Catalytic urea hydrolysis in the selective catalytic reduction of NO_x: catalyst screening and kinetics on anatase TiO₂ and ZrO₂

Abstract

The catalytic hydrolysis of urea was investigated under conditions relevant for the selective catalytic reduction of NO_x (urea-SCR). The hydrolysis activities of the tested catalysts coated on cordierite monoliths were in the order ZrO₂ > TiO₂ > Al₂O₃ > H-ZSM-5 > SiO₂. A comparison with isocyanic acid (HNCO) hydrolysis on the same catalysts showed that urea decomposition was much slower than HNCO hydrolysis; hence, catalytic urea thermolysis into NH₃ and HNCO is likely to be the rate-determining step in urea decomposition. Interestingly, a different order of catalyst activities was found in water-free experiments on urea thermolysis: TiO₂ > H-ZSM-5 ≈ Al₂O₃ > ZrO₂ > SiO₂. The widely accepted reaction pathway for urea decomposition, namely urea thermolysis followed by HNCO hydrolysis, seems to be valid on all the tested catalysts except ZrO₂: The high urea hydrolysis activity of the ZrO₂ catalyst compared to its low urea thermolysis activity suggested a different reaction pathway, in which water directly attacks adsorbed urea rather than adsorbed HNCO.

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Introduction

In the selective catalytic reduction of NO_x (SCR) with urea, the dosed urea solution decomposes in the hot exhaust gas and on the catalyst to yield the actual reducing agent ammonia (NH₃). According to the established mechanism, urea first thermolyzes into NH₃ and isocyanic acid (HNCO) in a non-catalytic reaction (1) and the intermediately formed HNCO then hydrolyzes catalytically (2). HNCO is relatively stable in the gas phase, hence non-catalytic HNCO hydrolysis can be neglected.¹ However, a large fraction of the dosed urea remains unreacted before it enters the catalyst.² Therefore, catalytic urea thermolysis should also be considered.

Thermodynamics show that urea hydrolysis into NH₃ and CO₂ is a significantly exergonic reaction at any temperature.³ On the other hand, an eventual equilibrium limitation of the urea thermolysis reaction into NH₃ and HNCO needs to be discussed. At 150 °C, the thermolysis of pure urea gas is endergonic with ΔG⁰ = 18.5 kJ mol⁻¹.³ Only above 260 °C, ΔG⁰ becomes negative.³ The hypothetical thermolysis of solid urea into gaseous NH₃ and HNCO is even more endergonic with ΔG⁰ = 54.7 kJ mol⁻¹ at 130 °C.³ However, these ΔG⁰ values do not represent the situation in the urea-SCR application, because the compounds involved in the SCR reaction are strongly diluted. The dilution corresponds to a low gas partial pressure, which shifts the urea thermolysis reaction to the product side. The thermodynamic equilibrium curve for the actual concentrations in our experiment is included in Fig. 5a. It shows that our catalyst screening was not biased by an equilibrium limitation of the urea thermolysis reaction.

Yet, most studies on urea decomposition in the SCR process have either focused on non-catalytic thermolysis,^4–13 or on catalytic HNCO hydrolysis.^1,14–21 Scattered information in the literature indicates that urea thermolysis (1) can be catalyzed as well,^22–25 but dynamic experiments, such as thermogravimetric analysis (TGA), have been prevailing so far. Dynamic experiments are difficult to use for kinetic evaluation due to mass and heat transfer phenomena. Only recently, we have reported about catalytic urea thermolysis (1) under steady state conditions over different metal oxides.²⁶ Also, a theoretical study of Todorova et al. (2011) proposes catalytic urea thermolysis. These authors calculated the reaction energy barriers for the hydrolysis of guanidine on anatase TiO₂ (101) using density functional theory (DFT) and found adsorbed urea as an intermediate on the TiO₂ surface.²⁷ From the adsorbed urea NH₃ can be eliminated, leaving HNCO on the catalyst surface. HNCO then hydrolyzes via carbamic acid as another intermediate.^21,27 Interestingly, an additional reaction pathway was proposed, where water directly attacks the adsorbed urea.²⁷ In this more direct reaction (3), carbamic acid is formed as an intermediate as well, but the reaction step of intermediate HNCO formation is skipped. Fig. 1 illustrates the two reaction pathways. The term “urea decomposition” is used to indicate any combination of reactions that consume urea.

Urea thermolysis: CO(NH₂)₂ → NH₃ + HNCO (1)

HNCO hydrolysis: HNCO + H₂O → NH₃ + CO₂ (2)

Direct urea hydrolysis: CO(NH₂)₂ + H₂O → 2NH₃ + CO₂ (3)

In the present study, we compare water-free (dry) catalytic urea thermolysis (1) on different catalysts with catalytic urea hydrolysis (wet), while keeping in mind that there may be two different reaction pathways for urea hydrolysis. Additionally, we used a pseudo first-order kinetic model to analyze urea thermolysis on all the catalysts tested and to analyze direct urea hydrolysis on ZrO₂.

Fig. 1 Reaction scheme of catalytic urea decomposition.

Experimental

Setup

Experiments were carried out using the lab-scale setup described in ref. 28–30 Briefly, the prepared urea solutions were constantly dosed by a gas-assisted spray nozzle into a tubular glass reactor with an inner diameter of 20.4 mm as shown in Fig. 2. Catalyst-coated cordierite monoliths were directly fitted into the reactor, 9 cm downstream of the spray nozzle, where the urea was well distributed on the reactor cross section. We did not observe catalyst deactivation or deposit formation in most of our experiments. Only dry experiments with the H-ZSM-5 catalyst induced slight deactivation. The absence of deposits is attributed to the high spray quality realized in our setup. The model gas was mixed from pure gases using electronic mass flow controllers (Brooks 5850S); water was provided pulsation-free by oxidation of hydrogen on a Pt-catalyst. Deviating from ref. 30 the reactor exit was heated to 170 °C, the extraction capillary to the FTIR spectrometer was heated to 150 °C and the glass wool condenser was removed.

Fig. 2 The reactor used to perform catalytic urea decomposition, with the thermocouples (T) upstream and downstream of a catalyst-coated monolith sample.

Catalyst coating and characterization

Cordierite monoliths (Corning) with a cell density of 600 cells per square inch (cpsi), a wall thickness of 0.15 mm,³¹ 216 channels, 20 mm long and 18 mm in diameter were coated with the catalyst powders shown in Table 1.

Table 1 Catalyst samples tested

Material	Producer, product name	BET surface/m^2 g^−1	Considered p/p0
Anatase TiO2	Crystal Global, DT-51	89	0.05–0.3
ZrO2	MEL Chemicals, XZO881	67	0.05–0.3
Al2O3	Condea, Disperal S	200	0.05–0.3
H-ZSM-5	Süd-Chemie, H-MFI 27	420	0.01–0.1
SiO2	Davison Catalysts, Davicat ® SI 1452	380	0.05–0.3

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The TiO₂ powder was suspended in a diluted aqueous NH₃ solution, the other powders in pure water. We improved the stability of the coating using a commercial colloidal silica binder (Ludox AS-40, 10% of the catalyst mass). The monoliths were calcined at 550 °C for 5 h. The specific surface areas of catalyst powders were measured by nitrogen physisorption on a Quantachrome Autosorb 1-c instrument. We prepared the samples for the physisorption measurements by calcining dried catalyst suspensions under the same conditions as the coated monoliths.

Measuring conditions

The tests were typically conducted at a space velocity (GHSV) of 91 000 h⁻¹ using the catalyst coated monoliths with a cell density of 600 cpsi shown in Table 2. The base feed of the reactor was composed of 10% O₂ and 0% or 5% H₂O in N₂, at a total flow rate of 500 L h⁻¹ at standard temperature and pressure (STP: 0 °C, 1013 mbar).

Table 2 Catalyst coated monoliths used (the GHSV was calculated at 500 L h⁻¹ at STP)

Monolith name	Active mass/mg	Length/cm	GHSV/h^−1	Loading/g L^−1
TiO2	45	2	91 000	8.3
TiO2b	85	1.06	170 000	29
TiO2c	540	6	30 000	33
TiO2d	24	2	91 000	4.4
ZrO2	52	2	91 000	9.4
Al2O3	64	2	91 000	12
H-ZSM-5	48	2	91 000	9.1
SiO2	55	2	91 000	10

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A concentration of 100 ppm urea in the gas phase was realized by spraying 4 wt% urea in ethanol (70.4 μL min⁻¹), 5 wt% urea in methanol (56.3 μL min⁻¹) or 15 wt% urea in water solution (14.5 μL min⁻¹). Urea solutions were prepared by dissolving urea (Merck, ≥99.5% purity) in ethanol (Merck, ≥99.9% purity), methanol (VWR, <0.05% water) or de-ionized water. Dosing of the organic solvents resulted in gas phase concentrations of 0.31% ethanol or 0.36% methanol. When urea in water solution was dosed, the water concentration in the base feed was reduced to maintain an effective water concentration of 5% in the reactor feed.

Non-catalytic urea decomposition could be neglected because very little urea conversion was observed in the empty reactor up to 200 °C (Fig. 3) and most of the data points considered for the Arrhenius analyses were below 200 °C.

Fig. 3 Urea conversion in the empty reactor.

In addition to dry and wet urea decomposition, we performed HNCO hydrolysis experiments (100 ppm HNCO, see the next section) and biuret hydrolysis experiments with 2% biuret (Fluka, p.a.) in de-ionized water solution (95.3 μL min⁻¹ liquid dosing rate, resulting in 50 ppm biuret).

HNCO generation for HNCO hydrolysis experiments

Urea was the reactant in most of the experiments. Additionally, experiments with HNCO as the reactant were carried out for comparison. HNCO was generated by thermolysis of cyanuric acid in a separate reactor as described in ref. 1. Cyanuric acid was sublimed at about 280 °C and thermolyzed downstream over stainless steel nuts at 380 °C. The carrier gas flow (dry N₂) through the HNCO generator was set to 20 L h⁻¹ at STP. After leaving the HNCO generator, its product gas was immediately mixed with the main gas stream (total gas flow = 500 L h⁻¹ at STP). The typical composition of the resulting gas mixture was 100 ppm HNCO, ≈15 ppm NH₃, ≈20 ppm CO₂ and ≈3 ppm CO. Before starting an experiment, the HNCO generator was allowed to stabilize for about 1 h, while its output was monitored by FTIR spectroscopy (bypassing the reactor containing the catalyst-coated monolith). After the experiment, the HNCO generator output was measured again. Typically, the HNCO concentration was stable within a few percent. From the HNCO generator output before and after the experiment, the values at any time during the experiment were calculated by linear interpolation.

Analytics

The gaseous compounds NH₃, CO₂, HNCO, H₂O, ethanol, methanol and urea were measured by Fourier transform infrared (FTIR) spectroscopy using a multi-component analysis method developed in-house.^26,30,32 The quantification of urea by FTIR spectroscopy was precise in the water-free experiments, but the presence of water induced scattering.²⁶ The product gas was also analyzed using a liquid-quench of the product gas, followed by HPLC analysis as described in ref. 30, 33, and 34. The quenching solution was the HPLC eluent itself: 5 mM phosphate buffer, set to pH 7 by NaOH. Three liquid samples were collected consecutively at a selected temperature at intervals of 10 min each. HPLC analysis was used to check the HNCO and urea concentration measured by FTIR spectroscopy and to measure urea in the presence of water. Furthermore, the HPLC method is capable of quantifying potential emissions of byproducts such as biuret, triuret, cyanuric acid and ammelide (Fig. 4),^33,34 but none of these byproducts were observed in the present study, except for the biuret decomposition experiments. The C-balance could usually be closed within ±5% accuracy by summing up the HNCO, urea and CO₂ amounts. The CO₂ yield in the dry experiments due to hydrolysis by water traces was mostly below 5%.

Fig. 4 Examples for urea decomposition byproducts that can be quantified by the applied HPCL method.

Calculation of product yields and urea conversions

The yields of the reaction products NH₃, HNCO and CO₂ were calculated based on the N- or C-balance. Using the concentrations was allowed, as all species were diluted to ≤200 ppm in the model exhaust gas matrix:

NH₃ yield = c(NH_3,out)/(2 × c(urea_in))

HNCO yield = c(HNCO_out)/c(urea_in)

CO₂ yield = c(CO_2,out)/c(urea_in)

urea slip = c(urea_out)/c(urea_in)

Urea conversion (X) was always calculated according to

X = 2 × NH₃ yield − CO₂ yield

Although it would have been easiest to calculate the urea conversion based on the urea slip, the equation X = 2 × NH₃ yield − CO₂ yield was used because of the following reasons: (a) NH₃ and CO₂ could be measured more precisely than urea; (b) a moderate relative measuring error of the large urea slip at low conversions would have resulted in a large relative error of the conversion rates, required for the Arrhenius analyses; (c) the urea conversion is still indicated irrespective of the presence of water (see explanation in the next two paragraphs) and (d) urea condensation in the reactor would have indicated wrong urea conversion.

For the dry experiments, it would have been acceptable to calculate the conversion based on the HNCO yield alone. However, intermediately formed HNCO, which is hydrolyzed due to water traces, is missing in the HNCO yield. A “true” HNCO yield, which is corrected for the HNCO missing due to HNCO hydrolysis, can be calculated by adding the CO₂ yield to the HNCO yield (provided that there is no direct urea hydrolysis). The HNCO quantification by FTIR spectroscopy is highly sensitive, but more vulnerable to systematic errors than the quantification of the other compounds because calibration of the reactive HNCO is difficult. The HNCO quantification by HPLC is robust, but time consuming and less sensitive than the quantification by FTIR spectroscopy. Hence, it was a better option to use the NH₃ yield to calculate the yield of the NH₃ produced by urea thermolysis alone (thermolysis-NH₃ yield), which is in principle equivalent to the “true” HNCO yield.

“true” HNCO yield: HNCO yield + CO₂ yield

thermolysis-NH₃ yield: 2 × NH₃ yield − CO₂ yield

If only urea thermolysis takes place, the HNCO yield, the “true” HNCO yield and the thermolysis-NH₃ yield are identical except for measuring errors. In the wet experiments, the chosen definition of the urea conversion based on the thermolysis-NH₃ yield is valid as well. In fact, the thermolysis-NH₃ yield indicates the urea conversion irrespective of the product selectivity. Even if direct urea hydrolysis (3) takes place, the thermolysis-NH₃ yield correctly indicates the urea conversion, but, of course, without indicating anymore how much NH₃ was produced by urea thermolysis. Anyway, if CO₂ is produced with a very high selectivity, the best option is to simply base the urea conversion on the CO₂ yield. We tested using the CO₂ yield to analyze urea hydrolysis on ZrO₂, but the result was very close to the result based on the thermolysis-NH₃ yield. Hence, we chose to use only one definition of the urea conversion. Using a uniform definition of the urea conversion is simple and provides excellent comparability between the results obtained under different conditions (dry or wet and different catalysts).

Results and discussion

Catalyst screening

Catalytic urea decomposition is a fast reaction. To realize reaction conditions with little diffusion limitation, small amounts of the different catalyst powders were coated on cordierite monoliths with a high cell density of 600 cpsi. The typical catalyst loading of about 50 mg corresponded to only 9 g L⁻¹, which is about 20 times lower than in a real application in exhaust gas aftertreatment. Also, we chose a rather low urea concentration of 100 ppm in the gas phase to allow measurements at low temperatures around 150 °C without urea condensation.

Fig. 5 shows the results of a catalyst screening for dry urea thermolysis (a and b) and urea hydrolysis in the presence of 5% water (c and d). Dry conditions, required for studying urea thermolysis without hydrolysis, were realized by using ethanol and methanol solutions of urea. Fig. 5a and b also compare urea thermolysis in the presence of ethanol, which was already reported elsewhere,²⁶ with thermolysis in the presence of methanol. Changing the solvent from ethanol to methanol influenced the thermolysis reaction only slightly, indicating that these water-free, polar solvents are suitable for studying the urea thermolysis only. In both cases, no compounds originating from side-reactions due to the presence of the solvent could be detected by FTIR spectroscopy and the mass balance could often be closed more precisely than in the hydrolysis experiments.²⁶

Fig. 5 Screening of catalysts for urea decomposition. (a, b) Dry experiments; solid lines: ethanol, dashed lines: methanol. (c, d) Hydrolysis with 5% water. (a) includes the NH₃ yield according to the thermodynamic equilibrium of the urea thermolysis reaction.³⁵

In the absence of water (Fig. 5a and b), the urea thermolysis products NH₃ and HNCO were formed with a high selectivity. Adding the urea slip (Fig. 5b) to the NH₃ or HNCO yield usually allowed for closing the mass balance. Only at the lowest temperatures, the urea slip decreased due to urea condensation.

Under hydrolysis conditions (Fig. 5c and d), NH₃ and CO₂ were the final products, however, only ZrO₂ always showed a high selectivity for CO₂. By contrast, TiO₂, Al₂O₃ and H-ZSM-5 showed significant local maxima in the HNCO yield at 165 °C, 180 °C and 190 °C, respectively. Local maxima of reaction intermediates are typical for sequent reactions. SiO₂ showed poor thermolysis and even lower hydrolysis activity. The hydrolysis activity was so low that even in the presence of water the HNCO yield increased steadily with increasing temperature (Fig. 5d).

In light of the very high HNCO hydrolysis rates on anatase TiO₂ reported by Hauck et al. (2007),¹⁴ the intermediate HNCO peaks shown in Fig. 5d may be surprising. Fig. 6 shows the hydrolysis of 100 ppm HNCO on some of the previously tested catalysts for urea hydrolysis. Indeed, we found HNCO hydrolysis to be very fast on TiO₂ and even much faster on ZrO₂. The HNCO slip on TiO₂ at 165 °C was only 3%, whereas a local maximum of 37% HNCO yield was observed at 165 °C during urea hydrolysis on the same TiO₂-coated monolith (Fig. 5d). Apparently, HNCO hydrolysis was strongly inhibited by the presence of urea.

Fig. 6 HNCO hydrolysis, using the same catalyst-coated monoliths as were used for the experiments shown in Fig. 5.

Fig. 7 shows a comparison of the HNCO yields observed in the presence of water and without water. To indicate the urea conversion irrespective of the product selectivity, Fig. 7 also shows the thermolysis-NH₃ yield. If only urea thermolysis takes place, the HNCO yield and the thermolysis-NH₃ yield are identical except for measuring errors. Due to the strong inhibition of the HNCO hydrolysis reaction by the presence of urea, the HNCO yield obtained on TiO₂ under hydrolysis conditions was only slightly lower than the HNCO yield without water up to 160 °C (Fig. 7a). Above 165 °C, the HNCO yield decreased again due to HNCO hydrolysis. On ZrO₂, the HNCO yield was much higher without water at most temperatures, however, at the lowest temperature investigated (160 °C), the difference was small (without water: 6%, with water: 5% HNCO yield, see Fig. 5).

Fig. 7 Dry and wet urea decomposition. The data are from the same experiments as those shown in Fig. 5. Dashed lines and empty symbols represent the NH₃-yield by thermolysis, solid lines and filled symbols show the FTIR data for HNCO. (a) TiO₂, (b) H-ZSM-5, (c) Al₂O₃ and (d) SiO₂. Dry experiments were conducted with urea in methanol solution (TiO₂ and H-ZSM-5) or urea in ethanol solution (Al₂O₃ and SiO₂).

Surprisingly, the HNCO yields obtained on the H-ZSM-5 (Fig. 7b) and on the Al₂O₃ (Fig. 7c) catalyst were higher in the presence of water than without water at some temperatures. On H-ZSM-5 at 180 °C, the HNCO yield increased dramatically from 33% without water to 93% with water. On SiO₂, the HNCO yield was higher with water over the whole temperature range (Fig. 7d).

The simplest explanation for the low HNCO yields in the dry experiments would be inhibition of urea thermolysis by the organic solvent used to dose the urea. However, it is plausible that the adsorption strength of the solvents on the metal oxide catalysts decreases with their polarity in the order H₂O > methanol > ethanol. Moreover, the concentration of the organic solvents was about 15 times lower than the H₂O concentration (H₂O: 5%, methanol: 0.36%, ethanol: 0.31%). Hence, the organic solvents are likely to inhibit urea adsorption less strongly than H₂O.

Another potential effect of the organic solvents that has to be checked is the presence or absence of side-reactions like condensation into ethers or substitution of the OH group with NH₃. Our results show that the organic solvents did not induce side-reactions because the C-balance of the dry urea thermolysis reaction could be well closed by summing up the HNCO, urea and CO₂ yields. Also, we did not observe solvent-related byproducts by FTIR spectroscopy or HPLC analysis.²⁶

Possibly, the presence of water on the catalyst surface accelerated the urea thermolysis by facilitating proton transfer reactions. This assumption of water assisting in the urea thermolysis is supported by a theoretical study by Alexandrova and Jorgensen (2007) that suggests water to act as a proton shuttle for the formation of the zwitterionic intermediate H₃N⁺C(O)N⁻H in the mechanism for urea thermolysis in aqueous solution.³⁶

Table 3 summarizes the catalyst activity results from Fig. 5 and 6. Interestingly, ZrO₂, which showed the highest hydrolysis activity, showed only low thermolysis activity (see also Fig. 9b). The high hydrolysis activity of the ZrO₂ catalyst in spite of its low thermolysis activity indicates that urea hydrolyzed directly (3) on the ZrO₂ surface without intermediate HNCO formation. Alternatively, if HNCO was formed on ZrO₂ during urea hydrolysis as a short lived intermediate, the low thermolysis activity of ZrO₂ might be due to a low proton transfer capability of the dry ZrO₂ surface. Anyway, since urea hydrolysis on the ZrO₂ catalyst always produced CO₂ with a high selectivity, a kinetic model with only one rate-determining step should be sufficient to describe urea hydrolysis on ZrO₂.

Table 3 Qualitative summary of the catalyst activities shown in Fig. 5 and 6

Catalyst	Urea thermolysis	Urea hydrolysis	HNCO hydrolysis
TiO2	High	High	High
ZrO2	Low	Very high	Very high
H-ZSM-5	Moderate	Low	Low
Al2O3	Moderate	Moderate	<TiO2^17,37
SiO2	Very low	Very low	Very low^17,37

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Kinetics

For comparison of the urea conversions on the different catalysts, we calculated pseudo first-order rate constants (k₁) using the equationwhere V* is the actual gas volume flow rate, W is the catalyst weight and X is the urea conversion. By using this equation, we assume a first-order dependency with respect to urea, whereas all the other reaction orders are approximated by zero. The same approach was used by Kleemann et al. (2000)¹ to calculate pseudo first-order rate constants of the HNCO hydrolysis (first-order with respect to HNCO, water present in excess). Fig. 8 shows the resulting Arrhenius plot of the urea thermolysis data presented in Fig. 5a.

Fig. 8 Arrhenius analysis of the urea thermolysis experiments presented in Fig. 5a.

To experimentally confirm that the urea conversion on the most active catalyst in our screening, TiO₂, was not limited by the thermodynamic equilibrium of the urea thermolysis reaction, we performed an experiment with increased active mass. As expected, increasing the active mass from 45 mg (curve “TiO₂” in Fig. 8) to 540 mg (curve “TiO₂c” in Fig. 8) on the monolith increased the HNCO yield from 22% to 56%, respectively. Since the pseudo first-order rate constants shown in Fig. 8 are normalized to the active mass, the rates for the “TiO₂c”-monolith were below those of the “TiO₂”-monolith. Still, the pseudo first-order rate constants at 150 °C were quite similar for the two TiO₂-coated monoliths due to the higher urea conversion on the “TiO₂c”-monolith. Below 150 °C, the pseudo first-order rate constants from the “TiO₂c”-monolith were probably lowered due to urea condensation, whereas, above 150 °C, they were lowered due to mass transport limitation.

To test the presence or absence of mass transfer limitation in the Arrhenius plot (Fig. 8), we calculated η_extDaII values using the equationwhere³⁸r_eff is the observed reaction rate, k_g is the gas phase mass transport coefficient, a is the geometric monolith surface area, T is the actual temperature and T₀ is the temperature for the tabulated value of the binary diffusion coefficient (D_1,2).

The η_extDaII value is a measure of the external mass transport limitation of a reaction (diffusion of gaseous urea in the monolith channels). A η_extDaII value below 0.1 means the mass transport limitation is insignificant.³⁸ The calculation of the η_extDaII values was based on the following assumptions:

• complete evaporation and mixing of the dosed urea with the model gas

• laminar gas flow

• diffusion length (d_g) = ¼ channel width

• The binary diffusion coefficient (D_1.2) of urea in the model gas was approximated by the binary diffusion coefficient of SO₂ in air: D_{urea,modelgas} ≈ D_SO2,air = 0.122 cm² s⁻¹ at 298 K.³⁸ The values for slightly different gas matrices do not differ significantly. The binary diffusion coefficient of SO₂ was chosen for the calculations with urea due to the quite similar molecular mass and due to the not completely different geometry.

• conversion at the catalyst exit used for calculating the concentration of gaseous urea (c_1,g)

Please note that using the conversion at the catalyst exit for calculating c_1,g means making a worst case assumption, because the low c_1,g at the catalyst exit leads to the assumption of a small urea concentration gradient. Another assumption made has to be discussed: the assumption of complete urea evaporation. Unfortunately, we could not directly measure if and where the urea aerosols evaporate in our reactor. However, the vapor pressure of urea is more than high enough to allow for complete urea evaporation under the conditions applied.^33,39 Indeed, several indications suggest that the high spray quality in our setup allowed for quantitative urea evaporation upstream of the catalyst.

• During method development, the performance of the SCR reaction with respect to NO_x reduction and NH₃ emissions at the reactor exit was found to be almost the same with sprayed urea solution as with NH₃ gas over a very broad parameter range, when the distance between the nozzle and catalyst was properly chosen.³⁰

• Fig. 5c shows the CO₂ yield obtained by urea hydrolysis in long (“TiO₂”) and short (“TiO₂b”) TiO₂-coated monoliths. The shorter monolith with the higher amount of the catalyst showed a higher CO₂ yield at low temperatures up to 180 °C, as expected for gaseous urea but not for urea aerosols. If evaporation of the urea aerosols would not have been complete upstream of the monolith, aerosols would have preferably slipped through the channels of the short monolith resulting in a lower conversion at all temperatures.

• In a first series of urea hydrolysis experiments, we placed a catalyst-coated monolith at a larger distance from the spray nozzle. Then we placed an inert cordierite foam between the spray nozzle and the catalyst. Neither of these measures increased the urea conversion. If the urea aerosols evaporated slowly, an inert structure upstream of the catalyst should have improved the urea evaporation, which should also have increased the urea conversion on the catalyst.

These observations suggest that most of the urea aerosols dosed by the spray nozzle evaporated upstream of the catalyst. If the urea aerosols did not evaporate, at least the aerosol slip through the catalyst-coated monoliths was insignificant.

Determination of apparent activation energies

Fig. 9a shows an Arrhenius plot of the urea thermolysis on anatase TiO₂ and calculated η_extDaII values for this experiment. An η_extDaII value below 0.1 indicates insignificant mass transport limitation.³⁸ The pseudo first-order rate constants of the urea hydrolysis were calculated likewise and plotted in Fig. 9a. We are aware of the fact that the simple kinetic model is not applicable for the two-step urea hydrolysis on TiO₂, where a significant amount of the intermediate HNCO can desorb into the gas phase. However, due to the definition of the conversion as X = 2 × NH₃ yield − CO₂ yield, the NH₃ produced by HNCO hydrolysis is subtracted and only the NH₃ produced by urea thermolysis is taken into account. Therefore, the pseudo first-order rate constants obtained for hydrolysis conditions (5% water) should be suitable for comparison. Comparison of the pseudo first-order rate constants shows that the hydrolysis was somewhat slower than thermolysis below 150 °C, whereas the hydrolysis was faster above 160 °C. Maybe, the hydrolysis was slower below 150 °C due to competitive adsorption of urea and water, whereas the hydrolysis was faster above 160 °C because the surface coverage was lowered by HNCO hydrolysis or because adsorbed water accelerated the thermolysis by acting as a proton shuttle.

Fig. 9 Arrhenius analyses of urea decomposition on anatase (a) TiO₂ and (b) ZrO₂. Urea in methanol solution was used for the dry experiments.

Based on the η_extDaII values and the apparent linearity of the Arrhenius plots in the temperature range between 150 °C and 170 °C, the urea thermolysis data points in this temperature range were used for a linear regression (highlighted in Fig. 9a). From the obtained regression line an apparent activation energy (E_a) of 90 kJ mol⁻¹ and a pre-exponential factor (A) of 6× 10¹³ s⁻¹ were calculated for urea thermolysis on TiO₂ (Fig. 9a).

The activity of the ZrO₂ catalyst (Fig. 9b) was analyzed likewise. If urea hydrolyzes in one step on ZrO₂ according to reaction (3) as proposed in the introduction, the kinetic model used for the urea thermolysis can be applied for urea hydrolysis on ZrO₂ as well. The corresponding values derived from Fig. 9b are E_a = 100 kJ mol⁻¹, A = 4 × 10¹³ s⁻¹ for the thermolysis (considering the data points at 180, 190, 195, 200, 205, 210 and 215 °C) and E_a = 70 kJ mol⁻¹, A = 8 × 10¹¹ s⁻¹ for the hydrolysis (considering the data points at 125, 130, 135 and 140 °C).

Biuret hydrolysis

Byproduct formation was not observed in the experiments reported above, however, byproducts may have accumulated slowly, which could cause catalyst deactivation after long-term operation. The most probable byproducts are biuret, cyanuric acid, triuret and ammelide.⁴⁰ Since only biuret has a significant solubility in water, 2% aqueous biuret solution was used in this study to investigate byproduct decomposition.

Fig. 10 shows the product distribution of biuret hydrolysis on TiO₂. Biuret hydrolyzed fairly well in a three-step reaction, showing local maxima of the formed intermediates urea and HNCO. The urea yield peaked at 190 °C and 0.29 mol/mol-biuret, whereas HNCO peaked at 200 °C and 0.23 mol/mol-biuret. When the water concentration in the gas phase was decreased from 5% to 1.4%, the urea peak increased to 0.36 mol/mol-biuret, whereas the HNCO peak shifted to 210 °C and increased to 0.52 mol/mol-biuret (not shown).

Fig. 10 Biuret hydrolysis on the TiO₂d-monolith sample.

In the empty reactor, higher biuret conversions were observed (not shown) than urea conversions shown in Fig. 3. This result is surprising because biuret is more stable against non-catalytic decomposition than urea.⁸ The high biuret conversion in the empty reactor was probably caused by the low vapor pressure of biuret. Unlike urea, which seemed to evaporate quickly and completely in our experiments, part of the biuret may have stuck to the reactor wall until it decomposed. In other words, the residence time and the heat transfer from the reactor wall to the reactant were increased for biuret compared to urea. The decomposition products urea, HNCO and NH₃ have higher vapor pressures than biuret and could therefore easily leave the reactor. As mentioned, the biuret conversion in the empty reactor was quite high. Still, the TiO₂ catalyst largely increased the biuret conversion. At 190 °C, where the urea yield with the TiO₂ catalyst peaked at 0.29 mol/mol-biuret, the yield in the empty reactor was only 0.08 mol/mol-biuret, which strongly indicates catalytic biuret decomposition.

Catalytic biuret decomposition activity is a highly welcome feature of a urea hydrolysis catalyst in an SCR system, because biuret that may form during low temperature operation will quickly hydrolyze at elevated temperatures. We have observed catalytic biuret decomposition before in a recent study on the formation and decomposition of byproducts on TiO₂.⁴⁰ In that study,⁴⁰ slow partial catalytic biuret hydrolysis yielding urea was observed at only 100 °C. At 160 °C, melamine was catalytically hydrolyzed into ammeline, ammelide and cyanuric acid. Cyanuric acid hydrolysis started at about 200 °C. Hence, temperatures around 200 °C should be sufficient to regenerate a TiO₂ catalyst after poisoning with biuret, cyanuric acid, ammelide, ammeline and melamine. Such temperatures are easily reached in urea-SCR applications. The decomposition of more stable byproducts like melem was not investigated in ref. 40 Melem is very unlikely to ever form on TiO₂, because melem can only be formed from melamine at very high temperatures around 500 °C,²² when melamine is already hydrolyzed in the presence of water. Notably, Al₂O₃ was reported to catalytically hydrolyze even melem.⁴¹ TiO₂, which showed a higher hydrolysis activity than Al₂O₃ in our study, may thus catalyze melem hydrolysis, too.

Conclusions

Experiments on the catalytic hydrolysis of 100 ppm urea to NH₃ and CO₂ in the gas phase showed the following order of hydrolysis activities: ZrO₂ > TiO₂ > Al₂O₃ > H-ZSM-5 > SiO₂. Dry experiments on the thermolysis of urea to HNCO and NH₃ revealed a different activity order: TiO₂ > H-ZSM-5 ≈ Al₂O₃ > ZrO₂ > SiO₂. For the most active catalysts TiO₂ and ZrO₂ the activation energies for hydrolysis and thermolysis were determined by means of an Arrhenius analysis of the apparent pseudo first-order rate constants:

Urea thermolysis on TiO₂: E_a = 90 kJ mol⁻¹

Urea thermolysis on ZrO₂: E_a = 100 kJ mol⁻¹

Urea hydrolysis on ZrO₂: E_a = 70 kJ mol⁻¹

The found activation energies for the thermolysis of urea are larger than the activation energy of 73 kJ mol⁻¹ reported for the hydrolysis of HNCO on anatase TiO₂ by Hauck et al. (2007).¹⁴ Consequently, our results clearly show that urea thermolysis is slower than HNCO hydrolysis (Fig. 5vs. Fig. 6) and urea thermolysis must be the rate-determining step in catalytic urea decomposition in the SCR process.

On TiO₂, the urea conversion in the absence of water was quite similar to the urea conversion with water (Fig. 9a). The main effect of the presence of water was a changed selectivity of the overall reaction due to HNCO hydrolysis. By contrast, the ZrO₂ catalyst showed a much lower urea conversion in the absence of water than with water, which indicates that urea hydrolyzes on ZrO₂ according to a different mechanism, in which water directly attacks adsorbed urea rather than adsorbed HNCO (Fig. 9b).

Appendix

a	diffusion surface area
A	pre-exponential factor in the Arrhenius equation
c	volume fraction for gas phase concentrations; mass fraction for liquid phase concentrations
c1,g	reactant concentration in the gas phase
dg	diffusion length
D1.2	binary diffusion coefficient
DaII	second Damköhler-number
Ea	apparent activation energy
reff	effective reaction rate
k1	apparent pseudo first-order reaction rate constant
kg	gas phase mass transport coefficient
ηext	efficiency factor, ηext=reff/r
R	gas constant
T	actual temperature
T0	temperature for the tabulated value of D1,2
V*	total flow rate at the actual temperature and pressure
W	catalyst weight
X	relative conversion

Acknowledgements

Funding by TOTAL (France) is gratefully acknowledged. TOTAL was not involved in the study design, collection, analysis or interpretation of data, nor was it involved in the preparation of the article.

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