air pollution control, such as in the catalyst industry and the nitric acid industry [1,2,3,4]. Due to the low temperature and high moisture content of the flue gas in such industries, the selective catalytic reduction (SCR) technology that is widely used in the power industry is difficult to apply in this field [5,6,7]. Wet denitration technologies are commonly adopted in this field, such as alkali absorption and urea absorption [8,9,10]. As urea is reducible and easy to cause to react with nitrite to generate nitrogen, a number of technologies and research reports related to denitrification using the urea absorption method have been investigated [11,12,13,14,15]. In particular, there have been many research reports on simultaneous desulfurization and denitrification using urea, which can even reach a higher level of desulfurization and denitrification at the same time [11,12]. Another key technology for the successful application of urea absorption denitrification is the use of a high gravity reactor. The high-gravity reactor is essentially composed of a rotor and a stator, rotating at high speed through a coaxial rotor. Under the action of shear, the gas phase is broken into a large number of tiny bubbles so as to achieve an efficient gas-liquid contact reaction [16]. The high-gravity reactor has good applications in the fields of separation enhancement, nanoparticle preparation and synthesis, and SO₂ removal [17,18]. A complete set of technology and equipment for nitrogen oxide removal and dust removal has been developed in China. This technology has achieved good results in the field of low-temperature denitration in the petrochemical industry and achieved the integration goal of high-efficiency denitrification, dust removal and desulfurization.

However, during the practical application of this technology, a series of secondary problems have emerged with the extension of the service time. The composition of calcined tail gas is unstable and contains catalyst dust because of the difference in the formula of hydrogenation catalyst. In addition, the denitrification process of tail gas is accompanied by side reactions such as urea hydrolysis, the formation of nitric acid and urea nitrate, resulting in a gradual reduction of the denitration capacity of urea absorption solution or even inability to use. At the same time, urea nitrate, as an energetic substance, has potential safety hazards. The successful solution to this problem would overcome an important problem that restricts environmental protection and up-to-standard production in the catalyst industry. It provides an important guarantee for the popularization and application of urea absorption denitrification technology in high-gravity beds.

Most of the research focuses on the adjustment of process parameters and the selection of oxidation additives. The research on the reaction mechanism of nitrogen oxide absorption by urea solution, especially in the occurrence of side reactions, is relatively lacking. For the mechanism of denitrification through the urea absorption method, on the main reaction path, a relatively consistent view has been formed on the key path, that nitrogen oxides are absorbed into nitrous acid, and then react with urea to generate N₂ and CO₂ [11,12,19]. The work of Lasalle et al. further gave the kinetic results of the reaction [19]. However, due to the limited application of this reaction and there being few studies of it, no detailed research into the mechanism of the side reaction, such as investigating the accumulation of nitrate or the impact of pH and nitrogen oxides on the side reaction, has been retrieved. In consideration of the lack of research in this field abroad, it is of great significance to study the side reaction of denitrification by the urea absorption method for the application of this technology.

Since urea nitrate has strong explosiveness and is also one of the precursors of explosive synthesis, it has a high risk [20,21,22]. Therefore, the formation of urea nitrate should be avoided under any circumstances. In order to avoid the production of urea nitrate in any case, such as the operation of the supergravity reactor and the disposal of the absorption liquid waste liquid, the formation path of urea nitrate was studied. Due to a large number of components in the supergravity reactor, it is necessary to eliminate them one by one when determining the conditions for the formation of urea nitrate. The workload is large, and it is difficult to efficiently determine the exact reaction mechanism. Starting from the crystal analysis of urea nitrate, we analyzed its crystal structure characteristics and preliminarily estimated the conditions for the formation of urea nitrate [23]. Whether urea nitrate was formed under different conditions was tested through experiments. Finally, the theoretical understanding of the formation of urea nitrate was formed, and the range of safe working process parameters to avoid the formation of urea nitrate was proposed. In addition, the following two issues are also studied in this paper: (1) does SO₂ promote or inhibit the denitrification reaction; (2) under what conditions are nitrate by-products easy to accumulate?

2. Results and Discussion

2.1. Determination of the Boundary Conditions for the Formation of Urea Nitrate

In order to investigate the effects of the concentration of urea, H⁺ and NO₃⁻ on the formation of urea nitrate, a series of experiments were designed based on the structural formula of urea nitrate. In five groups of experiments, precipitation was observed in the solution, as shown in Figure 1a. After adding a certain amount of nitric acid (HNO₃) into the urea solution, white crystals were formed. Combined with XRD pattern analysis, as shown in Figure 1b, the precipitated crystals were urea nitrate (JCPDS PDF #06-0332) [23].

Figure 2a shows the amount of HNO₃ required for precipitation under different urea concentrations. As the concentration of urea solution increased, the amount of HNO₃ required to produce urea nitrate precipitation decreased gradually. When the concentration of urea solution was 0.83 mol/L, 2.08 mol/L HNO₃ was needed to form precipitation. However, when the concentration of urea solution was increased to 3.33 mol/L, HNO₃ of 1.14 mol/L was added to precipitate in the solution. The following formula was obtained by linear fitting:

$$\left[{\mathrm{HNO}}_3\right]=-0.3461·\left[\mathrm{urea}\right]+2.2159$$

(1)

where [HNO₃] and [urea] are the concentrations of HNO₃ and urea solution, respectively. Based on this mathematical relationship, it is possible to predict the HNO₃ concentration required for the formation of urea nitrate precipitation under the conditions of different concentrations of urea solution, and then adjust the working conditions, such as the concentration of the absorption solution. Figure 2b shows the amount of HNO₃ required for the formation of urea nitrate precipitation under different NO₃⁻ concentrations. The amount of HNO₃ needed for urea nitrate precipitation reduced with the increase of NO₃⁻ concentration. Through data fitting, it can be seen that the two were in a linear relationship.

$$\left[{\mathrm{HNO}}_3\right]=-0.8019·\left[\mathrm{urea}\right]+2.5433$$

(2)

Each increase of 1 mol/L of NaNO₃ can reduce 1 mol/L of NO₃⁻ from other sources, that is, about 1 mol/L of HNO₃ can be reduced, so the slope was closed to −1. However, since HNO₃ also provided H⁺, the slope would be slightly larger than −1, and the fitting value was around −0.8.

Our findings lead us to conclude that the critical sedimentation condition of urea nitrate is closely related to the concentration of urea solution, H⁺ and NO₃⁻. The product of [urea], [H⁺] and [NO₃⁻] is shown in Figure 3. It can be found that the product of the three concentrations is approximately a constant. Taking into account the above observations and the crystal structure of urea nitrate, the formation of urea nitrate requires the carbonyl protonation of urea. The formation process of urea nitrate is displayed in the following Equations (4) and (5):

$$\mathrm{CO}{\left({\mathrm{NH}}_2\right)}_2\left(\mathrm{aq}\right)+{\mathrm{H}}^{+}\left(\mathrm{aq}\right)\leftrightarrow {\mathrm{C}}^{+}\left(\mathrm{OH}\right){\left({\mathrm{NH}}_2\right)}_2\left(\mathrm{aq}\right)$$

(3)

$${\mathrm{C}}^{+}\left(\mathrm{OH}\right){\left({\mathrm{NH}}_2\right)}_2\left(\mathrm{aq}\right)+{ \mathrm{NO}}_3^{-}\left(\mathrm{aq}\right)\leftrightarrow \mathrm{CO}\left({\mathrm{NH}}_2\right)·{\mathrm{HNO}}_3\left(\mathrm{s}\right)$$

(4)

where C⁺(OH)(NH₂)₂ is the product of carbonyl protonation to the hydroxyl of urea. In this process, the concentration of C⁺(OH)(NH₂)₂ is determined by the product of [urea] and [H⁺], which can be obtained from its acid-base equilibrium constant (Equation (6)).

$$\left[{\mathrm{C}}^{+}\left(\mathrm{OH}\right){\left({\mathrm{NH}}_2\right)}_2\right]=K_b·\left[\mathrm{CO}{\left({\mathrm{NH}}_2\right)}_2\right]·\left[{\mathrm{H}}^{+}\right]$$

(5)

where K_b is the base equilibrium constant of urea. the concentration product of C⁺(OH)(NH₂)₂ and NO₃⁻ could not be higher than the solubility product of urea nitrate ($K_{sp}^{\prime }$).

$$K_{sp}^{\prime }=\left[{\mathrm{C}}^{+}\left(\mathrm{OH}\right){\left({\mathrm{NH}}_2\right)}_2\right]·\left[{\mathrm{NO}}_3^{-}\right]$$

(6)

Substituting Equation (6) into Equation (7) to obtain Equation (8).

$$\left[\mathrm{CO}{\left({\mathrm{NH}}_2\right)}_2\right]·\left[{\mathrm{H}}^{+}\right]·\left[{\mathrm{NO}}_3^{-}\right]=\frac{K_{sp}^{\prime }}{K_b}=K_{sp}$$

(7)

Equation (8) also illustrates that the formation of urea nitrate is related to the product of [urea], [H⁺] and [NO₃⁻], coinciding with the observation from our experiments. These results indicate that only when the concentration product of the three solutes ([urea], [H⁺] and [NO₃⁻]) is high enough, it is possible to generate urea nitrate precipitation. The critical value of the concentration product, namely the solubility product $K_{sp}$, is about 0.87–1.22 mol³/L³.

In addition, we verified whether adding alkali could re-dissolve urea nitrate. As described earlier, the addition of concentrated HNO₃ to a concentrated urea solution resulted in urea nitrate precipitate. Whereas, when NaOH solution was added, the precipitate dissolved rapidly. This suggests that the formation of urea nitrate requires a certain pH, and that if the pH increases, the urea nitrate can be re-dissolved. Furthermore, when NaNO₃ was added to the concentrated urea aqueous solution, no precipitation was generated, regardless of the addition. It indicates that H⁺ is one of the necessary factors for the formation of urea nitrate. In other words, the formation of urea nitrate requires the simultaneous presence of high concentrations of urea, H⁺ and NO₃⁻. From the results we have obtained, one can conclude that, (1) if the formation of urea nitrate is observed in the supergravity urea absorption device, it means that the concentration of urea, H⁺ or NO₃⁻ is too high; (2) even if urea nitrate is formed, it can be re-dissolved as long as the pH of the absorption solution is increased; (3) the implication is that the formation of urea nitrate is not one of the main side reactions under normal working conditions.

2.2. Reaction Mechanism for NO_x Removal

In contrast to the alkali absorption of NO_x, the absorption path of NO_x using the urea reduction method is more complex, and no systematic study concerning the side reaction process of the urea reduction method has been published. The main and side reactions of the urea reduction method were systematically investigated. For the main reaction, it is necessary to confirm the flow transition paths of various nitrogen-containing substances in the chemical grid, so as to clarify the main reaction process. In particular, it’s important to confirm whether the optimal reaction range for the dissolution of NO_x and subsequent reaction with urea is in the same pH range. For the side reactions, gas-phase analysis and liquid-phase analysis were combined to track the gas-liquid two-phase changes in the process of the urea solution absorbing NO_x, thereby further exploring the side reaction mechanism, inhibiting or cutting off the side reaction paths, improving the efficiency of NO_x removal, and reducing the generation of by-products.

2.2.1. Concentration of Urea Solution

The concentration of urea solution is one of the influencing factors of the NO_x removal process [11], thus the influence of the urea concentration on the denitration process was investigated. Firstly, in order to eliminate the change in pH value of the absorption solution caused by the change of the urea concentration, the pH value of the urea solution with different concentrations was analyzed. As listed in Table S1, the pH value of the urea solution increased slightly from 7.06 to 7.58 with the increase of the urea concentration from 5 wt.% to 20 wt.%, showing a weak alkaline. This indicates that the urea concentration has no remarkable effect on the pH value of the absorption solution. Figure S1 shows the curve of NO, NO₂ and NO_x conversion rates over time with different urea concentrations. It can be seen that the reaction began to stabilize after 30 min, so each condition in the later experiment was stable for 60 min. As shown in Figure 4, the urea concentration exhibited no notable effect on NO_x removal efficiency when it was higher than 5 wt.%. The conversion of NO, NO₂ and NO_x was around 43%, 85% and 65%, respectively. This is because the urea concentration affects the removal of NO_x from both physical properties and chemical reactions. From the point of view of the chemical reaction, increasing the concentration of urea can accelerate the denitration reaction. However, in terms of physical properties, with the increase of the urea concentration, the viscosity of the urea solution increases, and the diffusion rate and solubility of NOx in the absorption solution decrease [9]. Herein, combining the actual operating conditions and experimental results, the 15 wt.% urea solution was employed in the latter.

2.2.2. Reaction Temperature

The diffusion, dissolution and reaction characteristics of the reactants or intermediate species in the urea absorption solution are closely related to the reaction temperature. Therefore, the denitration reaction was conducted at different temperatures. As shown in Figure 5, the NO conversion declined slightly, while the NO₂ conversion increased first and then dropped in the temperature range of 30–70 °C. Taken together, the NO_x conversion remained stable in the range of 30–50 °C, and it decreased at higher temperatures (50–70 °C). As the reaction temperature increases, on the one hand, the solubility of NO and NO₂ in the solution decreases, and the decomposition rate or the key intermediate HNO₂ accelerates, which are conducive to the absorption and removal of NO and NO₂. On the other hand, the NO oxidation rate increases, and the hydrolysis of the urea gradually strengthens, which are conducive to the absorption and removal of NO_x. Hence, the influence of reaction temperature on NO_x removal is the result of the comprehensive effects of diffusion, dissolution and reaction characteristics. When the reaction temperature exceeds 50 °C, various side reactions and negative effects dominate, leading to the reduction of NO_x removal efficiency.

2.2.3. Oxidation Degree of NO_x

Since NO is hardly soluble in water, the dissolution of NO_x in water can be divided into situations: one is the absorption of NO₂, and the other is the synergistic absorption of NO and NO₂ [24]. In the process of liquid-phase absorption of NO_x, NO_x in the gas phase must be dissolved in the absorption solution before they react with urea. Accordingly, the oxidation degree of NO_x shows a significant impact on the final absorption efficiency of NO_x. The influence of the oxidation degree of NO_x (NO₂/NO_x) on denitrification through the urea absorption method was investigated, and the results are shown in Figure 6. The NO_x removal efficiency improved with the increase of the NO_x oxidation degree in the gas phase, that is, the higher the NO₂ content was in the gas phase, the better the NO_x absorption efficiency was. With the oxidation degree of NO_x increasing from 30% to 95%, the NO₂ conversion and NO_x conversion increased gradually. Whereas, the NO conversion decreased, and when the oxidation degree of NO_x was higher than 50%, the NO conversion dropped markedly. Among them, when the oxidation degree of NO_x was 95%, the NO₂ conversion and NO_x conversion were as high as 97% and 84%, respectively, but the NO conversion dropped to a negative value of −162%. These results indicate that a great quantity of NO is generated in the denitrification reaction system of the urea solution. This is because, when the NO₂ in the gas phase is relatively excessive, a great deal of HNO₂ is unstable in the liquid phase, which can easily decompose and regenerate NO [25]. As shown in the reaction formula (9), it is generally considered that NO₂ and H₂O react to generate NO, and the essence of NO generation is the decomposition of HNO₂. From the results we have obtained, one can conclude that a high NO_x oxidation degree is conducive to the absorption and removal of NO_x using a urea solution.

$$3{\mathrm{NO}}_2+{\mathrm{H}}_2\mathrm{O}\to \mathrm{NO}+2{\mathrm{HNO}}_3$$

(8)

2.2.4. Initial pH of Urea Solution

The pH value of the absorption solution is one of the main factors affecting denitration efficiency. The pH value of a 15 wt.% urea solution is 7.54, which is alkalescent and close to neutral. HCl and NaOH were adopted to adjust the pH value of the absorption solution. On the one hand, it is an effect of the pH value on the gas dissolution process [11,12]. Nitrogen oxides are easy to dissolve under alkaline conditions. The higher the pH value is, the more conducive the solution is to the dissolution of NO₂ in NO_x, which promotes the dissolution of NO_x and the reaction rate. On the other hand, urea is easily activated under acidic conditions [19]. With the reduction of the pH value, the concentration of H⁺ elevates. H⁺ exhibits a considerable promoting effect on the hydrolysis of the urea solution. The hydrolysis of urea produces ammonium carbamate (NH₂COONH₄). NH₂COONH₄ is easier to make react with HNO₂, enhancing the reaction rate, thereby improving the removal efficiency. However, when the pH value of the absorption solution is too low, not only is the dissolution of NO_x is inhibited, but the decomposition rate of HNO₂ is also accelerated, which is not beneficial to the absorption and removal of NO_x.

Taking into account the complexity of the influence of pH on the denitration reaction through urea absorption, the denitration process was divided into two processes in the experimental design: the dissolution process and the reaction process. The complete absorption denitration process was expressed as (Dissolution + Reaction). In the Dissolution process, the simulated gas was introduced into the aqueous solution, and the absorption of the gas was monitored online via the flue gas analyzer. In the Reaction process, sodium nitrite (NaNO₂) was added to the urea solution, and N₂ was used as a carrier gas to analyze the gas components generated online. In the Dissolution + Reaction process, the simulated reaction gas was injected into the urea solution, and the composition changes in the gas phase were measured online. The removal efficiencies of NO, NO₂ and NO_x in the Dissolution process and the Dissolution + Reaction process at different pH values are displayed in Figure 7. As shown in Figure 7a, for the dissolution process, with the increase of the pH value of the aqueous solution, the dissolution efficiency of NO was significantly improved. The dissolution efficiency of NO₂ was also gradually enhanced, which was comprehensively reflected in the improvement of the NO_x dissolution efficiency. For the Dissolution + Reaction process, as displayed in Figure 7b, the conversions of NO and NO_x increased in the pH range of 0–12. The NO₂ conversion declined in the acidic range but accelerated in the alkaline, and pH = 7 was the turning point. The behavior of the NO, NO₂ and NO_x removal efficiencies makes us conclude that alkaline conditions are beneficial to gas-phase dissolution; nevertheless, acidic conditions are good for liquid-phase reaction.

The pH changes of the aqueous solution and urea solution before/after the Dissolution process and Dissolution + Reaction process are summarized in

Table 1. The ion concentration and pH value in the solution after absorption reaction.

Sample	NH4+ (mg/L)	NO2− (mg/L)	NO3− (mg/L)	pH
34%	5.80	12.35	1.13	6.29
42%	13.74	20.10	7.10	5.65
50%	11.10	4.14	--	5.82
70%	8.04	13.94	6.85	5.49
95%	10.38	10.92	6.25	6.09

and

Table 2. The pH value of the solution before/after Dissolution/Reaction process.

Sample	Dissolution		Dissolution + Reaction
	Fresh Solution	Used Solution	Fresh Solution	Used Solution
0	−0.38	−0.33	0.10	0.11
4	2.06	1.99	3.09	2.98
7	5.24	2.87	7.43	6.66
9	10.08	2.97	8.20	7.25
12	12.79	12.09	12.93	12.68

. After the dissolution and absorption denitrification reaction, the pH of the solution dropped to different degrees. Among them, the pH value of the aqueous solution in the range of 7–10 declined markedly after the dissolution process. This is mainly due to the formation of HNO₂ and HNO₃ after NO and NO₂ are dissolved in an aqueous solution, which leads to a decrease in the pH value after the Dissolution process. The concentration of NH₄⁺, NO₂⁻ and NO₃⁻ ions in the aqueous solution and urea solution after the Dissolution process and Dissolution + Reaction process are listed in

Table 3. The ion concentration of the solution after Dissolution/Reaction process (mg/L).

Sample	Dissolution		Dissolution + Reaction
	NO2−	NO3−	NO2−	NO3−	NH4+
4	3.44	43.24	2.58	8.28	43.8
7	3.70	24.01	14.81	8.42	22.6
9	6.72	46.65	18.36	10.24	13.2
12	69.34	10.66	23.92	4.13	1.2

. As for the Dissolution process, with the increase of the pH value of the aqueous solution, the concentration of NO₂⁻ increased from 3.44 to 6.72 mg/L. When the pH value was 12, the NO₂⁻ concentration increased significantly to 69.34 mg/L. Accordingly, the concentration of NO₃⁻ showed a downward trend on the whole. This results from NO₂⁻ being decomposed into NO₃⁻ at an accelerated rate under acidic conditions. When the pH is at a higher level, e.g., pH = 12, NO₂⁻ can be stably stored in the solution. When urea was added into the aqueous solution, the denitration reaction occurred, and the total accumulation of NO₂⁻ and NO₃⁻ declined. On the basis of the principle of mass conservation of the N element, it was speculated that part of the NOx to be removed would react with urea to generate N₂.

Furthermore, the Reaction process was also investigated under different pH values. The specific experimental operation was to add NaNO₂ into the urea solution and use N₂ as the carrier gas. Nitrite (HNO₂) is the key intermediate product during the denitrification reaction through the urea solution. Hence, the reaction process was simulated via adding NaNO₂. When NaNO₂ was added to the urea solution with pH = 0, a large amount of reddish-brown gas was produced immediately, and then the gas disappeared. HNO₂ only exists stably in a dilute aqueous solution. In a concentrated nitrite solution, nitrite will undergo disproportionation and decomposition reactions at the same time, generating the disproportionation products HNO₃ and NO, as well as the decomposition products N₂O₃. N₂O₃ will decompose into NO and NO₂ rapidly. Therefore, the reddish-brown gas should be NO₂ at the beginning of the reaction process. The NO and NO₂ produced in the Reaction process could reach 720 and 220 ppm, respectively, as shown in Figure 8a, and the corresponding CO₂ could reach ppm. No reddish-brown gas appeared when NaNO₂ was mixed with the urea solution with pH = 7. As shown in Figure 8b, the generation amounts of NO, NO₂ and CO₂ were 160, 100 and 350 ppm, respectively. There was a remarkable difference in CO₂ production between the two groups’ reactions. These results indicate that HNO₂ is the key intermediate, and the following reactions occur:

$$\mathrm{CO}{\left({\mathrm{NH}}_2\right)}_2+2{\mathrm{HNO}}_2\to 2{\mathrm{N}}_2+{\mathrm{CO}}_2+3{\mathrm{H}}_2\mathrm{O}$$

(9)

2.3. Simultaneous Removal of NO_x and SO₂

Urea is a strong reducing agent with weak alkalinity. Its aqueous solution has a high removal efficiency of SO₂. The desulfurization product is ammonium sulfate, which avoids the generation of large quantities of desulfurization gypsum and the occupation of land. Does SO₂ promote or inhibit the denitrification reaction in the absorption solution? From the current literature, there is no unified conclusion. Figure 9 shows the influence of 250, 600 and 1000 ppm SO₂ on the denitrification efficiency of the urea solution. The SO₂ removal rate could be maintained above 99% in the wake of the variety of SO₂ concentrations. SO₂ inhibited the denitrification reaction. The conversion of NO, NO₂ and NOx declined slightly. This is due to the small difference between the electrode potential of NO and that of SO₂, which compete with each other when they coexist. Therefore, when the initial concentration of SO₂ is increased, the probability of NO contacting with oxidant will inevitably be reduced, leading to the decrease of the NO removal rate. The ion concentration and pH value of the solution after the reaction process are displayed in

Table 4. The ion concentration and pH value of the solution after the absorption denitration process.

Sample	NH4+ (mg/L)	NO2− (mg/L)	NO3− (mg/L)	SO42− (mg/L)	pH
SO2 0	22.6	14.81	8.42	--	6.66
SO2 250 ppm	32.2	18.13	22.28	16.76	5.49
SO2 600 ppm	31.3	15.32	26.85	33.82	5.29
SO2 1000 ppm	38.3	--	46.02	53.47	3.64

. As the concentration of SO₂ increased from 0 to 1000 ppm, the pH value of the urea absorption solution dropped, and the accumulated NO₃⁻ and SO₄²⁻ increased. After absorption, SO₂ exists in the stable form of SO₄²⁻, and SO₃²⁻ is not detected in the absorption solution. In terms of the denitrification and desulfurization, the desulfurization efficiency of the urea solution is above 99%, and SO₂ will slightly inhibit denitrification. In general, a urea solution can realize simultaneous desulfurization and denitrification.

3. Experimental

The wet denitration of urea solution was evaluated in the laboratory’s self-made absorption equipment. The absorption solution was urea solution with a certain concentration, and the filling volume of the absorption bottle was 150 mL. In order to increase the contact area and enhance mass transfer of the gas-liquid two-phase, metal Pall rings with a diameter of 10 mm were filled into the absorption bottle. The temperature of the urea absorption solution was controlled by a thermostat water bath. The simulated flue gas was used in the experiment, and its specific components were: 500–1500 ppm NO (when used), 500–1500 ppm NO₂ (when used), 100–1000 ppm SO₂ (when used), 10 vol.% O₂, N₂ used as the balance gas, and a total gas flow rate of 100 mL/min. The concentrations of each component of the inlet and outlet gases were monitored online by a Fourier transform infrared (FTIR) spectrometer (Gasmet DX-4000). Each reaction was carried out for 70 min to ensure that the reaction reached equilibrium. The removal efficiency of NO/NO₂/NO_x was calculated using the following equation [26]:

$$N \mathrm{conversion}=\frac{C_N^{in}-C_N^{out}}{C_N^{in}}\times 100\%$$

(10)

where N refers to NO/NO₂/NO_x, and NO_x is the sum of NO and NO₂. Cⁱⁿ and C^out are the concentrations at the inlet and outlet, respectively.

The ion concentrations in the urea absorption solution were obtained with an ion chromatography system (761 Compact IC, Metrohm, Zurich, Switzerland) equipped with a conductivity detector. The C4 cation column and Metrosep A Supp 5 anion separation column were applied to separate anions (i.e., NO₂⁻, NO₃⁻, SO₃²⁻ and SO₄²⁻) and cations (i.e., NH₄⁺) in the solution. In order to protect the chromatographic column, each sample needed to be filtered to remove organic matter (i.e., urea) before testing. The pH value of the solution was analyzed using a pH meter (FE28-Standard, Mettler Toledo, Zurich, Switzerland). The urea concentration was characterized on a spectrophotometer (T6-1650E, PGENERAL, Beijing, China), and the wavelength of the light source was 420 nm.

4. Conclusions

The boundary conditions for the formation of urea nitrate in urea solution were determined, and the main and side reaction paths of the urea absorption denitrification process were revealed in this study. The product of [urea], [H⁺] and [NO₃⁻] concentrations was the critical condition for the formation of urea nitrate. Only when the product of the three concentrations in the absorption solution was higher than 0.87~1.22 mol³/L³ would the urea nitrate precipitation be released, which could preliminarily eliminate the generation of urea nitrate precipitation in the actual operation of the urea absorption unit. The urea concentration exhibited no considerable influence on denitrification in the range of 5–20 wt.%. The influence of the reaction temperature on NO_x removal was a comprehensive result, and 50 °C was a slightly preferred reaction temperature. The higher oxidation degree of the flue gas was conducive to the removal of NO_x. Alkaline conditions were beneficial to the dissolution process, while acidic conditions were beneficial to the reaction process. Therefore, the complete denitration process was the comprehensive embodiment of the dissolution and reaction processes. The desulfurization efficiency of the urea solution was above 99%, and SO₂ would slightly inhibit denitrification process.