Using Transient FTIR Spectroscopy to Probe Active Sites and

May 29, 2019 - With respect to standard SCR mechanisms, it has been well established ..... (19,20,28,36) Among these, δas(N–H) of L-NH3 at 1620 cmâ...
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Research Article Cite This: ACS Catal. 2019, 9, 6137−6145

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Using Transient FTIR Spectroscopy to Probe Active Sites and Reaction Intermediates for Selective Catalytic Reduction of NO on Cu/SSZ-13 Catalysts Yani Zhang,†,‡ Yue Peng,*,† Kezhi Li,† Shuai Liu,† Jianjun Chen,† Junhua Li,† Feng Gao,*,‡ and Charles H. F. Peden‡

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†

State Key Joint Laboratory of Environment Simulation and Pollution Control, School of Environment, Tsinghua University, Beijing 100084, China ‡ Institute for Integrated Catalysis, Pacific Northwest National Laboratory, Richland, Washington 99352, United States S Supporting Information *

ABSTRACT: A Cu/SSZ-13 catalyst containing predominately Z2Cu sites is prepared. Using FTIR spectroscopy, two nonsteady-state measurements, (1) continuous NO titration of an NH3 saturated catalyst and (2) intermittent NO on/off cycles during quasi-steady state (i.e., NO perturbation), are conducted to shed light on active sites, reaction intermediates, and possible reaction mechanisms. During continuous NO titration, a strong NH3 inhibition effect is found, where Cu active sites containing NH3 ligands are less active than Cu sites depleted of NH3 ligands. In the NO perturbation experiments, it is demonstrated that NO+NH3 interactions lead to formation of Brønsted acid sites, which further interact with NH3 to form NH4+. Perturbation measurements, using ND3 to replace NH3 in order to distinguish ammonia and nitrate vibrations, further show that NH4NO3 and other surface nitrates are not involved in the SCR process under typical low-temperature quasi-steady-state conditions. KEYWORDS: selective catalytic reduction, Cu/SSZ-13, transient kinetics, mechanism, DRIFTS

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INTRODUCTION Copper-exchanged SSZ-13 (Cu/SSZ-13) zeolite has been commercialized in a diesel vehicle after-treatment systems for selective catalytic reduction of NOx with ammonia (NH3− SCR).1−5 Cu/SSZ-13 adopts a small pore Chabazite (CHA) zeolite structure that contains catalytically active isolated Cuion sites, typically introduced via ion-exchange, for the SCR process. Compared with previously studied Cu-zeolite SCR catalysts such as Cu/ZSM-5 and Cu/Beta, this catalyst exhibits substantially improved SCR activity, hydrothermal stability, and sulfur resistance.6−8 Numerous experimental and theoretical studies have been conducted to investigate the mechanisms of the standard SCR reaction (4NO + 4NH3 + O2 → 4N2 + 6H2O) on the catalyst.9−14 With respect to standard SCR mechanisms, it has been well established that ammonia is much more strongly bound to the catalyst than NO under most reaction conditions.12,15−17 Via acid−base interactions, reactive NH3 first adsorbs onto Cu(II) Lewis acid sites (abbreviated as L-NH3) or Brønsted acid sites (as NH4+, abbreviated as B-NH3). Then, the Cu(II)-NH3 complexes react with NO to generate surface reaction intermediates which decompose into N2 and H2O, while divalent copper is reduced to monovalent copper.10,12−14 Depending on whether NO chemisorbs or only weakly interacts with the Cu(II)-NH3 complexes, standard SCR can © 2019 American Chemical Society

be regarded as a Langmuir−Hinshelwood type, or an Eley− Rideal type catalytic reaction, respectively. However, since the SCR process involves more than two reactants and since the active Cu-ion sites cycle between Cu(II) and Cu(I) oxidation states, SCR may be better categorized as a redox reaction, where the monovalent copper is oxidized back to divalent copper by O2 in order to complete the entire catalytic cycle.3 The mechanisms discussed above are very general, lacking many molecular-level details of the standard SCR chemistry. For example, even the nature of the reactive Cu(II)-NH3 complexes can be rather complex. In Cu/SSZ-13, active Cu(II) species can be present as either or both Z2Cu and ZCuOH (where Z represents a negative framework charge), depending on how the Cu(II) ions are charge-balanced by the zeolite framework.2,3,18 Correspondingly, Cu(II)-NH3 complexes can vary from fully ammonia coordinated [Cu(NH3)4]2+ and [Cu(OH)(NH3)3]+ species to partially ammonia coordinated moieties (e.g., [Cu(OL)4−x(NH3)x]2+ species), where OL represents framework oxygen atoms.16 Furthermore, it is expected that the relative populations of these complexes are dependent on catalyst compositions (Si/Al and Cu/Al ratios) Received: February 20, 2019 Revised: April 27, 2019 Published: May 29, 2019 6137

DOI: 10.1021/acscatal.9b00759 ACS Catal. 2019, 9, 6137−6145

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its semiquantitative nature (i.e., intensities of surface IR bands may not be directly proportional to concentrations of the corresponding surface species), DRIFTS does offer a few important advantages over other characterization techniques in terms of (1) being readily in situ/operando, (2) excellent ability for identifying and distinguishing various chemical species, and (3) good detection limits for low-concentration species. In the studies described here, we were particularly motivated to use time-resolved transient studies on the basis of very recent research by Marberger et al.17 that utilized a similar strategy to obtain some important new details of the SCR mechanism over Cu/SSZ-13 catalysts.

and reaction conditions (temperature and NO/NH3 partial pressures). Importantly, it can be anticipated that the various Cu(II)-NH3 complexes display different reactivity toward NO, a hypothesis that is difficult to probe using ensemble-averaging techniques under steady-state reaction conditions. Besides the Cu(II)-NH3 complexes described above, B-NH3 is another form of potentially reactive ammonia species. Some early titration measurements (i.e., using NO to titrate ammonia saturated Cu/SSZ-13 at relatively low temperatures, probed with FTIR) indicated that Cu(II)-NH3 complexes are much more reactive than B-NH3;19,20 the latter has been suggested, therefore, to participate in the SCR reaction at most as a reservoir of NH3. In fact, such “stored” ammonia is expected to influence SCR more significantly under nonsteady-state reaction conditions than at steady state, notably, when input concentrations of NH3 may fluctuate. Moreover, ammonia storage capacities for Brønsted acid sites will decrease as reaction temperatures rise, clearly leading to a decrease in potential mechanistic contributions from the B-NH3 reservoirs. Reaction intermediates that decompose to N2 have been extensively discussed in literature. Based on a charge balancing requirement for reaction intermediates, notably, if the N atom coming from NH3 has an oxidation state of −n, the other N atom (coming from NOx) should have an oxidation state of +n to facilitate N2 formation.21 This leads to two reactant activation scenarios: (1) NH3 is activated by Cu(II) via N− H bond cleavage to form a Cu(I)-NH2 complex and a H+. Note that this process is accompanied by an electron transfer from N to Cu such that the N atom now has a formal oxidation state of −2 in the resultant Cu(I)-NH2 complex. In this case, NO (with a +2 oxidation state for N) can directly interact with this complex to generate a Cu(I)-NH2NO intermediate, which can then rearrange and decompose to N2+H2O. Significantly, theoretical calculations indicate that NO-assisted NH 3 activation is energetically much more favorable than NH3 activation in the absence of NO.10,22 Alternatively, (2) NO is activated by Cu(II), forming a species with a +3 oxidation state for N that are reactive toward nonactivated NH3. The thusgenerated possible intermediates include NO+, NO2−, and HONO.23,24 In particular, theoretical calculations demonstrate that a ZCuOH active site can interact with NO to form HONO with rather low energy barriers. HONO can then react with NH3 to form an NH4NO2 intermediate that can readily decompose to N2+2H2O.13 It is important to note that, although NH2NO and NH4NO2 intermediates have been frequently proposed, particularly based on theoretical simulations,3,10,12,13 proof of their presence under typical reaction conditions via experimental approaches has been elusive because of their highly unstable nature. On the other hand, a surface species that is much more stable and, therefore, much more readily observable with spectroscopy (besides Cu-NH3 and B-NH3) is ammonium nitrate.19,25 While NH4NO3 is not expected to contribute to SCR at low temperatures because of its inertness, its dominance in the N−O vibrational regions, unfortunately, can severely complicate identification of other more important reaction intermediates (e.g., NH4NO2) via IR spectroscopy.19,25 In the present study, we aim to obtain further insights into standard SCR mechanisms and intermediates, particularly under nonsteady-state conditions, via time-resolved transient diffuse reflectance infrared Fourier transform spectroscopy (DRIFTS) combined with reactant isotope-labeling. Despite

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EXPERIMENTAL METHODS The Cu/SSZ-13 catalyst used here was prepared by a two-step ion-exchange method. First, a commercial H/SSZ-13 powder (Si/Al = 10, provided by BASF) was mixed with a 1 M NH4NO3 solution at 70 °C for 4 h under stirring to obtain NH4/SSZ-13. To ensure complete exchange, this process was repeated three times. Second, NH4/SSZ-13 was added into a 0.1 M Cu(CH3COO)2 solution to undergo Cu-ion exchange at 80 °C for 4 h under stirring. Finally, the solid was collected via filtration, washed, dried, and calcined at 550 °C for 4 h to obtain Cu/SSZ-13. Copper content of the catalyst was determined to be 1.40 wt % by inductively coupled plasmaoptical emission spectrometry (ICP-OES). This yields a Cu/Al ratio of ∼0.14. According to a theoretical Cu site compositional phase diagram versus Si/Al and Cu/Al ratios,12 this catalyst should contain predominately Z2Cu and minimal ZCuOH active sites. Furthermore, such ratios indicate that ∼70% of the Brønsted acid sites remain unchanged. Steady-state SCR reaction tests were performed in a fixedbed reaction system.26 First, 60 mg of the catalyst was pressed, crushed and sieved into 40−60 mesh sizes and placed inside a quartz tube with a 6 mm inner diameter. The gaseous reactant mixture consisted of 500 ppm of NH3, 500 ppm of NO, 5% O2 and balanced with N2. Water vapor (5%) was bubbled into the gas mixture along with the nitrogen carrier gas. The total flow rate was 300 mL/min. The estimated gaseous hourly space velocity was 100 000 h−1, and NOx conversions and N2 selectivity were calculated by following equations: NOx conversion =

[NO + NO2 ]inlet − [NO + NO2 ]outlet × 100% [NO + NO2 ]inlet

2[N2O] jij zyz jj1 − zz × 100% j z [ ] + [ ] − [ ] − [ ] NO NH NO NH 3 inlet 3 outlet { x inlet x outlet k

N2 selectivity =

In situ DRIFT spectra were obtained with a Nicolet 6700 FT-IR spectrometer, equipped with an MCT detector and a Harrick reaction cell with ZnSe windows. The catalyst powder (20−30 mg) was loosely packed above the copper grid of the sample holder to form a smooth flat surface within the cell. The mixed gases were passed through the catalyst from top to bottom. Prior to the tests, the sample was pretreated with 5% O2/N2 (100 mL/min) at 500 °C for 30 min. The sample was then cooled in the same flow to a few target temperatures and maintained at each one for 15 min for stabilization before background spectra acquisition. Reaction gases (100 mL/min, compositions detailed below) were then switched on, and the DRIFT spectra were recorded from 4000 to 400 cm−1 by accumulating 5 scans with a resolution of 4 cm−1. The time 6138

DOI: 10.1021/acscatal.9b00759 ACS Catal. 2019, 9, 6137−6145

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higher than 95% are readily achieved for the catalyst studied here. As also shown in the figure, N2 selectivity values near 100% are achieved at all temperatures tested. These results collectively indicate that (1) this catalyst is essentially free of CuOx cluster species that are active for the NH3 oxidation side reaction and that (2) isolated Cu(II) active sites are dominated by Z2Cu rather than ZCuOH; in particular, the latter display lower SCR selectivity than the former above ∼400 °C.35 These conclusions are also in line with the low Cu/Al ratio of ∼0.14 described above. Therefore, this catalyst is highly suitable for the fundamental studies to be described below, especially with respect to Z2Cu active species. The rest of the study focuses on the results of in situ DRIFTS experiments. At 250 °C, 500 ppm of NH3 was introduced to the Cu/SSZ13 until saturation as monitored with continuous DRIFT scans. The catalyst was then purged with N2 (100 mL/min) until stabilization of the DRIFT spectra. At this point, 500 ppm of NO/O2 (100 mL/min) was continuously introduced, and time series DRIFT spectra were acquired until NH 3 consumption was complete. Figure 2 presents five representa-

series spectra were obtained continuously with a scanner velocity of 0.9 s/spectrum. All infrared spectra were collected in the absence of water in the reactant feed because of its strong interference effects. It is assumed here that H2O, at most, only alters the rates but not the nature of the SCR reactions, and therefore, its absence in the feed does not affect the key conclusions of our DRIFTS investigations.27,28 To avoid overlapping of certain key IR bands, deuterium-labeled ND3 was used to replace NH3 in some measurements. ND3/N2 with 1% concentration was prepared by diluting pure ND3 with N2. In order to explore the dynamic changes of adsorbed species, IR band areas were integrated according to assignments given in Table 1 in order to monitor their variations as a Table 1. Integral Range and Assignments of Surface Species and Framework Vibrations Present in the Spectra Obtained during In Situ DRIFTS Experiments integral range (cm−1) 1640−1600 1500−1420 1650−1500 1280−1170 970−1070 1410−1480 890−910 935−960 2380−2480

assignments

reference

δas(N−H) of L-NH3 δas(N−H) of B-NH3 Cu-NO3− bidentate nitrate bridging nitrate free nitrate asymmetric TOT vibrations associated with Z2Cu asymmetric TOT vibrations associated with ZCuOH ND3

5 5 11,29 30 30 31 32,33 32,33 34

function of reaction time. It is important to note that, during the course of this study, these integrated areas are assumed to change linearly with concentrations of the surface species that they represent (i.e., obeying Beer’s law). However, this assumption is not strictly proved here and, therefore, should be treated only as an approximation.

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Figure 2. Selected DRIFT spectra acquired at different stages during titration of an NH3-saturated Cu/SSZ-13 catalyst by NO (500 ppm) + O2 (5%) at 250 °C.

RESULTS AND DISCUSSION Figure 1a presents NOx conversions as a function of temperature for standard SCR measured at a GHSV of 100 000 h−1. Between 250 and 550 °C, NOx conversions

Figure 1. (a) NH3−SCR activity and N2 selectivity of the Cu/SSZ-13 catalyst as a function of temperature in the range of 150−550 °C. (b) N2O yields during NH3−SCR as a function of temperature. The reaction gases consisted of 500 ppm of NH3, 500 ppm of NO, 5% O2 and 5% H2O. The total flow rate was 300 mL/min. 6139

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Upon NO+O2 admission, L-NH3 and B-NH3 are consumed; the surface chemistries are first described qualitatively. At 377s (second spectrum from the top), intensities of the L-NH3 feature at 1620 cm−1 and TOT vibrations at ∼950/900 cm−1 decrease as compared to those at time zero. Such changes demonstrate L-NH3 consumption. However, intensities of the B-NH3 feature at 1460 cm−1 and negative ν(OH) vibrations at ∼3600/3580 cm−1 remain invariant, indicating the absence of B-NH3 consumption. Note that similar titration studies by others also discovered earlier consumption of L-NH3 than BNH 3 . 19,20 At time 787s (the middle spectrum), the disappearance of TOT vibrations at ∼950/900 cm−1 indicates complete consumption of L-NH3. The substantially decreased intensities of the B-NH3 feature at 1460 cm−1 and negative ν(OH) vibrations at ∼3600/3580 cm−1 also demonstrate that B-NH3 has been partially consumed at this point. At 1198 s (second bottom spectrum), B-NH3 becomes essentially fully consumed. Interestingly, however, the ∼1620 cm−1 band does not completely disappear after complete consumption of LNH3; a weak band persists at ∼1624 cm−1 in the bottom three spectra. This new feature is assigned to a Cu-NO3− species,11,29 demonstrating that Cu active sites that are incompletely coordinated with NH3 ligands interact with NO+O2 to form surface nitrates in the presence of NO+O2. Note that the bottom two spectra also contain an extremely weak feature at ∼3650 cm−1. This band can be attributed to a ν(OH) vibration in ZCuOH.12,39,41 The detection of ZCuOH indicates that not all Cu sites interact with NO+O2 to form surface nitrates under this condition. The bottom two spectra also contain a band at 910 cm−1; the nature of this band is not clear to us. As such, we only tentatively assign this to a framework band. To obtain more quantitative information on the evolution of L-NH3 and B-NH3 during NO titration, their peak areas are integrated and ratioed against those at time zero and plotted in Figure 3a as a function of reaction time. The most obvious conclusion drawn here, as has been pointed out previously, is that L-NH3 has higher reactivity than B-NH3.19,20 For example, at a reaction time of ∼9 min where ∼80% L-NH3 has been consumed, B-NH3 consumption is barely initiated. Likewise, at a reaction time of ∼13 min where L-NH3 has been exhausted,

tive spectra collected at different stages of reaction. A more complete set of spectra with narrower time intervals is shown in Figure S1. At time zero (top spectrum, right before NO introduction), FTIR bands for both B-NH3 (3220, 3185, 1670, 1460 cm−1, marked with “B”) and L-NH3 (3380−3330, 3280, 1620 cm−1, marked with “L”) are detected.19,20,28,36 Among these, δas(N−H) of L-NH3 at 1620 cm−1 and δas(N−H) of BNH3 at 1460 cm−1 are well resolved and can be readily used for intensity comparisons, with the intensity ratio of the two features at time zero being ∼0.2. On the basis of three assumptions, notably that (1) each Brønsted acid site chemisorbs one NH3 molecule as NH4+; (2) L-NH3 is dominated by Cu(II)-NH3, with other Lewis acid sites, including extra-framework Al (EFAl), being too weak to chemisorb NH3 at the adsorption temperature of 250 °C; and (3) L-NH3 and B-NH3 have similar δas(N−H) sampling sensitivity factors, it is estimated that each Cu(II) site chemisorbs one NH3 molecule based on the catalyst composition (Cu/Al = 0.14). However, the third assumption above is likely not valid; even without considering geometry and symmetry differences of L-NH3 and B-NH3, B-NH3 has one more N−H bond than L-NH3. If, based simply on the number of N−H bonds, B-NH3 has a higher δas(N−H) sampling sensitivity factor, then each Cu(II) site chemisorbs more than one NH3 molecules. In a prior quantitative NH3TPD investigation, Luo et al.37 demonstrated that approximately two NH3 molecules adsorb on each Z2Cu active site and one NH3 molecule adsorbs on each ZCuOH active site at a temperature of 150 °C. While in the present study, the NH3 adsorption temperature was set at 250 °C, it should be noted that the thermocouple is placed beneath the sample (separated by a layer of quartz wool) in our DRIFTS cell design, so that the actual sample temperature may be somewhat lower. In any case, although we are not in a position to precisely quantify the numbers of NH3 molecules that adsorb on Z2Cu and ZCuOH sites, a good approximation is that 1−2 NH3 adsorb on Z2Cu and 0−1 on ZCuOH right before our NO titration experiments. Based on the signal intensity ratio and approximations discussed above, it is also estimated that BNH3 concentration is ∼3−6 times higher than L-NH3 at time zero. We note that NH3-TPD was also attempted here. Unfortunately at an ammonia adsorption temperature of 250 °C, desorption from L and B acid sites are not separable, precluding more accurate determination of these ratios. NH3 adsorption causes development of several negative going IR bands. In particular, bands at ∼3600 and 3585 cm−1 are attributed to consumption of Brønsted acid sites (i.e., NH3 + H+ = NH4+).20,38 Bands at ∼950, ∼900, and ∼820 cm−1 are perturbed framework vibrations (i.e., TOT vibrations) arising from Cu ions in the close vicinity of the zeolite framework.8,32,33,39,40 These bands are generated because of Cu(II)NH3 interactions, which “pull” Cu(II) ions away from positions at which the background spectra are taken, thus inducing “negative” bands as the perturbed vibrations are eliminated. Among these vibrations, the ∼950 cm−1 band is induced by ZCuOH repositioning, and the ∼900 cm−1 band is induced by Z2Cu repositioning. Relative sampling sensitivity factors for these two bands have not been established. In qualitative terms, the ∼950 cm−1 band is much weaker than the ∼900 cm−1 band at the same ZCuOH and Z2Cu concentrations,32 which is likely due to the higher rigidity of the 6-membered ring that Z2Cu occupies than the 8-membered ring that ZCuOH occupies.

Figure 3. Time evolution of selected signals during NO titration of an NH3-equilibrated Cu/SSZ-13 at 250 °C. Relative integrated peak areas are normalized with those at time zero (At/A0); (a) L-NH3 and B-NH3, (b) ZCuOH and Z2Cu. Condition: 500 ppm of NO, O2/N2 (5 vol %). 6140

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stage L-NH3 consumption phenomenon is no longer seen at 350 °C. The most likely explanation is that Cu sites with multiple NH3 ligands are no longer present at such a high NH3 adsorption temperature. In Figure 3b, for reaction times ≤ ∼8 min, the negativegoing peak at ∼950 cm−1 induced by ZCuOH repositioning appears to start to diminish somewhat earlier than that induced by Z2Cu repositioning. This may indicate that ZCuOH begins to reform before Z2Cu. However, since the ZCuOH induced TOT vibrations are quite weak (with correspondingly larger errors expected in peak area integration), and since 6- and 8membered rings have different rigidity, definitive conclusions based on analysis of the reformation of perturbed TOT vibrations may be difficult. Indeed, both negative-going FTIR peaks decay at similar rates after ∼8 min and become absent at ∼11 min; i.e., right before L-NH3 depletion. As shown in Figure 3a, B-NH3 becomes reactive at ∼11 min, right before the complete depletion of L-NH3. As demonstrated in Figure 2, Cu sites depleted of NH3 ligands interact with NO+O2 to form Cu-NO3− species. Therefore, it can be anticipated that in this final stage of titration, B-NH3 interacts with surface NOx species that are activated by Cu sites. On the basis of previous studies, two reaction pathways (i.e., R1 + R2, or R3 + R4) are likely:

B-NH3 is still clearly detectable. It is interesting to note that LNH3 consumption follows a nonlinear pattern with time. In the first ∼7 min, it reacts at a slower rate than that during the ∼7− 11 min period. The rate slows down again after ∼11 min due to depletion of L-NH3. On the basis of discussions above regarding the nature of Cu(II)-NH3 complexes, we argue that the rate difference before and after ∼7 min is due to reactivity difference of various Cu(II)-NH3 complexes; in particular, we suggest that complexes with higher NH3 coordination (e.g., [Cu(OL)2(NH3)2]2+) have reactivity lower than complexes with lower NH3 coordination (e.g., [Cu(OL)3(NH3)]2+). This argument is consistent with our current understanding about the reduction half-cycle of the redox SCR mechanism, where the Cu(II) → Cu(I) reduction half cycle proceeds with formation of NH2NO or NH4NO2 intermediates during which an electron is transferred from the reactants to Cu.10,13,14 At least two reasons can be considered for this NH3 inhibition effect: first, NH3 ligands spatially shield Cu centers from interacting with NO; and second, NH3 ligands destabilize the transition states involving electron transfer because of their electron-donating nature. We note that these postulations remain to be further verified (e.g., via theoretical simulations). The NH3 inhibition effect discussed above is well-supported by studies utilizing other techniques.17,42 The argument we made above regarding L-NH3 consumption rates before and after ∼7 min, however, may be alternatively explained by invoking geometry changes of NH3 ligands. In particular, after consumption of the first NH3 ligand bound to the Lewis site, the remaining NH3 ligand may adopt a different symmetry and, thus, have a different absorption cross-section. In this case, the faster signal intensity loss after ∼7 min may not be due to faster NH3 consumption at lower Cu-NH3 coordination. To gain more insights into this issue, Figure 3b plots changes in the negative-going TOT vibrational peaks at ∼900 and ∼950 cm−1 that are induced by Z2Cu and ZCuOH repositioning, respectively. In discussing these “negative” bands, it is again important to keep in mind that the “background” spectrum for all of these data is the oxidized catalyst before NH3 adsorption. Thus, upon exposure to saturation coverages of NH3, the perturbed TOT vibrations that result from strong interactions of Cu ions with framework oxygens (i.e., zeolite ion exchange sites) are eliminated because NH3 “solvates” the Cu ions and moves them away from these locations. This shows up as “negative” FTIR bands because, again, these spectra are ratioed to a background spectrum obtained just before NH 3 adsorption. As such, the disappearance of these “negative” bands also represents L-NH3 consumption during NO+O2 titration. Importantly, since these vibrations are framework modes of the SSZ-13 substrate, it is unlikely that geometry changes of NH3 ligands plays a role in their intensity variations. As is clearly shown in Figure 3b, TOT vibrations at ∼900 cm−1 change in a highly similar manner as the 1460 cm−1 LNH3 band; that is, it diminishes slowly in the first ∼7 min of NO+O2 titration but speeds up thereafter. This provides rather compelling evidence supporting our hypothesis above: Cu(II)NH3 complexes with higher NH3 coordination are indeed less reactive than complexes with lower coordination. Moreover, this argument is further corroborated by similar titration experiments carried out at 350 °C. Figure S2a displays selected time series spectra for NO titration of an NH3-saturated catalyst, and Figure S2b compares normalized L-NH3 and Z2Cu signals as a function of time for measurements at 250 and 350 °C. In these data, it is clearly evident that the two-

Cu−NO−3 + NH+4 → Cu 2 + + NH4NO3

(R1)

NH4NO3 + NO → NH4NO2 + NO2 → N2 + 2H 2O + NO2

(R2)

Cu−NO−3 + NO → Cu−NO−2 + NO2

(R3)

Cu−NO−2 + NH+4 → NH4NO2 → N2 + 2H 2O

(R4)

It is difficult to determine which of the two pathways dominate under our titration conditions. R2 has long been suggested as the rate-limiting step for fast SCR (i.e., NO + NO2 + 2NH3 = 2N2 + 3H2O) over zeolite and/or oxide-based SCR catalysts.43−46 Below ∼200 °C, NH4NO3 is known to be rather stable in small-pore zeolites.25,47,48 However, since R2 is strongly activated on Cu/SSZ-13,49 it may become fast enough to account for the rapid NH4+ consumption at 250 °C as shown in Figure 3a. On the other hand, R3 is known in proceed rapidly in the absence of NH3/NH4+;50 however, NO3− may preferentially react with NH3/NH4+ when they are present. To gain more insight into these possible pathways, it is useful to calculate the NH3 consumption rates. As shown in Figure 3a, the fast linear L-NH3 consumption period at 7−11 min comprises ∼57% of L-NH3 signal loss. This translates to an NH3 consumption rate of 3.2 × 10−3 s−1 assuming an NH3/ Cu ratio of 1 at time zero, or 6.4 × 10−3 s−1, and assuming an initial NH3/Cu ratio of 2. It is also shown from Figure 3a that B-NH3 intensity disappears between ∼11 and 15 min at a similar rate as L-NH3 between ∼7 and 11 min. Since B-NH3 concentrations at time zero are ∼3−6 times higher than LNH3, it can be estimated that the B-NH3 disappearance rates can reach up to ∼4 × 10−2 s−1. Recently, it has been demonstrated that R2, as the rate-limiting step in fast SCR over Cu/SSZ-13, is strongly activated with an apparent activation energy of ∼160 kJ/mol.49 By applying the Arrhenius

( )= (

expression ln

r1 r2

Ea R

1 T2

−

1 T1

), and reaction rates measured

at other temperatures, it is estimated that R2 proceeds at ∼3.8 6141

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ACS Catalysis × 10−2 s−1 at 250 °C, very similar to the titration rates estimated here. Therefore, we tentatively suggest that pathway R1+R2 dominates in the process of using NO+O2 to titrate BNH3 present in Cu/SSZ-13. Results shown above clearly demonstrate an ammonia inhibition effect under our titration conditions (i.e., NH3 consumption is the slowest when multiple NH3 ligands are chemisorbed on each Cu active site); the rate increases when only one NH3 ligand is chemisorbed on each Cu and increases even further when Cu sites are depleted of NH3 ligands. In the latter case, nitrate species become kinetically relevant. However, under a low-temperature steady-state or quasisteady-state standard SCR condition, where NH3 almost always bonds to Cu active sites, it is not clear whether nitrate species also play kinetically important roles. Next, infrared transient perturbation experiments were used to probe SCR under quasi-steady-state conditions. In this case, after a steadystate reaction condition is reached at 250 °C, the NO/O2 flow is turned on and off alternatively every 2 min (i.e., 4 min frequency for 1 full cycle) to perturb the reaction. As Figure 3a demonstrates, the first 2 min of NO titration of an NH3saturated catalyst only causes conversion of a small portion of adsorbed NH3. This indicates that in the transient perturbation experiments described as follows, such intervals are short enough that steady-state is only slightly perturbed. Our rationale regarding the efficacy of such pulsed experiments is that, even when the nitrate signals are weak and not resolved, if these species are kinetically relevant, by integrating the frequency ranges where these species appear, we may still be able to observe variations of the signal intensities with NO pulsing. Thus, the invariant signal intensities in these regions either suggest that (1) these nitrates are not present, or that (2) they are present but kinetically irrelevant. For this reaction system, possible surface NO3− species include Cu(II)-NO3−, free nitrate ion (i.e., NO3− in NH4NO3), bidentate nitrate, and bridging nitrate on support, with characteristic vibrations at 1500−1650, 1410−1480, 1170−1290, and 1000−1050 cm−1, respectively, according to literature assignments (Table 1). From N2O formation shown in Figure 1b, at least NH4NO3 is present under such measurements. It is readily seen from these vibrations, however, that Cu(II)-NO3− overlaps with L-NH3 (as is clearly shown in Figure 2, both species appear at ∼1620 cm−1) and free nitrate ion overlaps with B-NH3 (both appear at ∼1460 cm−1). Under our quasi-steady-state measurements, such nitrate signals may be readily overwhelmed by the much stronger ammonia signals. In order to overcome this problem, ND3 was used next to replace NH3. As shown in Figure S3, N− H vibrations occur at frequencies ∼1.37 times higher than N− D vibrations, sufficient to separate the overlapped ammonia/ ammonium ion and nitrate vibrations. Figure S4 presents time-resolved DRIFT spectra of the pulsed experiments, and Figure 4 presents time evolution of the selected signals. In this case, intensity of the 2380−2480 cm−1 band, assigned to L-ND3, changes periodically with NO pulses. This clearly demonstrates the kinetics relevance of LND3. In contrast, all possible nitrate bands do not show any variation during the course of the measurement. This strongly indicates that ammonium nitrate is kinetically insignificant under low-temperature steady-state standard SCR reaction conditions. Furthermore, results shown in Figure 4 also demonstrate that in addition to ammonium nitrate, all other possible nitrate species are equally irrelevant to standard SCR under typical low-temperature operating conditions. In

Figure 4. Time evolution of the selected DRIFT spectra during NO +O2+ND3/ND3 modulation experiments (frequency = 4 min) at 250 °C. Representation of the pulse sequence (ND3 in, NO in). Conditions: ND3 (500 ppm), NO (500 ppm), O2 (5 vol %) in N2.

particular, Cu active sites with NH3 ligands highly unlikely active NO in the same manner as “naked” Cu active sites (e.g., with the formation of NH3-Cu-NO3− intermediates). This finding corroborates a recent similar conclusion by Marberger et al.17 Finally, the only apparently kinetically relevant species during such perturbation measurements are described in more detail below. In this case, NH3 instead of ND3 is used. The time series spectra are displayed in Figure S5. As depicted in Figure 5, intensities of the 1620 cm−1 band (L-NH3) and the

Figure 5. Time evolution of the selected DRIFT spectra during NO +O2+NH3/NH3 modulation experiments at 250 °C (frequency = 4 min). Representation of the pulse sequence (NH3 in, NO+O2 in). Conditions: NH3 (500 ppm), NO (500 ppm), O2(5 vol %) in N2.

1460 cm−1 band (B-NH3) vary correspondingly with the pulsed admission of NO. In particular, the L-NH3 peak intensities decrease when NO is on and increase when it is off. This finding is fully consistent with Figure 3 where it is clearly demonstrated that NO reacts first with L-NH3. Note that the L-ND3 band shown in Figure 4 changes periodically in the same manner. Rather curiously, however, intensities for the 1460 cm−1 band increase with NO admission and decrease when NO is turned off. Since the 1460 cm−1 band is not associated with nitrates (Figure 4), we propose that the intensity increase of this band upon NO admission can be explained via the following processes

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Cu(II)−NH3 + NO → Cu(I) − NH 2NO + H+

(R5)

H+ + NH3 F NH+4

(R6) DOI: 10.1021/acscatal.9b00759 ACS Catal. 2019, 9, 6137−6145

Research Article

ACS Catalysis where R5 has been proposed by Paolucci et al.10 in their theoretical simulations, and R6 is well accepted to readily occur at 250 °C. In a recent study, Chen et al.51 also observed the formation of NH4+ intermediates resulting from the interaction of NO and NH3 on Fe-ZSM-5 catalysts using a combination of in situ electrical impedance spectroscopy and DRIFTS. Interestingly, intensities of the 1460 cm−1 band decrease when NO is turned off. This indicates a decrease of the Cu(I)/ NH4+ populations. Operando X-ray absorption near edge structure (XANES) studies have demonstrated, for the same catalyst and at the same reaction temperature, that the percentage of isolated Cu(I) relative to the total number of Cu ions at steady-state standard SCR is substantially higher than that at NH3 oxidation.10 Therefore, the 1460 cm−1 band intensity loss in the absence of NO in the feed found here is fully consistent with Cu(I) concentration drop found by XANES. Stoichiometrically, the chemistry can be described as follows: NH+4 + Cu(I) + O2 = Cu(II) +

1 N2 + 2H 2O 2

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AUTHOR INFORMATION

Corresponding Authors

*E-mail: [email protected]. *E-mail: [email protected]. ORCID

Yue Peng: 0000-0001-5772-3443 Junhua Li: 0000-0001-7249-0529 Feng Gao: 0000-0002-8450-3419 Charles H. F. Peden: 0000-0001-6754-9928 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was financially supported by the national key research and development program of China (2017YFC0211300 and 2016YFE0126600) and the China Postdoctoral Science Foundation (2017M620800). FG and CHFP gratefully acknowledge financial support for their participation from the US Department of Energy (DOE), Energy Efficiency and Renewable Energy, Vehicle Technologies Office. A portion of this work was performed in the Environmental Molecular Sciences Laboratory (EMSL), a national scientific user facility sponsored by the DOE’s Office of Biological and Environmental Research and located at Pacific Northwest National Laboratory (PNNL). PNNL is operated for the US DOE by Battelle.

(R7)

We note that although R7 is both charge- and massbalanced, this reaction is clearly not elementary. It is beyond the ability of in situ DRIFTS to provide further details.

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CONCLUSIONS The results show that additional details of the SCR reaction mechanism on a copper-exchanged CHA zeolite can be probed by time-resolved spectroscopic methods. In this work, using a Cu/SSZ-13 catalyst that contains active Cu sites primarily as Z2Cu, adsorbed NH3 species and the active NH3 sites participating in SCR reactions are clearly distinguished. LNH3 reacts with nitrogen oxides directly, presumably via a NO-assisted NH3 activation pathway as proposed previously. B-NH3, on the other hand, is much less active under lowtemperature steady-state conditions. Under such conditions, NH4NO3 is kinetically insignificant; other possible surface nitrate species are also kinetically insignificant and, in fact, may not even exist. However, when depleted of NH3 ligands, Cu active sites efficiently activate NO+O2 to form Cu-nitrates, and the latter species interacts with B-NH3 rather efficiently at 250 °C. For practical applications, this finding indicates that ammonia dosing strategies for steady-state and nonsteady-state operations should be different in order to optimize NOx removal efficiencies.

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spectra during NO+NH3+O2/NH3+O2 modulation experiments at 250 °C (T = 4 min) (PDF)

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REFERENCES

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acscatal.9b00759. Transient infrared spectra for NO titration of NH3saturated Cu/SSZ-13 at 250 °C; S2(a):Transient infrared spectra for NO titration of NH3-saturated Cu/ SSZ 13 at 350 °C; S2(b):The time evolution of L-NH3 and Z2Cu FTIR features during NO titration of NH3saturated Cu/SSZ-13 at 250 and 350 °C; S3: Infrared spectra obtained after NH3 (ND3) adsorption onto Cu/ SSZ-13 to saturation; S4: Time resolved DRIFT spectra during NO+ND3+O2/ND3+O2 modulation experiments at 250 °C (T = 4 min); S5: Time resolved DRIFT 6143

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ACS Catalysis

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